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IL FOUNDATION SERIES

CHEMISTRY

A Reliable Companion for JEE | NEET | Olympiads

IL Foundation Series - Chemistry Class 9 Legal Disclaimer This book is intended for educational purposes only. The information contained herein is provided on an “as-is” and “as-available” basis without any representations or warranties, express or implied. The authors (including any affiliated organizations) and publishers make no representations or warranties in relation to the accuracy, completeness, or suitability of the information contained in this book for any purpose. The authors (including any affiliated organizations) and publishers of the book have made reasonable efforts to ensure the accuracy and completeness of the content and information contained in this book. However, the authors (including any affiliated organizations) and publishers make no warranties or representations regarding the accuracy, completeness, or suitability for any purpose of the information contained in this book, including without limitation, any implied warranties of merchantability and fitness for a particular purpose, and non-infringement. The authors (including any affiliated organizations) and publishers disclaim any liability or responsibility for any errors, omissions, or inaccuracies in the content or information provided in this book. This book does not constitute legal, professional, or academic advice, and readers are encouraged to seek appropriate professional and academic advice before making any decisions based on the information contained in this book. The authors (including any affiliated organizations) and publishers disclaim any liability or responsibility for any decisions made based on the information provided in this book. The authors (including any affiliated organizations) and publishers disclaim any and all liability, loss, or risk incurred as a consequence, directly or indirectly, of the use and/or application of any of the contents or information contained in this book. The inclusion of any references or links to external sources does not imply endorsement or validation by the authors (including any affiliated organizations) and publishers of the same. All trademarks, service marks, trade names, and product names mentioned in this book are the property of their respective owners and are used for identification purposes only. No part of this publication may be reproduced, stored, or transmitted in any form or by any means, including without limitation, electronic, mechanical, photocopying, recording, or otherwise, without the prior written permission of the authors (including any affiliated organizations) and publishers. The authors (including any affiliated organizations) and publishers shall make commercially reasonable efforts to rectify any errors or omissions in the future editions of the book that may be brought to their notice from time to time. Subject to Hyderabad jurisdiction only. Copyright © 2025 Rankguru Technology Solutions Private Limited. All rights reserved. ISBN 978-81-985385-5-0 Second Edition

Contents 1.

Matter in our Surroundings

2. Is Matter Around Us Pure?

3. Atoms and Molecules

4. Structure of the Atom

112

5. Periodic Classification of Elements

160

1.1

MATTER IN OUR SURROUNDINGS

INTRODUCTION TO MATTER AROUND US

As we look at our surroundings, we see a large variety of things with different shapes, sizes, and textures. Everything in the universe can be classified as either matter or energy. The air we breathe, the food we eat, stones, clouds, stars, plants, and animals, even a small drop of water or a particle of sand - everything is matter.

Fig. 1.1 Different types of matter around us

All the things mentioned above occupy space and have mass. Early Indian philosophers classified matter in the form of five basic elements - the 'Panch Tatva' - air, earth, fire, sky, and water. According to them, everything living or non-living is made up of these five basic elements. Things that have mass and occupy space are called matter, or things that have both mass and volume are called matter.

MATTER IN OUR SURROUNDINGS

Matter is made up of particles. All matter can be broken up to get very small particles. Hence, we now conclude that all matter is made up of small particles. Modern-day scientists have evolved two types of classification of matter based on their physical properties and chemical nature. 1.1.1 Classification of matter Universe

Energy

Matter

Physical classification

Solids

Liquids

Gases

Chemical classification

Plasma

Pure Substances

Bose-Einstein condensate

Crystalline Amorphous

Elements

Metals

Non-Metals

Metalloids

Compounds

Mixtures

Homogeneous (True solutions)

Noble Gases

Organic

Acids

Heterogeneous

Colloids

Suspensions

Inorganic

Bases

Salts

Fig. 1.2 Classification of matter

Energy: Anything that is not matter and has the capacity to do work is called energy. For example, heat, light, electricity, and sound do not possess mass. So, they are not matter. However, they have the capacity to do work. Heat causes water to boil; light helps plants prepare their food from carbon dioxide and water (photosynthesis); electricity makes fans revolve and trains run; and sound causes your eardrum to vibrate so that you can hear. Thus, heat, light, electricity, and sound are forms of energy.

Sound

Heat

Light

Fig. 1.3 Forms of energy

Electricity

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1.2 PHYSICAL NATURE OF MATTER Matter is composed of small particles. Every matter is made up of certain particles which differ in shape, size, and nature from other types of matter. The particles which constitute a certain matter are so small that we cannot see them individually with the naked eye. What we see is an aggregate of small or tiny particles. Generally, matter is classified into six types: solid, liquid, gas, plasma, and Bose-Einstein condensate.

Liquid

Solid

Plasma

Gas

Bose-Einstein condensate Fig. 1.4 Different states of matter

Tiny particles of matter are so small that we can't see them with our eyes. We need super-strong microscopes to see just how tiny they are. These particles are way smaller than anything we come across in our daily lives. Even a speck of sand or a tiny drop of water is made up of a huge number of these little particles. Scientists call them atoms and molecules; they're like the building blocks for everything around us.

MATTER IN OUR SURROUNDINGS

1.3 CHARACTERISTICS OF PARTICLES OF MATTER 1.3.1 Space between particles

Gas

Liquid

Solid

Fig. 1.5 Space between the particles in solid, liquid, and gaseous states of matter

Particles of matter have some space between them. For example, when we prepare tea, coffee, or lemonade, particles from one substance go into the spaces between particles of the other. This demonstrates that there is ample space between the particles of matter. Order of particle spacing: Gas > Liquid > Solid 1.3.2 Motion in particles Particles of matter are always moving, which is known as kinetic energy. When the temperature goes up, the particles move at a faster rate. This means that as the temperature increases, the kinetic energy of the particles also increases. Types of motion in particles

1) Translatory motion: It occurs when particles move in straight lines and change direction without losing energy after interacting with another particle or the container's wall. When compared to liquids, translational motion is greatest in gases and least in solids. 2) Rotational motion: When particles travel about their own axis, this is known as rotational motion. This motion is comparable to the earth's rotation around its axis. In gases and liquids, the rotational motion will be quite high. 3) Vibrational motion: When particles move back and forth around a central point. Solids have the greatest amount of motion because the particles are held in a hard framework.

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1.3.3 Force of attraction in particles Particles of matter attract each other. This force of attraction varies in strength depending on the type of matter involved. Force of attraction in particles: Solid > Liquid > gas

Solid

Liquid

Gas

Kinetic Energy Particle Motions Force of Attraction Between Particles

Fig. 1.6 Force of attraction and particle motion in solid, liquid, and gas Types of forces of attraction in different states of matter

1) Gaseous state: The forces of attraction are very weak, and practically, the molecules move almost independently to their neighbour. 2) Liquid state: They have Van-Der-Waal forces of considerable strength or H-bonds 3) Solid state: The forces are strong, and they may be bonds like ionic, covalent, metallic bonds or even H-bonds. Sometimes, they are strong Van-Der-Waal forces (London dispersion forces, dipole-dipole forces) Note: The strength of these forces changes the physical state and physical properties of substances but not the chemical properties. 1.3.4 Diffusion in particles Molecules of gas are in constant motion. When two or more gases are placed in contact, their molecules mix spontaneously until a homogeneous mixture is formed between them. The mixing of gases occurs even against the gravitational force. The mixing of gases is an irreversible process. The spontaneous intermixing of two or more gases in all directions is called diffusion. Diffusion is an important characteristic property of gases. This phenomenon allows us to easily detect the fragrance of a flower or bad rotten egg smell of hydrogen sulphide. Gases from the air mix and dissolve in water. These gases, like oxygen and carbon dioxide, are crucial for the survival of underwater animals and plants.

MATTER IN OUR SURROUNDINGS

Every living thing needs to breathe to stay alive. Aquatic animals can breathe underwater because there's oxygen dissolved in the water. So, we can say that solids, liquids, and gases can all mix with liquids. The speed at which liquids mix is faster than solids. That's because, in liquids, the particles move around freely and have more space between them compared to particles in solids. Rate of diffusion: The amount of gas diffused in one second is called the rate of diffusion. The rate of diffusion is equal to the distance travelled by a gas through a tube of uniform cross-section in one second. A decrease in the pressure of a cylinder in one second is also called the rate of diffusion. The number of moles of the gas diffused in one unit of time is also called the rate of diffusion. The rate of diffusion/effusion depends upon several factors, i.e., molecular mass, temperature, pressure, and concentration. The rate of diffusion of a gas increases with an increase in the temperature, pressure and velocity of the gas.

High concentration

Low concentration

Fig. 1.7 Diffusion of particles from higher concentration to lower concentration

The rate of diffusion (r) is given as: r=

V t

, r=

n t

l t

Where V is the volume of gas diffused, t is the time taken for diffusion, n is the number of moles of the gas diffused in the time t and l is the distance travelled by the gas in time t. Units of the rate of diffusion

c.c.s-1 (Cubic centimetres per second) (or) l.s-1 (Litres per second) (or) mol.s-1 (moles per second) Applications: The phenomenon of diffusion has some important applications, which include: • Smell, odour, perfume, etc. are based on diffusion. • Ansil alarms in coal mines for the detection of marsh gas are based on diffusion. • Differences in the rate of diffusion are taken advantage of for the separation of components of the gaseous mixture. • Relative comparison of the rates of diffusion is useful in calculating molecular weights and density of gases and vapours. 6

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Other characteristics of particles of matter

1) Rigidity: The property by virtue of which a substance can retain its shape when force is applied to it. Solids have the property of rigidity. Whereas liquids and gases are not rigid. 2) Compressibility: The particles of matter have spaces between them, and they can be compressed by applying force (or) pressure. Gases have high compressibility due to large empty spaces between the molecules of the gas. Solids and liquids are not compressed. The gas we use at home for cooking in cylinders or the oxygen provided to hospitals also comes in cylinders, and it's called compressed gas. Nowadays, vehicles use something called compressed natural gas (CNG) as fuel. Because it can be squeezed a lot, we can fit a large amount of gas into a small cylinder, making it easy to transport. 3) Fluidity: Fluidity is the tendency of a substance to flow. Liquids and gases possess fluidity, whereas solids do not flow. 4) Filling a gas container: Gases have less force of attraction between their molecules, and therefore, they move in all directions with very high velocities and fill the container. 5) Shape: The definite geometrical arrangement of constituents of a substance determines its shape. Solids have definite shapes, whereas liquids acquire the shape of the container; gases also do not have a definite shape. 6) Kinetic energy: The energy possessed by the particles due to their motion is called kinetic energy. Gas molecules possess maximum kinetic energy, and solid molecules possess the least kinetic energy. The increasing order of kinetic energy is: Solids < Liquids < Gases. 7) Volume: Matter must occupy space. The space occupied by matter is called volume. 8) Mass: It represents the quantity of matter, which can be found by physical balance. 9) Density: The mass of the substance per unit volume is called density. Density =

Mass Volume

The density in different states decreases as Solids > Liquids > Gases Exception: Liquids generally have lower density than solids. However, ice has a lower density than water because it has an open cage-like structure. In this structure, many vacant spaces are left when water molecules are linked to ice. Because of the lower density of ice, it floats on water. 10) Weight: It represents the pull of gravity on matter. It can be measured with a spring balance. 11) Perception by physical senses: It is perceived by our physical senses, i.e., sense of touch, sense of hearing, sense of sight, sense of smell, and sense of taste.

MATTER IN OUR SURROUNDINGS

1.4 STATES OF MATTER When we look at the matter around us, we find it is in three forms. These are solid, liquid and gas. All human beings are outwardly solids. The water we drink daily is a liquid, while the air we breathe-in represents a gas. The three states of matter we have mentioned are because of the interparticle spaces in different kinds of matter. The state of matter refers to its nature or appearance. The properties in the bulks of these aggregates of molecules depend upon the strength of force holding these constituting units together, and roughly, there are three states of matter that make up matter, viz. solid, liquid and gas. A substance may attain either of these physical states under suitable conditions. Let us briefly discuss the characteristics of the three states of matter we have listed. 1.4.1 The solid state

Fig. 1.8 Arrangement of particles in the solid state

Can a rubber band change its shape when you stretch it? Is it considered a solid? What happens with sugar and salt when you put them in different jars? Do they take the shape of the jar, and does that mean they are solids? Think about a sponge. It's solid, but you can squeeze it. Why is that? A rubber band can stretch and take on different shapes when you pull it. Once you stop pulling, it goes back to its original shape. But be careful! If you pull too hard, it might break. Each tiny crystal of sugar or salt keeps its shape. It doesn't matter if you hold it in your hand, put it on a plate, or keep it in a jar-it stays the same shape. Imagine a sponge with tiny holes where air is trapped. When you press the sponge, the air comes out, and that's how we can squeeze and compress it. Solids are known for their hardness and rigid nature with definite shape and volume. The important characteristics of the solid state are: • Solids have definite shape and volume. • Solids are generally rigid. If some solid (such as rubber) changes its shape on the application of external force, then it regains its shape, upon the removal of force. • Solids can have any number of free surfaces.

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• The intermolecular spaces are very small. • The intermolecular forces are very large. • The dimensions of solids do not increase in large proportion to heating or cooling. • When two solids are kept in contact with one another, they do not mix with each other, i.e., diffuse. • The density of solids is generally high. Mass occupied by a solid per unit volume is obtained by dividing the mass of a particular solid by the volume occupied by that mass of the solid. The unit of density: kg/m3. Here, kg is the unit of mass while m3 is the unit of volume. • Solids can hardly be compressed by applying pressure. Examples: Rocks, stones, wood, sand, crystalline solids, metals like iron, copper, nickel, and ice, etc., are some typical examples of substances that are in the solid state. 1.4.2 The liquid state A liquid is a state of matter with definite mass and volume but no definite shape. The important characteristics of the liquid state of a substance are: • Liquids have a definite mass and volume. • Liquids do not have a definite shape but take the shape of the containing vessel. • The force of attraction between molecules of the liquids is less than that of solids. Thus, the liquids can flow.

Fig. 1.9 Arrangement of particles in the liquid state

• The intermolecular space in liquids is larger than in solids. Thus, liquids are slightly more compressible than solids. • Liquids have only one free surface. • The density of the liquids is relatively less than that of solids. • Liquids expand far more than solids on heating and contract far more on cooling.

MATTER IN OUR SURROUNDINGS

• The particles of two different liquids can diffuse in one another, depending upon the nature of the molecules of the liquids. • Milk and water particles diffuse in one another, but the particles of oil and water do not. • Liquids have fluidity and not rigidity. Examples: Water, milk, kerosene, petrol, alcohol, benzene, etc., are examples of the substances which exist in the liquid state. 1.4.3 The gaseous state The gaseous state is the simplest among the three states of matter. Throughout our lives, we remain immersed in the ocean of air, a mixture of gases. Nitrogen and oxygen account for more than 99% of the volume of dry air. The remaining 1% is largely argon with trace amounts of CO2, Ne, He, CH4, Kr, etc. We spend our life in the lowermost layer of the atmosphere (Troposphere), held to the surface of the Earth by gravitational force. When molecular interactions are weak, molecules do not cling together and are present in the gaseous state.

Fig. 1.10 Arrangement of particles in the gaseous state

Out of the three states of matter, the interparticle spaces are the maximum in the gaseous state. The interparticle forces which hold together different particles are minimal in the gaseous state. As a result, rigidity is the minimum, while fluidity is the maximum. A gas is a state of matter that has a definite mass but no definite shape or volume. The important properties of the gaseous state

• A gas contained in a vessel has a definite mass. • A gas can occupy the entire space of a given vessel in which it is enclosed. • Intermolecular spaces are very, very large as compared to solids and liquids. It is due to this reason that gases are highly compressible. • Intermolecular forces are negligible. It is due to this reason that they can fill the entire space. • Gases have no free surface. • Gases largely expand when heated. They largely contract when cooled. 10

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• The gases diffuse in one another rapidly to form a homogeneous mixture. This is due to large intermolecular spaces. • The density of the gases is extremely low compared to solids and liquids. • Gases do not keep their volume and are highly compressible. • The kinetic energy of the particles in the gaseous state is very high. • Gases exert pressure on their surroundings. Examples: Air is a common example of the gaseous state. It is a mixture of a number of gases, like nitrogen, oxygen, carbon dioxide, inert gases, etc. A few other examples are hydrogen, ammonia, sulphur dioxide, chlorine, etc. S. No

Property

Solids

Liquids

Gases

Shape

Have definite shape

Have no definite shape

Volume

Have definite volume

Can fill any available volume

Density

Have very high density

Have high density

Have a low density

Fluidity

Not fluid

Fluid

Diffusion

Almost none or very low

Diffusion is considerable

Diffusion is very rapid

Compressibility

Nearly incompressible

Only slightly compressible

Very compressible

Table 1.1 Important characteristics of solids, liquids, and gases

So far, we have studied solid, liquid and gas as the three states of matter. Scientists have discovered two more states of matter. These are Plasma and Bose-Einstein condensate. 1.4.4 Plasma state of matter We all know that the sun glows during the day while stars shine during the night. This is because of plasma is a mixture of free electrons and ions. Inside the sun and the stars, the temperature is very high. As a result, the atoms break, releasing electrons, and the residual particles are called ions, which carry a positive charge. The mixture of highly energetic electrons and ions is known as plasma and is responsible for the glow or shine. The fluorescent tubes and the neon sign bulbs also glow due to plasma. In the fluorescent tube, there is helium or some other gas, while neon is present in neon sign bulbs. As electricity flows through, the atoms of these gases break into charged electrons and ions. They constitute plasma, which glows. 11

MATTER IN OUR SURROUNDINGS

1.4.5 Bose-Einstein condensate The main work for the discovery of the fifth state of matter was done by Indian physicist Satyendra Nath Bose and well-known scientist Albert Einstein. But this was shown to exist when three American scientists Eric A Cornell, Wolfgang Ketterle, and Carl Wieman cooled certain gases of extremely low density to very low temperature, known as super low temperature. The particles which constitute this state of matter are often called Bose-Einstein condensate(BEC). All three scientists were awarded the Nobel Prize for their achievement.

1.5 CAN MATTER CHANGE ITS STATE? We have seen that matter exists in three different states. For a given substance, its state of matter is not permanent, i.e., a given state of matter can always be changed to other states of matter by altering the conditions of temperature and pressure. The phenomenon of change of matter from one state to another state and back to the original state by altering the conditions of temperature and pressure, etc., is called the interconversion of matter.

Solid n tio

sio

ifi

lid

tio

ion

sit

bli

Su De

Liquid

Vaporisation

Condensation

Gas

Fig. 1.11 Interconversion of matter

The obvious question that strikes everybody's mind is how to bring about a change of state. There are two ways to achieve this: by changing the temperature and by changing the pressure. Let us briefly study the effect of both of these factors. 1.5.1 Effect of change of temperature To study the effect of temperature in bringing about a change in state, consider a very popular and common example of ice. Consider a block of ice at 0° C placed in a beaker and heated. It changes to a liquid. Heat the water till it boils. It slowly gets converted to vapour (gas). From this observation, the solids convert into liquids, which in turn convert to gas when they are heated.

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Scales of measuring the temperature

We all know that temperature is generally recorded by a thermometer in which mercury is used. What about the scale of measuring the temperature? There are three scales on which temperature can be measured. These are known as Celsius scale (°C), Fahrenheit scale (°F) and Kelvin scale (K) • Thermometers with Celsius scale are calibrated from 0°C to 100°C. • Thermometers with Fahrenheit scale are calibrated from 32°F to 212°F. • Kelvin scale of temperature is the S.I. scale and is very common these days. Temperature on this scale is shown by the sign K. • The different scales are related to each other as: 9 0 F = ( 0 C) + 320 5 0 K = C - 273.15 (or 273.0 for conveniencce) Example 1: The room temperature on the Celsius scale is 25°C. Convert it into the other two scales of measurement. Solution: Temperature on Kelvin scale = 25 + 273 = 298 K. 9 × 25 + 32 = 770 F . 5 Example 2: The body temperature of a normal and healthy person is 98.4°F. What is the temperature on the Celsius scale? Temperature on Fahrenheit =

Solution: 9 0 ( C) = 0 F - 320 = (98.4 - 32) = 66.40 5 5 (or) (0 C) = 66.4 36.890 C 9 Melting

Fig. 1.12 Melting of ice 13

MATTER IN OUR SURROUNDINGS

Melting or fusion: The process by which a solid changes into a liquid state by absorbing heat energy is called melting or fusion. The melting point is the constant temperature at which a solid changes into a liquid state by absorbing heat energy. S.no

Solid

Melting point

Ice

0 °C

Sodium

97 °C

Sulphur

119 °C

Lead

327 °C

Zinc

420 °C

Iron

1535 °C

Table 1.2 Melting points of some common solid substances

Latent heat of fusion: The amount of heat energy that is needed to convert one kg of a solid into the liquid state without any rise in temperature. Example: The latent heat of the fusion of ice is 335 kJ.kg-1. Boiling Boiling or vapourisation: The process in which a liquid changes into a gaseous state by absorbing heat energy is called boiling or vapourisation.

Fig. 1.13 Boiling of water

Boiling point: The constant temperature at which a liquid rapidly changes into a gaseous state by absorbing heat energy is called boiling point.

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S.no

Liquid

Boiling point

Water

100 °C

Ethyl alcohol

78.3 °C

Benzene

80.2 °C

Mercury

357 °C

Table 1.3 Boiling points of some common solid substances

Latent heat of vapourisation: The amount of heat energy that is needed to convert one kg of a liquid at its boiling point temperature into its vapour state without any rise in temperature. The latent heat of vapourisation of water is 226 kJkg-1. Substance

Heat of Fusion (J/gm)

Heat of Vapourisation(J/gm)

Ethyl alcohol

104

854

Mercury

272

Water

334

225

Lead

871

Silver

234

Gold

1578

Table 1.4 Latent heat of certain substances Difference between gas and vapours

Gases: Substances that ordinarily exist in the gaseous state at room temperature are called gases. Examples: Hydrogen, nitrogen, oxygen, carbon dioxide and sulphur dioxide. Vapours: The gaseous substances which are obtained by heating those substances which exist as solids or liquids at room temperature are called vapours. Examples: Water vapour, mercury vapour and iodine vapour. Water and mercury ordinarily exist in the liquid state and iodine in the solid state. It may be noted that, basically, there is no difference between gases and vapours, as both are in the gaseous state.

MATTER IN OUR SURROUNDINGS

Test of purity based on melting point and boiling point

The melting point and boiling point of a pure substance are always constant and good measures to find the purity of the substance. For Example, the melting point of pure ice is 0°C, and the boiling point of pure water is 100°C at a pressure of 76 cm of Hg. The presence of impurities generally reduces the melting point of a pure solid, whereas the boiling point of a pure liquid tends to increase. Example: • The melting point of paraffin wax is not sharp and can be between 62°C and 65°C depending on the kind of compounds present in it. • If a few spoonfuls of salt are dissolved in pure water, then its boiling point becomes more than 100°C. 1.5.2 Effect of decrease in temperature on the physical state The process can be reversed if the temperature is lowered or decreased. The gas (or vapours) will be first converted to the liquid state, and then, on further cooling the liquid state will change to the solid state. • Condensation or liquefaction: The process by which a gas changes into a liquid state by giving out heat energy is called condensation or liquefaction. • Condensation point: The constant temperature at which a gas changes into a liquid state by giving out heat energy is called the condensation point. • Freezing or solidification: The process due to which a liquid changes into the solid state by giving out heat energy is called freezing or solidification. • Freezing point: The constant temperature at which liquid changes into a solid state by giving out heat energy is called the freezing point. • The numerical values of the melting point and freezing point, boiling point, and condensation point are equal. Thus, for a given substance Melting point = Freezing point Boiling point = Condensation point.   → Liquid State ←   → Vapour State • Solid state ←

Heat

Cool

bli ma

tio n

ti o pos i De

on ati ens nd n Co tio ora ap

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STATES OF WATER

Freezing Melting

Fig. 1.14 Effect of change of temperature on states of matter

Sublimation So far, we have studied that upon heating, a solid initially changes to a liquid state and then to a gaseous state when the temperature increases. The process of changing states can be reversed when the temperature is decreased. However, there are some exceptions. Certain solids directly change to the gaseous state upon heating without passing through the liquid state. Similarly, the gas changes back to the solid state without passing through the liquid state. This is called sublimation. The sublimation may be defined as the change of solid directly into the gaseous state without passing through the liquid state and back to the solid state when the temperature is lowered. Examples: Naphthalene, camphor, iodine, ammonium chloride, and dry ice (solid carbon dioxide) are some common examples of substances which undergo sublimation. Application of sublimation The process of sublimation can be used to purify the impure samples of substances which undergo sublimation and are associated with non-volatile impurities. An impure sample of naphthalene can be purified this way. The non-volatile impurities will not change into vapours. They will remain in the dish. Pure naphthalene can thus be recovered because of sublimation. 1.5.3 Effect of change in pressure We have learnt that the three states of matter differ with respect to interparticle spaces as well as interparticle forces. We have also discussed the effect of temperature in bringing about the change in state. In addition to temperature, pressure is another factor that can cause a change in the physical state of a substance. The increase in pressure also brings the particles of the substance closer. As a result, the interparticle spaces decrease. At the same time, the interparticle forces increase. This leads to a change in the physical state. To illustrate the effect of pressure, take a gas in a cylinder and apply pressure on it by a piston. 17

MATTER IN OUR SURROUNDINGS

At low pressure, the volume of the gas, as well as interparticle spaces, are very large. Under high pressure, the gas gets compressed. The interparticle spaces become less and interparticle forces become stronger. As a result, the gaseous state may change to the liquid state. Under very high pressure, there is a further decrease in volume. The interparticle forces become so strong that the liquid state may change to the solid state. The scale of measuring the pressure

Pressure is normally expressed in atmospheres. The pressure at sea level is 1 atmosphere and is regarded as normal atmospheric pressure. Name of the unit

Symbol

Pascal

Bar

bar

Atmosphere

atm

Torr

torr

Millimetre of Mercury

mm of Hg

Value

Table 1.5 Units of pressure

Evaporation During the study of the change of state, we have seen that a liquid changes to a gaseous state either by increasing the temperature or by decreasing the pressure. This process is known as evaporation. Evaporation may be defined as the phenomenon of change of liquid from its surface to the vapour state at any temperature below the boiling point of the liquid.

Fig 1.15 Evaporation of water 18

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Factors affecting evaporation The evaporation of liquids can be increased or accelerated by the following factors. An increase in surface area: Evaporation is a surface phenomenon. If the surface area is increased, the rate of evaporation increases. Example: Spreading of wet clothes for quick drying. An increase in temperature: With the rise in temperature, a greater number of particles get enough kinetic energy to go into the vapour state. A decrease in humidity: Humidity is the amount of water vapour in the air. The air around us cannot hold more than a definite amount of water vapour at a given temperature. If the amount of water in the air is already high, the rate of evaporation decreases. An increase in wind speed: It is a common observation that clothes dry faster on a windy day. With the increase in wind speed, the particles of water vapour move away with the wind as the amount of water vapour in the surroundings decreases. Nature of the liquid: We have so far discussed the external factors which influence the extent of evaporation. Apart from these, another factor of great importance is the nature of the liquid that is evaporating. We often see that alcohol evaporates at a faster rate than water. The boiling point of alcohol (350 K) is less than that of water (373 K). This means that interparticle forces of attraction in alcohol are less than in water. Therefore, alcohol will evaporate faster than water. Thus, we conclude that the lower the boiling point of a liquid, the more its tendency to change into vapours or to evaporate. Difference between boiling and evaporation

We have studied that both evaporation and boiling represent a change of state from liquid to gas or vapours. But still, they are different in certain aspects. Boiling

Evaporation

Fig 1.16 Boiling and evaporation of water

MATTER IN OUR SURROUNDINGS

Boiling

Evaporation

Boiling occurs only when the liquid is heated.

Evaporation of a liquid takes place on its own.

Boiling takes place at a specific temperature is known as the boiling point of the liquid.

Evaporation takes place at all temperatures below its boiling point.

Boiling occurs from the surface as well as from below the surface of the liquid.

Evaporation is a surface phenomenon and occurs only from the surface of the liquid.

No cooling is caused during boiling.

Cooling is always caused during evaporation.

S.no

Table 1.6 Difference between boiling and evaporation Dry ice

Solid carbon dioxide is called dry ice. It is stored under high pressure. Solid CO2 directly converts to a gaseous state by decreasing the pressure to 1 atm without changing into a liquid state. This is the reason that solid CO2 is known as dry ice. Application of dry ice: • A small quantity of CO2 will increase the growth rate of plants. Allow the dry ice to sublime near the plants for 10-15 mins daily. •

It is used to keep freezer contents frozen during breakdowns.

•

It can be used to protect post-harvested seeds and grains.

QUICK REVIEW • Anything which occupies space (volume), has mass and can be perceived by our physical senses is called matter. Examples: Stones, the air we breathe, the food we eat, the water we drink, stars, planets, etc. • Matter is made up of small particles. • Generally, matter is classified as solid, liquid, gas, plasma, Bose-Einstein condensate, and Fermi condensate. • The matter around us exists in three states - solid, liquid and gas. • The force of attraction between the particles is maximum in solids, intermediate in liquids and minimum in gases. • The spaces between the constituent particles and the kinetic energy of the particles are minimum in the case of solids, intermediate in liquids, and maximum in gases. • The arrangement of particles is most ordered in the case of solids; in the case of liquids, layers of particles can slip and slide over each other, while for gases, there is no order; particles just move about randomly. 20

IL Foundation Series Class 9

• The states of matter are inter-convertible. The state of matter can be changed by changing temperature (or) pressure. • Interconversion of matter: The phenomenon of change of matter from one state to another and back to the original state by altering the conditions of temperature and pressure, etc. • Sublimation is the change of solid state directly to the gaseous state without going through a liquid state, and vice versa. • Boiling is a bulk phenomenon; particles from the bulk (whole) of the liquid change into a vapour state. • Evaporation is a surface phenomenon in which particles from the surface gain enough energy to overcome the forces of attraction present in the liquid and change into a vapour state. • The rate of evaporation depends on the surface area exposed to the atmosphere, the temperature, the humidity, and the wind speed. • Evaporation causes a cooling effect. • Latent heat of vaporisation is the heat energy required to change 1 kg of a liquid to gas at atmospheric pressure at its boiling point. (2230 J/g, 533 cal/g, 22300 J/kg) • Latent heat of fusion is the amount of heat energy required to change 1kg of solid into liquid at its melting point. (334 J/g, 79.7cal/g, 33600 J/kg) • Diffusion is the process by which different substances mix because of the random motion of their molecules. • Condensation is the process of changing of state of a substance from its gaseous state to a liquid state at a particular temperature. • Freezing is the process of change of matter from the liquid to the solid state, at a particular temperature. • Melting is the process of changing of a solid substance to its liquid state at a particular temperature.

WORKSHEET - 1 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER I.

Introduction to matter in our surroundings

1. Assertion (A): Anything which occupies space (volume) and has mass, can be perceived by our physical senses is called add Reason (R): Matter is made up of large particles. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. 21

MATTER IN OUR SURROUNDINGS

c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 2. Which of the following is not a matter? a. Feeling warmth

b. Smoke

c. Humidity

d. Water

3. As per the definition of matter, which of the following is not matter? a. Water

b. Light

c. Sugar

d. Oxygen

4. Matter has: a. No mass but occupies space b. Mass but occupies no space c. Mass and occupies space d. No mass and occupies no space 5. The SI unit of density is: a. Kilogram per metre cube

b. Gram per kilometre cube

c. Kilogram per kilometre cube

d. Gram per metre cube

6. The SI unit of volume is: a. Cubic metre

b. Cubic kilometre

c. Cubic litre

d. Cubic kilogram

7. In how many states is the matter classified normally?

II.

a. 1

b. 2

c. 3

d. 4

States of matter and properties of solids, liquids, and gases

1. Which of the following statements do not go with the liquid state? a. Particles are loosely packed in the liquid state. b. Fluidity is the maximum in the liquid state. c. Liquids cannot be compressed (or) slightly compressed. d. Liquids conform to the shape of the container in which they are placed. 2. The physical state of matter which can be easily compressed is: a. Solid

b. Liquid

c. Gas

d. None of these

3. Intermolecular forces of attraction are least effective in:

a. Solids

b. Gases

c. Liquids

d. Plasma

IL Foundation Series Class 9

4. The state of matter which is found to be more stable at lower temperatures: a. Solid

b. Liquid

c. Gas

d. None of these

5. A substance has neither a fixed shape nor a fixed volume; which physical state is represented by this statement? a. Solid

b. Liquid

c. Gas

d. None of these

6. Gases can exert pressure in: a. One direction

b. Three directions

c. Four directions

d. All directions

7. The ability of a substance to decrease its volume when force is applied is called: a. Expansibility

b. Compressibility

c. Diffusion

d. All

8. Which among the following substances has the highest density? a. Wood

b. Sponge

c. Coal

d. Stone

9. Gases have: a. Definite shape

b. Definite volume

c. Indefinite volume

d. None

10. Assertion (A): The force of attraction between the particles is maximum in solids, intermediate in liquids and minimum in gases. Reason (R): Liquids have highest rate of diffusion. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 11. Assertion (A): Diffusion is the process by which different substances mix as a result of the random motion of their molecules. Reason (R): The arrangement of particles is most ordered in the case of solids. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 23

MATTER IN OUR SURROUNDINGS

12. Assertion (A): Solid CO2 is called dry ice. Reason (R): The atmospheric gas can be liquefied by cooling under pressure. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 13. Assertion (A): The physical state of water at 25°C is liquid. Reason (R): The physical state of water at 0°C is liquid (or) solid. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. III. Changes in states of matter

1. Evaporation is called: a. Surface phenomenon

b. Bulk phenomenon

c. Both surface and bulk phenomenon

d. Unique phenomenon

2. During the evaporation process, the heat is: a. Absorbed

b. Evolved

c. First absorbed, then evolved

d. Initially evolved and then absorbed

3. The physical state of water at 273 K is: a. Solid

b. Liquid

c. Gas

d. Both a & b

4. The melting point of ice is: a. 273.16 K

b. 373.16 K

c. 283.16 K

d. 263.16 K

5. Boiling of a liquid takes place at: a. A fixed temperature lower than its boiling point b. A fixed temperature and normal atmospheric pressure c. A fixed temperature higher than its boiling point d. A fixed temperature and higher atmospheric pressure 6. Boiling process is a: a. Surface phenomenon

IL Foundation Series Class 9

b. Bulk phenomenon c. Both surface and bulk phenomenon d. Rare phenomenon 7. Evaporation of a liquid occurs at: a. At any temperature below boiling point. b. It’s boiling point c. All fixed temperatures lower than the boiling point d. Fusion 8. When the liquid starts boiling, further heat energy which is supplied: a. Is lost to the surrounding as such b. Increases the temperature of the liquid c. Increases the kinetic energy of the particles in the liquid d. Is absorbed as latent heat of vaporization by the liquid 9. The amount of heat energy required to convert 1 gm of ice into water is called: a. Latent heat of vaporization

b. Latent heat of fusion

c. Specific heat capacity

d. All

10. The temperature above which a gas cannot be liquified is called: a. Critical temperature

b. Critical pressure

c. Triple point

d. Boiling point

11. The relation between boiling point of liquid and pressure is: a. P = B.P.

c. P ∝ 1/(B.P.)

b. P ∝ B.P. d. All

12. Assertion (A): Sublimation is the change of a gaseous state directly to the solid state without going through the liquid state. Reason (R): The state of matter can be changed by changing temperature (or) pressure. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 13. Assertion (A): Evaporation causes a heating effect. Reason (R): The rate of evaporation depends on the surface area exposed to the atmosphere. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. 25

MATTER IN OUR SURROUNDINGS

c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 14. Assertion (A): Latent heat of vaporisation is the heat energy required to change 1kg of a liquid to gas at atmospheric pressure. Reason (R): Latent heat of fusion is the amount of heat energy required to change 1kg of solid into liquid. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 15. Assertion (A): Melting is the process of changing a solid substance to its liquid state at a particular temperature. Reason (R): Vapour is the gaseous form of a substance which normally exists as a solid (or) liquid. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 16. When a substance is undergoing a phase change, its temperature: a. Remains constant

b. Decreases

c. Increases

d. Decreases then increases

17. The latent heat of vaporization is the energy required to change a substance from: a. A solid to a liquid

b. A liquid to a solid

c. A liquid to a gas

d. A gas to a liquid

18. Which of the following statements is true regarding the latent heat of fusion? a. It is the heat required to change a liquid into a gas. b. It is the heat required to change a solid into a liquid. c. It is the heat required to raise the temperature of a substance. d. It is the heat required to change a gas into a liquid. 19. When water is heated to a temperature 'X', it gets converted into steam at temperature 'X' by a process called 'R', and when steam at temperature 'X' is cooled, it gets reconverted into liquid at the same temperature 'X' by a process called 'S'. What is the name of the energy absorbed during the process 'R'? 26

IL Foundation Series Class 9

a. Latent heat of fusion

b. Latent heat of vaporization

c. Specific heat

d. Adsorption

20. While determining the melting point of ice, the thermometer is immersed in the beaker containing crushed ice. When heating the beaker on a low flame, what would happen to the temperature? a. An increase in temperature during the melting of ice. b. A decrease in temperature during the melting of ice. c. A decrease first and then an increase in the temperature during the melting of ice. d. The temperature remains constant during the melting of ice.

WORKSHEET - 2 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER 1. Choose the correct statement from the following: a. The volume of gas expands on heating. b. Two gases cannot diffuse into each other. c. Gas is converted into solid; it is called condensation. d. Gases cannot diffuse in solids. 2. Observe the given table carefully: Substance

Melting point

-129°C

98°C

1540°C

3550°C

Choose the correct statement related to A, B, C, D: i. Substance 'A' can exist as a gas at room temperature, whereas substance 'B' is a liquid. ii. The intermolecular distance between particles of 'C' is more than 'D'. iii. The substance 'B' has less force of attraction between its particles than the particles of C. iv. Generally, the density of 'C' is more than the density of 'D'. a. i, ii, iii only

b. ii, iii, iv only

c. i, iii, iv only

d. ii, iii only

MATTER IN OUR SURROUNDINGS

3. Which of the following is not characteristic of particles of matter? a. Particles in the matter are minimal in size. b. There are no spaces between the particles in the substance. c. The particles in the matter are constantly moving. d. Particles in the matter are attracted to each other. 4. Which of the following is not a measurable property of a gas? a. Amount of gas

b. Shape of gas

c. Temperature of gas

d. Density of gas

5. How does evaporation cause cooling? a. The heat from the surroundings is absorbed by the liquid during evaporation, causing the surroundings to cool down. b. The heat from the surroundings is released by the liquid during evaporation, causing the surroundings to cool down. c. The heat from the surroundings is absorbed by the gas during evaporation, causing the surroundings to cool down. d. The heat from the surroundings is released by the gas during evaporation, causing the surroundings to cool down. 6. Arrange in the order indicated for solids, liquids, and gases: Decreasing order: space between the particles a. Solids, liquids, gases

b. Liquids, solids, gases

c. Gases, liquids, solids

d. Solids, gases, liquids

7. The state of matter in which the interparticle attraction is weak and the inter-particle space is so large that the particles become free to move randomly in the entire space is known as: a. Solid

b. Liquid

c. Gas

d. Both A and B

8. A 1° C rise in temperature is equal to a rise of: a. 1 °F

b. 9/5 °F

c. 5/9 °F

d. 33 °F

9. Which one of the following statements is not correct about the three states of matter, i.e., solid, liquid and gas? a. Molecules of a solid possess the least energy, whereas those of a gas possess the highest energy. b. The density of solids is the highest, whereas that of gases is the lowest. c. Gases like liquids possess definite volumes. 28

IL Foundation Series Class 9

d. Molecules of a solid possess vibratory motion. 10. Assertion (A): Gases are highly compressible. Reason (R): Large intermolecular space between the gas molecules. a. Both A and R are true, and R is the correct explanation for A. b. Both A and R are true, and R is not the correct explanation for A. c. A is correct, but R is incorrect. d. A is incorrect, but R is correct. 11. Statement (A): Anything which occupies space and has no mass is called matter. Statement (B): Solids can be compressed. Statement (C): The intermolecular spaces in liquids are more than in gases. a. All the statements A, B and C are correct. b. All the statements A, B and C are incorrect. c. A, B are correct, but C is incorrect. d. A, B are incorrect, but C is correct. 12. Which of the following is not matter? a. Air

b. Feeling of cold

c. Dust

d. Humidity

13. Which of the following are the main components of the universe? a. Matter

b. Energy

c. Argon

d. Both a & b

14. Which of the following are different forms of energy? a. Heat

b. Light

c. Electricity

d. All

15. The space occupied by the matter is called its: a. Mass

b. Volume

c. Weight

d. All

16. Assertion (A): Particle motion always increases with the rise in the temperature. Reason (R): Upon heating, the kinetic energy of the particles increases, and that leads to an increase in their motion. a. Both A and R are correct and R is correct explantion of A. b. Both A and R are correct, but R is not correct explantion of A. c. A is true, but R is false. d. A is false, but R is true. 29

MATTER IN OUR SURROUNDINGS

17. Statement (A): The interparticle force of attraction depends on the space present between the particles of matter. Statement (B): If the space between the particles is less, the force of attraction is more and vice versa. a. Both A and B are true. b. Both A and B are false. c. A is true, but B is false. d. A is false, but B is true. 18. Statement (A): Mass represents the quantity of matter which can be found by physical balance. Statement (B): Weight represents the pull of gravity on matter, which can be found by spring balance. Statement (C): Sense of touch, sense of hearing, and sense of sight etc are perception by physical senses. a. All the above statements are correct b. All the above statements are incorrect c. A, B are correct, and C is incorrect d. A, B are incorrect, and C is correct 19. Assertion (A): The component of the universe that is useful for doing some work is called energy. Reason (R): Heat is the form of energy, helps plants to prepare their food from CO2 and water. a. Both A and R are true and R is the correct explanation of A b. Both A and R are true and R is not correct explanation of A c. A is correct and R is incorrect d. A is incorrect and R is correct. 20. Assertion (A): Gases diffuse very rapidly. Reason (R): The interparticle spaces are very large and interparticle forces are quite weak. a. Both A and R are correct and R is correct explantion of A. b. Both A and R are correct but R is not correct explantion of A. c. A is true, R is false. d. A is false, R is true.

IL Foundation Series Class 9

21. Statement (A): Gases exert pressure Statement (B): Gases are generally very light. a. Both A and R are true and R is the correct explanation of A b. Both A and R are true and R is not correct explanation of A c. A is correct and R is incorrect d. A is incorrect and R is correct. 22. Assertion (A): The kinetic energy of the particles in the gaseous state is very high. Reason (R): A gas can occupy the entire space of a given vessel in which it is enclosed. a. Both A and R are true and R is the correct explanation of A b. Both A and R are true and R is not correct explanation of A c. A is correct and R is incorrect d. A is incorrect and R is correct. 23. Statement (A): Solids have definite shape and volume. Statement (B): The density of solids is generally high. Statement (C): Liquids have fluidity but not rigidity. a. All the above statements are correct b. All the above statements are incorrect c. A, B are correct, and C is incorrect d. A, B are incorrect, and C is correct. 24. The order of boiling point of water, pentane and benzene is a. Water > Pentane > Benzene b. Pentane > Water > Benzene c. Water > Benzene > Pentane d. Benzene > Pentane > Water 25. Which of the following substance is in a gaseous state at room temperature? a. Water vapour

b. Alcohol vapour

c. Mercury vapour

d. Chlorine

26. The solid room fresheners are based on the properties of a. Evaporation

b. Sublimation

c. Decantation

d. Sedimentation

27. The melting point temperature of the solid state and freezing point temperature of the liquid state of the same substance are a. Same

b. Different 31

MATTER IN OUR SURROUNDINGS

c. May vary slightly

d. None

28. Assertion (A): The room temperature in Celsius scale is 25 °C which is equal to 77 °F on Fahrenheit scale.

9 0 C 32 5 a. Both A and R are true and R is the correct explanation of A

Reason (R): Temperature on Fahrenheit

b. Both A and R are true and R is not correct explanation of A c. A is correct and R is incorrect d. A is incorrect and R is correct. 29. Statement (A): The constant temperature at which a solid changes into a liquid state by absorbing heat energy is called melting point. Statement (B): The constant temperature at which a liquid changes into a gaseous state by absorbing heat energy is called boiling point. Statement (C): The melting and boiling points of a pure substance are always constant and are a good measure of finding the purity of the substance. a. All the above statements are correct b. All the above statements are incorrect c. A, B are correct, and C is incorrect d. A, B are incorrect, and C is correct 30. During evaporation of a liquid a. The temperature of the liquid falls b. The temperature of the liquid rises c. The temperature of the liquid remains unchanged d. All statements are wrong 31. A liquid is kept in an open china dish. The evaporation of liquid can be accelerated a. By keeping the dish open b. By blowing air on the liquid c. By keeping the dish under a running fan d. All are correct 32. The amount of water vapour present in the atmosphere is called a. Heredity

b. rigidity

c. humidity

d. all

33. Assertion (A): Pressure is normally expressed in atmospheres. Reason (R): 1 atmosphere is equal to 76 cm of Hg (or) 760 mm of Hg. a. Both A and R are true and R is the correct explanation of A 32

IL Foundation Series Class 9

b. Both A and R are true and R is not correct explanation of A c. A is correct and R is incorrect d. A is incorrect and R is correct. 34. Statement (A): Under high pressure, the gas gets compressed. Statement (B): At low pressure, the volume of the gas as well as inter particle spaces are very large. Statement (C): Under high pressure, the interparticle forces become strong and gas may change to a liquid state. a. All the above statements are correct b. All the above statements are incorrect c. A, B are correct, and C is incorrect d. A, B are incorrect, and C is correct. 35. During evaporation particles of a liquid change into vapours only a. From the surface

b. From the bulk

c. From both surface and bulk

d. Neither from surface nor from bulk

36. In the Plasma state a. Ions and electrons co-exist

b. Atoms and molecules co-exist

c. Atoms and ions co-exist

d. molecules and protons co-exist

37. Assertion (A): Alcohol evaporates faster than water. Reason (R): The interparticle forces of attraction in alcohol are less than in water. a. Both A and R are true and R is the correct explanation of A b. Both A and R are true and R is not correct explanation of A c. A is correct and R is incorrect d. A is incorrect and R is correct.

IS MATTER AROUND US PURE?

2.1 INTRODUCTION How can we tell if the milk, ghee, butter, salt, spices, mineral water, or juice we purchase from the store is pure?

Fig. 2.1 Grocery items

Have you ever noticed the term 'pure' on the labels of food items? To the average person, pure suggests no contamination. However, for a scientist, these items are mixtures of different substances and are not completely pure. Take milk, for instance; it is a combination of water, fat, proteins, and more. When a scientist claims something is pure, it means all the particles in that substance are identical in their chemical nature. A pure substance consists of only one type of particle, making it a singular form of matter. Looking around, we see that most of the matter we encounter is a blend of two or more pure components such as seawater, minerals, and soil are all mixtures. 2.1.1 Classification of pure substances In the previous classes, you have studied the scientific definition of matter. You have also studied how matter may be classified into three states and the changes it could undergo. But matter can also be classified chemically. Matter can be classified either into pure substance (an element or a compound) or impure substance, which can also termed as a mixture of various substances.

IL Foundation Series Class 9

Matter Impure substances (Mixtures)

Pure substances Compounds

Elements

Organic Metals (Na, Ca, etc.)

Inorganic Non-metals (O2, C12, etc.)

Homogeneous mixtures (True solutions)

Metalloids (As, Sb, etc.)

Heterogeneous mixture Colloids

Suspensions

Noble gases (He, Ne, etc.)

Fig. 2.2 Chemical classification of matter

2.2 TYPES OF PURE SUBSTANCES Pure substances exist in two primary forms: elements and compounds. An element represents a fundamental form of matter that remains indivisible through chemical reactions, resisting the breakdown into simpler substances. On the other hand, a compound is a substance formed by the chemical combination of two or more distinct types of elements, maintaining a consistent and fixed proportion. 2.2.1 Elements Robert Boyle is credited as the pioneer who introduced the term element in 1661. Following him, Antoine Laurent Lavoisier, a French chemist (1743-94), is recognized for providing the first experimentally practical definition of an element. According to Lavoisier, an element is considered a fundamental form of matter that cannot be decomposed into simpler substances through chemical reactions. An element is defined as a basic form of matter that cannot be broken down into simpler substances by any chemical reaction. An element is a pure substance made up of only one type of particle called an atom. An element cannot be broken down or converted into anything simpler by itself through a physical or chemical change. Likewise, it cannot be formed from simpler substances. So, elements are the basic substances from which all other substances can be made. Examples: Hydrogen, oxygen, nitrogen, copper, silver, etc. Atom: An atom is defined as the smallest particle of an element which takes part in the chemical reaction. An element is made up of the same kind of atoms. Elements are monoatomic, diatomic, triatomic, polyatomic etc.

IS MATTER AROUND US PURE?

Monoatomic elements - He, Ne, Ar, Kr, Xe, Rn, and metals Di-atomic elements - H2, O2, N2, F2, Cl2, Br2, I2, etc. Tetra-atomic elements - P4, As4, Sb4, etc. Octa-atomic element - S8 Molecule: The smallest particle of an element or compound that has independent existence and can retain all its properties is called a molecule. Examples: H2, O2, N2, H2O, CO2, etc. As the above are made up of only one kind of atoms, they are called homoatomic or homonuclear molecules. Further, depending upon whether the molecule contains only one, two, three, etc. atoms, they are called monoatomic, diatomic, triatomic, etc., respectively. For example, He, Ne, Ar, etc. are monoatomic, H2, O2, N2, etc. are diatomic, ozone (O3) is triatomic, P4 is tetra atomic, and S8 is octa atomic (polyatomic) molecules. Characteristics of an element

• Each element consists of only a single type of atom. No two elements contain the same kind of atoms. • An element is composed either of individual identical atoms or of molecules made up of these atoms. • An element is a pure or homogeneous substance. • It has its own fixed boiling and melting points. • An atom is the smallest particle of an element taking part in a chemical reaction. • An element cannot be broken down into simpler substances by any chemical or physical means. (The only exception in this case is when a nuclear reaction takes place.) • An element may chemically react with another element or compound. • An element may be a metal, non-metal, metalloid or noble gas. • An element may occur in a solid, liquid or gaseous state. The elements may be regarded as the building units of the universe. Scientists believe that 92 different kinds of elements exist in nature. Others have been prepared in the laboratory by artificial means during nuclear research, and the total number of elements known to man so far is 118. While most elements occur in nature in combination with each other, some of them, like oxygen, nitrogen, copper, gold, and so on, also occur in a free state, that is, in an uncombined state. In the Earth's crust, comprising the outer portion of the Earth, oceans, and air, we find that oxygen is present to the maximum extent. Next comes silicon, which is an essential constituent of the rocks. The approximate average percentage by weight of the various elements of the Earth's crust is given below: 36

IL Foundation Series Class 9

Element

% by weight

Element

% by weight

Oxygen

48.85

Potassium

2.33

Silicon

26.03

Magnesium

2.11

Aluminium

7.28

Hydrogen

0.97

Iron

4.12

Titanium

0.41

Calcium

2.18

Chlorine

0.20

Sodium

2.33

Carbon

0.09

Other elements

1.00

Table 2.1 The weight of the various elements of the Earth's crust

Elements found in the human body The elements that biochemists have found to be most important in the human body are oxygen 65%, carbon -18%, nitrogen -3.2%, calcium -1.5%, phosphorus - 1%, and others. Classification of elements

Elements are classified on the basis of their properties and also on the basis of their sub-divisions. Based on their sub-divisions, elements may be solids, liquids or gases. Mercury, bromine, caesium, and galium can exist as liquids at 30° C. Hydrogen, nitrogen, oxygen, chlorine, fluorine, helium, neon, argon, krypton, xenon, and radon exist as gasses at room temperature. The remaining elements are solids. Metals elements which generally

Non-metals elements which generally

Metalloids elements which generally

Noble gas elements which

i) Have metallic lustre

i) Do not have lustre

i) Have properties mid-way between metals and nonmetals

i) Have properties are that are gaseous in nature

ii) Are good conductors of heat and electricity

ii) Are bad conductors of heat and electricity

ii) Contain one kind of atom and are monoatomic

ii) Chemically inert

iii) Are malleable (beaten into sheets) and ductile (drawn into wires)

iii) Are not malleable or ductile

iii) Occur in a free state in traces in the atmosphere

IS MATTER AROUND US PURE?

Metals elements which generally

Non-metals elements which generally

iv) Contain one kind of atoms and are monoatomic

iv) Contain one kind of atoms and are mono-atomic or di-atomic or polyatomic

Examples: Sodium, potassium, iron, lead, etc.

Examples: Gaseous, hydrogen, chlorine, oxygen, etc.

Exceptions:

i) Mercury (Hg) is a liquid metal at room temperature

i) Iodine and graphite are lustrous

ii) Zinc (Zn) is nonmalleable and nonductile

ii) Carbon fibre is ductile but not malleable

iii) Tungsten (W) is a poor conductor of electricity

iii) Graphite is a good conductor of electricity

Metalloids elements which generally

Noble gas elements which iv) Contain one type of atoms and are monoatomic

Examples: Germanium (Ge) arsenic (As) antimony (Sb) bismuth (B)

Examples: Helium (He) neon (Ne) argon (Ar) krypton (Kr) xenon (Xe)

Table 2.2 Classification of elements based on their properties

2.2.2 Compounds A compound is a substance composed of two or more different elements, chemically combined with one another in a fixed proportion by weight. The constituents of a compound can be separated only by chemical reactions but not by physical changes. Example: Water contains oxygen and hydrogen. Similarly, carbon dioxide contains carbon and oxygen. However, the properties of a compound are different from the properties of its constituent elements. For example, sodium is a soft metal that reacts violently with water. A person who swallows sodium would die. Chlorine is a greenish-yellow gas which has an overpowering smell. Chlorine, too, is poisonous in nature. But a substance which is composed of these two elements, sodium and chlorine, forming a white, crystalline compound, sodium chloride, is non-poisonous and is used for cooking food. 38

IL Foundation Series Class 9

Note: However, if the same elements are combined chemically, they give more than one compound. For example, if hydrogen and oxygen are combined, they form two different compounds, i.e., water (H2O) and hydrogen peroxide (H2O2) under different conditions. In water, two hydrogen atoms are combined with one atom of oxygen, whereas in hydrogen peroxide, two atoms of oxygen combine with two atoms of hydrogen. Molecules of compounds: They are made up of atoms of different elements and hence are called heteroatomic or heteronuclear molecules. They may be diatomic, triatomic, etc., depending upon the number of atoms present in one molecule of the compound. For example, HF, HCl, HBr, and HI, are diatomic, H2O, CO2, etc., are triatomic, and NH3 and PH3 are tetratomic. Atomicity: The number of atoms present in one molecule of an element or compound is called atomicity. Examples: Atomicity of helium (He), hydrogen (H2), ozone (O3), phosphorus (P4), and sulphur (S8) are 1, 2, 3, 4, and 8 respectively. Atomicity of water (H2O), ammonia (NH3 ), and methane (CH4) are 3, 4, and 5 respectively. Characteristics of a compound

•

Homogeneity: A compound has a homogeneous composition. In other words, all samples of a compound have identical physical and chemical properties. For example, if one molecule of water contains a combination of oxygen and hydrogen in each pattern, then all the other molecules of water will also contain hydrogen and oxygen in the same pattern.

•

Combination: A compound contains atoms of two or more elements combined by chemical forces. For example, Iron (II) sulphide (FeS) is a compound of iron and sulphur. When a magnet is placed near FeS, its iron is not attracted by the magnet. Similarly, sulphur present in FeS is not soluble in its solvent viz, carbon disulphide (CS2).

•

Separability: The components of a compound cannot be separated by physical means, though their components can be separated by chemical means. For example, water (H2O) is a compound made up of hydrogen and oxygen. Hydrogen and oxygen, which are components of water, cannot be separated by physical methods. However, they can be separated by a chemical process called electrolysis, in which electric current is passed through acidified water.

•

Proportion: The elements in a compound are present in a fixed ratio(proportion) by mass. If the proportion or ratio is changed, a different compound is formed. For example, the ratio of iron to sulphur in iron sulphide is 7:4. That means iron sulphide contains 7 parts of iron combined with 4 parts of sulphur by mass.

•

Similarly, the chemical combination of two H atoms and one O atom gives us the water molecule H2O. But if you combine 2H atoms and 2O atoms, it will give us a hydrogen peroxide molecule.

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•

Energy change: The formation of a compound involves the absorption or release of energy in the form of heat, light, electricity, etc. For example, carbon dioxide is formed from carbon and oxygen with the liberation of heat (exothermic reaction). C + O2 → CO2 + heat energy

On the other hand, nitric oxide is formed from nitrogen and oxygen with the absorption of heat (endothermic reaction). N2 + O2 + heat energy → 2NO The process of photosynthesis proceeds with the absorption of light energy (endothermic process). 6CO2 + 6H2 O + light energy → C6H12O6 + 6O2

Element

Atoms of the element argon exist on their own.

Oxygen atoms join in pairs. Argon and oxygen are elements.

Compound

Mixture

O Carbon and oxygen atoms are joined together in carbon dioxide.

N O

Air is a mixture of elements and compounds.

Fig. 2.3 Different types of substances

2.3 MIXTURE AND ITS TYPES A mixture is a matter composed of two or more substances (elements, compounds, or both) whose particles are in contact but are not chemically combined, and each of the components still exhibits its own characteristics and properties. The properties of a mixture of common salt and water will have different densities, boiling points, etc., depending on the amounts of the two components present in it. These properties can be varied by adding water or salt. But the properties of water and salt are not lost by doing so. Some of the mixtures that are available at home for daily use are toothpaste or tooth powder, lime juice, tea, coffee, kerosene oil, cough syrups, oranges, squash, etc. 40

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Pure substance

Mixture Fig. 2.4 Mixture

2.3.1 Characteristics of mixtures 1. A mixture may be homogeneous or heterogeneous: The components of a homogeneous mixture are mixed uniformly, and hence its composition is uniform throughout the mixture. Alloys are examples of a homogeneous mixture. Bronze is an example of an alloy, which is a mixture of Cu(80%), Zn(2%), and Sn(18%). The constituents of a heterogeneous mixture are not mixed uniformly, and hence, its composition is not uniform throughout the mixture. For example, gunpowder, which is a mixture of charcoal, sulphur, and nitre, is an example of a heterogeneous mixture. 2. Separability: The components of a mixture can be separated by physical means. For instance, in a mixture of iron and sulphur, iron can be separated with the help of a magnet or by adding HCl to the mixture, a colourless and odourless gas (H2 ) is evolved. The gas burns with a blue flame and is extinguished with a pop sound. This gas is generated by the action of HCl on Fe (iron). Sulphur remains unreacted.

Fe + S + 2HCl  → FeCl2 + H2 + S(unreacted) mixture

(OR) Sulphur can be separated from the mixture by adding carbon disulphide (CS2) to the mixture when sulphur gets dissolved in CS2 and Fe remains as such. A mixture consists of two or more substances that exist together without any chemical force acting upon them. A mixture can have its components in varying proportions. Properties: A mixture has no definite set of properties. The individual components decide the properties of a mixture because the particles of the different components are not chemically combined. For example, Fe is attracted by a magnet and gets dissolved in an acid even when it is mixed with sulphur. The melting point or the boiling point of the mixture is not fixed. It depends on the proportion of the components present in it. For example, the boiling point of a solution of sugar in water depends on the amount of sugar; the more sugar, the higher the boiling point of the solution. Solids, liquids,

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and gases can be mixed in different combinations to form either a homogenous or heterogeneous mixture. There are two types of mixtures: homogeneous mixture and heterogeneous mixture.

HOMOGENEOUS HOMOGENEOUS

Fig. 2.5 Types of mixtures

2.3.2 Homogeneous mixtures A homogeneous mixture has the same properties and characteristics in all the parts of its volume. For example, alcohol mixed with water. They can be mixed well and miscible with each other. Also, milk with water or sugar with milk and so on. These are also known as true solutions.

Sugar

Water

Sugar solution (Homogeneous mixture)

Fig. 2.6 Homogeneous mixture

2.3.3 Heterogeneous mixtures A heterogeneous mixture is that which does not have the same properties throughout its bulk. For example, when oil and water are mixed, oil, being lighter than water, forms the upper layer, while water, being denser, forms the lower layer. They do not mix with each other as they are immiscible. Sand mixed with salt would be another such example.

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Sand

Water

No solution (Heterogeneous mixture)

Fig. 2.7 Heterogeneous mixture

Note: Alloys are like mixtures of two or more metals or a metal with a non-metal. You can't easily separate them into their parts using normal methods. Even though they can't be split apart physically, we still call them mixtures. This is because alloys act like a mix of their original parts and can have different amounts of each. Take brass, for example, which has about 30% zinc and 70% copper. It shows some features of both zinc and copper, making it useful for different things.

Fig. 2.8 Classification of mixture

Based on the particle size of the substance, the solutions may be divided into three types. These are 1. True solutions 2. Suspensions 3. Colloidal solutions A true solution is a homogeneous solution which contains small solute particles (molecules or ions) dispersed throughout a solvent. For example, the solution of sodium chloride in water. The particle size is less than 1 nm. The particles of a solute in a true solution are invisible even under a microscope, and its particles can pass through ordinary filter paper as well as through animal membranes.

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2.4 PHYSICAL AND CHEMICAL CHANGES Physical properties encompass observable and measurable characteristics such as colour, hardness, rigidity, fluidity, density, melting point, boiling point, and more. These attributes define the distinctive features of a substance without involving any chemical changes.

Combustion

Rotting

Melting

Shredding

Rusting

Digestion

Boiling

Chopping

Physical Changes Chemical Changes Fig. 2.9 Chemical changes and physical changes

When substances change from one form to another, like from ice to water or water vapour, it's called a physical change. This happens without any alteration in what the substance is made of or its chemical nature. Even though ice, water, and water vapour may look different and have different properties, they are made up of the same chemicals. Now, let's talk about water and cooking oil. Both are liquids, but they have different characteristics. They smell different, and oil can catch fire, while water puts out flames. This is because of their chemical properties. When oil burns, it undergoes a chemical change. In a chemical change, substances react with each other, leading to a new chemical composition. This kind of change is often called a chemical reaction. When a candle burns, it goes through both physical and chemical changes. Physical changes involve the way something looks or feels, while chemical changes lead to the creation of entirely new substances. 2.4.1 Evaporation During the study of the change of state, we have seen that a liquid changes to a gaseous state either by increasing the temperature or by decreasing the pressure. This process is known as evaporation. Evaporation may be defined as the phenomenon of change of liquid from its surface to the vapour state at any temperature below the boiling point of the liquid.

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Factors affecting evaporation

The evaporation of liquids can be increased or accelerated by the following factors. • An increase in surface area: Evaporation is a surface phenomenon. If the surface area is increased, the rate of evaporation increases. Example: Spreading of wet clothes for quick drying. • An increase in temperature: With the increase in temperature, a greater number of particles get enough kinetic energy to go into the vapour state. • A decrease in humidity: Humidity is the amount of water vapour present in the air. The air around us cannot hold more than a definite amount of water vapour at a given temperature. If the amount of water in the air is already high, the rate of evaporation decreases. • An increase in wind speed: It is a common observation that clothes dry faster on a windy day. With the increase in wind speed, the particles of water vapour move away with the wind as a decrease in the amount of water vapour in the surroundings. • Nature of the liquid: We have so far discussed the external factors which influence the extent of evaporation. Apart from these, another factor which is of great importance is the nature of the liquid which is evaporating. We often see that alcohol evaporates at a faster rate than water. In fact, the boiling point of alcohol (350 K) is less than that of water (373 K). This means that interparticle forces of attraction in alcohol are less than in water. Therefore, alcohol will evaporate faster than water. Thus, we conclude that the lower the boiling point of a liquid, the more its tendency to change into vapours or to evaporate.

2.5 SOLUTION AND ITS PROPERTIES A solution is like a mix where everything is spread out evenly. Think about lemonade or soda - those are examples of solutions. Usually, we think of solutions as liquids mixed with something solid, liquid, or gas. But did you know we can also have solutions that are solid, like when different metals mix (we call that an alloy), or even solutions that are gases, like the air we breathe? In a solution, everything is mixed so well that it all looks the same, even at the tiniest particle level. For example, when you drink lemonade, it tastes the same from the first sip to the last. That's because the sugar or salt in it is spread out evenly throughout the whole drink. Example: If a sugar lump is dipped in a beaker of water, the lump disintegrates and, within a short time, disappears into the liquid phase. In this process, the molecules of the sugar leave the crystal structure of the solid and become uniformly dispersed throughout the water, thus producing a complete mixture of the two substances. Thus, it is a solution of sugar in water.

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Solute: The component present in a lesser quantity in a binary solution is referred to as solute. Solvent: The component present in a larger quantity, generally in any solution, is called the solvent. Ternary solutions are composed of three components.

Fig. 2.10 Components of solution

Examples: i) Carbonated drinks like soda water are examples of gas-in-liquid solutions. They consist of carbon dioxide (gas) as the solute and water (liquid) as the solvent. ii) Air is a blend of gases, forming a gas-in-gas mixture. It is a uniform mix of various gases, primarily oxygen (21%) and nitrogen (78%), with other gases present in very small amounts. 2.5.1 Properties of a solution • A solution is a type of mixture where everything is evenly spread out, making it look the same throughout. • The tiny particles in a solution are so small, less than 1 nanometer (10-9 meters), that we can't see them with our eyes. • Because these particles are incredibly small, they don't make the light passing through the solution visible. The path of light remains invisible. • Unlike some mixtures, you can't separate the particles in a solution using a process called filtration. Also, the particles don't sink down when left alone, making a solution stable. 2.5.2 Classification of solutions Solutions can be classified based on the physical state of the solute and solvent. Since there are three states of matter, there are theoretically nine possible types of solutions. Three types are possible

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when a liquid is the solvent since the solute may be a gas, a liquid or a solid. Similarly, three types are possible when the solvent is a gas and three when the solvent is a solid. Type of the solution

Gaseous solutions

Liquid solutions

Solid solutions

Solute

Solvent

Example

Gas

Mixture of gases

Liquid

Gas

Chloroform mixed with nitrogen gas

Solid

Gas

Camphor in nitrogen gas

Gas

Liquid

Oxygen in water, soda water

Liquid

Alcohol in water

Solid

Liquid

Glucose in water, sugar in water

Gas

Solid

Hydrogen adsorbed on palladium

Liquid

Solid

Mercury in gold (Amalgam)

Solid

Metal alloys (Brass, Zn and Cu)

Table 2.3 Various types of solutions with examples Classification of solutions based on the quantity of solute present in a given solution

1. Saturated solution: A solution in which no more solute can be dissolved at a given temperature and pressure is called a saturated solution. 2. Unsaturated solution: A solution in which more solute can be dissolved at the same temperature and pressure is called an unsaturated solution. 3. Supersaturated solution: A solution in which more solute is dissolved than its saturated level by increasing temperature or pressure is called a supersaturated solution.

Fig. 2.11 Types of solutions

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Classification of solutions based on the relative amount of solute present in a given solution

1. D ilute solutions: A solution containing a relatively very small quantity of solute is called a dilute solution. 2. C oncentrated solution: A solution containing a relatively very large quantity of solute is called a concentrated solution. Classification of solutions depending on the solvent

1. Aqueous solutions: In this type, water acts as the solvent. 2. N on-aqueous solutions: In this type, a substance other than water (alcohol, CHCl3, CCl4, C6H6) acts as the solvent. 2.5.3 Solubility The solubility of a substance is its maximum amount that can be dissolved in 100 g of solvent at a specified temperature. Solubility =

Mass of solute Mass of solvent

×100

Solubility depends upon: • Nature of solute and solvent • Temperature • Pressure Factors affecting solubility

Factors affecting the solubility of a solid in a liquid: • Nature of solute and solvent: Every solid does not dissolve in each liquid. Sodium chloride and sugar dissolve readily in water. However, naphthalene and anthracene do not dissolve in water. Naphthalene and anthracene dissolve readily in benzene. However, sodium chloride and sugar do not dissolve in benzene. It is observed that polar solutes dissolve in polar solvents and non-polar solutes in non-polar solvents. In general, a solute dissolves in a solvent if the intermolecular interactions are similar in the two, or we may say like dissolves like. • Effect of temperature: The solubility of a solid in a liquid is significantly affected by temperature changes. Consider the following equilibrium. Solute + Solvent ⇌ Solution

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In general, if in a nearly saturated solution, the dissolution process is endothermic (Hsol>zero), the solubility should increase with the rise in temperature and if it is exothermic (Hsol < zero). The solubility should decrease. • Effect of pressure: Pressure does not have any significant effect on the solubility of solids in liquids. This is so because solids and liquids are highly incompressible and practically remain unaffected by changes in pressure. Solubility of gas in liquid: All gases are soluble in water as well as in other liquids to a greater or lesser extent. The solubility of a gas in liquids depends upon the following factors. 1. Nature of the gas 2. Nature of the solvent 3. Temperature 4. Pressure i) Nature of the gas and nature of the solvent: Generally, the gases which can be easily liquified are more soluble in common solvents. Example: CO2 is more soluble than hydrogen or oxygen in water. The gases which are capable of forming ions in aqueous solutions are much more soluble in water than in other solvents. Example: Gases like hydrogen chloride (HCl) and ammonia (NH3 ) are highly soluble in water but not in organic solvents in which they do not ionise. ii) Effect of temperature: The solubility of gases in liquids decreases with the rise in temperature. When dissolved, the gas molecules are present in the liquid phase, and the process of dissolution can be considered similar to condensation, and heat is evolved in this process. If dissolution is an exothermic process, the solubility should decrease with the increase in temperature. When a solution of a gas is heated, the gas is usually expelled. iii) Effect of pressure: The solubility of gases in liquids increases with the increase of pressure.

Fig. 2.12 Effect of pressure on solubility

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2.6 SUSPENSION AND ITS PROPERTIES Suspension is a heterogeneous mixture that contains small insoluble particles. The particle size is more than 100 nm. For example, dirt particles in water. The particles of a suspension may not be visible to the naked eye but are visible under a microscope. The particles of a suspension can neither pass through an ordinary filter paper nor through the animal membrane. The mixture still has undissolved particles settled at the bottom with some particles still floating.

Water

Sand Fig. 2.13 Suspension

Suspension

2.6.1 Properties of a suspension The properties of a suspension are: • Suspension is characterized as a heterogeneous mixture of visible particles. These particles are observable to the naked eye. • When a beam of light passes through a suspension, the particles scatter the light, making their path visible. • In undisturbed conditions, the solute particles in a suspension settle down, rendering it unstable. • Separation of these particles from the mixture can be achieved through the process of filtration. Once the particles settle, the suspension breaks and it no longer scatters light.

2.7 WHAT IS A COLLOIDAL SOLUTION? Thomas Graham, in 1861, observed that certain solutes, such as starch, glue, gelatine, etc., could not pass through the parchment membrane while ordinary solutes, such as sodium chloride, urea, sugar, etc., could easily do so. Graham called the former solutes colloids (Greek, kollo meaning glue) while the latter were called crystalloids. However, the above classification of solutes into crystalloids and colloids proved unsatisfactory because a particular substance would be crystalloid in one solvent and a colloidal in the other. For example, in an aqueous solution, NaCl is a crystalloid, while in benzene, it behaves as a colloid. Similarly, soap is a typical colloid in water, but it acts as a crystalloid in alcohol. Further studies of the behaviour of the solutes have shown that the nature of the substance, 50

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whether colloid or crystalloid, depends upon the particle size. When the size of the particles is between 10-9 m(1 nm) to 10-7 m(100 nm ), it behaves like a colloid, and in case it is less than this range, it gives the characteristics of a crystalloid. Thus, colloid is not a substance, but it is a state of a substance which depends upon the molecular size. A colloidal solution is a heterogeneous solution which contains particles of intermediate size. For example, milk. The particles of a colloidal solution have diameters between 1 to 1000 nm. Such particles cannot be normally seen with the naked eye. However, the light reflected by them can be seen under an ultramicroscope. The particles of a colloidal solution can pass through ordinary filter paper but not through animal membranes. In a colloid, the dispersed phase may consist of particles of a single macromolecule (such as synthetic polymer or protein) or an aggregate of many atoms, molecules, or ions. Colloidal particles have an enormous surface area per unit mass. For example, consider a cube having each side as 1 cm. It has a total surface area of 6 cm2 because it has six faces, and each face has an area of 1 cm2. Now, if it were divided equally into 1012 cubes, the cubes would be the size of large colloidal particles and have a total surface area of 60,000 cm2 or 6 m2. This enormous area is responsible for some special properties of colloids, which will be learned in this unit. The dispersed particles have the ability to disperse a visible light beam. This phenomenon, known as the Tyndall effect, is named after the scientist who first identified it. The Tyndall effect is also noticeable when a narrow beam of light enters a room through a small aperture, as it is caused by the scattering of light by airborne particles such as dust and smoke. Additionally, the Tyndall effect can be witnessed when sunlight traverses through the canopy of a dense forest. In this environment, mist carries minute water droplets that function as colloidal particles dispersed in the air.

Flashlight Solution

Colloid

Suspension

Fig. 2.14 Tyndall effect

2.7.1 Phases of colloids and their classification 1. Dispersed phase: It is the component present in a small proportion and is just like a solute in a solution. For example, in the colloidal solution of silver in water, the former acts as a dispersed phase. 2. Dispersion medium: It is generally a component present in excess and is just like a solvent in a solution. In the above example, water acts as a dispersion medium. Thus, the particles of the

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dispersed phase are distributed in the dispersion medium. Out of solid, liquid, and gas, each one can act as a dispersed phase and dispersion medium, leading to eight types of colloidal systems. Classification of colloids The colloids are classified on the basis of the following criteria : 1. Physical state of the dispersed phase and the dispersion medium. 2. Nature of interactions between the dispersed phase and the dispersion medium. 3. Type of particles of the dispersed phase. A. Classification based on the physical state of the dispersed phase and dispersion medium. Depending upon the physical state of the dispersed phase and dispersion medium, whether these are solids, liquids or gases, eight types of colloidal systems are possible. Examples of the various types of colloids and their typical names are given in the table below. It may be noted that a gas mixed with another gas forms a homogeneous mixture, and therefore, it is not a colloidal system.

S.No..

Dispersion Phase

Dispersion medium

Type of colloid

Example

Solid

Solid sol

Alloys, coloured glasses, gemstones, ruby glass

Solid

Liquid

Sol

Paints, cell fluids, starch dispersed in water

Solid

Gas

Aerosol

Liquid

Solid

Gel

Jelly, butter, cheese, boot polish, curd

Liquid

Emulsion

Milk, hair cream, emulsified oils, medicines, mayonnaise

Liquid

Gas

Aerosol

Mist, fog, cloud, insecticide spray

Gas

Solid

Solid sol

Pumice stone, foam rubber

Gas

Liquid

Foam

Soap leather, froth, whipped cream, soda water

gold sol

Table 2.4 Types of colloidal systems

It is clear from the table that many common commercial products and natural objects are colloids. For example, whipped cream is a colloidal system (foam), a gas dispersed in a liquid. Out of the different types of colloids, the most common are sols (solids in liquids), gels (liquids in solids) and emulsions liquids in liquids. The important distinguishing features of the three types of solutions are as follows:

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S. No.

Property

Nature

Particle size

Suspension

Colloid solution

Heterogeneous

(or

)

True solution Homogeneous

)

(or

)

Separation by i) Ordinary filtration

Possible

Not possible

ii) Ultra-filtration

Possible

Not possible

Setting of particles

Settle under gravity

Settle only on centrifugation

Do not settle

Visibility

Particles visible to naked eye or under a microscope

Scattering of light by the particles is observed under ultra-microscope

Particles are invisible

Appearance

Opaque

Generally transparent

Transparent

Tyndall effect

Shows

Does not show

Diffusion of particles

Does not diffuse

Diffuses slowly

Diffuses rapidly

Brownian movement

May show

Shows

Negligible

Table 2.5 Properties of a colloid solution, suspension solution, and true solution

Thus, colloidal solutions are intermediate between true solutions and suspensions. In other words, the size of dispersed particles in colloidal solutions is more than that of solute particles in a true solution and smaller than that of suspension. The size of different solutions is sometimes expressed in other units also, as given below: True solutions

Colloids

Suspensions

Relation

Table 2.6 Size (diameter) of particles in different units

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2.8 CONCENTRATION OF SOLUTION Different things dissolve differently in the same liquid when it's at the same temperature. Now, when we talk about the concentration of a solution, we mean how much of the stuff we dissolved (like salt or sugar) is in a certain amount of the liquid. It's like asking- How much salt or sugar is in this cup of water? Concentration is just a way of answering that question. And guess what? There are a few different ways we can talk about this concentration stuff. Today, we'll keep it simple and learn about three ways to do that. The concentration of a solution is defined as the amount of solute present in a unit volume of the solution. The concentration of a solution is expressed in several ways. 1. Mass percentage (Weight percentage) (w/w) 2. Volume percentage (v/v) 3. Mass by volume percentage (w/v) 4. Parts per million (ppm) 5. Molarity 6. Molality 7. Mole fraction 2.8.1 Weight percentage, volume percentage with questions Mass/Weight percentage (W %)

The mass of solute (in grams) present in 100 grams of solution is called mass percentage. Let 'w' grams of solute be added to 'W' grams of solvent. Mass of a solution = (w + W) grams. mass percentage = w% =

Mass of solute Mass of solution w w+W

× 100

Note: Mass percentage has no units. It is independent of temperature. Example: 5 grams of sodium carbonate are present in 120 grams of water. Calculate the mass percentage. Solution: Mass of solute, w = 5 g

Mass of solution, (w + W = 5 + 120 = 125 g)

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Mass percentage, w % = w%=

5 125

× 100

w+W

× 100

w%=4 Volume percentage (V%)

The volume of solute (in ml ) present in 100ml of solution is called volume percentage. Let 'v' ml of solute be added to 'V' ml of solvent. Then, the volume of solution = (v + V)ml. Volume percentage, v % = v% =

Volume of solute Volume of solution v v+V

× 100

×100

Note: Volume percentage has no units. It depends on temperature. Example: 15 ml of hexane is mixed with 45 ml of heptane. Calculate the volume percentage of this solution. Solution: Volume of solute, v = 15 ml

Volume of solvent, V= 45 ml

Volume of solution = v + V

= 15+45

= 60 ml

Volume of percentage, V% =

V% =

V% = 25

15 15+45

× 100

v+V

× 100

Mass by volume percentage (w/v)

It is the mass of the solute dissolved in 100ml of the solution. It is commonly used in medicine and pharmacy. (w/v)% = (w/v)% =

Mass of solute Volume of solution w v

× 100

Note: It depends on temperature.

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Parts per million: When a solute is present in trace quantities, it is convenient to express the concentration in parts per million (ppm). Parts per million =

Number of parts of the component Total number of parts of all components of the solution

× 106

• A litre of seawater (which weighs 1030 g ) contains about 6 × 10-3 g of dissolved oxygen. Such a small concentration is also expressed as 5.8 g per 106 g(5.8 ppm) of seawater. • The concentration of pollutants in water or atmosphere is often expressed in terms of μgml-1 or ppm. 2.8.2 Moles, mole-volume relation Quite often, we use the unit dozen to represent 12 articles, irrespective of their nature. For example, one dozen books mean 12 books, whereas one dozen apples mean 12 apples. In a similar way, chemists use the unit mole for counting atoms, molecules, ions, etc. A mole is a collection of 6.023 × 1023 particles. Thus, a mole represents 6.023 × 1023 particles. The actual meaning of a mole is a heap, but here, it represents the number of atoms or molecules or particles present in a fixed amount of the substance. This fixed amount of the substance is known as atomic mass or molecular mass for molecules. Relative atomic mass

Relative atomic weight or atomic mass of any element is the number which indicates how many times its atom is heavier than 1/12th of the weight of a CARBON-12 atom. (Or) It is the ratio of the mass of one atom of an element to 1/12th of the mass of one atom of the CARBON-12 isotope. Atomic mass of an element =

Mass of one atom of an element 1/12

part of mass of an atom of C-12 isotope

× 106

• Here, we have to remember one thing strictly. We cannot measure atomic mass accurately and directly, so we have to use a reference. Even though we have a number of references, nowadays, we consider CARBON-12 as the reference as it gives an accurate atomic weight. • Atomic mass can be expressed as 1 a.m.u (or) one unified mass (u) (or) one Avogram (or) one Aston (or) one Dalton. • 1 a.m.u. = 1.66 × 10-24 g = 1.66 × 10-27 kg Gram atom or gram atomic mass

Gram atomic mass is defined as the mass of Avogadro number of atoms of an element. (Or) The amount of substance is equal to the atomic weight of an element expressed in grams.

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Example: Atomic weight of oxygen =16 a.m.u. One gram atom of oxygen = 16 g 16 g of oxygen atoms = 6.023 × 1023 number of oxygen atoms. Relative molecular mass

It is defined as the average relative mass of a molecule as compared to the 1/12th mass of a CARBON-12 atom. (In other words, molecular mass indicates how many times a molecule of the substance is heavier than 1/12th of the mass of a CARBON-12 atom. It is expressed in a.m.u. Determination of molecular weight

i) Homoatomic molecules: Molecular weight = Atomic weight × Atomicity Example: Phosphorus (P4) = 31 × 4 = 124, Ozone (O3) = 16 × 3 = 48 etc. ii) Heteroatomic molecules: The molecular weight of a heteroatomic molecule is equal to the sum of atomic weights of atoms of all the elements present in one molecule. Example: Na2CO3 = 2 × 23 + 12 + 3 × 16 = 106. Gram molecule or gram molecular mass

A substance with a one-gram molecular weight is termed a one-gram molecule. (or) The amount of substance equal to the molecular weight of that substance expressed in grams is called a gram molecule or gram molecular mass. Example: Molecular weight of oxygen = 32 a.m.u. One gram molecule of oxygen = 32 g of oxygen Gram molecular weight of oxygen =32 g Mole

One mole is defined as the amount of the substance that contains as many particles as atoms exactly present in 12 g of the C - 12 Isotope. (Or) The amount of the substance which contains the same number of elementary particles (atoms, molecules, ions or electrons) as the number of atoms present in 12 g of Carbon (C-12). Example: Let us calculate the number of atoms present in 12 g of C-12 Isotope. One C-12 atom weighs 1.9926 × 10-23 g Number of atoms in 12 g per mole of C - 12 isotope =

12 g/mole 1.9926 × 10-23 g/atom

= 6.023 × 1023 atoms/mole 57

IS MATTER AROUND US PURE?

This number of entities in one mole of substance is known as Avogadro's number. It is denoted by N or NA. • Avogadro's number = 6.023 × 1023 number of particles (atoms or molecules or ions or electrons). This number is named in honour of the Italian scientist Avogadro. It is used as a reference for most of the calculations and equations found in chemistry. Applications of the mole concept

• The chemical formula represents one mole of the substance. • The formula mass in grams represents the mass of one mole of that substance. • One mole of any substance contains 6.023 × 1023 particles. • At STP, one mole of any gas occupies 22.4 litres of volume Note: STP is a shorthand way of representing the standard temperature, i.e., 0° C or 273 K and a standard pressure of 1 atmosphere. Now, we will discuss these rules in detail. Rule 1: The chemical formula represents a mole of that substance. Remember that any number placed to the left of a chemical symbol or formula is the coefficient. This number (integer, decimal or scientific notation) tells us the number of moles of that substance. Example: Pb→1 mole of lead atoms 3 Pb→3 moles of lead atoms 1.5 Cl- → 1.5 moles of chloride ions

2 CaCl2 → 2 moles of calcium chloride units (molecules) Rule 2: The formula mass in grams represents the mass of one mole of that substance. The formula mass of a compound is the sum of the atomic masses of all the atoms present in that compound. If formula mass is expressed in grams, then it is called Gram formula mass. The following example is given to demonstrate how to find the formula mass or formula unit mass. Example: CaCO3 (Calcium carbonate)

Ca → 1 × 40.1 = 40.1 C →1 × 12.0 = 12.0

O → 3 × 16.0 = 48.0 100.1 amu = 100 amu ( 3 significant figures) So, the formula unit mass of CaCO3 is 100 a.m.u Then, the gram formula mass of CaCO3 is 100 grams

IL Foundation Series Class 9

5.00 NaCl → 5 moles of sodium chloride →5.00 × 58.5 = 292 g (rounded) 2.5 H2 SO4 → 2.5 moles of sulphuric acid →2.5 × 98.1 = 245 g (rounded) Rule 3: One mole of any substance contains 6.023 × 1023 particles Here, particles mean atoms, molecules, ions, electrons, or just about anything that we might need to work with. Remember, just as there are 12 items in a dozen, 6.023 × 1023 particles are present in a mole. Example 1: HNO3→1 mole of nitric acid →1.00 × 63.0 = 63.0 grams, 6.023 × 1023 molecules of nitric acid. Example 2: 3.00 K→3.00 moles of potassium atoms,

3.00 × 39.1 = 117 grams of potassium,

3.00 × 6.023 × 1023 = 1.81 × 1024 potassium atoms.

Rule 4: At STP, one mole of any gas occupies 22.4 litres of volume. This is called gram molar volume (GMV) STP means standard temperature and pressure. STP is a shorthand way of representing the standard temperature, i.e., 0° C or 273 K and a standard pressure of 1 atm. 1 atm = 76 cm of Hg or 760 mm of Hg This rule is most commonly used while studying the gas laws. Suppose you have 4 g of helium gas. This represents one mole of helium (see 2nd rule). These 6.023 × 1023 atoms of helium would occupy 22.4 litres of volume at STP. This large volume would be fully occupied when the temperature is 0° C(273 K) and the pressure is 1 atmosphere. A change in the temperature or pressure would change the volume occupied by the gas. Sample calculations on the mole concept

Formula to find the number of moles

How many moles of sodium carbonate are there in 4.5 moles of Na2CO3? Answer: Remember, the coefficient in front of an element or compound tells you the number of moles you have. Of course, we are dealing with 4.50 moles of sodium carbonate. Remember that the coefficient can be a whole number, a decimal, or a number in scientific notation, and that the number of scientific figures in that coefficient indicates the detailed precision needed in your final answer.

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Conversion of moles to grams

2.00 moles of Ca(OH)2 would represent how many grams? Answer: Remember that 1 mole of a compound is represented by the formula mass of that compound. Also, 1 mole of an element is equal to its atomic mass. To solve this problem, first, we have to calculate the gram formula mass. And then, multiply the number by the number of moles we have. To calculate gram formula mass

First, list out the elements in the formula along with their number (hint: use the subscripts).

Then, multiply that number by the atomic mass of that element.

Add those masses, and you get the gram formula mass.

Remember, to get your final answer, you must multiply the formula mass by the number of m oles.

1) Calculate the gram formula mass of Ca(OH)2 1(Ca)→1 × 40.1 = 40.1 2 (O) → 2 × 16.0 = 32.0

2(H) → 2 × 1.01 = 2.02 Gram formula mass is 74.1 g( rounded ) 2) Calculate the gram formula mass of 2.00 Ca(OH)2 2.00 × 74.1 g = 148 g (again rounded to 3 significant figures) 3) Conversion of grams to moles 48.5 grams of CaCO3= moles of calcium carbonate. Remember that first, you must find the gram formula mass of the compound. 4) Calculate formula mass Ca → 1 × 40.1 = 40.1 C → 1 × 12.0 = 12.0

3(O) → 3 × 16.0 = 48.0

5) Gram formula mass of CaCO3→100 g (rounded) 100 g of CaCO3-1 mole 48.5 g CaCO3→

48.5 g CaCO3 100 g CaCO3

= 0.485 moles of CaCO3

The gram units are cancelled, leaving mole as the proper unit. 6) Conversion of moles to particles (atoms, molecules, ions) 4.20 moles of hydrogen fluoride = molecules of HF

IL Foundation Series Class 9

Remember that 1 mole of anything has 6.023 × 1023 particles. So, to answer this problem, we just multiply 4.20 with 6.023 × 1023 and get the answer 2.53 × 1024 molecules. The general formula used to calculate: The number of molecules = number of moles × Avogadro number = n × NA = n × 6.023 × 1023 2.8.3 Molarity, molality with questions Molarity (M)

The number of moles of a solute dissolved in one litre of the solution is known as the molarity (M) of the solution. Let 'V' litres of a solution contain 'n' moles of solute dissolved in it. Then, the molarity of the solution, M = n/V(lit) But, n =

Mass of solute (w) Gram Molecular mass of solute

Therefore, M =

W GMW

1 Vlit

If the volume is expressed in ml, then M =

W GMW

1000 Vml

Units: moles/lit: It depends on temperature. The molarity of the solution decreases with the increase in temperature, since the volume of the solution is directly proportional to temperature. Example: Calculate the molarity of 6.3 g of oxalic acid (H2C2O4.2H2O) present in 500 ml of solution. (Molecular weight of oxalic acid is 126). Solution: M=

W GMW

1000 Vml

W = 6.3 g GMW = 126 g Vml = 500ml M=?

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∵ M=

GMW 6.3

126

1000 Vml 1000 500

M = 0.1 Molality (m)

The molality of a given solution (m) is defined as the number of moles of the solute present in one kg (or) 1000 g of the solvent. Let 'n' moles of the solute be dissolved in 1 kg of the solvent, Then, the molality of the solution is m=

n Wkg

But, n =

∴m=

Mass of solute (w) Gram Molecular mass of solute(GMW)

W GMW

1 Wkg

If the mass of the solvent is expressed in grams then, m =

W GMW

1000 Wg

Units: moles/kg Since the quantities of the solute and the solvent are expressed in weights, the molality does not change with the change in temperature. ∴ Molality is independent of temperature. Example: Calculate the molality (m) of a solution containing 0.365 g of HCl in 100 g of O2. Solution: Molality of the solution, m =

W GMW

Mass of the solute, w = 0.365 g Mass of the solvent, W = 100 g GMW of HCl = 36.5 g

1000 Wg

IL Foundation Series Class 9

∴m =

0.365 36.5

1000 100

m = 1/10 = 0.1 moles/kg 2.8.4 Avogadro's law At the same temperature and pressure, equal volumes of all the gases contain an equal number of moles (n) (or) molecules (N). Vα n (or) VαN V n V1 n1

=K =

V2 n2

If V1 = V2 then, n1= n2 2.8.5 Vapour density Vapour density or relative vapour density (VD) is the ratio of the density of gas to the density of hydrogen gas. It has no units. VD =

Density of gas Densityof hydrogen gas

2 × vapour density = Molecular mass Formulae of mole concept: Number of moles of an element = Number of moles of a compound =

Weight of element Gram atomic weight Weight of compound Gram molecular weight

GMW(or) GAW

S. No.

Chemical substance

No. of chemical units

One mole of hydrogen atoms (H)

6.023 × 1023 atoms of hydrogen

1.008 g

One mole of hydrogen molecules (H2)

6.023 × 1023 molecules of hydrogen

2.016 g

One mole of oxygen atoms (O)

6.023 × 1023 atoms of oxygen

16 g

One mole of oxygen molecules (O2)

6.023 × 1023 molecules of oxygen

32 g

Table 2.7 Chemical units of hydrogen and oxygen 63

IS MATTER AROUND US PURE?

It is important to note that while using the unit mole, it is also necessary to specify the chemical unit. For example, 1 mole of hydrogen atoms

6.023 × 1023 atoms of hydrogen

1 mole of hydrogen molecules

6.023 × 1023 molecules of hydrogen

1 mole of carbon dioxide

6.023 × 1023 molecules of carbon dioxide

1 mole of electrons

6.023 × 1023 electrons

1 mole of sodium ions (Na+ )

6.023 × 1023 Na+ ions

Example 1: What is the mass (in grams) of a single atom of chlorine?

(Atomic mass of chlorine = 35.5 )

Solution: 6.023 × 1023 atoms of Cl = gram atomic mass of Cl = 35.5 g 35.5 g of Cl = -6.023 × 1023 of Cl atoms? = 1 Cl atom Mass of one Cl atom = (1 × 35.5)/(6.023 × 1023 ) = 5.9 × 10-23 g Example 2: Calculate the number of moles of atoms in 5.75 g of sodium (atomic mass of sodium = 23 ) Solution: 1 mole of sodium atoms = gram atomic mass of sodium = 23 g 23 g of sodium =1 mol of sodium atom 5.75 g of sodium = 5.75/23 = 0.25 moles (or) According to the formula, Number of moles of an element = Weight of an element/Gram atomic weight = 5.75/23 = 0.25 moles Example 3: How many grams of each of the following elements must be taken to get 1 mole of the element? a) Sodium

b) Chlorine

c) Copper

Solution: The mass of 1 mole of an element is its atomic mass expressed in grams, and the atomic masses of sodium, chlorine and copper are 23 g, 35.5 g and 63.5 g, respectively. Hence, 23 g of sodium, 71 g of chlorine and 63.5 g of copper must be taken to get one mole of each of these elements.

IL Foundation Series Class 9

2.8.6 Specific gravity The density of a substance provides information about the proximity of its particles. A higher density indicates that the particles are more tightly packed together. The density of gas at STP is defined as the mass of gas per unit volume. Density of gas = Units: g/cc or kg/m3

Gram molecular mass Mass = Gram molar volume (22.4 lit) Volume

This unit is big, and chemists usually talk about density using grams per cubic centimetre (g/cm3), where mass is measured in grams and volume in cubic centimetres. Specific Gravity =

Density of the object Density of water

Specific gravity is a dimensionless measure that compares the density of a substance to the density of a reference substance, typically water, at a specified temperature. It provides insight into how much heavier or lighter a given substance is compared to an equal volume of water. The specific gravity of a substance is calculated by dividing its density by the density of the reference substance. Since specific gravity is a ratio, it has no units. This property is commonly used in various industries, such as in determining the buoyancy of objects in fluids and assessing the concentration of solutions. 2.8.7 Equivalent weight The equivalent weight of an element is defined as a number which denotes the number of parts by weight of the element required to combine with or displace 8 parts by weight of oxygen or 1.008 parts by weight of hydrogen or 35.5 parts by weight of chlorine. For example, 23 grams of sodium combines with 8 grams of oxygen to form sodium oxide. Similarly, 23 tons of sodium combines with 8 tons of oxygen to form sodium oxide. In general, 23 parts by weight of sodium combine with 8 such parts by weight of oxygen to form the oxide. The number 23 is known as the equivalent weight of sodium, i.e., the equivalent weight of an element is a number that denotes the number of parts by weight of it combining with 8 parts by weight of oxygen. The analysis of sodium hydride (NaH) shows that 23 parts by weight of sodium combine with 1.008 such parts by weight of hydrogen. Again, common salt or sodium chloride is found to contain 35.5 parts by weight of chlorine in union with 23 such parts by weight of sodium. In all the above examples, the number 23 is termed as the equivalent weight of sodium. And the numbers 8 for oxygen, 1.008 for hydrogen, and 35.5 for chlorine are referred to as the equivalent weights of the respective elements.

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Atomic weight of the element = Valency of the element Equivalent weight of the element (or) Equivalent weight =

Atomic weight of the element Valency of the element

Elements showing more than one valency exhibit different equivalent weights. Example: Equivalent weight of iron in ferrous (Fe+2 ) compounds is 56/2 = 28, whereas its equivalent weight in ferric (Fe+3) compounds is 56/3 = 18.66. Equivalent weight of an element in terms of electrons

Consider the following reaction. Mg + Cl2 → MgCl2 1 atom of Mg loses 2 electrons to become Mg2+ ions Mg → Mg2+ + 2e1 mole of Mg→2 moles of electrons 24 g of Mg→2 N electrons 12 g of Mg→1 N electrons Thus, the equivalent weight of an element is that weight of the element which loses or gains the Avogadro number (N) of electrons. Gram equivalent weight Equivalent weight expressed in grams is known as gram equivalent weight. Number of gram equivalents or equivalents of a substance in a given weight =

Weight of the substance in grams Equivalent weight of the substance in grams

Example: One equivalent of chlorine = 35.5 g Two equivalents of chlorine = 71.0 g Three equivalents of chlorine = 106.5 g

IL Foundation Series Class 9

2.8.8 Mole fraction The ratio between the number of moles of a component of the solution to the total number of moles of all components of the solution is called the mole fraction of that component. Suppose 'A' is a solute and 'B' is a solvent in a binary solution. Let the number of moles of A be nA and that of B be nB in the solution. Total number of moles in the solutions = nA + nB Then, the mole fraction of solute A is given by, nA

XA =

nA+nB

Similarly, the mole fraction of the solvent B in a solution is nB

XB =

nA+nB

The sum of the mole fractions of all the components in the solution is equal to unity. ∴XA + XB = 1 For a solution containing n1, n2, n3,…, etc., moles of the various components, the mole fraction of ith component can be written as Xi =

ni n1+n2+n3+.....

Mole fraction in dilute binary solution

The mass of the solute (wA) is much less compared to that of the solvent (WB). Under these conditions number of moles of the solute (nA) is neglected in comparison with the number of moles of solvent (nB). Then, the mole fraction of the solute XA = XA= ⇒XA=

nA nA+nB nA nB

becomes,

(∵nA<<nB )

wA.WB WA.wB

Here, wA= Mass of solute WA= Gram molecular mass of solute WB= Mass of solvent wB= Gram molecular mass of solvent

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Example: 3.65 g of HCl is dissolved in 16.2 g of water. Calculate the mole fractions of HCl and water. Solution: Mass of HCl, wA= 3.65 g Gram Molecular mass of HCl, WA= 3 6.5 g Mass of water, wB=16.2 g Gram Molecular mass of water, WB= 18 g Number of moles of HCl, nA = Number of moles of H2O, nB =

wA WA WB wB

∵ Mole fraction of HCl, XA =

Mole fraction of H2O, XB =

= =

3.65 36.5 16.2 18

= 1/10 = 0.1 = 0.9

nA nA+nB 0.1 0.1+ 0.9 nB

nA+nB

= 0.1 0.9 0.1+ 0.9

Verification: XHCl + XH2O = 0.1 + 0.9 = 1

XHCl + XH2O = 1

= 0.9

QUICK REVIEW • Pure substances exist as either elements or compounds. Elements are fundamental forms of matter resistant to breakdown through chemical reactions. On the other hand, compounds consist of two or more distinct types of elements chemically bonded in a fixed proportion. • Properties of a compound differ from those of its constituent elements. In contrast, a mixture exhibits the properties inherent in its constituent elements or compounds. • A mixture is comprised of multiple substances, which can include elements and/or compounds, blended in varying proportions. Mixtures can undergo separation through various techniques to isolate pure substances effectively. • Colloids, another type of heterogeneous mixture, contain particles too small to be seen with the naked eye but large enough to scatter light. Colloids find practical applications in both industrial processes and daily life, with dispersed particles termed the dispersed phase and the medium holding them called the dispersion medium. • A solution is a homogeneous mixture of two or more non-reacting components. Formation of solution is a process. • A solution whose molar concentration is known as a standard solution.

IL Foundation Series Class 9

• The ideal method of expressing concentration is molality (m). • The commonly used method of expressing concentration is molarity (M). • Mass by volume percentage (w/v) indicates the mass of solute in 100 ml of solution. • Mass by weight (w/w) indicates the mass of solute in 100 g of solution. • The number of millimoles of the solute present in Vml of the solution is given as product in V litres of solution is given as M × V. • When a solution is diluted, its molarity decreases. • V1 M1=V2 M2 where V1= Volume of the solution before dilution, M1= Molarity of the solution before dilution, V2= Volume of the solution after dilution and M2= Molarity of the solution after dilution. • Molarity (M) =

W GMW

1000 V

, W = weight of the solute in grams and,

V = volume of the solution in millilitres. • Equivalent weight of a substance expressed in grams is known as gram-equivalent weight or gram equivalent or equivalent. • Molality (m) =

W GMW

1000 Wkg

• Mole fraction of the solute = Xsolute = n2 n1+ n2

n1 n1 + n2

. Mole fraction of the solvent = Xsolvent =

, where n1 and n2 are the number of moles of solute and solvent.

• For a binary solution, Xsolute + Xsolvent = 1. • Weight percentage, molality, and mole fraction are independent of temperature.

WORKSHEET - 1 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER I.

Pure substances

1. Which of the following is a pure substance? a. Milk

b. Seawater

c. 24-carat gold

d. Soil

c. Water

d. Soil

c. Compound

d. Mixture

2. In the given, which one is in the elemental form? a. CO

b. N2

3. What is the impure substance? a. Element

b. Molecule

IS MATTER AROUND US PURE?

4. Which property does not describe a compound? a. It is a pure substance. b. It is mixed in any proportion by mass. c. It cannot be separated into constituents by physical means. d. It's composed of two or more elements. 5. Suspension comes under a. Element

b. Compound

c. Homogeneous mixture

d. Heterogeneous mixture

6. True solution comes under

II.

a. Heterogeneous mixture

b. Homogeneous mixture

c. Compound

d. Molecule

Impure substances

1. The number of components present in a binary solution is a. 1

b. 2

d. 4

2. Among the following mixtures, the solution is a. Sand in water

b. Chalk in water

c. Sugar in water

d. Water in kerosene

3. 10 g of NaCl is dissolved in one litre of water, then the solute is a. NaCl

b. Water

c. Either a or b

d. Both a and b

4. The component present in lesser quantity in a binary solution is a. Solvent

b. Solute

c. Solution

d. Either a or b

5. If air is taken as a binary solution, the solvent is a. N2

b. O2

c. CO2

6. The size of colloidal particles is in the range of

a. 0.1-1 nm

b. 1 nm-100 nm

c. 100 nm-1000 nm

d. 1000-10000 nm

d. H2

IL Foundation Series Class 9

7. Statement (A): The solution contains only one component. Statement (B): Starch in water is an example of a homogeneous mixture. Statement (C): The solution is a homogeneous mixture. a. All the above statements are correct

b. All the above statements are incorrect

c. A and B are correct, and C is incorrect

d. A and B are incorrect, and C is correct.

8. Assertion (A): The solution of sugar and common salt in water is an example of ternary solution. Reason (R): Ternary solution contains three components.

a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 9. Among the following, the example of liquid in solid type solution is a. Mercury in gold

b. Alcohol in water

c. Camphor in air

d. Moisture in air

10. The solution of soda water is an example for a. Gas in gas

b. Gas in liquid

c. Liquid in gas

d. Solid in liquid

11. Among the following, the example of solid in solid type solution is a. Sugar in water

b. Camphor in air

c. Brass

d. Mercury in gold

12. Among the following, the example of solid in gas type solution is a. Camphor in air

b. Sugar in water

c. Mercury in gold

d. Moisture in air

13. Which of the following acts as a solvent other than water a. C2H5OH

b. CHCl3

c. CCl4

d. All of these

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14. Statement A: If a solution contains the maximum quantity of solute, then it is called an unsaturated solution. Statement B: If a solution contains less than the maximum quantity of solute, then it is called a saturated solution. Statement C: If a solution contains more than the maximum quantity of solute, then it is called a supersaturated solution. a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct 15. Assertion (A): Sugar in water is an example of an aqueous solution. Reason (R): In aqueous solutions, water acts as a solvent. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct III. Solutions

1. Mass percentage is equal to a.

mass of solute mass of solution × 100

mass of solute volume of solution × 100

volume of solute volume of solution × 100

volume of solute mass of solution × 100

2. Volume percentage is equal to

mass of solute volume of solution × 100

volume of solute mass of solution × 100

mass of solute mass of solution × 100

volume of solute volume of solution × 100

IL Foundation Series Class 9

3. 4 ml of alcohol is present in 36 ml of water. Then, the volume percentage of the solution is a. 10

b. 20

d. 40

4. 10 g of sugar is present in 190 g of water. Then, the mass percentage of the solution is a. 20

b. 10

d. 15

5. Atmospheric pollution in cities is expressed in a. Mass percentage

b. Volume percentage

c. ppm

d. Mass by volume percentage

6. 10 g of sodium carbonate is present in 120 g of its aqueous solution. Then, the mass percentage is

a. 8.33

b. 9.33

c. 4.00

d. 10.00

7. 10 ml of hexane is mixed with 4 0ml of heptane. Then, the volume percentage of the solution is a. 25

b. 15

c. 20

d. 10

8. Which concentration method is commonly used in medicine and pharmacy? a. (w/W)%

b. (v/V)%

c. (w/V)%

d. Molarity

9. Statement (A): Mass percentage has no units. Statement (B): Volume percentage has no units. Statement (C): The concentration of a solution is the amount of solute present in unit volume of the solution. a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct 10. Statement (A): Mass percentage (w/W)% is independent of temperature. Statement (B): Volume percentage (v/V) % is depends on temperature. Statement (C): Mass by volume percentage (w/V) depends on temperature. a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct

IS MATTER AROUND US PURE?

11. Statement (A): Mass percentage has no units. It is independent of temperature. Statement (B): Volume percentage has no units. It depends on the temperature. Statement (C): The concentration of pollutants in water or atmosphere is often expressed in terms of μgml-1 or ppm.

a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct 12. The unit of molarity is a. g/lit

b. moles/lit

c. equivalents/lit

d. moles/kg

13. The amount of Na2CO3 present in 100 ml of 0.1M Na2CO3 solution is a. 2.12 g

b. 0.53 g

c. 1.06 g

d. 5.3 g

14. The number of moles of the solute present in 600 ml of 0.05M solution is a. 0.01

b. 0.02

c. 0.03

d. 0.04

15. The volume of water that must be added to 100 ml of 0.5M NaOH solution to get 0.2M solution is

a. 50 ml

b. 100 ml

c. 150 ml

d. 200 ml

16. The number of particles present in a mole a. 6.023 × 1024

b. 6.023 × 1023

c. 6.023 × 10-23

d. 6.023 × 10-24

17. Gram molecular weight of oxygen is a. 16 a.m.u

b. 32 a.m.u

16 g

d. 32 g

18. The element taken as a reference in measuring atomic mass is a. C-14

b. C-12

c. C-13

19. The weight of one C-12 atom is

a. 1.9926 × 1023 g

b. 1.9926 × 10-23 kg

c. 1.9926 × 1023 kg

d. 1.992 6× 10-23 g

d. None of these

IL Foundation Series Class 9

20. One a.m.u equals to a. 1.66 × 1027 kg

b. 1.66 × 10-27 g

c. 1.66 × 1027 g

d. 1.66 × 10-27 kg

21. Assertion (A): Avogadro number is denoted by NA. Reason (R): Gram molecular weight of oxygen is 32 a.m.u. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 22. Statement (A): C -12 is the reference for measuring atomic mass as it gives an accurate atomic mass. Statement (B): Atomic mass can be expressed as one unified mass Statement (C): One gram atom of oxygen mass is 16 a.m.u. a. All the above statements are correct b. All the above statements are incorrect c. A and B are correct, but C is incorrect d. A and B are incorrect, but C is correct

WORKSHEET - 2 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER 1. The molarity of a solution is M. To decrease the molarity of this solution to M/2 a. The weight of the solute is to be doubled b. The weight of the solvent is to be doubled c. The volume of the solvent is to be doubled d. The volume of the solution is to be doubled 2. The concentration units which depend on temperature would be a. Normality

b. Mass-mass percentage

c. Molality

d. Mole fraction

IS MATTER AROUND US PURE?

3. Statement (A): Molarity is dependent on temperature. Statement (B): Molarity increases with decreasing in temperature. Statement (C): Normality is independent on temperature. a. All the above statements are correct b. All the above statements are incorrect c. A and B are correct, but C is incorrect d. A and B are incorrect, and C is correct 4. Assertion (A): Molarity of a solution decreases with an increase in temperature. Reason (R): As the temperature increases, the volume of the solution increases. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 5. The molarity of pure water is a. 55.6

b. 50

c. 100

d. 18

6. One molal solution is one that contains a. 1 g of the solute in 1000 g of solvent b. 1 gram mole of solute in 100 ml of solution c. 1 gram mole of solute in 1000 lit of solution d. 1 gram mole of solute in 1000 g of solvent 7. The units of molality are a. mole/lit

b. mole/ml

c. mole/kg

d. gram equivalent/kg

8. 0.1 gram mole of urea is dissolved in 100 g of water. The molality of the solution is a. 0.1 m

b. 0.01 m

c. 0.001 m

d. 1.0 m

9. 16 g of methanol is present in 100 ml of the solution. If the density of the solution is 0.96 g/ml. Then, the molality of the solution is a. 6.25 m

b. 5.20 m

c. 5.75 m

d. 5 m

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10. 6 moles of a solute is present in 100 moles of a solution. Then, the mole fraction of the solute is a. 0.6

b. 0.06

c. 0.03

d. 0.94

11. 60 g of urea is dissolved in 90 g of water. Then, the mole fraction of the solute is a. 1/5

b. 1/6

c. 5/6

d. 4/5

12. Statement (A): Molarity always equals to its molality. Statement (B): The amount of solvent in 1M aqueous solution is more than in 1 m aqueous solution. Statement (C): At 25° C, for a given solution M = m, then at 50° C, the correct relation is M < m. a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct 13. Assertion (A): Molality is independent of temperature. Reason (R): There is no volume factor in the expression of molality. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 14. Assertion (A): Mole fraction has no units. Reason (R): Mole fraction is a ratio of the number of moles of solute to the number of moles of solvent. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 15. The number of moles of a solute per kilogram of a solvent is called a. Molarity

b. Molality

c. Normality

d. Formality

IS MATTER AROUND US PURE?

16. The molality of a solute in the solution is the a. Mass of solute per dm3 of solution

b. Amount of solute per dm3 of solution

c. Amount of solute per kg of solution

d. Amount of solute per kg of solvent

17. Assertion (A): One molar aqueous solution is always more concentrated than one molal aqueous solution. Reason (R): The amount of solvent in 1M solution is less than in 1 m solution. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 18. Among the following, naphthalene is soluble in a. Water

b. Alcohol

c. Benzene

d. Chloroform

19. Among the following substances, solubility in water decreases with an increase in temperature of

a. NaCl

b. KNO3

c. NH4Cl

d. Ce2(SO4)3.9H2O

20. 2.5 g of NaCl is dissolved in 25 g of water. Then, the solubility of NaCl in its solutions is a. 10 g

b. 1 g

c. 0.1 g

d. 100 g

21. If the pressure applied to a gas increases, then the mass of the gas dissolved is a. Decreases

b. Increases

c. No change

d. Forms product

22. Statement A: The solubility of gases in liquids increases with the increase of pressure. Statement B: The solubility of gases in liquids decreases with the increase in temperature. Statement C: The solubility of gases in liquids decreases with the increase of pressure. a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct

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23. Assertion (A): Solubility of KNO3 in water at 1atm pressure is greater than the solubility of KNO3 in water at 0.1 atm pressure. Reason (R): Solubility of solids in liquids is unaffected by pressure. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 24. Assertion (A): NaCl dissolves in water but not in kerosene. Reason (R): Like dissolves like. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 25. The solubility of gas in a liquid increases with a. Increase in temperature

b. Reduction of gas pressure

c. Decrease in temperature

d. Amount of liquid taken

26. One mole of any substance is equal to a. Gram atom

b. Gram molecule

c. Avogadro number of entities

d. All of the above

27. The volume of any gas at STP is a. 22.4 lit

b. 2.24 lit

c. 11.2 lit

d. 1.12 lit

28. The weight of one mole of atoms of an element is called a. Gram molecular weight

b. Gram atomic weight

c. Molecular weight

d. Atomic weight

29. Statement (A): One mole of any substance contains 6.023 × 1023 particles Statement (B): The gram formula mass of CaCO3 is 100 g Statement (C): 6.023 × 1023 particles of MgCl2 mass is 95 a.m.u a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect

IS MATTER AROUND US PURE?

c. A and B are correct, and C is incorrect d. A and B are incorrect, and C is correct 30. 7.4 moles of iron weight equals a. 111.1 g

b. 212.2 g

c. 414.4 g

d. 112 g

c. 44 × 1023

d. 9.4 × 1023

31. The number of atoms in 7.4 moles of iron is a. 44.4 × 1023

b. 22. 2 × 1023

32. Assertion (A): The formula mass in grams represents the mass of one mole of that substance. Reason (R): Gram formula unit mass of a substance contains Avogadro's number of entities. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 33. Statement (A): The formula mass in grams represents the mass of one mole of that substance. Statement (B): One mole of any substance contains 6.023 × 1023 particles. Statement (C): At STP, one mole of any gas occupies 22.4 litres of volume. a. All the statements A, B, and C are correct b. All the statements A, B, and C are incorrect c. A and B, are correct and C is incorrect d. A and B, are incorrect and C is correct 34. The molar volume of any gas at STP is a. 22,400 lit

b. 22,400 ml

c. 22.4 lit

d. Both b and c

c. 2.0 × 1024

d. 1.07 × 1022

35. The number of atoms in 1 gram of iron is a. 2 × 1022

b. 1.07 × 1020

36. Assertion (A): Avogadro's number = 6.02 3 × 1023 atoms or molecules or ions or electrons. Reason (R): Standard temperature is 0° C or 273 K, and standard pressure is 1 atmosphere. a. Both A and R are true, and R is the correct explanation of A b. Both A and R are true, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 80

ATOMS AND MOLECULES

3.1 INTRODUCTION An Indian philosopher, Maharishi Kanad, proposed that matter (padarth) can be continuously divided into the smallest form of particles. He identified these as the smallest particles beyond which further division would not be possible and called them paramanu. Around the same era, Greek philosopher Democritus called these indivisible particles as atoms (atom means indivisible).

Fig. 3.1 Maharishi Kanad, Democritus

The foundation of all this was based on philosophical ideas. Until the 18th century, there weren't many chances to test these ideas through experiments. At that time, experimental work was limited, making it hard to prove these early thoughts. By the end of the 18th century, scientists recognised the difference between elements and compounds. Elements are substances having molecules of the same type of atoms, i.e., elements comprised of identical atoms. A compound is composed of two or more elements chemically combined with each other in a definite proportion by weight.

ATOMS AND MOLECULES

3.2 LAWS OF CHEMICAL COMBINATION The combination of two or more elements in a chemical reaction is governed by some combination laws. Due to these laws, changes often occur in the matter. These laws are known as the laws of chemical combination. The following are some of the laws from the different laws of combination1. Law of conservation of mass 2. Law of constant proportions 3. Law of multiple proportions 3.2.1 Law of conservation of mass When a chemical reaction occurs, does the mass of the elements involved in the reaction change? No! French chemist Antoine Lavoisier created a law regarding the change of mass in a chemical reaction. His given law states, "Mass can neither be created nor destroyed in a chemical reaction". That means during chemical changes, the sum of masses of all the reactants remains equal to the mass of products. That isThe sum of masses of all the reactants = The sum of masses of all the products.

LAW OF CONSERVATION OF MASS String Small test tube

Sodium hydroxide solution Copper sulfate solution

Sodium sulphate solution

Small test tube

Mix Thoroughly

Copper hydroxide precipitate

67.25g

One pan electric balance

Fig. 3.2 Law of conservation of mass

3.2.2 Law of constant proportions The law of constant proportions was established by Joseph L. Proust. This law is also known as the law of definite proportions. In a chemical substance, the elements are always present in definite proportions by mass.

IL Foundation Series Class 9

For example, • Water obtained from any source contains hydrogen and oxygen combined in the ratio of 1: 8 by weight. • Sodium chloride obtained from rock salt or seawater always contains sodium and chlorine in the ratio of 23: 35.5 by weight. • Carbon dioxide is prepared using two methods, but the ratio of carbon and oxygen is always 3:8. 3.2.3 Law of multiple proportions It states that, when two elements combine to form more than one compound, the several weights of the first element that combine with a fixed weight of the second are in the ratio of their small whole number. OR Different weights of an element that combine with a fixed weight of the other element bear a simple numerical ratio. For example, i. Carbon and oxygen combine to form carbon monoxide and carbon dioxide. 2C

→

2CO

2×12g 32g 6g C

8g +

12g

32g

16g

→

CO2

In carbon monoxide, exactly 8 g of oxygen is combined with 6 g of carbon. Whereas, in CO2, 16 g (exactly twice 8 g) of oxygen is combined with 6 g of carbon. The weights of oxygen that combine with a fixed weight of carbon (6 g), are in the ratio of 1:2. ii. Hydrogen and oxygen combine to form H2O and H2O2 H2O = 1: 8 H2O2 = 1:16 The weights of oxygen that combine with a fixed weight of hydrogen (2 g), are in the ratio of 1:2.

ATOMS AND MOLECULES

Weight of nitrogen

Weight of oxygen

Wt. of nitrogen=28 fixed

Weight of oxygen which combines with a fixed weight of nitrogen

N2O

NO(N2O2 )

N2O3

N2O4

N2O5

Oxides of nitrogen

Table 3.1 Oxides of nitrogen

From the above observations, weights of oxygen, which combine with a fixed weight of nitrogen (28 gm), are in the ratio of 16:32:48:64:80 or 1:2:3:4:5.

3.3 WHAT IS AN ATOM? The word 'atom' is derived from the Greek word 'a-tomio' (which means indivisible). According to Dalton's atomic theory, all matter, whether an element, a compound, or a mixture, is composed of small particles called atoms. An atom is the smallest particle of the element that can retain all its chemical properties. More than millions of atoms, when stacked, would make a layer barely as thick as this sheet of paper. Atomic radius is measured in nanometres. 1 109

= 1 nm or 1 m = 109 nm

1 nm = 10-9 m 3.3.1 Modern-day symbols of atoms of different elements In daily life, abbreviations are generally written in place of lengthy names to save time and labour. For example, RBI represents the Reserve Bank of India, and ATM indicates Automatic Teller Machine. Similarly, elements are represented by symbols. Definition: The shorthand notation of an element is called a symbol. For example, oxygen is represented by the letter 'O'. Symbols of elements with a single letter

For some elements, the first letter of their English names represents their symbols. For example, nitrogen (N).

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Nitrogen 14.007 Fig. 3.3 Symbol for nitrogen Name of the element

Symbol

Hydrogen

Oxygen

Nitrogen

Carbon

Fluorine

Table 3.2 Some elements and their symbols

Similarly, Boron, Sulphur, Phosphorus, Iodine, and Uranium are represented by the letters B, S, P, I, and U, respectively. Symbol of elements with two letters

Two-letter symbols are used to represent most of the elements. The list of such elements is given below. The first letter is always written in upper case, and the second letter is always in lower case. For example, the symbol of cobalt is Co but not CO because CO represents a molecule of carbon monoxide.

Fig. 3.4 Symbol for aluminium

ATOMS AND MOLECULES

Element

Symbol

Element

Symbol

Aluminium

Helium

Argon

Lithium

Arsenic

Magnesium

Astatine

Manganese

Barium

Molybdenum

Beryllium

Neon

Bismuth

Nickel

Bromine

Palladium

Table 3.3 Some elements and their two-letter symbols

There are some elements whose symbol is derived from their Latin names. Some of them are given below. Element

Latin name

Symbol

Antimony

Stibium

Copper

Cuprum

Gold

Aurum

Iron

Ferrum

Lead

Plumbum

Mercury

Hydragyrum

Potassium

Kalium

Silver

Argentum

Sodium

Natrium

Tin

Stannum

Tungsten

Wolfram (German Name)

Table 3.4 Elements with their Latin names and symbols

Tungsten 183.84 Fig. 3.5 Symbol for Tungsten

IL Foundation Series Class 9

There are some elements that are named after scientists. Element

Scientist's name

Symbol

Curium

Madam Curie

Einsteinium

Albert Einstein

Fermium

Enrico Fermi

Einsteinium

Nobelium

Alfred Nobel

252

Mendelevium

Mendeleev

Fig. 3.6 Symbol for Einsteinium

Table 3.5 Elements with scientists' names and symbols

Some elements are named after countries and laboratories. Element

Country and Laboratory

Symbol

Berkelium

City of Berkely

Californium

University of California

Polonium

Poland

Americium

America

Ruthenium

Russia

Germanium

Germany

Fig 3.7 Symbol for Polonium

Table 3.6 Elements with the names of countries and laboratories

Some elements are named after certain planets. Elements

Name of the Planet

Symbol

Uranium

Uranus

Neptunium

Neptune

Plutonium

Pluto

Table 3.7 Elements with the names of planets 87

ATOMS AND MOLECULES

3.3.2 Atomic mass An individual atom is very small and assumed to be spherical. The radius of an atom is in the order of 10-10 m. If ten grams of iron or ten grams of copper is taken and made into 1023 small pieces of the same size, we may get each fine particle as an atom. The real mass of an individual atom is as small as about 10-26 kg. The real masses are not practicable to measure. Hence, relative masses are used. The concept of relative masses of atoms was first introduced by Dalton by taking water molecules. A water molecule has two H atoms and one O atom. If the mass of an H atom is taken as one unit, then that of an O atom will be 16 units. The unit of relative mass is the atomic mass unit. It is familiarly known as 'amu', but now denoted by 'u' and is called unified mass. The present reference of relative masses is a 12C isotope. The relative mass of 12C is taken as 12amu. 1 amu =

Mass of 12C 12

Hence, 'amu' is defined as one-twelfth of the mass of a 12C isotope: Uniﬁed atomic mass unit 1.660538782(83) X 10-27 kg

C612

/12

1 amu = 931.46 MeV/C2

Fig. 3.8 Unified atomic mass unit

Example: If the relative mass of He is taken as one unit, what is that of calcium? Solution: The relative masses of He and Ca are 4 and 40, respectively. Compared to He as 1, that of Ca is 40/4=10 Atomic weight is stated as the number of times the weight of one atom is heavier than the unified mass. Atomic Weight =

Weight of one atom of an element 1/12 × weight of 12C

Atomic weight is also called atomic mass, and has no units since it is relative. Though the real masses of atoms are very small, we now have sophisticated techniques, like mass spectrometry, for the determination of atomic masses fairly accurately. When we use atomic masses of elements in calculations, we use average atomic masses. 88

IL Foundation Series Class 9

Element

Atomic mass (U)

Hydrogen

Carbon

12U

Nitrogen

14U

Table 3.8 Atomic masses of a few elements

3.3.3 Average atomic mass Many naturally occurring elements exist in more than one form and are called isotopes. When we take into account the existence of isotopes and their per cent abundances, the average mass of that element can be computed. The average isotopic mass of chlorine is 35.5. This is because, naturally, chlorine is available in two stable isotopic forms, 35Cl and 37Cl, in the approximate mass ratio of 3:1. Example: Neon is naturally available as 20Ne and 22Ne with an average atomic mass of 20.2. Calculate the relative abundance of the heavier isotope. Solution: L et the percentage abundance of 22Ne be x. The percentage abundance of 20 Ne is 100-x. 20.2 =

(100-x)20+22x 100

Percent abundance of 22Ne = 10% Atomic masses are fractional, but mass numbers are whole numbers. A mass number is usually given for elements. It is also given for the fundamental particles. The mass numbers of electrons, protons, and neutrons are 0, 1, and 1, respectively. The mass number of an element denotes the sum of the number of protons and neutrons present in the nucleus of its atom. 3.3.4 Molecular mass Molecular mass is the sum of masses of all sub-atomic particles present in a molecule. It is generally calculated as the sum of the atomic masses of the elements constituting the molecule. Molecular weight is also measured relatively. It has no units.

Molecular Weight =

Weight of one mole of substance 1/12 × weight of 12C

The molecular weight of water is the sum of the atomic weight of O and twice the atomic weight of H. It is equal to 18.02 u. 89

ATOMS AND MOLECULES

Ionic substances do not contain discrete molecules as their constituent units. For example, in sodium chloride, sodium cations (Na+ )and chloride anions (Cl-) are present in the ratio 1:1. For such substances, the formula mass is calculated instead of molecular mass. The formula mass of sodium chloride is the sum of atomic masses of sodium and chlorine. It is equal to 58.2u. SOLVED EXAMPLES Example 1: What is the formula mass of gypsum? Solution: The formula of gypsum is CaSO4.2H2 O Formula mass of CaSO4.2H2O = (1×40) + (1×32) + (4×16) + (2×18) = 172u Example 2: Calculate the mass of a fructose molecule. Solution: The molecular formula of fructose is C6H12O6 Molecular mass of C6H12O6 = (6×12u) + (12×1u) + (6×16u) = 180u While atomic masses are determined using Dulong and Petit's law, molecular masses are determined by Cannizzaro's method and the vapour density method. If the specific heat of a metal is given in calg-1, the product of the atomic mass of the element and specific heat is approximately 6.4. Atomic mass =

6.4 Specific heat

If the valency of an element and its equivalent mass are known, the product gives atomic mass. Atomic mass = Valency × Equivalent mass In Cannizzaro's method, atomic mass is defined as the smallest mass of the element present in the molecular mass of any of its compounds. In this method, the molecular mass of a number of compounds in which the element is present is determined. Each compound is analysed, and the mass of an element is obtained from the molecular mass of each compound. The lowest mass of the element from different compounds is taken as the atomic mass. If the vapour density of a volatile substance is obtained, molecular weight is taken as twice the vapour density. Molecular mass =2 × Vapour density In Victor Mayer's method, the volume of air displaced by a fixed mass of a volatile substance is determined. The volume is translated to standard temperature and pressure conditions. The mass corresponding to 22400 cc of air gives the molecular mass.

IL Foundation Series Class 9

Molecular mass =

22400× Weight of the substance Volume of air displaced at STP

The molecular weight of substances can also be determined based on the following methods: 1. Raoult's law method 2. Depression in the freezing point method 3. Elevation in boiling point method 4. Osmotic pressure method 5. Measurement of density, etc. The mass spectroscopic method, however, is the most accurate experimental method for the molecular masses. The molecular weight of a homogenous mixture is called effective molecular weight. If the number of moles of different substances in the mixture is n1,n2,..., then, Effective molecular weight of the mixture =

(n1 M1+n2 M2+⋯..) (n1+n2+ ....)

3.4 WHAT IS A MOLECULE? A molecule is the smallest particle of an element or a compound capable of independent existence under ordinary conditions. It shows all the properties of the substance. 3.4.1 Molecules of elements The number of atoms present in one molecule of an element is called the atomicity of that element. • Monoatomic elements - Helium (He), Argon (Ar)... etc. • Diatomic elements - Oxygen, Hydrogen, Nitrogen... etc. • Triatomic element - Ozone • Tetratomic element - Phosphorus (P4) • Polyatomic elements - Sulphur (S8), etc.

ATOMS AND MOLECULES

O3 O

Oxygen

Ozone

Sulphur

Fig. 3.9 Molecules of oxygen, ozone, and sulphur

3.4.2 Molecules of compounds Atoms of different elements join together in definite proportions to form molecules of compounds. For example - Water - H2O, Ammonia - NH3, Carbon dioxide - CO2

Fig 3.10 Water molecule

3.4.3 What is an ion? Definition: Charged species that are formed by losing (or) gaining electrons are known as ions. Ions may contain a single atom or a group of atoms. Types of Ions

Ions are of two types based on the nature of the charge they carry. 1. Electropositive ion or Cation: An ion having a positive charge on it is known as an electropositive ion or a cation. For example: K+, Ag+, Pb+2, etc. 2. Electronegative ion or Anion: The ion having a negative charge on it is known as an electronegative ion or an anion. For example: Cl-, O-2, SO-2 , etc. 4

IL Foundation Series Class 9

3.5 WRITING CHEMICAL FORMULAE Definition: The chemical formula of a compound is a symbolic representation of its composition. Importance: A Chemical formula gives the exact number of atoms of the same or different elements present in a chemical substance. To write chemical formulae, the symbols of elements and the combining capacities of elements are required. Note: Here, the combining capacity of an element indicates the number of bonds formed by that element. For example, Hydrogen

Chlorine

Its valency is 1.

Its combining capacity is 1.

Number of bonds formed by it is 1.

H-Cl: The combining capacity of hydrogen and chlorine is 1. So, only 1 bond is formed between H and Cl . Rules for writing chemical formulae

• The valencies or charges on the ion must be balanced. • In writing formulae of a compound, the metal comes first, followed by the non-metal. 3.5.1 Formulae of simple compounds Criss-Cross method to write the formulae for compounds

1) First, we write the constituent elements in the compound. 2) Then, write their valencies. 3) After that, we must cross over the valencies of the combining atoms.

ATOMS AND MOLECULES

Examples: 1. Formula of hydrogen chloride: Cl Symbol H Hydrogen chloride (HCl) Valency 1

2. Formula of hydrogen Sulphide: S

Symbol H

Hydrogen sulphide (H2S) 2

Valency 1

3. Formula of carbon tetrachloride:

Symbol C

Carbon tetrachloride (CCl4) Valency 4

4. Formula of magnesium chloride: Symbol Mg

Valency 2

3.6

Cl Magnesium chloride (MgCl2) 1

MOLECULAR MASS AND MOLE CONCEPT

3.6.1 Molecular mass The Molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. So, we can simply use the molecular formula to calculate the molecular mass. For example, The molecular mass of water (CH4) will be the sum of the atomic mass of four atoms of hydrogen and the atomic mass of one atom of carbon. Molecular mass = Atomic mass of four hydrogen atoms + Atomic mass of one carbon atom

= 4 × 1 + 1 × 12

IL Foundation Series Class 9

= 16 u

3.6.2 Formula unit mass The formula unit mass of a substance is the sum of the atomic masses of all atoms in a formula unit of a compound. Note: The word formula unit mass is used for those substances whose constituent particles are ions. For Example. Formula unit mass of sodium chloride (NaCl) = 1 × 23 + 1 × 35.5 = 58.5U Formula unit mass of CaCl2 = Atomic mass of Ca + 2 × atomic mass of Cl

= 40 + 2 × 35.5 = 40 + 71 = 111U

3.6.3 Mole concept The actual meaning of a mole is a heap, but here, it represents the number of atoms or molecules or particles present in a fixed amount of the substance. This fixed amount of the substance is known as atomic mass for atoms and molecular mass for molecules. (Or) One mole is defined as the amount of the substance that contains as many particles as atoms exactly present in 12 g of the C-12 isotope. (Or) The amount of the substance that contains the same number of elementary particles (atoms, molecules, ions or electrons) as the number of atoms present in 12g of Carbon (C-12). For example: Let us calculate the number of atoms present in 12g of C-12 isotope. One C-12 atom weighs 1.9926×10-23 g Number of atoms in 12g per mole of C-12 isotope =

12 g/mol 1.9926×10-23 g/atom

= 6.023 × 1023 atoms/mol

This number of entities in one mole of a substance is known as Avogadro's number. It is denoted by 'N’ or 'NA’. Avogadro's number = 6.023 × 1023 number of particles (atoms or molecules or ions or electrons). This number is named in honour of the Italian scientist Avogadro. It is used as a reference for most of the calculations and equations found in chemistry. Quite often, we use the unit dozen to represent 12 articles, irrespective of their nature. For example, one dozen books mean 12 books, whereas one dozen apples mean 12 apples. Similarly, chemists use

ATOMS AND MOLECULES

the unit mole to count atoms, molecules, ions, etc. So, a mole is a collection of 6.023 × 1023 particles. A mole represents 6.023 × 1023 particles. 1 mole of carbon atoms 6.022 ◊ 1023 atoms of C 6.022 ◊ 1023 atoms of H 6.022 ◊ 1023 number of that particle 6.022 ◊ 1023 numbers of molecules

1 mole of hydrogen atoms 1 mole of any particle (atoms, molecules, icon)

12g of carbon 1g of H atoms

1 mole of molecules Molecular mass in grams

Fig. 3.11 Different representations of a mole

3.7 EMPIRICAL FORMULA AND MOLECULAR FORMULA A formula represents the chemical composition of the substance. There are three kinds of formulae of compounds. They are 1. Empirical formula 2. Molecular formula 3.7.1. Empirical formula The empirical formula is the simplest formula of a chemical substance. It is the formula of the substance which gives the relative whole number ratio of atoms of each element present in the molecule of a substance. For example, CH is the empirical formula of benzene. This formula indicates that the ratio of the number of 'C' and the number of 'H' atoms is 1:1. It also indicates that the ratio of the mass of 'C' and 'H' is 12:1. It suggests that the weight percentages of carbon and hydrogen in benzene are 92.3% and 7.69%, respectively. Compounds with the same empirical formula possess the same percentage composition of elements. The empirical formula can be determined from the mass percentages of various elements present in a compound.

IL Foundation Series Class 9

The sequence of steps in the determination of empirical formula are:

1. The weight percentage (or weight) of each constituent element is taken. 2. The per cent weight of each constituent is to be divided by its atomic weight. 3. The simplest whole number ratio of the values of step (2) is to be obtained. This may be done by dividing all values with the smallest among them. 4. The whole number is written as a suffix against the symbol of each element. 5. Ionic substances have no molecules. They are represented only by their empirical formula.

SOLVED EXAMPLES Examples 1: An organic compound contains carbon, hydrogen, oxygen, and nitrogen in the weight ratio 3 : 1 : 8 : 3.5. Calculate its empirical formula. Solution: 3

3.5

atio of atoms, C:H:O:N= R : : : = 0.25 : 1 : 0.5 : 0.25 = 1:4:2:1. The empirical formula is 12 1 16 14 CH4O2N Examples 2: The combustion of 0.202 g of a carbon compound gave 0.361 g of carbon dioxide and 0.147 g of water. Determine the empirical formula of the compound. Solution: %C =

0.361×12×100 0.202×44 48.76

C:H:O =

= 48.76%; %H= :

8.07 1

43.17 16

0.147×2×100 0.202×18

= 8.07%; %O = 100-56.83 = 43.17

= 4.06 : 8.07 : 2.7 = 3:6:2

The empirical formula of the compound is C3H6 O2. 3.7.2 Molecular formula The molecular formula of a compound is one which expresses the actual number of atoms of each element present in one molecule. The molecular formula of benzene is C6H6. This indicates that the benzene molecule has six carbon atoms and six hydrogen atoms. Molecular formula = n × Empirical formula, Here, n is the number of empirical formula units repeated in the molecular formula. The number of repeating empirical units is obtained as a ratio of the molecular mass and the mass represented by the empirical formula. n=

Molecular formula mass Empirical formula mass

ATOMS AND MOLECULES

The empirical formula may be the same for different compounds. Formaldehyde, acetic acid, lactic acid, and glucose all have the same empirical formula CH2O. The difference is n = 1 for formaldehyde, n = 2 for acetic acid, n = 3 for lactic acid, and n = 6 for glucose. The empirical formula is the same for all alkenes or cycloalkanes, but the individual members differ in their molecular weights and structural formula. The difference between empirical and molecular formulas is illustrated with some examples in Table. The empirical and molecular formula of some substances: Chemical substance

Empirical formula

Molecular formula

Methane

CH4

Ethane

CH3

C2H6

Acetylene

C2H2

Benzene

C6H6

Oxalic acid

HCO2

H2C2O4

Glucose

CH2 O

C6H12O6

Sucrose

C12H22O11

Hydrogen peroxide

H2O2

Sodium carbonate

Na2CO3

Table 3.9 Empirical and molecular formula of some substances

SOLVED EXAMPLES Example 1: A brominated alkane, on analysis, gave 12.8% carbon and 2.1%H. If its vapour density is 93.95, what is its molecular formula? Solution: Per cent weight of bromine =100-(%C+%H)=100-14.9=85.1 C:H:Br =

12.8 2.1 12

85.1 80

=1:2:1

Empirical formula is CH2Br Molecular weight = 2 × V.D. = 2 × 93.95 = 187.9 n = 187.9/94 = 2 Molecular formula = 2 × CH2Br = C2H4Br2 Example 2: A compound has a molar mass of 98.98 g. The analysis gave 24.27%C, 4.07%H, and 98

IL Foundation Series Class 9

71.6%Cl. Determine the molecular formula of the alkane from which the chlorinated compound is obtained. Solution: C : H : Cl =

24.27 12

4.07 1

71.6 35.5

=1:2:1

Empirical formula is CH2Cl Molecular formula mass

Number of repeating empirical units in the molecular formula =

Empirical formula mass

98.98 49.5

=2 Molecular formula is C2H4Cl2 The formula of the alkane from which the halogenated compound is obtained = C2H6.

3.8 LIMITING REAGENT Many times, the reactions are carried out when the reactants are not present in the amount required by a balanced chemical reaction. In such cases, one reactant may be in excess when compared to the other. The reactant that is present in the lesser amount gets consumed after some time, and then no further reaction occurs even though the other reactant is present in the large amount. Hence, the reactant that gets consumed limits the amount of product formed and is, therefore, called the limiting reagent. If a reaction involves two or more reactants, then the reactant consumed first, limiting the amount of product, is called a limiting reagent or limiting reactant. The best way of finding the limiting reagent is to find out the amount of product formed using both the reactants. The reactant that gives the smaller amount of products is called a limiting reagent. Example: 2H2(g) + O2(g) → 2H2 O(l) If two moles each of H2 and O2 react, then the limiting reagent is H2, as it is completely consumed in the reaction. So, a limiting reagent is the reactant that is entirely consumed during the completion of the reaction. A reactant that is not completely consumed is often called an excess reactant. In the above example, the excess reactant is oxygen.

ATOMS AND MOLECULES

SOLVED EXAMPLES Example 1: 0.5 mole BaCl2 is mixed with 0.2 mole Na3PO4. The maximum number of moles of Ba3(PO4)2 that can be formed is Solution: 3BaCl2 + 2Na3PO4 → Ba3(PO4)2 + 6NaCl Number of moles of Ba3(PO4)2 formed by BaCl2 = 1/3 × 0.5 = 0.166 Number of moles of Ba3(PO4 )2 formed by 0.2 mole Na3PO4 = 1/2 × 0.2 = 0.1 Then, Na3PO4 will be the limiting reactant, and the actual amount of product will be 0.1 mole. Example 2: In the following reaction: 4NH3(g)+ 5O2(g) → 4NO(g) + 6H2O(l) when 1 mole of ammonia and 1 mole of O2 are mixed, then the number of moles of NO formed will be Solution: 1 mole of NH3 (on complete reaction) gives 1 mole NO. Similarly, 1 mole of O2 (on complete reaction) gives 4/5, i.e., 0.8 mole NO. Thus, O2 will be the limiting reactant, and the actual amount of NO formed in the reaction will be 0.8 mole.

QUICK REVIEW • All chemical reactions between elements take place according to certain laws. These are called the law of conservation of mass, the law of definite proportions, the law of multiple proportions, and the law of combining volumes, which are important. • According to the law of conservation of mass, in a chemical reaction, the total mass of the products is equal to the total mass of the reactants or mass is neither created nor destroyed in a chemical reaction. • The law of conservation of mass holds good for all chemical reactions, but not for nuclear reactions. • According to the law of definite proportions, a given chemical substance (compound) always contains the same elements combined in a fixed proportion by weight. • Shorthand notation of an element is called a symbol. For example, oxygen is represented by the letter 'O'. • Symbols of elements with a single letter: For some elements, the first letter of their English names represents their symbols. • The atomic weight or atomic mass of an element is a relative mass, and is expressed in atomic mass unit or atomic weight units. 100

IL Foundation Series Class 9

• The SI unit of atomic weights is 'u'. • The latest standard for determining atomic masses is 6C12, which is assigned the mass of an atom. One a.m.u. is 1/12th part of the mass 6C12 atom. • One a.m.u is also known as one Dalton or one Aston. • The molecular weight (or) molecular mass is also a relative mass expressed in a.m.u. • One mole of carbon is 12g of graphite, and one mole of oxygen is 32g of oxygen. • One atomic mass unit is a mass unit equal to exactly one-twelfth (1/12th) mass of one atom of carbon-12. • Atomic mass of an element =

Mass of one atom of an element 1/12th part of mass of an atom of C-12 isotope

• Ions: Charged species that are formed by losing or gaining electrons are known as ions. • Cation: An ion having a positive charge on it is known as an electropositive ion or a cation. • Anion: The ion having a negative charge on it is known as an electronegative ion or an anion. • Simple ion: It is an ion which contains one or more atoms of the same element. • Compound ion: Ions consist of two or more atoms of different elements to form a single unit. • An ionic compound is a combination of two or more simple compound ions. • Chemical formula: "The chemical formula of a compound is a symbolic representation of its composition". • Importance of chemical formula: "A Chemical formula gives the exact number of atoms of the same or different elements present in a chemical substance." • In writing formulae of a compound, the metal comes first, followed by the non-metal. • The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. • The formula unit mass of a substance is the sum of the atomic masses of all atoms in a formula unit of a compound. • Mole: A mole represents 6.023 × 1023 particles. • The actual meaning of a mole is a heap, but here, it represents the number of atoms or molecules or particles present in a fixed amount of the substance. • Mole: The amount of the substance that contains as many particles as atoms exactly present in 12 g of the C-12 Isotope. • Avogadro’s number (NA): The number of entities in one mole of a substance is known as “Avogadro’s number”. • The chemical formula represents one mole of the substance. 101

ATOMS AND MOLECULES

• The formula mass in grams represents the mass of one mole of that substance. • One mole of any substance contains 6.023 × 1023 particles. • Limiting Reagent: The reactant that is present in the lesser amount gets consumed after some time, and then no further reaction occurs even though the other reactant is present in the larger amount; this is called a limiting reagent.

WORKSHEET - 1 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER I. Introduction to atoms, molecules, and laws of chemical combination

1. The law that deals with the relation between the mass of the reactants and the products in physical changes or chemical reactions is the a. Law of constant composition

b. Law of conservation of mass

c. Law of multiple proportions

d. None of the above

2. The law that deals with the composition of the various elements present in a compound is the a. Law of multiple proportions

b. Law of conservation of mass

c. Law of constant composition

d. None of these

3. According to the law of constant composition, a pure chemical compound always consists of the same elements combined together in a fixed proportion by a. Weight

b. Volume

c. Density

d. Pressure

4. "Elements always combine in a fixed ratio of their atomic weights." This statement is based on a. Law of conservation of mass

b. Law of conservation of energy

c. Law of multiple proportions

d. Law of definite proportions

5. Statement (A): Stoichiometric equation obeys the law of conservation of mass. Statement (B): The weight of reactants is equal to the weight of products in a balanced chemical reaction. Statement (C): The law of conservation of mass is given by Lavoisier. a. All the above statements are correct. b. All the above statements are incorrect. c. A and B are correct, but C is incorrect. d. A and B are incorrect, but C is correct 102

IL Foundation Series Class 9

6. If 32 g of oxygen reacts with 32 g of sulphur, then find the weight of the SO2 formed according to the law of conservation of mass. a. 32 g

b. 64 g

c. 128 g

d. 16 g

7. Among the following, which is the element that has fractional atomic mass? a. Carbon

b. Magnesium

c. Chlorine

d. Calcium

c. 32

d. 15

8. The atomic mass of sulphur is a. 16

b. 31

9. Which of the following elements has the same molecular mass as its atomic mass? a. Nitrogen

b. Neon

c. Oxygen

d. Chlorine

10. Which of the following contains the maximum number of molecules? a. 1 g CH4

b. 1 g CO2

c. 1 g of H2

d. 1 g of N2

11. The atomicity of ozone, sulphur, phosphorus, and argon are respectively a. 8, 3, 4, and 1

b. 1, 3, 4, and 8

c. 4, 1, 8, and 3

d. 3, 8, 4, and 1

II. What is an atom?

1. Which of the following has the maximum number of atoms? a. 18 g of CO2

b. 18 g of H2O

c. 18 g of O2

d. 18 g of CH4

2. The symbol of a metal element that is used in making thermometers is a. Ag

b. Hg

c. Mg

d. Sg

3. Arrange the following in the order of increasing mass. (Atomic mass of O = 16u, Cu = 63u, N = 14u ). I) One atom of oxygen II) One atom of nitrogen III) 1×10-10 mole of oxygen gas IV) 1×10-10 mole of copper a. II < I < III < IV

b. I < II < III < IV

c. III < II < IV < I

d. IV < II < III < I

103

ATOMS AND MOLECULES

4. The formula of calcium sulphate is a. Ca2SO4

b. Ca(SO4)2

c. CaSO4

d. CaSO3

b. 1.66×10-27 g

c. 1.66×1027 g

d. 1.66×10-27 kg

5. One a.m.u equals to a. 1.66×1027 kg 6. Match the following: Column - I

Column - II

A) 1 a.m.u in grams

1. Mole

B) This is a relative quantity

2. Avogadro number

C) 1 mole of carbon contains these many atoms

3. Relative atomic mass

D) 12 g of C-12

4. 1.66×10-24

a. A- 4, B-3, C-2, D-1

b. A- 4, B-3, C-1, D-2

c. A- 3, B- 4, C-2, D-1

d. A- 3, B- 2, C-1, D-2

III. What is a molecule?

1. The number of molecules in 3 moles of CO2 are a. 18×1023

b. 6×1023

c. 12×1023

d. 3×1023

b. 11×10-23 g

c. 44×10-23 g

d. 22×10-23 g

2. 3 molecules of CO2 weigh a. 33×10-23 g

3. The number of ions in 2.2×10-3 moles of Ca(OH)2 are a. 11×1020

b. 14.2×1020

c. 1020

d. 39.6×1020

c. 12.046×1023

d. 24×1023

4. The total number of atoms in one gram atom is a. 6.023×1023

b. 3.0115×1023

5. The total number of atoms present in 1 mole of CO2 is a. N

b. 2N

c. 3N

d. 4N

c. 44 kg

d. 44 mg

6. The weight of one carbon dioxide molecule is a. 44 g

b. 44 a.m.u

7. The charge present in one mole of electrons is

104

a. 96,500 coulomb

b. 1 Faraday

c. both a and b

d. none of these

IL Foundation Series Class 9

8. Identify the gas whose 2 gram molecules weigh 32 g a. He

b. O2

c. CH4

d. SO2

9. Identify the gas whose 2 gram molecules weigh 64 g a. He

b. O2

c. CH4

d. SO2

IV. Writing chemical formulae, molecular mass, and mole concept

1. The chemical formula represents _______ of that substance. a. 1 mole

b. 10 a.m.u

c. 1 g

d. All of the above

2. In a chemical symbol or formula, the coefficient put in front represents that a. Atomic mass number

b. Atomicity

c. No of moles

d. Charge of atom

3. The number of particles present in a mole a. 6.023×1024

b. 6.023×1023

c. 6.023×10-23

d. 6.023×10-24

c. 16 g

d. 32 g

4. Gram molecular weight of oxygen is a. 16 a.m.u

b. 32 a.m.u

5. The element taken as a reference in measuring atomic mass is a. C-14

b. C-12

c. C-13

d. None of these

6. The weight of one C-12 atom is a. 1.9926×1023 g

b. 1.9926×10-23 kg

c. 1.9926×1023 kg

d. 1.9926×10-23 g

V. Empirical and molecular formula

1. X and Y are two different elements having their atomic masses in a 1:2 ratio. The compound formed by the combination of X and Y contains 50% of X by weight. The empirical formula of the compound is a. X2Y

b. XY2

c. XY

d. X4Y

105

ATOMS AND MOLECULES

2. The empirical formula of a gaseous compound is CH2. The density of the compound is 1.25gm/ lit. at S.T.P. The molecular formula of the compound is X a. C2 H4

b. C3 H6

c. C6 H12

d. C4 H8

3. The percentages of C, H, and N in an organic compound are 40%,13.3%, and 46.7%, respectively. Then, the empirical formula is a. C3H13N3

b. CH2N

c. CH4N

d. CH6N

4. In a compound, C, H, and N atoms are present in a 9 : 1 : 35 ratio by weight. The molecular weight of the compound is 108. The molecular formula of the compound is a. C2H6N2

b. C3H4N

c. C6H8N2

d. C9H12N3

5. A certain compound contains Calcium, Carbon, and Nitrogen in the mass ratio 20:6:14. The empirical formula of the compound is a. CaCN

b. CaC2N

c. Ca(CN)2

d. CaCN2

6. A compound contains carbon and hydrogen in the mass ratio of 3 : 1. The formula of the compound is a. CH2

b. CH3

c. CH4

d. C2 H6

7. A compound contains 90% C and 10% H. The empirical formula of the compound is a. C8H10

b. C15H30

c. C3H4

d. C15H40

8. The empirical formula of a compound is CH2O. Its molecular weight is 120. The molecular formula of the compound is a. C3H6O3

b. C4H8O4

c. C2H4 O2

d. C6H12O6

VI. Limiting reagent

1. The mass of Na2 CO3 required to prepare 500 ml of 0.1 M solution is a. 10.6 g

b. 5.3 g

c. 2.65 g

d. 7.95 g

2. 'X' grams of Calcium carbonate was completely burnt in the air. The weight of the solid residue formed is 28 g. What is the value of 'X' (in grams) a. 44

b. 200

c. 150

d. 50

3. 'X' litres of carbon monoxide is present at STP. It is completely oxidised to CO2. The volume of CO2 formed is 11.207 litres at STP. What is the value of 'X' in litres? a. 22.414

106

b. 11.207

c. 5.6035

d. 44.828

IL Foundation Series Class 9

4. How many litres of oxygen (at STP) is required for the complete combustion of 39 gms of liquid benzene? (Atomic weights : C = 12, H = 1, O = 16 ) a. 84

b. 22.4

c. 42

d. 11.2

5. What is the volume (in lit) of carbon dioxide liberated at STP when 2.12 grams of sodium carbonate (mol.wt. = 106) is treated with excess dilute HCl? a. 2.28

b. 0.448

c. 44.8

d. 22.4

WORKSHEET - 2 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER 1. The number of atoms in 1 mole of iron are a. 6.023×1023

b. 3.0715×1023

c. 12.046×1023

d. 24×1023

c. 18×1023

d. 24×1023

b. 12×1023

c. 18×1023

d. 24×1023

b. 212.2 g

d. 112 g

2. The number of atoms of carbon in 3 moles of CO2 is a. 6×1023

b. 12×1023

3. The number of molecules in 2 moles of CO2 are a. 6×1023 4. 7.4 moles of iron weigh a. 111.1 g

414.4 g

5. One mole of KNO3 contains a. 1 mole of potassium

b. Three moles of oxygen

c. Two moles of Nitrogen

d. Both (a) and (b)

6. The number of atoms in 1 gram of iron are a. 2 × 1022

b. 1.07 × 1020

c. 2.0 × 1024

d. 1.07 × 1022

b. 93×10-23 g

c. 9.33×10-23 g

d. 40×1023 g

c. 5.2

d. 2.1

7. 1 atom of iron weighs a. 8×10-23 g

8. The number of moles in 2×1024 atoms of iron are a. 3.3

b. 4.5

107

ATOMS AND MOLECULES

9. The number of moles in 4×1031 atoms of iron are a. 6.6×107

b. 3.3×107

c. 1014

d. 8×1014

c. 17×1016 g

d. 14×1016 g

c. 30×1023

d. 60×1023

10. 6.4×1031 atoms of iron weigh a. 64×108 g

b. 59.73×108 g

11. The number of atoms in 140 grams of iron are a. 45×1023

b. 15×1023

12. Fe2(SO4)3 is the chemical formula of: a. Ferrous sulphate

b. Ferric sulphate

c. Iron sulphate

d. Iron sulphide

13. The name of the compound Ca(NO3)2 are a.

Calcium nitrate

b. Carbon nitrate

c. Calcium nitride

d. Carbon nitride

14. 2.2×10-3 moles of Ca(OH)2 weigh a. 1.628 g

b. 0.01628 g

c. 0.1628 g

d. 162.8 g

c. 44×1023

d. 9.4×1023

b. 1.5 N

c. N

d. 2 N

b. 300 g

c. 396 g

d. 400 g

15. The number of atoms in 7.4 moles of iron are a. 44.4×1023

b. 22.2×1023

16. The number of ions in one mole of Ca(OH)2 are a. 3 N 17. 4 moles of Ca(OH)2 weigh a. 296 g

18. Which of the following will represent the correct chemical formula for aluminium sulphate and calcium carbonate, respectively? a.

Al2(SO4)3 and CaCO3

b. CaPO4 and AlCl3

c. Ca(HCO3)2 and AlSO4

d. Al(CO)3 and CaCl2

19. 3.1 moles of CO2 weigh a. 136.4 g

108

b. 144.2 g

c. 186.2 g

d. 190.5 g

IL Foundation Series Class 9

20. The number of moles in 6 grams of KClO3 are a. 1.2

b. 4.8

c. 48

d. 0.048

c. 245 g

d. 402 g

c. 0.5

d. 0,4

b. 6.25 g

c. 9.75 g

d. 8.25 g

b. 3

c. 4

d. 5

21. Two moles of KClO3 weigh a. 142.5 g

b. 122.5 g

22. The number of moles in 20.2 g of KNO3 a. 0.1

b. 0.2

23. The weight of potassium in 0.25 moles of KNO3 is a. 7.25 g 24. The atomicity of NO2 is a. 2

25. The number of moles of water in 488 grams of BaCl2.2H2O are a. 1

b. 2

c. 3

d. 4

26. The names of the compounds Ba(NO3)2, K2CO3, and SiO2 are: a. Barium nitrate, potassium carbonate, and silicon dioxide b. Barium nitrite, potassium carbon, and silicon oxide c. Beryllium nitrate, potassium carbonate, and silicon-oxygen d. Barium nitrate, Carbon potassium, and silicon dioxide 27. The formula of a chloride of a metal M is MCl3. The formula of the phosphate of metal M will be: a. MPO4

b. M2PO4

c. M3PO4

d. 2(PO4)3

28. Suppose a chemist has chosen 1020 as the number of particles in a mole. Then, the molecular mass of oxygen is a. 5.32 × 10-3

b. 53.2 × 10-3

c. 0.532 × 10-3

d. 532 × 10-4

b. 15×1023

c. 30×1023

d. 60×1023

b. 44 g

c. 56 g

d. 40 g

29. The number of atoms in 140 grams of iron are a. 45 × 1023 30. 1 mole of CO2 weighs a. 22 g

109

ATOMS AND MOLECULES

31. The number of water molecules present in a drop of water weighing 0.05 mg are a.

1.666 × 1016

b. 16.66 × 1018

c. 1.666 × 1017

d. 16.73 × 1017

c. 101

d. 404

32. The mass of 1 mole of KNO3 in grams is a. 202

b. 303

33. If a 5-star chocolate of 10 rupees contains 3.42 g of sugar, then the number of sugar molecules present in two 5-star chocolates are a. 12.046 × 1019

b. 12.046 × 1020

c. 12.046 × 1021

d. 12.046 × 1022

34. The number of moles in 1.6×104 grams of Ca(OH)2 are a. 2.162

b. 21.62

c. 0.2162

d. 216.2

c. 44.8

d. 33.6

35. The volume occupied by 4g of H2 at STP is a. 22.4

b. 11.2

36. The number of moles of KNO3 in 12×1023 ion pairs of KNO3 are a. 1

b. 2

c. 3

d. 4

37. A compound has 20% of nitrogen by weight. If one molecule of the compound contains two nitrogen atoms, the molecular weight of the compound are a. 35

b. 70

c. 140

d. 280

38. The empirical formula of an organic compound is CH2O. Its vapour density is 45. The molecular formula of the compound is a. CH2O

b. C2H4O2

c. C3H6O3

d. C6H12O6

c. CH2O

d. CHO

39. The empirical formula of acetic acid is a. CH3-COOH

b. C2H4O

40. The empirical formula weight of a compound containing carbon and hydrogen is 13. The molecule of the compound is 39 times heavier than a molecule of hydrogen. The molecular formula of the compound is a. CH

110

b. C3H3

c. C13H13

d. C6H6

IL Foundation Series Class 9

41. An organic compound containing carbon, hydrogen, and oxygen has 52.20% carbon and 13.04% hydrogen. The vapour density of the compound is 23. Its molecular formula will be a. C2H6O

b. C3H8O

c. C4H8O

d. C5H10O

42. An alkane has a C/H ratio (by mass) of 5.1428. Its molecular formula is a. C5H12

b. C6H14

c. C8H18

d. C7H16

43. 0.262 g of a substance, on combustion, gave 0.361 g of CO2 and 0.147 g of H2O. What is the empirical formula of the substance? a. CH2O

b. C3H6O

c. C3H6O2

d. C2H6O2

44. 2.76 g of silver carbonate on strong ignition leaves a residue weighing a. 2.48 g

b. 2.16 g

c. 2.32 g

d. 2.84 g

45. The number of moles of CO2 produced when 3 moles of HCl react with the excess of CaCO3 is a. 1

b. 1.5

c. 2

d. 2.5

46. 6 g of Mg reacts with the excess of an acid. The amount of hydrogen produced would be a. 0.5 g

b. 1 g

c. 2 g

d. 4 g

47. What volume of H2 at NTP is required to convert 2.8 g of N2 into NH3 ? a. 2240 ml

b. 22400 ml

c. 6.72 lit

d. 224 lit

48. What is the volume (lit) of oxygen required at STP to completely convert 1.5 moles of sulphur into sulphur dioxide a. 11.2

b. 22.4

c. 33.6

d. 44.8

111

STRUCTURE OF THE ATOM

4.1

THE STRUCTURE OF AN ATOM

4.1.1 Introduction Frictional studies indicated the electrical nature of matter; for example, when substances like glass or ebonite are rubbed with silk or fur, they generate electricity. However, the greatest breakthrough in the atomic theory of matter came to light with the discovery of Faraday's laws of electrolysis, which proved the electrical nature of matter. Faraday suggested that there is some relationship between matter and electricity. The electrical nature of matter was further supported by Thomson's experiment on the electrical conduction of gases through discharge tubes. Year

Landmark

1896

J.J. Thomson's discovery of the electron

1909

Rutherford's nuclear atom

1913

Mosley's determination of atomic number

1913

Bohr's atom

1921

Bohr - Bury scheme of electronic arrangement

1932

Chadwick's discovery of the neutron.

Table 4.1 The main landmarks in the evolution of the atomic structure Discharge tube experiment

The discovery of electrons was a result of the study of the electric discharge (tube) in the discharge tube. Air at very low pressure

Discharge tube

- - - - - -

Green glow

Cathode rays Anode

Cathode To vacuum pump

High voltage generator

Fig. 4.1 Production of cathode rays 112

IL Foundation Series Class 9

Properties of cathode rays

• They travel in straight lines away from the cathode and cast shadows of opaque objects placed in their path. • Cathode rays cause mechanical motion of a small pin-wheel placed in their path. Thus, they possess kinetic energy and must be material particles. • They produce fluorescence (a glow) when they strike the glass wall of the discharge tube. • They heat a metal foil to incandescence, upon which they impinge. • Cathode rays produce X-rays when they strike a metallic target. • The speed of cathode rays is less than the speed of light. • In an electric field, cathode rays are deflected towards the positive end, which indicates the presence of a negative charge in cathode rays. In a magnetic field, cathode rays are deflected towards the positive end out of the plane of a paper. Properties of anode rays

• They travel in a straight line in a direction opposite to the cathode. • In an electric field, anode rays are deflected towards the negative end, which indicates the presence of a positive charge in anode rays. In a magnetic field, anode rays are deflected towards the south pole. • The charge-to-mass ratio (e/m) of positive particles varies with the nature of the gas placed in the discharge tube. • They possess a mass much greater than that of an electron. • They cause fluorescence when striking on zinc sulphide. 4.1.2 Thomson’s model of an atom An atom may be considered a sphere of positive charge in which the electrons are distributed uniformly. This model of an atom is known as the plum pudding or watermelon model. This model could explain the electrical neutrality of an atom but failed to explain the observations of Rutherford's α-ray scattering experiment. Thomson's model

• Thomson proposed the first atomic model in 1904. • According to him, an atom contains an equal number of negative and positive charges, so that as a whole, an atom is electrically neutral. • Thomson believed that the electrons are embedded in the positively charged atomic mass like the seeds embedded in the fibrous mass of a watermelon fruit. • The whole positive charge and mass of the atom are uniformly distributed inside it. 113

STRUCTURE OF THE ATOM

Thomson’s Atomic Model

Atom model

Watermelon Positive charge

- -

-+ - + - ++ + - - +-

Electron

Fig. 4.2 Thomson’s atomic model Limitations of J.J Thomson's model of an atom

• J .J Thomson attributed the mass of an atom to the electrons and protons, which are evenly spread throughout the atom. This does not meet with Rutherford's observations, who concluded that an atom's mass is concentrated in a very small space called the nucleus. • H owever, this model also failed to explain how the positively charged particles are shielded from the negatively charged particles, and that the electric charges, when squeezed in a small volume, are expected to repel each other. Rutherford’s Atomic Model (Planetary Model) 4.1.3 Rutherford’s model of an atom

α-particles

movable screen gold foil

Source of alpha rays

radioactive substance (polonium) deﬂected α-particles lead plate

ZnS screen

The scattering pattern on a gold foil

Rutherford alpha ray scattering experiment

Fig. 4.3 Rutherford's atomic model (planetary model) Alpha-particle scattering experiment

• Ernest Rutherford designed an experiment to check the validity of Thomson's atomic model. • A stream of high energy α-particles (nuclei of helium atoms) from a radioactive source was directed at a thin foil (thickness nearly 100 nm) of gold metal. The thin gold foil had a circular fluorescent zinc sulphide screen around it. Whenever an α-particle struck the screen, a tiny flash of light was produced at that point. 114

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• α-rays consist of 42 He 2+ .

In this α-ray scattering experiment, the following observations were made: • Most of the α-particles passed through the gold foil undeflected.

• A small fraction of the alpha particles were deflected by small angles. • A very few alpha particles (1 in 20,000) bounced back, i.e., were deflected by an angle of nearly 180°. 4.1.4 Conclusion from alpha-particle scattering From this experiment, Rutherford concluded that: • Most of the space inside the atom is empty (as most of the alpha particles passed through the gold foil undeflected). • Very few α-particles were deflected from their path, indicating that the positive charge of the atom occupies very little space at the centre. On the basis of his experiment, Rutherford put forward the nuclear model of an atom, which has the following features (postulates): • There is a positively charged centre in the atom called the nucleus. Almost all the mass of an atom is present in the nucleus. • The electrons revolve around the nucleus like the planets revolve around the sun. • The size of the nucleus is very small compared to the size of the atom. The size of the nucleus is in the order of 10-13 cm, and the size of the atom is in the order of 10-8 cm. 4.1.5 Drawbacks of Rutherford’s model • He could not explain the stability of the atom. Any particle moving in a circular path would undergo acceleration. According to the electromagnetic theory, during acceleration, a charged particle radiates energy continuously. Thus, the revolving electron would lose energy and finally fall into the nucleus. If this were so, the atom would be highly unstable, and hence, matter would not exist. But we know that matter exists and atoms are stable. • He could not explain the line spectrum of an atom. If the electron loses energy continuously, then the atomic spectra should be continuous. Experimentally, atomic spectra are made up of discrete spectral lines.

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STRUCTURE OF THE ATOM

Ernest Rutherford was born in New Zealand but went to Cambridge University in England to pursue his Ph.D. in 1895. In 1899, he moved to McGill University in Canada. In 1907, Rutherford moved from Canada to Manchester University in England, and there, he performed the experiments that gave us the modern view of the atom. Rutherford was awarded the 1908 Nobel Prize in Chemistry for his investigations on the chemistry of radioactive substances. In a series of experiments analysing the radiation emitted by elements such as Uranium, he discovered that such radiation consisted of three components: alpha, beta and gamma rays. E. Rutherford

4.1.6 Bohr's model of the atom Rutherford's atomic model could not explain electromagnetic radiation and, therefore, could not account for the existence of spectral lines even for the simplest atom, i.e., the hydrogen spectrum. To seek a theoretical explanation of the existence of spectral lines and their regularities and to also solve the conflict between the conclusions of conventional mechanics and the laws of electrodynamics, Bohr, in 1913, put forward a theory to improve Rutherford’s model of the structure of an atom. His theory was based on Max Planck's principles about the absorption and emission of radiation by an electron. To overcome the drawbacks of Rutherford's model of an atom, Niels Bohr, a brilliant Danish Physicists pointed out that the old laws of physics did not work in the submicroscopic world of the atom. He closely studied the behaviour of electrons, radiations and atomic spectra. In 1913, Bohr proposed a new model of the atom based on the modern quantum theory of energy. With his theoretical model, he was able to explain why an orbiting electron did not collapse into the nucleus and how the atomic spectra were caused by the radiations emitted when electrons moved from one orbit to the other. Bohr's postulates

1. Electrons revolve around the nucleus in specified circular paths called orbits or shells. These orbits are numbered as 1, 2, 3, 4.., or represented as K, L, M, N,.., respectively, and are represented by the symbol n. 2. Each orbit is associated with a definite amount of energy. Hence, these orbits are also called energy levels. 3. As long as the electron revolves in a particular orbit, the electron neither gains nor loses energy. Therefore, these orbits are called stationary orbits or main energy states, and the electrons are said to be in stationary energy states. 4. The energy associated with the energy levels increases with the increase in n value. 5. When an electron absorbs energy, it jumps from a lower energy level to a higher energy level, and by emitting energy, it jumps from a higher energy level to a lower energy level. This absorption and emission take place in the form of electromagnetic radiation. 116

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Let E1 be the energy of the lower energy orbit, and E2 be the energy of the higher energy orbit. Then, E2 - E1 = ΔE = hϑ, where h = Planck's constant and ϑ = Frequency of radiation, h = 6.6256×10-34 joule.sec (or) h = 6.6256×10-27 erg.sec ΔE is the energy difference between the two orbits. 6. The angular momentum of an electron revolving in a particular orbit is quantised and is an integral multiple of h/2π. This is known as the quantisation of angular momentum. h

7. The angular momentum is given by the formula mvr = n , which is called 'Bohr's quantum 2π condition.' Here, n is an integer (n = 1, 2, 3...); this is also called a 'principal quantum number.' m = mass of the electron, v= velocity of the electron r = distance of the electron from the nucleus, h= Planck's constant Niels Bohr, a Danish physicist, received his Ph.D. from the University of Copenhagen in 1911. He then spent a year with J.J. Thomson and Ernest Rutherford in England. In 1913, he returned to Copenhagen, where he remained for the rest of his life. In 1920, he was named the Director of the Institute of Theoretical Physics. After the First World War, Bohr worked energetically to make peaceful use of atomic energy. He received the first Atoms for Peace award in 1957. Bohr was awarded the Nobel Prize in Physics in 1922. Niels Bohr

4.1.7 Fundamental particles of an atom The three fundamental particles present in an atom are electron, proton and neutron. Property

Electron

Proton

Neutron

1) Charge

-Ve

+Ve

no charge

2) Notation

9.11 × 10-28 g

1.672×10-24 g

1.675×10-24 g

(or)

0.000548 amu

1.007277 amu

1.008665 amu

-1.602×10-19 coulombs

1.602×10-19 coulombs

(or)

-4.8×10-10 esu

4.8×10-10 esu

-1

3) Mass

4) Charge 5) Relative charge

Table 4.2 Fundamental particles in an atom 117

STRUCTURE OF THE ATOM

In a neutral atom, the number of positive charges is equal to the number of negative charges (or the number of protons is equal to the number of electrons).

Oxygen Atom electron <10-16 cm proton (neutron) ~10-13 cm

atom ~10-8 cm

nucleus ~10-12 cm

Electrons

Second Electron Shell

First Electron Shell

+ +

quark ~10-16 cm

Nucleus

Fig. 4.4 Fundamental particles of an atom

4.2

ELECTRON DISTRIBUTION IN DIFFERENT ORBITS (SHELLS)

The electrons in an atom revolve around the nucleus in definite paths called orbits or shells. These orbits are numbered 1, 2, 3, 4… and are represented as K, L, M, N..., respectively and are denoted by n. Each orbit consists of a subshell. The first orbit (n = 1, K- shell) has only one subshell and is named 1s subshell. It can accommodate a maximum of two electrons only. The second orbit (n = 2, L- Shell) has two types of subshells, which are named 2 s and 2p. The third orbit (n = 3, M- Shell) has three types of subshells, which are named 3s, 3p and 3d. The fourth orbit (n = 4, N- Shell) has four types of subshells, which are named 4s, 4p, 4d and 4f. The capacity of the s, p, d and f subshells to hold electrons present in any subshell is 2, 6, 10 and 14, respectively.

5 4 3 2 1

K L MN O

Fig. 4.5 Electronic distribution in different orbits

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Note: The number of subshells in the nth shell is given by n. The number of orbitals in the nth shell is given by n2. The number of electrons in the nth shell is given by 2n2.

4.3

VALENCY

Definition: Valency is the combining capacity of an element. (OR) Modern definition of Valency: According to the new concept, valency can be defined as the number of electrons that are lost or gained or shared with one atom of an element to acquire the stable configuration of the nearest noble gas element. (OR) The number of hydrogen atoms (or) chlorine atoms (or) double the number of oxygen atoms with which one atom of an element combines is called its valency. Element

Atomic number

Electronic configuration

2,8,1

2,8,2

2,8,3

: 1 Valency

Valency

Capacity of an atom to give, accept or share electrons to achieve the octet state

Needs 1e-

Hydrogen Valency

Oxygen (2,6) Fig. 4.6 Valency

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STRUCTURE OF THE ATOM

Examples: Element

with hydrogen

with chlorine

with oxygen

Valency

NaH

NaCl

Na2O

MgH2

MgCl2

MgO

AlH3

AlCl3

Al2O3

Table 4.3 Compound formation Valency of metals

Number of valence electrons present in a metal atom is its valency. Generally, metals have 1, 2 or 3 valence electrons. So, their valencies are 1, 2 or 3, respectively.

Examples:

Fig. 4.7 Sodium element Valency of Non-Metals

Valency of non-metals = (8 - number of valence electrons) Generally, non-metals possess 4, 5, 6 or 7 valence electrons. Thus, their valencies are 8-4, 8-5, 8-6, 8-7, i.e., 4, 3, 2 and 1, respectively.

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Examples: Element

Atomic number

Electronic configuration

2,4

2,5

2,6

2,7

Valency

8-4=4

8-5=3

8-6=2

8-7=1

Compounds example

CH4

NH3

H2O

Table 4.4 Non-metal compounds Name of the element

Symbol

Atomic number

Lithium

Distribution of electrons | Valency K

Beryllium

Boron

Carbon

4(8-4)

Nitrogen

3(8-5)

Oxygen

2(8-6)

Fluorine

1(8-7)

Neon

0(8-8)

Table 4.5 Distribution of electrons

4.4

ELECTRONIC CONFIGURATION OF ELEMENTS

The systematic arrangement of electrons in various atomic orbitals in the increasing order of their energies is known as electronic configuration. The increasing order of various atomic orbitals is as follows: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p To write the electronic configuration easily, we use Moeller's diagram, which is shown on the next page.

121

STRUCTURE OF THE ATOM l=0 l=1 l=2 l=3 n=1

n=2

n=3

n=4

n=5

n=6

n=7

n=8

Fig. 4.8 Moeller's diagram Electronic configuration of elements

122

Atomic Number

Element

Electronic Configuration

→

1s1

→

1s2

→

1s2 2s1

→

1s2 2s2

→

1s2 2s2 2p1

→

1s2 2s2 2p2

→

1s2 2s2 2p3

→

1s2 2s2 2p4

→

1s2 2s2 2p5

→

1s2 2s2 2p6

→

1s2 2s2 2p6 3s1

→

1s2 2s2 2p6 3s2

→

1s2 2s2 2p6 3s2 3p1

→

1s2 2s2 2p6 3s2 3p2

→

1s2 2s2 2p6 3s2 3p3

IL Foundation Series Class 9

Atomic Number

Element

Electronic Configuration

→

1s2 2s2 2p6 3s2 3p4

→

1s2 2s2 2p6 3s2 3p5

→

1s2 2s2 2p6 3s2 3p6

→

1s2 2s2 2p6 3s2 3p6 4s1

→

1s2 2s2 2p6 3s2 3p6 4s2

Table 4.6 Electronic configuration of the first 20 elements

4.5

ATOMIC NUMBER AND MASS NUMBER

4.5.1 Atomic number It is defined as the number of protons (positive charges) present inside the nucleus or the number of electrons (negative charges) present outside the nucleus in a neutral atom. The atomic number is represented by the symbol Z. mass number electrical charge

24 12

atomic number

Fig. 4.9 Magnesium element

4.5.2 Mass number It is defined as the sum of the number of protons and neutrons present in the nucleus of an atom. It is represented by A. Protons and neutrons together are called nucleons.

123

STRUCTURE OF THE ATOM

Atomic Mass on the Periodic Table Atomic Number

Symbol

Na 22.99

Atomic Mass

Fig. 4.10 Atomic mass of the sodium element

An element can be represented as z XA, where X is an element, A is the mass number, and Z is the atomic number. A = Number of protons + Number of neutrons ⇒ A = Z + n ⇒ n = A - Z 4.5.3 The atomic numbers and atomic masses of the first 30 elements Element

Symbol

Atomic number

Atomic mass (Rounded values)

Hydrogen

Helium

Lithium

Beryllium

Boron

Carbon

Nitrogen

Oxygen

Fluorine

Neon

Sodium

Magnesium

S.no

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Element

Symbol

Atomic number

Atomic mass (Rounded values)

Aluminium

Silicon

Phosphorus

Sulphur

Chlorine

35.5

Argon

Potassium

Calcium

Scandium

Titanium

Vanadium

Chromium

Manganese

Iron

Cobalt

Nickel

Copper

64.5

Zinc

S.no

Table 4.7 Atomic numbers and atomic masses of the first 30 elements

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STRUCTURE OF THE ATOM

4.6

ISOTOPES, ISOBARS, ISOTONES

4.6.1 Isotopes The atoms of the same element with the same atomic number but different mass numbers are called isotopes. Examples: 1) H11 , H12 , H13 35 37 2) Cl17 , Cl17 235 238 3) U 92 , U 92

Isotopes differ in number of neutrons. They have the same chemical properties but differ in physical and radioactive properties. Hydrogen (H)

Deuterium (D)

Tritium (T)

Electron

Proton

The Three0Isotopes of1Hydrogen2 Neutron Table 4.8 Isotopes of hydrogen

Proton P+

Proton N-

N-

P+

Neutron

P+

Neutron

N-

Proton

Electron

Protium 1 H 1

Deuterium 2 H 1

Tritium 3 H 1

Fig. 4.11 Isotopes of hydrogen

4.6.2 Isobars • The atoms of different elements with the same mass number but different atomic numbers are called isobars. • Isobars have different numbers of electrons, protons and neutrons. They differ in chemical, physical and radioactive properties.

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Examples: 14 1) C14 6 , N7 40 2) Ar1840 , Ca 20

131

131 Isobars

54 Protons 77 Neutrons

53 Protons 78 Neutrons same

131

Fig. 4.12 Isobars

4.6.3 Isotones • Isotones are the atoms of different elements which have the same number of neutrons. Examples: 30 32 1) Si14 , P1531 , S16 16 2) C14 6 , O8 23 24 3) Na11 , Mg12

• They differ in physical and chemical properties. Isotones X A=131 N=53 Z=78

Same

Y A=127 N=53 Z=74

Proton Neutron

Fig. 4.13 Isotones

127

STRUCTURE OF THE ATOM

4.6.4 Isodiapheres • The atoms of different elements have different atomic numbers and mass numbers, but the difference between the numbers of neutrons and protons in the nucleus is the same, and they are called isodiapheres. Example: 90

Th234 and 92U238

The difference between the neutron number (N) and proton number (Z) in both nuclei is the same N-Z=54.

Fig. 4.14 Isodiapheres

4.7

BOHR'S THEORY OF HYDROGEN ATOM

Let us consider a system of hydrogen-like atoms having one proton of charge +Ze. The electron with charge - e revolves around this nucleus in an orbit of radius r. Let v be the tangential velocity and m be the mass of the revolving electron.

Fig. 4.15 Electron revolving in hydrogen atom

Evidently, on the revolving electron, the following two types of forces act: • The centrifugal force, due to the motion of the electron, tends to take the electron away from its orbit. It is equal to +mv2/r and acts outwards from the nucleus. 128

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• The centripetal force (also called the electrostatic force of attraction) between the revolving electron and the nucleus, tends to pull the electron towards the nucleus. It is given by Ze

Coulomb's inverse square law and is, therefore, equal to -e × 2 r nucleus.

-Ze2 r2

. It acts towards the

In order for the electron to continue revolving in its orbit, these two forces, which act in the opposite directions to each other must balance each other, i.e., -Ze2

-mv2 r

or v2 =

Ze2

or v =

According to Bohr's quantisation rule: nh

mvr =

2π

2πmr

Squaring both sides, we get: v2 = Comparing equation (1) and (2), we get: Ze2 mr

r= rn =

n2h2

4π2m2r2 n2h2

4π2m2Ze2 n2h2

4π2mZe2

In the above expression, h,π,m and e are constants. Therefore, K =

= constant K = 0.529 × 10-10 m = 0.529A0

4π2me2

rn = K Z or rn = 0.529A0 × Z rn =

r0 × n2 Z

Thus, the radius of the atom is directly proportional to the number n of energy level and inversely proportional to Z (atomic number). Thus, the plot of r versus n2 shows a linear increase. If n is a constant, then r∝

1 Z

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STRUCTURE OF THE ATOM

4.8

ENERGY OF AN ELECTRON IN THE nth ORBIT (En)

The total energy of an electron in the nth orbit is equal to the sum of the kinetic energy and potential energy, i.e.,

We know that for an electron revolving in an orbit: mv2 r

(Centrifugal force) =

Ze2 r2

(or) mv2 r

Substituting value (ii) in equation (i), we get:

Substituting the value of r in equation (iii), we get:

130

Ze2 r2

(Centripetal force)

IL Foundation Series Class 9

atom If Z is a constant, then E∝-

1 n2

Therefore, the energy of an electron increases as the number of orbits increases.

If n is a constant, then

, i.e

Expression for velocity of electron (Vn ) in the nth orbit. We know that: equ (i) Substituting the value of

where,

in above equation (i), we get:

constant =2.18×106 m s-1.

131

STRUCTURE OF THE ATOM

If Z is a constant, then vα

1 n

Therefore, velocity decreases with an increase in the number of orbits. If n is a constant, then v α Z.

V 1 n

V Z

Therefore, velocity increases with the increasing atomic number. S. No.

Radius

Velocity

Energy

Wavelength

1. 2. 3. 4.

constant

5. constant

constant

Time period 6.

constant Table 4.9 Relationship and important formulae related to Bohr's theory

4.8.1 Significance of negative value of energy The energy of an electron at an infinite distance from the nucleus is arbitrarily assumed to be zero. This state is called the zero-energy state. When an electron moves and comes under the influence of the nucleus, it does some work and spends its energy in this process. Thus, the energy of the electron decreases, and it becomes less than zero, i.e., it acquires a negative value.

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When the electron is at an infinite distance from the nucleus, its energy is zero. When the electron jumps from the orbit at an infinite distance to a given orbit, its energy decreases from zero to a negative value.

4.9

IONISATION ENERGY

The minimum amount of energy required to remove an electron from the outermost orbit to an infinite distance is known as ionisation energy or ionisation potential. Ionisation energy = -( first energy of the orbit)

(for Hydrogen atom) (or)

The minimum amount of energy required to remove an electron from the outermost orbit of an isolated neutral gaseous atom. H(g) + IE → H+(g) + eAlternatively, ionisation energy corresponds to the energy difference between the ground state of an atom and the excited state, i.e., I.E = E∞ - En = E(ground state) (∵E∞=0) For H-atom, (I.E )H = E1 (H) = -2.17 × 10-18 Jatom-1 For H-like species (He+, Li+, etc.) (I.E)H.like =

I.EH n2

[1eV=1.6×10-19 J] where n is the first excited state. Note: As the principal quantum number increases, the difference of energy between the consecutive energy levels decreases. The energy of an electron in any orbit is calculated by En=

-2π2me4 z2 n2 h2

If E1, E2, E3, E4, and E5 are energies of electrons in the first orbit ( n = 1), second orbit (n = 2), third orbit (n = 3), fourth orbit(n = 4), fifth orbit(n = 5), respectively, then:

E5 = -0.54 eV

n=5

E4 = -0.85 eV

n=4

133

STRUCTURE OF THE ATOM

E3 = -1.5 eV

n=3

E2 = -3.4 eV

n=2

E1 = -13.6 eV

n=1

The difference in the energy levels is: E2-E1 = -3.4-(-13.6) =

10.2eV

E3-E2 = -1.5-(-3.4)

1.9eV

E4-E3 = -0.84-(-1.5) =

0.66eV

E5 -E4 = -0.54-(-0.84) =

0.30eV

All the calculations indicate that the gap between the consecutive energy levels decreases with the increase in the value of n.

4.10 QUANTUM NUMBERS It specifies the location and energy of an electron in an atom and is a measure of the effective volume of the electron cloud. Quantum numbers are proposed in order to know the complete information regarding the position of an electron in an atom. There are four quantum numbers, they are: 1. Principal quantum number 2. Azimuthal quantum number 3. Magnetic quantum number 4. Spin quantum number 4.10.1 Principal quantum number (n) • This quantum number was introduced by Niels Bohr. • It explains about • Origin of lines in the spectrum of hydrogen. • Number of the orbit • Size of the orbit • Energy of the orbit 134

IL Foundation Series Class 9

The value of n is taken as 1, 2, 3… and represented as K, L, M… 0.53 × n2

The radius of the nth orbit = The energy of an electron =

Z -13.6 n2

Å

× Z2 eV/ atom

The number of electrons in any orbit is given by the formula 2n2. For example: Orbit number

Name of the orbit

Number of electrons (Maximum)

Table 4.10 Number of electrons in orbits

4.10.2 Azimuthal quantum number (l) Orbital

The three-dimensional space around the nucleus in an atom where the probability of finding an electron is at its maximum is called an orbital. For an atomic orbital, the electron has a fixed energy, and its angular momentum is fixed. The angular momentum of orbitals is quantised like the energies, i.e., it can have specific values. Like the principal quantum number, which is used to specify the energies of an electron in various states, the orbital angular momentum quantum number is used to specify the orbital angular momentum. This quantum number was introduced by Sommerfeld. h 2π Example: When l = 0, there is no nodal plane in the wave function, and the orbital is spherically symmetrical about the nucleus. Since the total number of nodes is equal to n, l can have positive integral values from 0 to (n-1). Orbital angular momentum (L)= l (l + 1)

Example: If n = 5, l can be 0, 1, 2, 3, 4. The azimuthal quantum number tells us about the shape of the electron cloud. The greater the value of l, the more diffused the electron cloud is.

135

STRUCTURE OF THE ATOM

It gives information about the subshells and shapes of orbitals. It tells us about the sub-energy levels (subshells). The values of l depend on the values of n, and l takes the values from 0 to n-1. Example: When n = 1( K shell), l = 0 and is known as the s subshell.

When n = 2 (L shell), l = 0, 1 and are called the s and p subshells, respectively.

When n = 3 ( M shell), l = 0, 1 and 2 are called the s, p, and d subshells, respectively.

The number of subshells in any given shell is given by the formula n.

Example: The first shell contains only one subshell.

The second shell contains two subshells.

The number of electrons in a given subshell = 2 (2l+1).

4.10.3 Number of electrons in filled shells and subshells Shell

Subshell

Number of electrons in filled shell

Number of electrons in filled subshell

K (n=1)

s (l=0)

L (n=2)

s (l=0)

2 8

M (n=3)

p (l=1)

s (l=0)

p (l=1)

N (n=4)

d (l=2)

s (l=0)

p (l=1)

6 32

d (l=2)

f (l=3)

Table 4.11 Number of electrons in shells and subshells

4.10.4 Magnetic quantum number (m) • It was introduced by Lande. • It indicates the spatial orientation of the orbitals or the presence of orbitals in a subshell.

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• Its values depend on the value of l, and it can take only integral values ranging from -l to +l. • The total number of m values indicates the number of spatial orientations of the orbital, i.e., (2l+1). • The number of orbitals in a given orbit = n2 Name of the subshell

r values

Values of m

Number of orbitals

Name of orbitals

-1 (or) +1

0 p

+1( or )-1 -2 (or) +2 -1 (or) +1 0

+1 (or) -1

-2 (or) +2 Table 4.12 Values of m in orbitals

Note: For pz orbital and d Z 2 orbital, the m value is zero. • All the orbitals in a given subshell are energetically identical and are called degenerate orbitals. The positive and negative signs indicate the orientations of orbitals. • Purpose: To explain the Zeeman and Stark effects. • Zeeman Effect: The splitting of spectral lines in a strong magnetic field is called the Zeeman Effect. • Stark effect: The splitting of spectral lines in a strong electric field is called the Stark effect. • Space quantisation: To explain the Zeeman Effect, it is proposed that the orbital can assume only certain orientations in space relative to the direction of the external field. This is known as space quantisation.

137

STRUCTURE OF THE ATOM

4.10.5 Spin quantum number (s) Three quantum numbers are necessary to describe the spatial distribution of electrons in atoms. To describe an electron in an atom completely, a fourth quantum number, s, also called the spin quantum number, must be specified. This is because every electron is associated with a magnetic moment, which is quantised in one of the two possible orientations, either parallel or opposite to an applied magnetic field. When spectral lines of H, Li, Na, and K, etc., were observed by means of an instrument with high resolving power, each line of the spectral series was found to consist of a pair of lines (known as doublets or double-line structures). It should be understood clearly that this doublet is different from fine structure, which consists of closely spaced fine lines (not widely separated ones). To account for these doublets, Uhlenbeck and Goud Smith suggested that the electron, while moving around the nucleus in an orbit, also rotates (or spins) about its own axis either in a clockwise direction or in an anti-clockwise direction. s can have two values, +1/2 (corresponding to the spin of the electron in the clock-wise direction) or - 1/2 (corresponding to the spin of the electron in the anti-clockwise direction). The clockwise and anti-clockwise spins are represented as ↑ and ↓ Two electrons with the same sign of spin quantum numbers are said to have parallel spins (shown as ↑↑ ), while those having opposite signs of spin quantum numbers are said to have opposite spins or anti-parallel spins (↑↓ ). The electrons having opposite spins are called paired-up electrons. Purpose: It gives information about the spin of an electron.

Fig. 4.16 Types of electron spin

4.11 SHAPES OF ORBITALS Orbital: The three-dimensional space around the nucleus in an atom, where the probability of

finding an electron is at its maximum, is called an orbital.

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Shape of the s-orbital: The orbital in which the electrons with the quantum numbers n = 1 and l = 0 are present, which is called the 1s orbital. Similarly, the electrons having quantum numbers n = 2 and l = 0 are present in the 2s orbital. Thus, all the s-orbitals have l = 0, while n can have values 1, 2, 3, 4 and so on. The s-orbital is spherical. It has a spherical symmetry.

Fig. 4.17 Shape of s-orbital

Shape of the p-orbital: The orbital in which the electrons with quantum numbers n = 2 and l =1 are present is called the 2p orbital. Similarly, the orbital with quantum numbers n = 3 and l = 1 is called 3p orbital. For all p-orbitals, l = 1, and n = 2, 3… In each principal quantum number (except the first orbit), there are three p-orbitals, and they are mutually perpendicular to one another and oriented along the three axes. Each p-orbital has a dumbbell shape. The probability of finding an electron in the nucleus is zero. At a certain distance from the nucleus, the probability is at its maximum. A p-orbital has one nodal plane, and it consists of two lobes on either side as the nodal plane extends along the axis. For the px orbital, the nodal plane is the YZ plane. For the py & pz orbitals, the nodal planes are the XZ and XY planes, respectively.

Fig. 4.18 Shape of the p-orbital

139

STRUCTURE OF THE ATOM

Shapes of the d-orbital: The orbitals having electrons with quantum numbers n = 3 and l = 2 are called 3d orbitals.

Fig. 4.19 Shape of the d-orbital is a double dumbbell

There are five d-orbitals. They are: dxy, dyz, dxz, d x2 − y 2 , d Z 2 . The first four are double dumbbells in shape. Each has four lobes. d - orbital

Plane

dxy

between the X and Y axes

dyz

between the Y and Z axes

dzx

between the Z and X axes

d x2 − y 2

along the X and Y axes

dZ 2

along Z axis with a dumbbell shape. It contains a ring (torus, collar or tyre) along the XY plane

Table 4.13 d-orbitals SOLVED EXAMPLES Example 1: s-orbitals have no direction. Comment. Solution: For s -orbital, ψ=f(r) This expression only has distance ( r ) but no direction. Example 2: How many peaks and radial nodes are present in the radial probability curve of the 3s orbital? Solution: 3 s orbital curve has 3 peaks. Number of radial nodes = n - l - 1 = 2

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Example 3: Calculate the radial distance between the two peaks in the radial probability curve of the 2 s orbital. Solution: The distance of the first peak is 0.53 Å. The distance of the 2nd peak is 2.6 Å. The distance between two peaks = 2.6 - 0.53 = 2.07 Å. Example 4: How many nodal planes and nodal regions are present in the 3p orbital? Solution: The number of nodal planes for any p-orbital is 1. The number of nodal regions = n - 2 = 3 - 2 = 1. Each d-orbital has ( n - l - 1 ) number of radial nodes.

The nodal planes are given by l.

S. No.

Orbital

Number of

radical nodes (n-l-1)

nodal planes (l)

Total number of nodes (n-1)

In the radical probability distribution curves number of minimas

number of maximas

Table 4.14 Nodes and nodal planes of orbitals 141

STRUCTURE OF THE ATOM

The five orbitals present in a given d-sublevel will have the same energy in their ground state. Like p-orbitals, d-orbitals are similar in shape, but their energy and size differ as the value of n increases. The radial probability distribution curve for 3d orbital is shown in the figure below.

Fig. 4.20 Radial distribution curves for 3d

4.12

ARRANGEMENT OF ELECTRONS IN ORBITALS

The distribution and arrangement of electrons in the main shells, subshells and orbitals of an atom are called the electronic configuration of the element. The electronic configuration of different atoms can be best represented in two ways: a) nlx method b) Orbital diagram method nlx method: The nlx method is a way of representing the electronic arrangements of an atom. Here, n is the principal quantum number, which represents the main energy level. l is orbital quantum number expressed in terms of s, p, d, or f and x is the number of electrons in the given value of l. The nlx method is illustrated in the following example. The atomic number of neon is 10Ne (Z=10)=1 s2 2 s2 2p6 This means that there are 2 electrons in the s orbital of the first energy level. 2 electrons in the s orbital, and 6 electrons in the p orbitals of the second energy level. The total value of x is equal to the atomic number. The filling of electrons into the orbitals of different atoms takes place according to the Aufbau principle, Pauli's exclusion principle, Hund's rule of maximum multiplicity and the relative energies of the orbitals.

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IL Foundation Series Class 9

Fig. 4.21 Atomic number

4.12.1 Aufbau principle As much as possible, electrons will go to the orbital with the lowest energy. (OR) The Aufbau principle states that the electrons tend to occupy the available orbitals of minimum energy in the ground state of an atom. The order of increasing energy of atomic orbitals is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s… The sequence in which the electrons occupy various orbitals can be easily remembered with the help of Moeller's diagram, as shown below: l=0

l=1 l=2 l=3

n=1

n=2

n=3

n=4

n=5

n=6

n=7

n=8

Fig. 4.22 Moeller's diagram 143

STRUCTURE OF THE ATOM

The sequence in which the electrons occupy various orbitals can be found with the help of the (n+l ) rule, where n is the principal quantum number, and l is the azimuthal quantum number. The energy of an orbital is given by (n+l) value. As the (n+l) value increases, the energy of the orbital also increases. The orbital with the lowest (n+l) value is filled first. When two or more orbitals have the same (n+l) value, the one with the lowest n value is preferred for filling. Consider the two orbitals, 3d and 4s. The (n+l ) value of 3d = 3+2 = 5, and of 4s = 4+0 = 4. Since 4s has the lower (n+l ) value, it is filled before filling 3d. Consider the orbitals 3d, 4p, and 5s. The (n+l ) value of 3d = 3+2 = 5. The (n+l ) value of 4p = 4+1 = 5. The (n+l ) value of 5s = 5+0 = 5. 4.12.2 Pauli's exclusion principle It is impossible for two electrons with the same spin quantum number to be in the same orbital. (OR) No two electrons in an atom can have the same set of values for all the four quantum numbers. This means that two electrons in an atom may have the same n, l and m but differ in spin quantum number. In an orbital, one electron has a clockwise spin, while the other has an anti-clockwise spin. An orbital can hold a maximum of two electrons with opposite spins. For example, a helium atom has two electrons in its 1s orbit. The quantum numbers for the first electron in a helium atom are n = 1, l = 0, m = 0 and s = +1/2 The quantum numbers for the second electron in n helium atom are n = 1, l = 0, m = 0 and s = -1/2 The two electrons can have the same value for n, l and m but differ in s. The maximum capacity of a main energy shell is equal to 2n2 electrons, and of a subshell is equal to 4l + 2. 4.12.3 Hund's Rule The pairing of electrons in the orbitals belonging to the same subshell (p, d or f ) does not take place until each orbital belonging to that subshell has got one electron each, i.e., it's singly occupied. Since, there are three p, five d, and seven f-orbitals, the pairing of electrons will start in the p, d and f orbitals with the entry of 4th, 6th, and 8th electron, respectively. It has been observed that a halffilled and fully filled degenerate set of orbitals acquire extra stability due to their symmetry. To explain Hund's rule of maximum multiplicity, the following examples are taken: a) Boron's atomic number is 5. Its electronic configuration is:

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IL Foundation Series Class 9

1s2

2s2

1s2

2s2

2p1

2px1

2py0 2pz0

Here, each orbital is shown by one box. b) Similarly for carbon,

1s2

2s2

1s2

2s2

2p2

2px1

2py1

2pz0

the configuration in p-orbitals is 2px1,2py1 but not 2px2. A given subshell has degenerate orbitals, i.e., orbitals of the same energy. c) For nitrogen, 1s2

2s2

1s2

2s2

2p3

Nitrogen 2px1

2py1

2pz1

the configuration in p-orbitals is 2px 1,2py1 and 2pz1 but not 2px2 and 2py1. d) For oxygen,

1s2

2s2

1s2

2s2

2p4

Oxygen 2px2

2py1

2pz1

Sometimes, the electrons are shown by their respective numbers instead of simple arrows. This is to make the Hund's rule easier to understand. The unpaired electrons in the degenerate orbitals are indicated by arrows, all written in one direction to represent the parallel spin. 4.12.4 Configuration of ions Ions are of two types. Anions are negatively charged ions and are formed by adding one or more electrons to the neutral atom. Cations are positively charged ions and are formed by removing one or more electrons from the neutral atom. The electronic configuration of anions is written using the

145

STRUCTURE OF THE ATOM

same method as writing the configuration of neutral atoms. Oxide ion is O2- and is formed by adding two electrons to a neutral oxygen atom. The configuration of oxide (10 electrons) is 1s2 2s2 2p6. The chloride ion is Cl- and is formed by adding one electron to a neutral chlorine atom. The configuration of chloride (18 electrons) is 1s2 2s2 2p6 3s2 3p6. The electronic configuration of cations is also written using a similar method. For the transition metal cations, the electrons in the ultimate shell are to be removed first. The ferrous ion is Fe2+ and is formed by removing two electrons from a neutral iron atom. The electronic configuration of Fe2+ is 1s2 2s2 2p6 3s2 3p6 3d6. A set of atoms or ions with the same electronic configuration are called isoelectronic species. An isoelectronic group of 2 electrons is H-, He, Li+ ,and Be2+. An isoelectronic group of 10 electrons is N3-, O2-, F-, Ne, Na+, Mg2+, and Al3+. An isoelectronic group of 18 electrons is P3-, S2-, Cl-, Ar, K+, Ca2+, and Sc3+. An isoelectronic group of 28 electrons is Cu+, and Zn2+.

QUICK REVIEW • Properties of cathode rays: 1. They heat up a metal foil to incandescence, upon which they impinge. 2. Cathode rays produce X-rays when they strike a metallic target. 3. The speed of cathode rays is less than the speed of light. • Properties of anode rays: 1. They travel in a straight line in the direction opposite to the cathode. 2. They possess a mass much higher than that of an electron. 3. They cause fluorescence when they strike on zinc sulphide. • Specific charge (e/m ratio): The charge-to-mass ratio is known as specific charge. The specific charge for cathode rays is the same for different gases, but the specific charge for anode rays is different for different gases present in the discharge tube. • Structure of α-ray: 24 He2+

• According to Rutherford, most of the α-particles passed through the gold foil undeflected. • The size of the nucleus is very small compared to the size of the atom. The size of the nucleus is in the order of 10-13 cm, and the size of the atom is in the order of 10-8 cm. • Bohr's postulates: Electrons revolve around the nucleus in specified circular paths called orbits or shells. These orbits are numbered as 1, 2, 3, 4…, or represented as K, L, M, N..., respectively and are represented by the symbol n. 146

IL Foundation Series Class 9

• Atomic number: The number of protons present in the nucleus, denoted by Z. • Mass number: The sum of the protons and neutrons in the nucleus, denoted by A. • Isotopes: The atoms of the same element which have the same atomic number but different mass numbers. • Isobars: The atoms of different elements which have the same mass number but different atomic numbers. • Isotones: The atoms of different elements which have the same number of neutrons. • Isodiapher: The difference between the numbers of neutrons and protons in the nucleus of two different atoms is the same. • The principal quantum number denotes energy, and the azimuthal quantum number denotes shape, and the magnetic quantum number denotes the spatial orientation of an orbital. • Orbitals with the same energy and shape but different orientations are called degenerate orbitals. The number of subshells in a given shell is n, the number of orbitals is n2, and the maximum number of electrons filled is 2n2. • Pauli's exclusion principle states that no two electrons of an atom have all the same four quantum numbers. • The shape of the s-orbital is spherical,the p-orbital is dumb-bell, the d-orbital is double dumbbell and the f-orbital is a tetra dumb-bell. • An orbital can hold maximum of two electrons, and these electrons must have opposite spins. • The differentiating electron always enters into the orbital with the least energy among the available orbitals. • The sequence of energies of orbitals is 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p <... This sequence is obtained from Moeller's diagram and using the (n+l) rule. • Hund's rule states that the pairing starts only if all orbitals in the same subshell are singly filled with electrons. • The distribution and arrangement of electrons in the main shells, subshells and orbitals of an atom are called the electronic configuration of the element. • The atoms or ions with unpaired electrons are paramagnetic, and those with no unpaired electrons are diamagnetic. • If n is a number of unpaired electrons, magnetic moment μ is given as μ= n(n + 2) BM . • The atoms or ions with the same electronic configuration are called isoelectronic species. 147

STRUCTURE OF THE ATOM

WORKSHEET - 1 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER I. Structure of the atom

1. Choose a value for the charge-to-mass ratio of electrons from the options. a. 1.758820 ◊ 1011 C/Kg

b. 1.738820 ◊ 1011 C/Kg

c. 1.758820 ◊ 1011 C Kg

d. 1.708820 ◊ 1011 C/Kg

2. Statement A: J.J. Thomson said that the atom consisted of uniformly distributed positive and negative charged particles. Statement B: Cathode rays produce X-rays when they strike a metallic target. Statement C: Anode rays travel in a straight line in a direction opposite to the cathode. N O

All three statements are correct.

b. A and B are correct, but C is incorrect.

All the statements are incorrect.

d. A and B are incorrect, but C is correct.

3. Rutherford's atomic model is known as the: a.

Planetary model

b. Apple pie model

Plum pudding model

d. Watermelon model

4. The size of the nucleus is in the order of: a.

10-13 cm

b. 10-8 m

1Å

d. Both a and c

5. Rutherford's model of the atom accounts for the following: a. Scattering of alpha particles by metal b. Stability of the electron orbits foils c. Stability of the atom d. Line spectra of light elements 6. The average distance of an electron from the nucleus in an atom is of the order of: a.

1 cm

b. 10-8 cm

10-13 cm

7. Gold foil is selected for α rays scattering experiment because of:

148

d. 10-6 cm

IL Foundation Series Class 9

The very thin gold plate with a thickness equal to 1000 atoms can be made

b. The internal structure, subatomic particles, and nucleus can be observed easily c. High malleability capacity of gold d. All 8. According to Bohr's atomic model, the electrons revolve in: a.

Circular orbits

b. Elliptical orbits

Stationary orbits

d. Both a and c

9. In Bohr's atomic mode, l orbits are represented by: a.

K, L, M, N

b. 1, 2, 3, 4

A, B, C, D

d. Both a and b

10. Assertion (A): In stationary orbits, the energy of electrons is constant. Reason (R): Electrons do not revolve in atoms. a.

Both A and R are correct, and R is the correct explanation for A.

b. Both A and R are correct, but R is not the correct explanation for A. c.

A is correct, but R is incorrect.

d. A is incorrect, but R is correct. 11. Statement A: The angular momentum of an electron in its orbits is quantised. Statement B: While revolving in a stationary orbit, electrons do not lose energy. Statement C: mvr = nh/2π. a.

All the given statements are correct.

b. A and B are correct, but C is incorrect.

All the given statements are incorrect.

d. A and B are incorrect, but C is correct.

12. The fundamental particle with no charge is called: a. Electron

b. Proton

Neutron

d. Positron

149

STRUCTURE OF THE ATOM

II. Electronic distribution, valency, configuration of elements

1. The number of electrons present in the penultimate shell of nitrogen is: a. 2

b. 3

d. 8

3s1

d. 3 s2

d. 6

c. 3

d. 1

c. 3

d. 4

2. The valence electronic configuration of calcium is: a. 4s1

b. 4 s2

3. The valency of carbon in CH4 is: a. 1

b. 2

4. The valency of aluminium in Al2 O3 is: a. 6

b. 2

5. The combining capacity of oxygen in water is: a. 1

b. 2

6. An element has X valence electrons in its outermost orbit, its valency may be: a. equal to X

b. equal to 8 - X

both a & b

d. none of these

7. Identify the possible number of valence electrons for metals: a. 4

b. 5

d. 1

d. All

III. Atomic number and mass number

1. The number of protons present in nitrogen is: a. 14

b. 7

2. Among the following, the element with an odd number as their atomic number is/are: a. C

b. Mg

d. 0

Z-A

d. A×Z2

d. 15

3. In an atom, the number of neutrons is equal to: a. A+Z

b. A-Z

4. The atomic number of Argon is: a. 2

b. 10

5. Among the following, the elements with an even number of neutrons is/are: a. Na

b. Al

d. All

6. The atomic number is equal to the:

150

a. number of neutrons in the nucleus

b. sum of protons and neutrons

c. number of protons in the nucleus

d. atomic mass of the element

IL Foundation Series Class 9

7. A and Z can be: a. negative

b. fractional

zero

d. whole number

8. The number of protons, electrons and neutrons in 80 35 Br are ___________, respectively. a. 35, 35, 80

b. 35, 35, 45

80, 80, 35

d. 45, 45, 35

IV. Isotopes, isobars and isotones

1. The number of nucleons in the isotope of an atom z Xm are: a. m

b. z

m+z

d. m-z

Potassium

d. Aluminium

2. Which of the following elements contains 20 neutrons? a. Sodium

b. Magnesium

3. Two nuclides, X and Y, are isotonic to each other with mass numbers 70 and 72, respectively. If the atomic number of X is 34, then that of Y would be: a. 32

b. 34

d. 38

4. Natural chlorine has 2 isotopes 1735Cl and 1737Cl in the ratio of 1:3. Calculate the atomic mass of natural chlorine: a. 35 a.m.u

b. 37 a.m.u

35.5 a.m.u

d. 34 a.m.u

5. Which one of the following is an isobar of 6 C14 ? a. 6C13 b. 6C12 c. 7N14 d. 7N15 6. Two nuclides, A and B, are isoneutronic. Their mass numbers are 76 and 77, respectively. If the atomic number of A is 32 , then the atomic number of B will be: a. 33 b. 34 c. 32

d. 30

V. Bohr's theory of hydrogen atom, energy of the electron in the nth orbit (En), and ionisation energy

1. The correct order of energy of the shells is: a.

K>L>M>N

b. L>K>M>N

c. N>K>L>M

d. N > M > L > K

c. 1.058 Å

d. 0.176 Å

d. 10.2 eV

2. The radius of the K shell in Li+2 is: a.

0.529 Å

b. 2.116 Å

3. The energy difference between 5th and 4th orbits is: a.

0.66 eV

b. 1.9 eV

0.30 eV

4. According to Bohr's theory, the radius of the nth orbit of the hydrogen atom is: a.

0.529×n2 Å

b. 0.529/n2 Å

0.529×n Å

d. 0.529/n Å

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STRUCTURE OF THE ATOM

5. The energy of an electron in the nth orbit of the hydrogen atom is: a.

(-2π2 Z2 me4)/(n2 h2 )

b. (-4π2 me4)/(n2 h2 )

(-n2 h2)/(2π2 kme4 )

d. (-n2 h2)/(4π2 kme4 )

6. Statement A: The potential energy of an electron = 2× total energy. Statement B: The kinetic energy of an electron = - total energy Statement C: rn=(n2 h2)/(4π2 me2 Z) a.

All the above statements are correct.

b. All the above statements are incorrect.

A and B are correct, but C is incorrect.

d. A and B are incorrect, but C is correct.

7. Which state of the triply ionised Beryllium (Be+++) has the same orbital radius as that of the ground state of hydrogen? a.

n=3

b. n=4

n=1

d. n=2

8. The ionisation energy of a hydrogen atom is 13.6eV. The energy required to ionise a hydrogen atom in which the electron is in the second orbit from the nucleus is (eV): a.

-13.6

b. 3.4

10.2

d. 13.6

9. The ionisation energy of a hydrogen atom is 13.6eV. Following Bohr's theory, the energy corresponding to a transition between the 3rd and 4th orbit is: a.

3.40 eV

b. 1.51 eV

0.85 eV

d. 0.66 eV

VI. Quantum numbers

1. The maximum number of electrons that can be accommodated in the M shell are: a.

b. 8

c. 18

d. 32

2. The maximum number of electrons in an orbital is: a. 1 b. 2 c. 3 3. Assertion (A): The principal quantum number is denoted by the letter n.

d. 4

Reason (R): The principal quantum number gives information about the size and energy of the a.

orbit. Both A and R are correct, and R is the correct explanation for A.

b. Both A and R are correct, but R is not the correct explanation for A. c.

A is incorrect, but R is correct.

d. A is correct, but R is incorrect.

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IL Foundation Series Class 9

4. The azimuthal quantum number for the last electron in the sodium atom is: a.

b. 2

d. 3

5. The maximum number of electrons that can be filled into all the orbitals corresponding to the azimuthal quantum number, l = 3, is: a.

b. 6

d. 14

6. For the f-subshell, l = 3, then the total number of m values will be: a.

b. 5

d. 4

c. 9

d. 3

7. If n = 3, then the total number of m values will be: a.

b. 14

8. The s value for the anticlock-wise direction is: a. +1/2 b. -1/2 c. Both (a) & (b) d. None 9. The values of four quantum numbers of a valence electron of an element are: n = 4, l = 0, m=0 and s = +1/2. The element is: a. K b. Fe

d. Sc

VII. Shapes of orbitals and arrangement of electrons in orbitals

1. The shape of d-orbital is: a.

dumbbell

b. spherical

double dumbbell

d. complex structure

2. The shape of f-orbital is: a.

double dumbbell

b. dumbbell

spherical

d. complex structure

3. The d-orbital, which is different in shape from the remaining d-orbitals, is: dZ 2

d x2 − y 2

dxy

d. dzx

4. The spherically symmetrical orbital is: a. s b. p c. d d. f 5. As much as possible, electrons will go into the orbital with the lowest energy. This statement is based on: a. Aufbau principle

b. Pauli's exclusion principle

d. Bohr's rule

Hund's rule

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STRUCTURE OF THE ATOM

6. Statement A: The (n+l) value of 3d is 5. Statement B: The (n+l) value of 5p is 6. Statement C: The (n+l) value of 5d is 4. a.

All the above statements are correct.

b. A and B are correct, but C is incorrect.

All the above statements are incorrect.

d. A and B are incorrect, but C is correct.

WORKSHEET - 2 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER

1. If the specific charge of a proton (e/m) is 9.6×107 C/kg, then that for an α-particle will be: a.

2.4×107 C/kg

b. 4.8×107 C/kg

c. 19.2×107 C/kg

d. 38.4×107 C/kg

2. Match the following: Column – I

Column - II

A. Cathode rays possess kinetic energy

1. When they strike the metallic target

B. Cathode rays produce X – rays

2. In the presence of an electric field

C. Cathode rays deflect towards the positive end 3. Cathode rays cause mechanical motion A B C a.

3 1 2

b. 1 2 3

2 3 1

d. 1 3 2

3. Match the following. Set A

Set B

A) Size of the nucleus

1) Doubly charged helium

B) Size of the atom

2) Fluorescent screen

C) α-particle

3) 10-8 cm

D) Zinc sulphide screen

4) 10-13 cm

A B C D a. 4 3 1 2

b. 4 2 3 1

c. 4 3 2 1

d. 4 1 2 3

4. According to Bohr's atomic model, mvr = a.

154

n h/2π

b. nh/π

h/4π

d. h/2nπ

IL Foundation Series Class 9

5. The value of Planck's constant is: a.

6.6×10-32 gm2 sec

b. 6.6×10-34 kg m2 sec-1

c. 6.6×10-32 kg msec-1 d. 6.6×10-34 g m2 sec 6. During the transition of electrons from the orbits of higher energy to the orbit of lower energy, the energy is: a. absorbed

b. radiated

d. none of these

either (a) or (b)

7. In Bohr's model of an atom, which of the following is an integral multiple of h/2π? a.

Kinetic energy

b. Radius of an atom

c. Potential energy

d. Angular momentum

8. Bohr's model can explain: a.

The spectrum of hydrogen atoms only

b. The spectrum of an atom or ion containing one electron only c.

The spectrum of hydrogen molecule

d. The solar spectrum 9. The difference in angular momentum of a revolving electron in the 3rd to that of the 6th orbit is: a.

6 h/π

b. 4.5 h/π

3 h/π

d. 1.5 h/π

(n-2)

d. All

10. The valence shell is represented as: a. n

b. (n-1)

11. The increasing order of energy of the orbitals 3s, 3p, 3d, 4 s is: a.

3s <3p <3d < 4s

b. 3s < 3p < 4s < 3d

3d <4s <3p < 3s

d. 3s = 3p = 3d < 4s

12. 1s2 2s2 2p6 3s2 3p5 is the electronic configuration of: a. Cl

b. Al

c. P

d. Ca

13. Match the following. Column-1

Column-2

A. Ultimate shell

1. n-1

B. Penultimate shell

2. n-2

C. Antipenultimate shell

3. n

A B C a.

3 1 2

b. 2 1 3

2 3 1

d. 3 2 1 155

STRUCTURE OF THE ATOM

14. Match the following. Column - I (Shell)

Column - II (Subshell)

A. K shell

1. s, p, d, f

B. L shell

2. s, p, d

C. M shell

3. s, p

D. N shell

4. s

A B C D a. 4 3 2 1

b. 3 4 1 2

c. 4 2 1 3

d. 4 3 1 2

15. The valency of copper in cupric chloride is: a. 2

b. 1

c. 3

d. 4

16. Match the following. Elements

Valency

A. Beryllium

1. 0

B. Boron

2. 4

C. Carbon

3. 3

D. Neon

4. 2 A B C D

a. 4 3 2 1

b. 4 2 3 1

c. 4 2 1 3

d. 4 1 2 3

17. The nucleus of an atom contains: a.

Electrons and protons

b. Electrons and beta particles

Protons and neutrons

d. Protons and alpha particles

18. The isotopes of neutral atoms of an element differ in: a. Atomic number

b. Mass number

d. Chemical properties

Number of electrons

19. The nucleus of tritium consists of:

156

1 proton +1 neutron

b. 1 proton +3 neutrons

1 proton + zero neutrons

d. 1 proton +2 neutrons

IL Foundation Series Class 9

20. An atom differs from its ion in: a. Nuclear charge

b. Mass number

c. Number of electrons d. Number of neutrons 21. The ratio of the neutrons present in a carbon atom and silicon atom with mass numbers 12 and 28, is: a. 07:03 b. 3:7 c. 1:2 d. 2:1 22. An oxide of nitrogen has a molecular weight of 30. The total number of electrons in one molecule of the compound is: a. 15

b. 30

c. 45

d. 60

23. The maximum sum of the number of neutrons and protons in an isotope of hydrogen is: a. 6

b. 5

c. 4

(JEE)

d. 3

24. In two elements, z 1 AM1 and z 2 B M 2 M1 ≠ M2 and Z1 ≠ Z2, but M1-Z1= M2-Z2. These elements are: ,

a. Isotonic

b. Isobaric

Isotopic

d. Isosonic

25. If two neutrons are added to an element X, then it will get converted to its: a. Isotope

b. Isotone

d. None of the above

Isobar

26. The radius of Bohr's first orbit is: a. 0.529 Å

b. 5.29 Å

0.0529 Å

d. 0.529×10-8 Å

27. In case of the hydrogen atom, the energy of the electron in the nth orbit is given by: a.

(-313.6)/n2 Joule / mole

b. (-313.6)/n2 Kcal/mole

(-2.8×10-18)/n2 Cal/atm

d. (-2.8×10-18)/n2 J/mole

28. If the radius of Bohr's first orbit is given as 4 A0; then the radius of Bohr's fourth orbit is: a. 4 Å b. 8 Å c. 16 Å d. 64 Å 29. The radius of Bohr's first orbit is r0. The radius of the electron in the first orbit of a singly ionised helium atom is: a. r0/2

b. 2r0

4r0

d. r0/4

30. According to Bohr's theory, the radius of an electron in an orbit is proportional to: a. z2 n2 b. z2/n2 c. z2/n d. n2/Z 31. The radius of the first orbit of the electron in a hydrogen atom is 0.53 A. So, the radius of the third orbit will be: a. 2.12 Å

b. 4.77 Å

1.06 Å

d. 1.59 Å

157

STRUCTURE OF THE ATOM

32. If Ep and Ek are the potential and kinetic energies of the electron in a stationary orbit in the hydrogen atom, then the value of Ep/Ek is: a. -2 b. 2

c. -1

d. 1

33. The number of orbitals in the quantum level n = 4 is: a. 4

b. 9

c. 16

d. 18

34. Statement (A): The Azimuthal quantum number specifies the shapes of orbitals. Statement (B ): l values are from 0 to n-1. Statement (C): l = 0 represents the s-orbital. a.

All the above statements are correct.

b. A and B are correct, but C is incorrect.

All the above statements are incorrect.

d. A and B are incorrect, but C is correct.

35. Statement (A): m values depend on l values. Statement (B): For d subshells, the total number of m values is 5. Statement (C ): The magnetic quantum number indicates the size of the orbit. a.

All the above statements are correct.

b. A and B are correct, but C is incorrect.

All the above statements are incorrect.

d. A and B are incorrect, but C is correct.

36. For the pz orbital, conventionally m is: a. -2

b. 2

c. 0

d. None of these

37. If the value of the principle quantum number is 3 , the total possible values for the magnetic quantum number will be: a. 5

b. 9

c. 8

d. 10

38. An electron has a magnetic quantum number 3 , then its principal quantum number is: a. 3

b. 2

c. 1

d. 4

c. 3

d. 1

39. The number of electrons with +1/2 spin in n = 3 is: a. 18

b. 9

40. The impossible set of quantum numbers for an electron is: a. n = 2, l = 0, m = 0, s = +1/2

b. n = 2, l = 1, m = 0, s = +1/2

d. n = 3, l = 1, m = -1, s = -1/2

n = 2, l = 0, m = 1, s = -1/2

41. The number of electrons with -1/2 spin in n = 4 is: a. 32

b. 16

c. 8

d. 9

42. Which of the following elements has the least number of M-shell electrons: a.

158

b. Mn

d. Sc

IL Foundation Series Class 9

43. The total number of p electrons present in the phosphorus atom is: a. 9

b. 2

c. 8

d. 3

44. The valence electronic configuration of an element with atomic number 23 is: a.

3d5

b. 3 d3 4 s2

3d2 4s1 4p1

d. 3d2 4s2 4p1

45. Mg2+ and Al3+ have the same: a.

Protons

b. Neutrons

Electronic configuration

d. Neutrons + protons

46. The number of unpaired electrons in the valence shell of silicon is: a. 2

b. 3

c. 1

d. 0

47. Among the following ions, which are isoelectronic with V3+ ? a.

Sc+1

b. Mn+7

Ni+2

d. Fe+3

48. No two electrons in an orbital can have a parallel spin. This statement emerges from: a.

Hund's rule

b. Aufbau principle

Pauli's exclusion principle

d. Heisenberg's uncertainity principle

159

PERIODIC CLASSIFICATION OF ELEMENTS

5.1 EARLY ATTEMPTS AT THE CLASSIFICATION OF ELEMENTS 5.1.1 Introduction to early attempts It is difficult to study individually the chemistry of more than one hundred elements known today and their innumerable compounds. The experimental data regarding elements and their compounds can only be systematised if proper classification is done. The basic object of classification is to arrange the facts regarding elements and their compounds in such a way that we may have greater control over their characteristics with less possible effort. The best classification would be the one that puts together those elements that resemble in most respects and separate each other. At present, 118 elements are known. Of these elements, 92 are available in the elemental form, and the remaining 26 elements are human-made. The classification may help to study the elements better and to correlate the properties of elements with some fundamental properties that are characteristic of all the elements. 5.1.2 Döbereiner’s triads In 1829, John Wolfgang Dobereiner, a German scientist, was the first to consider the idea of trends among the properties of elements. He noted the physical and chemical properties of several sets of three elements, arranged in the increased order of their atomic weights called 'triads'. The law of triads states, 'When elements are arranged in order of their increasing atomic mass, the

mass of the middle element is approximately the arithmetical mean of the remaining two elements of the triad.'

Elements of triad

Symbol

Atomic mass

Lithium

7.0

Sodium

23.0

Potassium

39.0

Chlorine

35.5

Bromine

80.0

Iodine

127.0

Calcium

Strontium

87.5

Barium

137

Table 5.1 Examples of Dobereiner's triads

160

Arithmetic mean 7 + 39 = 23.0 2 35.5 + 127 = 81.25 2 40 + 137 = 88.5 2

IL Foundation Series Class 9

Limitations of Dobereiner’s system •

Few elements were known at the time of Dobereiner. The law of triads seemed to work only for a few elements. Quite a large number of similar elements could not be grouped into triads. For example, Fe, Mn, Ni, Co, Zn, and Cu are similar elements but cannot be placed in the triads.

•

Similarly, it was possible that quite dissimilar elements could be grouped into triads. For example, Carbon (12), Nitrogen (14), and Oxygen (16) can form a triad, but their properties are entirely different from each other.

5.1.3 Newland’s law of octaves The English chemist John Alexander Newlands (1865), a lover of music, propounded the law of octaves. He arranged many of the known elements in the increasing order of their atomic masses and noticed that the properties of every eighth element were a kind of repetition of the first element, just like the eighth note of octave in music, Indian as well as Western. Based on this observation, Newland formulated a law stating that when the known elements were arranged in the increasing order of their atomic mass, every eighth element had properties similar to the first one. Octaves of music

Newland’s arrangement

Indian

Western

Li (7.0) Na (23.0) K (39.0)

Be (9.0) Mg (24.0) Ca (40.0)

B (11.0) Al (27.0) Cr (52.7)

C (12.0) Si (28.0) Ti (48.0)

N (14.0) P (31.0) Mn (55.0)

O (16.0) S (32.0) Fe (56.0)

F (19.0) Cl (35.5) Br (80.0)

Table 5.2 Newland’s arrangement of elements

From Newland's classification, a very important conclusion was made that there is some systematic relationship between the order of atomic mass and repetition of properties, which gives rise to a new term, 'periodicity'.

5.2 MENDELEEV’S PERIODIC TABLE 5.2.1 Introduction to Mendeleev’s periodic table In 1869, Dmitri Ivanovich Mendeleev (father of periodic classification), an eminent Russian chemist, was the first to arrange all elements successfully. After a thorough study of elements and their properties, he proposed that when elements are arranged in the increasing order of their atomic masses, the elements with similar properties appear at regular intervals. He fully recognised the significance of periodicity and used a broader range of physical and chemical properties to classify the elements.

161

PERIODIC CLASSIFICATION OF ELEMENTS

Mendeleev was responsible for publishing the periodic law for the first time. It states as follows. 'The properties of the elements are a periodic function of their atomic weights'. The main points of periodic law, as stated by Mendeleev in his original published paper, are as follows: 1. I f the elements are arranged according to their atomic masses, they exhibit an evident periodicity of properties. 2. The elements with similar properties have either almost the same atomic weights, e.g., Fe(56), Co(59), Ni(59), or increase regularly, e.g., K(39), Rb(85), Cs(133), Ca(40), Sr(99), Ba(137). 3. The magnitude of atomic weights determines the character of the element. 4. The arrangement of elements in the order of their atomic weights corresponds to their 'valencies' and their properties. 5. The elements with low atomic weights were found to be widely distributed in nature and possess sharply defined properties. 6. Many yet unknown elements may be discovered, and their characteristic properties can be foretold from their atomic weights. Elements with similar properties were arranged in vertical columns called 'groups'. The horizontal rows were called 'periods'. Series 1 2 3 4 5 6

162

III

R2O

R2O3

9.4

27.3

35.5

Fe Co Ni Cu

56 59 59 63

Ru Rh Pd Ag

100

104 104 106 108

RO2/RH4 R2O5/RH3

VII

VIII

R2O7/RH

RO4

H 1

IL Foundation Series Class 9

108

112

113

118

122

127

133

137

138

140

178

180

182

184

199

200

204

207

208

7 8

11 12

Th 231

Os Ir Pt Au 195 197 198 199

U 240

Table 5.3 Mendeleev’s original periodic table of elements (1871)

Mendeleev arranged all the known 63 elements in the increasing order of their atomic masses. This arrangement showed that the elements having similar chemical properties came directly under one another in the same group. In this periodic table, Mendeleev left a few gaps (shown by ?) as unknown elements could be placed there. These missing elements are called 'Eka' elements. Some of these elements include Scandium (Sc) - Eka Boron, Gallium (Ga) - Eka Aluminium, and Germanium (Ge) - Eka Silicon. Noble gases were not known at that time. When they were discovered towards the end of the nineteenth century, they were easily accommodated as a separate column in the periodic table. Description of Mendeleev’s periodic table

Mendeleev's modern periodic table consists of vertical columns called 'groups' or 'families' and horizontal rows called 'periods'. The groups are marked from 0 to VIII. The groups I to VII are divided into subgroups A and B. Groups I to VII are termed normal groups. In the VIII group, three similar elements are placed together in one place, known as transition triads - Fe, Co, Ni; Ru, Rh, Pd; Os, Ir, Pt; Hs (Uno), Mt (Une), Uun. The first three periods of Mendeleev’s periodic table are called short periods, and the other periods are known as long periods. The elements present in the short periods are referred to as typical elements. The long periods of Mendeleev’s periodic table consist of two rows of elements, and each row of elements is called a series.

163

PERIODIC CLASSIFICATION OF ELEMENTS

5.2.2 Achievements of Mendeleev’s periodic table 1. Classification of elements

Mendeleev’s periodic system is more elaborate and is far superior to all earlier classifications. He fully recognised the importance of periodicity and utilised a wide range of physical and chemical characteristics as well as the formulae and the properties of the compounds formed by the elements in classifying elements. 2. Prediction of new elements

Several gaps were left in Mendeleev’s periodic table for unknown elements. He even predicted the properties of these unknown elements, helping scientists discover the unknown elements more readily and accurately. Thus, the periodic system hastened the discovery of many new elements. The following table shows a comparison of properties predicted by Mendeleev for the elements and those found experimentally after their discovery. Name of the element

Property

Eka Boron

Scandium

Atomic weight

43.80

Specific gravity

3.5

3.864

Formula of oxide

(Eka B)2 O3

Sc2O3

Formula of sulphate

(Eka B)2 (SO4)3

Sc2(SO4)3

Eka Aluminium

Gallium

Atomic weight

69.90

Specific gravity

5.90

5.94

Formula of oxide

(Eka Al)2O3

Ga2O3

Formula of chloride

(Eka Al)Cl3

GaCl3

Solubility in acids and alkalis

Dissolves slowly in both

Eka Silicon

Germanium

Atomic weight

72.6

Specific gravity

5.50

5.47

Valency

Formula of oxide

(Eka Si)O2

GeO2

Formula of chloride

(Eka Si)Cl4

GeCl4

Isolation

By the reduction of its oxide By the reduction of GeO2 with (or) its fluoro complex with Na C or of K2GeF6 with Na Table 5.4 Comparison of properties of elements

164

IL Foundation Series Class 9

3. Correction of atomic masses

The atomic mass of an element is related to the equivalent mass of the element by the formula. Atomic mass = Equivalent mass ◊ Valency The valency of the element can be known from its position in the periodic table. For example, the atomic mass of beryllium was corrected from 13.5 to 9 (4.55 × 2 = 9.10) . Similarly, the atomic masses of In, Au, Pt, etc. were also corrected. 5.2.3 Limitations of Mendeleev’s classification Despite its great usefulness in the study of various elements, Mendeleev’s periodic table suffers from the following limitations. 1. Position of hydrogen

he position of hydrogen in the periodic table is uncertain. It is sometimes placed in the first group T and sometimes in the seventh group, as hydrogen shows a resemblance with alkali metals and also with halogens. 2. Grouping of elements

ertain elements which show similar properties have been separated in the periodic table. For C example, Ba and Pb resemble each other in so many respects but have been placed in the second and fourth groups, respectively. Similarly, Ag and Tl, which have similar properties, have been placed in the first and third groups, respectively. Certain chemically dissimilar elements have been grouped together. For example, the elements Cu, Ag, and Au were grouped with IA group elements such as Li, Na, and K, which have quite dissimilar properties. 3. Anomalous pairs of elements

or certain pairs of elements, the chemical properties observed were not in apparent agreement with F the positions allotted according to atomic weights. In Mendeleev’s periodic table, four pairs of elements are in the reverse order of their atomic masses. These elements are called anomalous pairs of elements. These areElement

Atomic number

Atomic weight

Argon (Ar)

Potassium(K)

Cobalt(Co)

58.9

Nickel (Ni)

58.6

Tellurium (Te)

127.60

Iodine (I)

126.90

Thorium (Th)

232

Protactinium (Pa)

231

Table 5.5 Anomalous pairs of elements 165

PERIODIC CLASSIFICATION OF ELEMENTS

4. Position of isotopes

I sotopes are atoms of the same elements having different atomic mass but same atomic number. 1 2 3 For example, Hydrogen has three Isotopes - 1H , 1H , and 1H . According to Mendeleev’s periodic law, these elements should be placed at three separate places in the periodic table. However, isotopes have not been given separate places in the periodic table.

5.3 THE MODERN PERIODIC TABLE The empirical evolution of the periodic table reached its peak in 1913 when Henry Moseley showed that atomic number is a more fundamental property of an element than its atomic weight. The position of an element in the periodic table depends on its atomic number, and the reason for the anomalies in the original periodic table becomes clear at once. 5.3.1 Periodic table Moseley modified Mendeleev’s periodic law and stated, "The physical and chemical properties of the elements are periodic functions of their atomic numbers". The atomic number is equal to the nuclear charge or the number of electrons in the neutral atom. Further, it was recognised that the periodic law is essentially the consequence of the periodic variation in electronic configurations. The configurations determine the properties of elements and their compounds and are the basis for the modern periodic law.

Periodic Table of the Elements (Long Form) (Representing Electron Conﬁguration)

Main group Elements p-subshell is gradually ﬁlled up

Main group Elements s-subshell is gradually ﬁlled up Group Period

3 4 5 6 7

H 1s

1 2

IA 1

IIA

1s1 3

2s1

3s1

3s2 20

IIIB

6s1

6s2

5d16s2

Hf 104

5d76s2 109

7s1

7s2

6d17s2

6d27s2

6d37s2

6d47s2

6d57s2

6d67s2

6d77s2

5d96s1 110

6d97s1

4s24p6 54

5s25p5 85

5s25p6 86

6s26p4 116

4s24p5

5s25p4

6s26p3 115

Ar 3s23p6 36

6s26p2 114

Cl 3s23p5

4s24p4

5s25p3 83

Ne 2s22p6 18

6s26p1 113

5s25p2 82

F 2s22p5

3s23p4

4s24p3 51

1s1 10

5d106s2 112

4s24p2

5s25p1

O 2s22p4

3s23p3 33

5d106s1 111

4d105s2 80

VIIA 9

3s23p2

4s24p1 49

4d105s1 79

5d66s2 108

4d10 78

Zn 3d104s2 48

4d85s1

5d56s2 107

Cu 3d104s1 47

N 2s22p3 15

VIA 8

2s22p2

3s23p1 31

VA 7

IIB 30

3d84s2 46

5d46s2 106

Ru 4d75s1 76

3d74s2 45

4d55s6 75

3d64s2 44

4d55s1

5d36s2 105

3d54s2

Mo 74

Mn 43

4d45s1 73

5d26s2 **

4d25s2 72

IB 26

3d44s2 42

VIIB 25

3d34s2 41

4d15s2 57

VIB 24

3d24s2 40

5s2 56

VB 23

3d14s2 39

5s1

IVB 22

4s2

IVA 6

VIIIB 21

IIIA 5

2s22p1

4s1 37

2s2 12

0 (zero)

Transition Elements d-subshell is gradually ﬁlled up

Atomic number Symbol Valence-shell Conﬁguration

6s26p5 117

6s26p6 118

Uut

Uup

Uus

Uuo

6d107s1

6d107s2

7s27p1

7s27p2

7s27p3

7s27p4

7s27p5

7s27p6

Inner-Transition Elements

f-subshell is gradually ﬁlled up 58

*Lanthanides

**Actinides

4f26s2 90

4f36s2 91

4f46s2 92

4f56s2 93

4f66s2 94

Eu 4f76s2 95

Gd 4f75d16s2 96

Tb 4f96s2 97

4f106s2 98

4f116s2 99

4f126s2 100

4f136s2

4f146s2 102

101

Lu 4f145d16s2 103

6d27s2

5f26d17s2

5f36d17s2

5f46d17s2

5f67s2

5f77s2

5f76d17s2

5f97s2

5f107s2

5f117s2

5f127s2

5f137s2

5f147s2

5f146d17s2

Fig. 5.1 Periodic table of elements 166

IL Foundation Series Class 9

The modern periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers or their electronic configurations. The original form of the periodic table has since been modified as a result of the structural elucidation of atoms and the discovery of noble gas elements. Numerous forms of the periodic table have been devised from time to time. Moseley constituted the periodic table by unfolding Mendeleev’s table, in which the elements are arranged according to the atomic numbers. This is closely similar to the arrangement of the Bohr's arrangement. The most convenient version of the periodic table was constructed by Bohr based on the modern periodic law, and the elements were arranged in the order of their electronic configurations. 5.3.2 Position of elements in the modern periodic table It is very difficult to study and remember the properties of all the elements. So, to study the properties of elements in a better way and to know the relation between one element and another, a classification of elements is necessary. The vertical columns of the periodic table are called groups or families. There are eighteen groups. Among these groups, eight are important. These groups are 1, 2, 13, 14, 15, 16, 17 and 18. A part of group 3 elements are separately shown at the bottom of the periodic table in the form of Lanthanides and Actinides. Fourteen elements coming after lanthanum are called lanthanides. Lanthanides are from Cerium (Z = 58) to Lutetium. (Z = 71). Fourteen elements coming after actinium are called actinides. Actinides are from Thorium (Z = 90) to Lawrencium. (Z = 103). Lanthanides are commonly called rare earths. Most of the actinides are mainly synthetic elements. Hydrogen is the only element that can be placed in two different groups of the periodic table: group 1 and group 17. Groups : Groups number

Common name

IA Group elements

Alkali metals

IIA Group elements

Alkaline earth metals

IIIA Group elements

Boron family

IVA Group elements

Carbon family

VA Group elements

Nitrogen family (pnictogens)

VIA Group elements

Oxygen family (Chalcogens)

VIIA Group elements

Halogens

VIIIA Group elements

Noble gases / Inert gases / Aerogens / Zero group elements Table 5.6 Groups in a periodic table

Periods: There are 7 horizontal rows in the periodic table. These are called periods. 1. First Period: It contains only two elements: hydrogen and helium. So, it is called the 'shortest period’.

167

PERIODIC CLASSIFICATION OF ELEMENTS

2. Second Period: It contains 8 elements: Li, Be, B, C, N, O, F and Ne. 3. Third Period: It contains 8 elements: Na, Mg, Al, Si, P, S, Cl, and Ar. The second and the third periods are called 'Short periods’. 4. Fourth Period: It contains 18 elements from K to Kr (Z = 19 to 36). 5. Fifth Period: It contains 18 elements from Rb to Xe (Z = 37 to 54). The fourth and fifth periods are called 'Long periods’. 6. Sixth Period: It contains 32 elements from Cs to Rn and is called the 'longest period’ ( Z = 55 to 86 ). 7. Seventh period: It contains 32 elements from Fr to Og and is called the 'Longest period’ ( Z = 87 to 118 ). Based on the differentiating electron, the periodic table is divided into four main blocks - s, p, d, and f. i) Elements of groups 1 and 2 are s - s-block elements. ii) Elements of groups 13, 14, 15, 16, 17, and 18 are p - block elements. iii) Elements of groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12 are d- block elements. iv) The f-block elements comprise two horizontal rows placed at the bottom of the periodic table. s-block 1

1s1

p-block

1s1

3s1

5s1

Ba 6s

Fr 7s1

Ra 7s2

* * * *

La Lu

89 103

Ac Lr

104

105

14 1 1 2 2 5f6d 6d 7s7s

106

6d27s2

107

6d37s2

108

6d47s2

5d76s2 109

6d57s2

6d67s2

5d96s1

6s26p2

6s26p3

6d107s1

f-block 57

5d 6s 1

4f 6s

6d17s2

4f 6s

6d27s2

Nd 4

5f26d17s2

5f36d17s2

4f 6s

5f46d17s2

4f 6s

5f67s2

4f 5d 6s 7

Am 5f77s2

4f 6s

5f76d17s2

Dy 10

5f97s2

4f 6s

5f107s2

4f 6s

100

5f117s2

Tm 4f 6s

102

101

Fm 5f127s2

4f146s2

Md 5f137s2

5f147s2

5s25p5 85

6s26p4

4s24p6

5s25p4

6d97s1

3s23p6 36

4s24p5

5s25p3 83

6s26p1

Fig. 5.2 Different blocks in a periodic table

168

5s25p2 82

5d106s2

3s23p5

4s24p4

111

6d77s2

5s25p1 81

5d106s1

110

4d105s2 80

2s22p6

3s23p4

4s24p3

4s24p2

4d105s1 79

5d66s2

4d10

4s24p1

2s22p5

3s23p3

3d104s2

5d56s2

3s23p2

3d104s1

4d85s1 77

5d46s2

3d84s2

4d75s1 76

5d36s2

3d74s2

4d55s6 75

5d26s2

1 2 2

4d55s1

3d64s2

4f5d 5d 6s6s 14 1

4d45s1 73

3d54s2

2s22p4 16

3s23p1 27

3d44s2

4d25s2

57 71

3d34s2

4d15s2

5s2

3d24s2

3d14s2

4s2

2s22p3

d-block

3s2

2s22p2

2s22p1

2s2

2s1

5s25p6 86

6s26p5

6s26p6

IL Foundation Series Class 9

5.3.3 Trends in the modern periodic table Atoms of the elements present in the same group of the periodic table have similar outer shell configurations. Repetition of similar valence shell configuration is the cause of periodicity in properties. In the periodic table, the properties of elements change gradually with a change in their electronic configurations. This trend repeats itself at regular intervals. This repetition of a character is called "periodicity" or periodic trends. Some of the periodic trends are listed below•

Valency

•

Atomic size

•

Metallic and Non-metallic property

Valency

Valency is the combining capacity of an element. (OR) According to the new concept, valency may be defined as "The number of electrons that are lost or gained or shared with one atom of that element to acquire the stable configuration of the nearest noble gas element''. The valency of metals is given by the number of valence electrons present in an atom. The valency of non-metals is given by subtracting the number of valence electrons present in an atom from 8, i.e., (8-number of valence electrons). The combining capacity of elements is often compared to that of hydrogen, chlorine or oxygen. Valency is defined as the number of hydrogen atoms, the number of chlorine atoms, or twice the number of oxygen atoms with which one atom of the element combines. The maximum valency of an element is its group number. The minimum valency is zero. The highest valency is exhibited by Os or Xe. The value is 8. In the second period, the highest valency of 4 is exhibited by carbon, and in the third period, the highest valency of 7 is exhibited by chlorine. Atomic size

Atomic radius is the distance between the centre of the atomic nucleus and the electron cloud of the outermost energy level. Atomic radius is also commonly referred to as atomic size. But atomic size is truly regarded as the diameter of the atom. Variation of atomic radii in period

The atomic radii decrease from left to right along a period in the periodic table. In a period, the atomic number increases, and distinguishing electrons enter the same outer shell; hence, nuclear charge increases. This increases the attraction between the nucleus and the extranuclear electrons as the number of orbitals remains the same. Due to this, all electrons in orbitals are pulled closer to the nucleus. This goes on from atom to atom in a period. The atomic radius of inert gas is shown 169

PERIODIC CLASSIFICATION OF ELEMENTS

to be the largest in a period because of its van der Waals radius, which is generally larger than the covalent radii. Example : 2nd period

Li > Be > B > C > N> O > F < Ne

Atomic radii (in Å) 1.23, 0.89, 0.82, 0.77, 0.75, 0.73, 0.72, 1.60 Variation of atomic radii in a group

In a group from top to bottom, the atomic number increases, and the valence shell increases; hence, the atomic radii increase. Ex: Li (1.23), Na (1.57), K (2.03), Rb (2.16), Cs (2.35). The increase in size is due to the presence of extra energy shells in the elements as we go down the group. Metallic and non-metallic properties

Metallic: It is the tendency to lose electrons and form a positive ion. It decreases over a period and increases on moving down a group. The metals with low melting and boiling points are - Cs, Ga, Na, K, etc. The metals with high melting and boiling points are - W, Hg, Fe, etc. It increases along a period and decreases on moving down a group due to an increase in size.

5.4

ELECTRONIC CONFIGURATIONS AND TYPES OF ELEMENTS: S-, P-, D-, F-BLOCKS

In this part, we’ll see how the way atoms are arranged in the periodic table is directly related to how their electrons are arranged. Based on the entry of the differentiating electron into the sub-shells, elements are classified into four blocks. They are: 1. s-block elements,

2. p-block elements,

3. d-block elements and

4. f-block elements.

S-block is present on the left side, the p-block is on the right, the d-block is in the middle, and the f-block is at the bottom of the long form of the periodic table. Different blocks of the elements in the long form of the periodic table are given in the following figure.

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4 5

s-block

p-block d-block

6 7

f-block

Fig. 5.3 s, p, d, and f-block elements in the long form of the periodic table

Atomic number of aluminium is 13. The thirteenth electron of aluminium enters into the 3p-orbital. Hence, aluminium is called a p-block element. Similarly, the atomic number of titanium is 22, with the configuration of the differentiating electron 3d2. Titanium is called a d-block element. He, due to the 1s2 configuration, belongs to the 's’ block but is placed in the zero group due to its inert nature. 5.4.1 The s-block elements The elements in which the differentiating electron enters the s-orbital of the outermost shell are called s-block elements. The first two elements of each period belong to the s-block. Groups 1 and 2 (alkali and alkaline earth metals) constitute the s-block. These elements are located on the left side of the long form of the periodic table. The general electronic configuration of s-block elements is ns 1− 2 . Characteristic properties of s-block elements 1. The elements are highly electropositive and are soft metals with lower densities. 2. They are very good reducing agents. 3. They have low melting and boiling points. 4. They are very reactive and form ionic substances, except lithium and beryllium. 5. They exhibit an oxidation state of +1 or +2. 6. They impart characteristic colours to the flame. 5.4.2 The p-block elements The elements in which the differentiating electron enters the p-orbitals of the outermost shell are called p-block elements. Groups 13, 14, 15, 16, 17 and 18 (IIIA to VIIA and Zero groups) constitute 171

PERIODIC CLASSIFICATION OF ELEMENTS

the p-block. The group 18 element that is misplaced in the p-block is helium. These elements are located on the right side of the periodic table. The p block elements together with s block elements (except zero group) are referred to as representative elements or main group elements. The general electronic configuration of p-block elements is ns 2np1-6 or ns 2np x ( x = 1 to 6) . Characteristic properties of p-block elements 1. The elements include all metalloids, most of the non-metals and some metals. 2. All gaseous elements (except H2 and He) are p-block elements. 3. Most of these elements are highly electronegative and have high electron gain enthalpy. 4. Some elements are good oxidising agents. Some of them act as reducing agents. 5. Except for group 18, these elements are very reactive. They mostly form covalent compounds, e.g., Cl2 , O 2 , HCl, HCl etc., though ionic halides, oxides, sulphides, nitrides, etc., are also known to be formed by the p-block elements. 5.4.3 The d-block elements (transition elements) The elements in which the differentiating electron enters the d-orbitals of the penultimate shell are called d-block elements. Groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12 constitute the d-block. These elements are located in the middle, between the s-block and p-block of the long form of the periodic table. They have the properties intermediate to those of the s-block and p-block elements. They are referred to as transition elements. The d-block contains four series of elements. The 3d, 4d and 5d series are filled with ten elements each. The 6d- series is incomplete and has only eight elements. Zinc, cadmium and mercury, which have the electronic configuration (n − 1)d 10 ns 2 , do not show most of the properties of transition elements. The general electronic configuration of d-block elements is (n-1)d1-10ns1or2. Characteristic properties of d-block elements 1. The elements are all electropositive and are metals. 2. They are all solids except mercury, which is a liquid at room temperature. 3. Most of the elements possess catalytic activity. E.g., i. In the preparation of NH3 by Haber’s process, the catalyst is Fe, and the promoter is Mo. ii. In the preparation of H2SO4 by the contact process, the catalyst is platinised asbestos. 4. They form cations with different magnitudes of the charge. 5. They form ionic as well as covalent compounds. 6. They form complex compounds. e.g. Cu ( NH 3 )4  SO 4 , K 4 [ Fe(CN)6 ] 7. They form alloys and interstitial compounds.

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They mostly form coloured ions and exhibit variable valency ( e.g., Fe , Fe = BM for Fe +2 , µ 5.92 BM for Fe +3 ) ( µ 4.90 +2

) and paramagnetism.

5.4.4 The f-Block elements (inner transition elements) The elements in which the differentiating electron enters the f-orbitals of the anti-penultimate shell are called f-block elements. A part of group 3 constitutes the f-block. These elements are located at the bottom of the periodic table. There are two series of f-block elements. The first series follows lanthanum (Z = 57) and is called lanthanides [Ce(Z = 58) to Lu(Z = 71)]. The second series follows actinium (Z = 89) and is called actinides [Th(Z = 90) to Lr(Z = 103)]. The general electronic configuration of f-block elements is(n-2)f1-14(n-1)d0 or 1ns2. La and Ac are d-block elements, but lanthanides and actinides are f-block elements. Characteristic properties of f-block elements 1. These elements are heavy metals with high density. They form coloured ions and complexes with paramagnetic properties similar to d-block elements. 2. They are naturally available in very small quantities and are called rare earths. 3. Trans-uranic elements (Z > 92) are all synthetic. 4. They also form complexes and interstitial compounds. 5. They show a great deal of similarity among themselves in their properties. 6. Actinide elements are radioactive. 7. Actinides show a greater number of oxidation states compared to lanthanides. 5.4.5 Metals, nonmetals, and metalloids In addition to displaying the classification of elements into s, p, d- and f-blocks, other broad classifications of elements are given based on their properties. Elements can also be divided into metals, nonmetals, and metalloids based on their properties. Among the elements in the periodic table, more than 75% are metals. These metals appear on the left side of the periodic table. Metals are usually solids at room temperature. They usually have high melting and boiling points. They are good conductors of heat and electricity. Metals are malleable and ductile. About one dozen elements are nonmetals. They are placed at the top right-hand side of the periodic table. Nonmetals are usually gases at room temperature. Some of them are also solids with low melting and boiling points. These are poor conductors of heat and electricity (except graphite). Most of the non-metallic solids are brittle and are neither malleable nor ductile. Some other elements exhibit both metallic and non-metallic properties. Such elements are called metalloids. Metalloids are placed in the p-block. Examples include Ge, As, Sb, Te, etc. Metals usually react with dilute acids, while non-metals react with alkali.

173

PERIODIC CLASSIFICATION OF ELEMENTS

5.5 DIAGONAL RELATIONSHIP Definition: A characteristic analogy exists in elements of the 2nd and 3rd periods. The first element of a group resembles closely with the second element of the next successive group. This phenomenon is known as a diagonal relationship. Therefore, the first three elements of the second period (Li, Be, and B) not only show resemblance with the elements of their groups but also show resemblance with the elements diagonally placed in the higher groups.

Period

2nd Period 3rd Period

I Li Na

Groups

II Be Mg

Fig. 5.4 Diagonal relationship

III B Al

IV C Si

Lithium shows a close resemblance with magnesium, beryllium with aluminium and boron with silicon. The resemblance, however, disappears after these pairs. On moving far away in the period or in the group, the relationship disappears since the elements located diagonally show large differences in their atomic radii because the increasing or decreasing trend in atomic radii does not occur in the same order.

QUICK REVIEW

174

•

The first periodic table was constructed by Mendeleev. The periodic law is the physical and chemical properties are periodic functions of their atomic weights.

•

The elements with low atomic weights that were found to be widely distributed in nature are referred to as typical elements which are present in three short periods of Mendeleev’s table.

•

VIII group of the Mendeleev table contains three triads.

•

Newland’s Law of Octaves states that when the known elements were arranged in the increasing order of their atomic mass, every eighth element had properties similar to the first one.

•

There are 18 vertical columns in the long form, called groups and 7 horizontal rows called periods.

•

The first period is the shortest period; it has only 2 elements. The second and third periods are short periods with eight elements each. The fourth and fifth periods are long periods with eighteen elements each. The sixth period is the longest period, with 32 elements. The seventh period is incomplete.

•

The period number indicates the valence shell, and the group number in Roman letters denotes the number of electrons in the outermost shell. However, elements of the zero group generally have eight electrons in the outermost shell.

IL Foundation Series Class 9

•

The atomic radius is defined as the distance between the centre of the nucleus and the outermost shell where electrons are present.

•

Atomic radius increases down the group because of the presence of new shells.

•

In any period, in general, atomic radius is the least for halogen and highest for noble gas elements. Radius is highest for the element 'Cs’ and least for the element 'H’.

•

Elements in a group generally show the same valency. Across a period, the valence increases unit by unit. Transition metals show variable valency.

•

The maximum valency of an element is its group number.

•

The minimum valency is zero. The highest valency is exhibited by Os or Xe. The value is 8. In the second period, the highest valency of 4 is exhibited by carbon, and in the third period, the highest valency of 7 is exhibited by chlorine.

•

The tendency of an element to lose an electron is called metallic nature. It increases down the group as the size increases and across a period decreases as the size decreases.

•

Metallic nature increases and non-metallic nature decreases down the group. Metallic nature decreases across a period.

WORKSHEET - 1 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER l. Early attempts at the classification of the elements

1. The maximum number of elements available in elemental form is: a. 26

b. 92

c. 111

d.118

2. Which of the following is the Dobereiner triad? a. Li, Na, K

b. Fe, Co, Ni

c. Ru, Rh, Pd

d. Os, Ir, Pt

3. Identify the set of elements showing similar properties according to Newlands a. F and Cl II.

b. Al and Si

c. Mg and K

d. All

Mendeleev's periodic table

1. According to Mendeleev’s periodic law, the physical and chemical properties of the elements are periodic functions of their a. atomic number

b. atomic mass

c. both a and b

d. none

2. Anomalous pair among the following is a. Boron–silicon

b. Beryllium-indium

c. Aluminium–gallium

d. cobalt–nickel

3. The triad not present in Group VIII of Mendeleev’s table is 175

PERIODIC CLASSIFICATION OF ELEMENTS

a. Li, Na, K

b. Fe, Co, Ni

c. Ru, Rh, Pd

d. Os, Ir, Pt

4. Which of the following pairs of atomic numbers represents elements belonging to the same group? a. 11,20

b. 12,30

c. 13,31

d. 14,33

5. As per the modern periodic law, the physical and chemical properties of elements are periodic functions of their a. Atomic number

b. Mole number

c. Atomic weight

d. Atomic size

6. Taking chemical properties into consideration, the atomic weights of the following elements were corrected a. Te and I

b. Ar and K

c. Co and Ni

d. Be, Au, and Indium

7. Which one of the following has the most negative electron affinity, and which one has the least negative electron affinity? a. F, Cl

b. Cl, F

c. Cl, P

d. Cl, S

8. Which of the following pairs is against Mendeleev’s periodic law? a. Cr, Mn

b. Cu, Zn

c. Te, I

d. Na, Mg

9. Which of the following sets of elements are known as transition triads? a. Fe, Co, Ni

b. Ru, Rh, Pd

c. Os, Ir, Pt

d. all

10. Which of the following is not a defect in Mendeleev’s periodic table? a. Position of hydrogen

b. Grouping of elements

c. Anomalous pairs

d. Common valency of elements in a group

III. Modern periodic table

1. In a periodic table, the vertical columns and horizontal rows are respectively called has a. Groups, periods

b. Periods, groups

c. Groups, groups

d. Periods, periods.

c. Eka silicon

d. Eka mercury

c. Na < Li < K

d. Fe3+ < Fe 2+ < Fe 4+

2. The element Sc is known as a. Eka aluminium

b. Eka boron

3. The correct order of atomic radii is a. N < Be < B b. F− < O 2− < N 3−

4. According to the Periodic Law of elements, the variation in properties of elements is related to their a. Nuclear neutron-proton number ratios

b. Atomic masses

c. Nuclear masses

d. Atomic number

5. The metallic character of elements with an increase in atomic size. 176

IL Foundation Series Class 9

a. Increases

b. Decreases

c. Remains the same

d. Increases first, then decreases

6. The increasing order of the atomic of Si, S, Na, Mg and Al is a. S < Si < Al < Mg < Na

b. Na < Al < Mg < S < Si

c. Na < Mg < Si < Al < S

d. Na < Mg < Al < Si < S

7. The increasing order of atomic radii of the following Group 13 elements is a. Al < Ga < In < Tl

b. Ga < Al < In < Tl

c. Al < In < Ga < Tl

d. Al < Ga < Tl < In

8. Rare earths are generally a. Actinides

b. f-block elements

c. Inner transition elements

d. Lanthanides

9. Lanthanum element with z = 57 belongs to a. s-block

b. p-block

c. d-block

d. f-block

10. In the periodic table, transition elements begin with a. Scandium

b. Zinc

c. Copper

d. Mercury

11. Identify the correctly matched set among the following a. scandium, d-block-representative element b. anthanum, d-block-inner transition element c. cerium, f-block-transition element

d. actinium, d-block-transition element

IV. Electronic configurations and type of elements: s-,p-,d-,f-blocks

1. The general electronic configuration of d-block elements is a. ns1− 2 (n − 1)d 1−10

b. ns 2 (n − 1)d 1 (n − 2) f 1−14

c. ns1− 2 (n − 1)d 1−9

d. ns1− 2 np 6 (n − 1)d 1−10

2. Configuration that does not denote a transition element a. 3d14s2

b. 3d104s1

c. 3d104s24p2

d. 3d84s2

3. Inert gas element which has a different valence shell configuration a. Xe

b. Ne

c. Kr

d. He

c. 58 to 71

d. 90 to 103

4. Atomic numbers of actinides are a. 57 to 71

b. 80 to 103

5. The period in which s-block, p-block and d-block elements are present a. 1

b. 6

c. 7

d. 3

6. Elements of p-block are 177

PERIODIC CLASSIFICATION OF ELEMENTS

a. Only non-metals

b. Only metalloids

c. Metalloids and non-metals

d. Metalloids, non-metals and metals

7. Following are some statements about the modern periodic table i) It consists of s, p, d, and f blocks. ii) The energy levels filling order in 6th period is 6s, 4f, 5d, and 6p. iii) IIIA group contains the maximum number of elements. a. Only i & ii are correct

b. Only i is correct

c. Only ii & iii are correct

d. All are correct

8. The electronic configuration of an element is 1s22s2, and the position of that element in the periodic table is a. Second period, IIA Group

b. Second period, IA group

c. Third period, II A Group

d. First period, II A group

9. The general electronic configuration of VA group elements is a. ns 2

b. ns 2 np 5

c. ns 2 np1

d. ns 2 np 3

10. The outermost electronic configuration of the element is ns2np6. This element belongs to a. Alkali metals

b. Halogens

c. Carbon family

d. Noble gases

V. Diagonal relationship

1. A diagonal relationship is shown by a. All elements with their diagonally opposite elements b. All elements of 3rd and 4th periods c. Some of the elements of 2nd and 3rd periods d. Elements of d-block 2. Aluminium is diagonally related to a. Li

b. Si

c. Be

d. B

c. Li-Mg

d. B-Al

c. Be

d. B

c. S

d. Na

3. Which pair is different from the others? a. Na-K

b. Ca-Mg

4. Silicon is diagonally related to a. Li

b. Al

5. Li is diagonally related to a. Mg

178

b. Al

IL Foundation Series Class 9

WORKSHEET - 2 MULTIPLE CHOICE QUESTIONS WITH SINGLE CORRECT ANSWER 1. The correct arrangement of increasing atomic radius among Na, K, Mg, and Rb is a. Mg < K < Na < Rb

b. Mg < Na < K < Rb

c. Mg < Na < Rb < K

d. Na < K < Rb < Mg

2. Which of the following has the largest atomic radius? a. Al

b. Si

c. Cl d. Na

3. The representative elements get the nearest inert gas configuration by a. Losing electrons

b. Gaining electrons

c. Sharing electrons

d. Losing, gaining or sharing electrons

4. According to the Periodic Law of elements, the variation in properties of elements is related to their [AIEEE] a. Nuclear neutron-proton number ratios

b. Atomic masses

c. Nuclear masses

d. Atomic number

5. In transition elements, which are the shells that are incompletely filled? a. Ultimate shell only

b. Penultimate shell only

c. Both ultimate and penultimate shells

d. Outermost three shells [EAMCET 1993]

6. In Lanthanides, the differentiating electron enters the a. d-subshell

b. f-subshell

c. p-subshell

d. s-subshell [EAMCET 1993]

7. The non-transition metal among the following is a. Ag

b. Zn

c. Os

d. Pt [EAMCET 1996]

8. The 100th numbered element is named in honour of a. Einstein

b. Bohr

c. Fermi

d. Curie

9. The element which belongs to the 3rd period and IVA group of the periodic table is a. Silicon

b. Carbon

c. Germanium

d. Tin

10. The number of periods present in the long form of the periodic table is a. 6

b. 7

c. 8

[EAMCET 1999]

d. 18

11. A pair of elements with the following atomic numbers have the same chemical properties a. 13 & 22

b. 3 & 11

c. 4 & 24

d. 2&1

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PERIODIC CLASSIFICATION OF ELEMENTS

12. The element Z = 114 has been discovered recently. Which of the following family groups and electronic configurations will it belong to? [NEET 2017] a. Halogen family [Rn]5f146 d107 s27p5

b. Carbon family [Rn]5f146 d107 s27p2

c. Oxygen family [Rn]5f146 d107 s27p4

d. Nitrogen family [Rn]5f146 d107 s27p6

13. Element X forms a chloride with the formula XCl4, which is a solid with a low melting point. X would most likely be in the same group of the periodic table as a. Sodium (Na)

b. Magnesium (Mg)

c. Aluminium (Al)

d. Silicon (Si)

14. Which of the following pairs of atomic numbers represent s-block elements? a.7 and 15

b. 6 and 12

c. 9 and 17

d. 3 and 12

15. The most metallic element amongst Li, K, Mg, C, Al, and S is a. Al

b. K

c. Mg

d. S

16. An element X has 12 protons and 12 electrons, and another element Y has 12 protons and 10 electrons. Which of the following is the correct option regarding their atomic radii? a. Element X will have larger atomic radii b. Element Y will have smaller atomic radii c. The atomic radii of both elements will be almost similar d. Atomic radii of X and Y cannot be compared 17. What is the chemical formula of the carbonate compound of element X, where element X has a valency of hydrogen? a. XCO2

b. X2CO2

c. X2CO3

d. XCO3

18. Elements have been arranged in the following sequence according to Newland’s law of octaves, as shown below. Which of the elements given below will show similar properties? F, Na, Mg, Al, Si, P, S, Cl, K a. Na and Cl

b. F and K

c. Mg and Cl

d. Both Na and K, F and Cl

19. An element X belongs to the 3rd period and group 2. The nature of an element is a. Metalloids

b. Noble gas

c. Metal

d. Non-metal

20. Which of the following is the mass of the third element if element A has an atomic mass of 14 and middle element B has an atomic mass of 31 in Dobereiner’s triads? a. 74.90

b. 22.5

c. 48

d. 17

21. What will be the formula and the nature of bonding of its chloride of an element X placed in group 14? a. XCl2, ionic

180

b. XCl4, covalent

c. XCl4, metallic

d. XCl2, coordinate

IL Foundation Series Class 9

22. The electronic configuration of an element X is 2, 8, 8, 2. What are its period number, group number, and valency, respectively? a. 4, 2, and 2

b. 2, 2, and 4

c. 2, 4, and 2

d. 4, 2, and 4

23. Which among the following groups of elements can be called Dobereiner’s triads? Group A: Na, Si, Cl

Group B: Be, Mg, Ca

a. Group A only

b. Group B only

c. Both groups A and B

d. Neither group A nor group B

24. Which one of the following elements exhibits the maximum number of valence electrons? a. Al

b. Si

c. Na

d. P

25. Which of the following is the outermost shell for elements of period 2? a. N shell

b. L shell

c. M shell

d. K shell

26. An element which is an essential constituent of all organic compounds belongs to a. Group 16

b. Group 14

c. Group 1

d. Group 15

27. The elements A, B, C, D, and E have atomic numbers 9, 11, 17, 12, and 13, respectively. Which pair of elements belong to the same group? a. A and B

b. B and D

c. A and C

d. D and E

28. Which of the following statements about the modern periodic table is correct? a. It has 7 horizontal rows known as Groups. b. It has 18 horizontal rows known as Periods. c. It has 18 vertical columns known as Groups. d. It has 7 vertical columns known as Periods. 29. Which of the following statement(s) about the modern periodic table are incorrect? i) The elements in the modern periodic table are arranged on the basis of their decreasing atomic number. ii) The elements in the modern periodic table are arranged on the basis of their increasing atomic masses. iii) Isotopes are placed in adjoining group(s) in the periodic table. iv) The elements in the modern periodic table are arranged on the basis of their increasing atomic number. a. iv only

b. i, ii, and iii

c. i only

d. i, ii, and iv

181

PERIODIC CLASSIFICATION OF ELEMENTS

30. The electronic configuration of the atom of element X is 2,8,4. In the modern periodic table, the element X is placed in a. 2nd group

b. 4th group

c. 14th group

d. 8th group

31. Five elements A, B, C, D, and E have atomic numbers 2, 3, 7, 10, and 18, respectively. The elements which belong to the same period of the periodic table are a. A, B, C

b. B, C, D

c. A, D, E

d. B, D, E

32. Most of the non-metals are present in the long form of the periodic table in a.

p-block

b. f-block

c. d-block

d. s-block

33. A set of elements with the following atomic numbers belong to the same group a. 9, 16, 35, 3

b. 12, 20, 4, 38

c. 11, 19, 27, 5

d. 24, 47, 42, 55

c. Hg

d. Graphite

34. A non-metal having metallic properties is a. Al

b. Ag

35. Which of the following statements is correct? Statement 1: 4Be and 9F belong to the same period and, 9F is bigger in size. Statement 2: 4Be and 9F belong to the same period and, 4Be is bigger in size. Statement 3: 19K and 20Ca belong to the same period and, 19K is bigger in size. Statement 4: 19K and 20Ca belong to the same period and, 20Ca is bigger in size. a. Statement 2 is correct.

b. Statement 3 is correct.

c. Both statements, 2 and 3, are correct.

d. Both statements, 2 and 3, are incorrect.

36. Assertion (A): Rare earths are generally lanthanides. Reason (R): Lanthanides are transuranic elements. a. Both A and R are correct, and R is the correct explanation of A b. Both A and R are correct, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 37. Assertion (A): Chemistry of actinides is more complicated than lanthanides. Reason (R): Actinides elements are radioactive. a. Both A and R are correct, and R is the correct explanation of A b. Both A and R are correct, and R is not the correct explanation of A c. A is correct, and R is incorrect d. A is incorrect, and R is correct 182

IL Foundation Series Class 9

38. Assertion (A): 2nd and 3rd period elements have a diagonal relationship. Reason (R): The first element of a group closely resembles the second element of the next successive group. a. Both A and R are correct, and R is the correct explanation of A. b. Both A and R are correct, and R is not the correct explanation of A. c. A is correct, and R is incorrect. d. A is incorrect, but R is correct 39. Assertion (A): Caesium is highly reactive. Reason (R): Caesium is highly electropositive in nature. a. Both A and R are correct, and R is the correct explanation of A. b. Both A and R are correct, and R is not the correct explanation of A. c. A is correct, and R is incorrect d. A is incorrect, and R is correct 40. Which of the following pairs has both members from the same group of the periodic table? a. Mg-Ba

b. Mg-Na

c. Mg-Cu

d. Mg-Cl

183

ANSWER KEY 1: MATTER IN OUR SURROUNDINGS Worksheet 1 I. Introduction to matter in our surroundings 1. c 2. a 3. b 4. c 5. a 6. a 7. c II. States of matter & properties of solids, liquids, and gases 1. b 2. c 3. b 4. a 5. c 6. d 7. b 8. d 9. c 10. c 11. b 12. b 13. b III. Changes in states of matter 1. a 2. a 3. d 4. a 5. b 6. b 7. a 8. d 9. b 10. a 11. b 12. d 13. d 14. b 15. b 16. a 17. c 18. b 19. b 20. d Worksheet 2 1. a 2. d 3. b 4. b 5. a 6. c 7. c 8. b 9. c 10. a 11. b 12. b 13. d 14. d 15. b 16. a 17. a 18. a 19. b 20. a 21. b 22. b 23. a 24. c 25. d 26. b 27. a 28. a 29. a 30. a 31. d 32. c 33. a 34. a 35. a 36. a 37. a 2: IS MATTER AROUND US PURE? Worksheet 1 I. Pure substances 1. c 2. b 3. d 4. b 5. d 6. b II. Impure substances 1. b 2. c 3. a 4. b 5. a 6. b 7. d 8. a 9. a 10. b 11. c 12. a 13. d 14. d 15. a III. Solutions 1. a 2. d 3. a 4. c 5. c 6. a 7. c 8. c 9. a 10. a 11. a 12. b 13. c 14. c 15. c 16. b 17. d 18. b 19. d 20. d 21. c 22. c

184

Worksheet 2 1. d 2. a 3. c 6. d 7. c 8. d 11. b 12. d 13. a 16. d 17. a 18. c 21. b 22. c 23. d 26. d 27. a 28. b 31. a 32. a 33. a 36. b

4. a 9. a 14. c 19. d 24. a 29. a 34. d

5. a 10. b 15. b 20. a 25. c 30. c 35. d

3: ATOMS AND MOLECULES Worksheet 1 I. Introduction to atoms, molecules, and laws of chemical combination 1. b 2. c 3. a 4. d 5. a 6. b 7. c 8. c 9. b 10. c 11. d II. What is an atom? 1. d 2. b 3. a 4. c 5. d 6. a III. What is a molecule? 1. a 2. d 3. d 4. a 5. c 6. b 7. c 8. c 9. b IV. Writing chemical formulae, molecular mass, and mole concept 1. a 2. c 3. b 4. d 5. b 6. d V. 1. a 2. a 3. c 4. c 5. d 6. c 7. c 8. b VI. Limiting reagent 1. b 2. d 3. b 4. a 5. b Worksheet 2 1. a 2. c 3. b 6. d 7. c 8. a 11. b 12. b 13. a 16. a 17. a 18. a 21. c 22. b 23. c 26. a 27. a 28. a

4. c 9. a 14. c 19. a 24. b 29. b

5. d 10. b 15. a 20. d 25. d 30. b

ANSWER KEY 31. d 36. b 41. a 46. a

32. c 37. c 42. b 47. c

33. c 38. c 43. a 48. c

34. d 35. c 39. c 40. d 44. b 45. b

4: STRUCTURE OF THE ATOM Worksheet 1 I. Structure of the atom 1. a 2. a 3. a 4. a 5. a 6. b 7. d 8. d 9. d 10. c 11. a 12. c II. Electronic distribution, valency, configuration of elements 1. a 2. b 3. c 4. c 5. b 6. c 7. d III. Atomic number and mass number 1. b 2. c 3. b 4. c 5. d 6. c 7. d 8. b IV. Isotopes, isobars and isotones 1. a 2. c 3. c 4. c 5. c 6. a V. Bohr's theory of hydrogen atom, energy of the electron in the nth orbit (En ) & ionization energy 1. d 2. d 3. c 4. a 5. a 6. a 7. d 8. b 9. d VI. Quantum numbers 1. c 2. b 3. a 4. c 5. d 6. a 7. c 8. b 9. a VII. Shapes of orbitals and arrangement of electrons in orbitals 1. c 2. d 3. b 4. a 5. a 6. b Worksheet 2 1. b 2. a 3. a 6. b 7. d 8. b 11. b 12. a 13. a 16. a 17. c 18. b 21. b 22. a 23. d 26. a 27. b 28. d

4. a 9. d 14. a 19. d 24. a 29. a

31. b 36. c 41. b 46. a

32. a 37. b 42. a 47. a

33. c 34. a 35. b 38. d 39. b 40. c 43. a 44. b 45. c 48. c

5: PERIODIC CLASSIFICATION OF ELEMENTS Worksheet 1 I. Early attempts at the classification of the elements 1. d 2. a 3. a II. Mendeleev’s periodic table 1. b 2. d 3. a 4. c 5. c 6. d 7. c 8. c 9. d 10. d III. Modern periodic table 1. a 2. b 3. b 4. d 5. a 6. a 7. b 8. d 9. c 10. a 11. d IV. Electronic configurations and types of Elements: s-,p-,d-,f-blocks 1. a 2. c 3. d 4. d 5. b 6. d 7. a 8. a 9. d 10. d V. Diagonal relationship 1. c 2. c 3. c 4. d 5. a Worksheet 2 1. b 2. d 3. d 6. b 7. b 8. c 11. b 12. b 13. d 16. a 17. c 18. d 21. b 22. a 23. b 26. b 27. c 28. c 31. b 32. a 33. b 36. c 37. b 38. a

4. a 9. a 14. d 19. c 24. d 29. b 34. d 39. a

5. c 10. b 15. b 20. c 25. b 30. c 35. c 40. a

5. b 10. a 15. a 20. c 25. a 30. d

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