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Chapter 5 PHYSICAL AND CHEMICAL CHANGES. 5.1 Physical and Chemical Changes. Changes that occur in matter are classified as: a) physical change b) chemical change Physical change. • is a change that alters the form or appearance of material that does not convert the material into new substances. • the chemical composition of the material produced remains the same. • a physical change usually involves little or no change in energy. • examples of physical changes are : a) melting ice b) evaporation of water c) solubility of salt in water Chemical change. • is a change in matter that forms one or more new substances. • the chemical properties and composition of the new substance are different from those of the original. • the chemical change that occurs usually involves absorption or release of heat. (sometimes light energy is also produced) • the new substances formed usually cannot be changed back to the original material physically. • examples of chemical changes are : a) burning paper b) change in colour of sliced apples c) rusting of iron 5.2 Heat Changes in Chemical Reactions. • heat change occurs in most of the chemical reactions. • all occurring chemical reactions involve energy transformation. • when a chemical reaction occurs, heat energy is absorbed or released. • heat energy is absorbed to break the bond in a compound. • conversely, heat energy is released when the bond is formed. • chemical reaction can be classified into two type : • exothermic reactions • endothermic reactions


Comparison between physical and chemical changes. Similarities - both experience physical changes like appearances. Physical changes No Same Normally yes Usually requires a little energy Same

Differences Formation of new substance Properties of composition of reactants and products Reversible change Energy requirement Mass of reactants and products

Chemical changes Yes Different Normally no Usually requires a lot energy Different

Comparison between Endothermic and Exothermic reactions. Definition Reactant Energy content

Exothermic reaction Reactions in which heat energy is released to the surroundings. When this reaction occurs, the reactant will lose heat to the surroundings. The total energy content of the product is less than the total energy content of the reactant. The energy transfer can be shown in an energy level diagram.

Endothermic reaction Reactions in which heat energy is absorbed from the surroundings. When this reaction occurs, the reactant will gain heat from the surroundings. The total energy content of the product is more than the total energy content of the reactant. The energy transfer can be shown in an energy level diagram.

Reactant

Product Product

Surrounding temperature Contents in container

The surrounding temperature is raised. The contents in the container become hot.

Reactant

The surrounding temperature is lowered. The contents of the container become cool.

Chemical reactions in industry. • usually occur under optimum conditions so that the time of reaction is very short and the cost involved is minimal. • two important chemical process in industry are : a) Haber process which produces ammonia. b) Contact process which produces sulphuric acid • in the Haber process, ammonia is produced from the mixture of nitrogen and hydrogen. Nitrogen + Hydrogen •

Haber process is a reversible reaction.

Ammonia


• • •

when ammonia is produced, bond formation occurs between the atoms of nitrogen and hydrogen, a lot of heat is released to the surroundings. in the contact process, sulphuric acid is produced through three stages. in the first stage, sulphur is burnt in the air to produce sulphur dioxide gas. Sulphur + Oxygen

• •

Sulphur Oxide

heat is released to form the bond when sulphur oxide is produced. in the second stage, sulphur dioxide gas reacts with oxygen to form sulphur trioxide gas at temperature of 400-500 oC and pressure of 1 atmosphere. Vanadium (V) oxide is used as a catalyst in this reaction. Sulphur Dioxide + Oxygen

• • • •

Sulphur Trioxide

this reaction is reversible. heat is released to form the bond when sulphur trioxide is produce. at the third stage, the reaction which occurs involves two steps. the reaction at both steps releases heat to form the chemical bonds when oleum and sulphuric acid are produce. Sulphur Trioxide + Concentrated Sulphuric Acid Oleum + Water Sulphuric Acid

Oleum

just like at the second stage, optimum conditions are maintained at this stage. Haber Process Nitrogen

Hydrogen

Mixture of nitrogen and hydrogen is compressed

Ammonia gas

temperature : 450oC pressure : 200 atm catalyst : iron filings

condensed

Ammonia (liquid)

Contact Process Sulphur combustion

Sulphur Dioxide

Sulphur temperature : 400- 500oC Trioxide pressure : 1 atm catalyst : vanadium(V) oxide dissolved in concentrated sulphuric acid

Oxygen Sulphuric acid

diluted with water

Oleum Liquid


5.3 Reactivity Series of Metals. Reactivity of Metals with Water. • some metals react with water more vigorously than others. • metals like sodium and potassium react very vigorously with water. • metals like calcium react less vigorously with water as compared to sodium and potassium. • reactivity of metals with water can be represented by the following equations : Reactive metal + water metal hydroxide + hydrogen ex : Sodium + water sodium hydroxide + hydrogen •

metals like magnesium, aluminium, zinc, and iron react less vigorously with water. (this metals only react vigorously with water) Metal + steam ex : Magnesium + steam

metal oxide + hydrogen magnesium oxide + hydrogen

metals like lead, copper, silver, and gold do not react with water and steam.

Reactivity of Metals with Acid. • some metals react with dilute acid to produce salt and release hydrogen gas. Metal + dilute acid salt + hydrogen ex : Magnesium + dilute hydrochloric acid magnesium chloride + hydrogen • •

reactive metals like magnesium, aluminium, zinc, and iron react vigorously with dilute acid. metals like copper silver and mercury do not react with dilute acid because these metals are not reactive.

Reactivity of Metals with Oxygen. • when a metal is heated in oxygen, it combines with oxygen to form metal oxide. Metal + oxygen ex : Magnesium + oxygen • •

metal oxide magnesium oxide

metals burn in oxygen with different reactivity. when a metal is heated, its reactivity can be determined through the brightness of the flame that is produced.


Reactive of metal can be determined by the brightness of its flame. Metal Flame Very reactive Ignite brightly Reactive Glows brightly Less reactive Glows dimly Comparison of reactivity of metals with water, dilute acid and oxygen. Metal Reactivity of metals when reacting with Water Dilute acid Oxygen Potassium React with cold water Sodium Calcium React with dilute Magnesium acid React with steam Aluminium Zinc Reactivity of metals Iron in decreasing order Tin Lead React with hot Do not react with dilute acid cold water or steam Copper Do not react with dilute acid Silver Gold Reactivity Series of Metals. • reactivity series of metals is a series that shows the order of metal reactivity. • the series is formed based on metals reactivity with oxygen. • metals that react vigorously with oxygen are placed at the top of the series. • metals that react less vigorously are placed at the bottom of the series. Reactivity series of metals Potassium Sodium Calcium Magnesium Aluminium Zinc Iron Tin Lead Copper Silver Gold

reactivity of metals in decreasing order


Position of carbon in reactivity series of metals. • although carbon is a non-metallic element, it reacts with excess oxygen to form carbon dioxide. • if carbon is more reactive than metal X, a bright flame or glow will be seen when a mixture of carbon and oxide of metal X is heated. Oxide of metal X + carbon

metal X + carbon dioxide heated

for example, carbon can eliminate zinc from zinc oxide. therefore, the position of carbon is higher than that of zinc in the reactivity series of metals. Zinc oxide + carbon

zinc + carbon dioxide

if carbon is less reactive than metal Y, a flame or glow is not seen when a mixture of carbon dioxide of metal Y is heated. Carbon + oxide of metal Y

• •

for example, no reaction takes place when a mixture of carbon and aluminium oxide is heated. therefore, the position of carbon is lower than that of aluminium in the reactivity series of metals. Aluminum oxide + carbon

no reaction

no reaction

by conducting reactions between carbon and oxides of metals, the position of carbon in the reactivity series of metals can be determined. Potassium Sodium Magnesium Aluminium Carbon

reactivity of metals in decreasing order

Zinc Iron Tin Lead Copper Silver Gold position of carbon in the reactivity series of metals.


5.4 Application of The Concepts of Reactivity Series of Metals. Relationship between the position of metals in the reactivity series and the method of metal extraction. • knowledge of the position of metals in the reactivity series of metal can be applied in the method of extracting metals from their ores. • there are two methods of extracting metal : (a) reduction by carbon (b) electrolysis on smelting metal ore • carbon is used in the extraction process because (a) it is cheap (b) easily obtained (c) the side product of carbon (carbon dioxide) during the extraction process is a type of gas which is easily eliminated. Extracting metal from metal ore. • in nature, metal tends to react with oxygen to form metal oxide. (ore) • metal in the form of ore does not have much use and needs to be extracted. • the method of extracting metal depends on the position of the metal in the reactivity series of metals. Extraction of tin. • tin ore (cassiterite) is tin oxide which exists in the earth’s crust. • tin oxide is washed with water to remove dirt. • then, tin oxide is roasted to remove impurities like sulphur and oil. • after that, tin oxide is extracted by heating tin oxide with carbon and limestone in a high temperature blast furnace. • the function of the limestone is to remove impurities. • during heating, carbon which is more reactive than tin removes oxygen from the tin oxide to produce pure tin and carbon dioxide. Tin oxide + carbon (tin ore) • •

tin + carbon dioxide

molten tin is poured into moulds to form tin ingots. at the same time, the limestone )calcium carbonate) decomposes into quicklime (calcium oxide) which reacts with impurities to form slag, i.e. dirt which is unwanted.

Importance of reactivity. • the importance of reactivity series of metals : i. Reactivity series of metals enables the reactivity of metals to be compared. Metals with higher positions in the series are more reactive than those below them.


ii. The series is used to determine whether a reaction can occur. For example, sodium has a higher position than iron in the series. This means that sodium is more reactive than iron. Therefore, sodium can remove oxygen from iron oxide. iii. Knowledge of reactivity series of metals can be applied in choosing the method of metal extraction from its ore. 5.5 Electrolysis. • is a dissociation process of chemical substances in aqueous solution or molten state to its constituents by using electricity. • a dry cell or battery supplies electricity to dissociate chemical substances to their constituents. • electrical energy changes to chemical energy in electrolysis. electrical energy  chemical energy dry cell/battery

ammeter

rheostat

A +

-

anode

cathode

cathode

+

+ -

cation

-

+

anion

arrangement of apparatus for electrolysis process

electrolyte


Electrolysis. • is usually used in electrolysis because carbon is inert and does not take part in reaction. • during electrolysis, cation moves towards the cathode while anion moves towards the anode. • at the cathode, cation receives electron from the cathode and is discharged to form a neutral atom. positive ion + electron  neutral atom •

at the anode, anion releases electron and is discharged to form a neutral atom. negative ion  neutral atom + electron

• •

discharge is a charge neutralisation process in ions to form neutral atoms. electrolysis of copper(II) chloride solution. o copper(II) ion with positive charge will attract to cathode to discharge as a copper metal. o chloride ion will attract to anode to discharge as a chlorine gas. o at anode, chloride ions lose of electrons, greenish-yellow bubbles of chlorine gas are released. o at the cathode, copper(II) ion receives electron, brown copper metal is deposited. o the blue colour of copper(II) chloride solution fades. copper(II) chloride

copper + chlorine electrolysis

Electrolysis of molten lead(II) bromide. • lead(II) bromide in a crucible is heated until it melts. • two carbon electrodes are put in the molten lead(II) bromide. • at the anode, brown vapour is released, i.e. bromine vapour. • at the cathode, shiny grey solid is produced, i.e. lead. • this is because lead(II) ion and bromide ion move freely when lead(II) bromide is melted. • during electrolysis, lead(II) ion which is positively charged moves towards the cathode to receive electrons and form lead. lead(II) ion + electrons •

lead

bromide ion which is negatively charged moves towards the anode to release electron and form bromine vapour.


bromine ion •

bromine + electron

thus, electrolysis of lead(II) bromide produces lead and bromine gas.

Uses of electrolysis in industry. Electrolysis is widely used in industry for the following purposes : a) extraction of metals b) purification of metals c) electroplating of metals Extraction of metals. • metals that are more reactive than carbon are extracted from their ores by electrolysis. • extraction of aluminium from bauxite : i. molten aluminium oxide (bauxite) and carbon electrodes are used in the extraction of aluminium. ii. the steps are as follows : 1. Aluminium oxide with cryolite is heated until it melts. The function of cryolite is to lower the melting point of aluminium oxide. 2. Aluminium oxide dissociates into aluminium ions (cation) and oxide ions (anion). iii. at the cathode, aluminium ions receive electrons and are discharged. molten aluminium forms and settles at the base of the electrolytic cell. aluminium ions + electrons  aluminium ions iv. at the anode, oxide ions release electrons and are discharged. oxygen atoms are formed. the combination of two oxygen atoms forms an oxygen molecule. thus, oxygen gas is released. oxide ion  oxygen atom + electrons Purification of metals. • metals can be purified through electrolysis. • in this process, the impure metal becomes the anode while the pure metal becomes the cathode. • electrolyte is a salt solution of that respective metal.


purification of copper :

A +

-

impure copper plate

pure copper plate

copper(II) sulphate solution purification of metal through electrolysis. i. anode is an impure copper plate while cathode is a pure copper plate. copper(II) sulphate solution is used as electrolyte. ii. at the anode, the impure copper plate will dissolve to form copper(II) ion. copper atom  copper(II) ion + electrons a. impurities will be left at the base of the beaker when the impure copper plate dissolves. b. the impure copper plate will become thinner after a while. iii. at the cathode, copper(II) ion will move towards the cathode to receive electrons and is discharged. copper metal is formed. copper(II) ion + electrons  copper atom eventually, the cathode will become thicker because pure copper sediment will settle on it. thus, the copper is purified. Electroplating of metals. • in electroplating process,  the metal used for electroplating becomes the anode.  the object to be plated becomes the cathode.  the electrolyte is a salt solution of that metal. • during electrolysis, the anode dissolves to form metallic ions.


• • •

these ions then move towards the cathode and settle as a thin layer of metal. thus, the metallic object (cathode) is coated with a thin layer of metal from anode. electroplating iron nail with copper :

A +

-

copper plate

iron nail

copper(II) sulphate solution electroplating iron nail with cooper i. the surface of the iron nail and the copper metal are rubbed with a sandpaper. ii. the iron nail and the copper metal are immersed in copper(II) sulphate solution as shown above. (the circuit is completed) iii. at the anode, the copper metal becomes thinner, this is because copper atoms at the anode release electrons to form copper cation. copper atom  copper(II) ion + electrons iv. at the cathode, the surface of the iron nail is coated with a brown copper layer. this is because copper ions in the solution move towards the cathode to receive electrons and are discharged. copper metal is formed. copper(II) ion + electrons  copper atom v. the copper metal formed settles on the surface of the iron nail. •

the aims of electroplating are to 1. prevent the metal from corrosion or rusting 2. make the metal look more attractive


5.6 Production of Electrical Energy From Chemical Reactions. Production of Electrical Energy by simple voltaic cell. • electrical energy can be produced from chemical reactions. • simple voltaic cell consists of two different metals, or one of it is carbon, that is immersed into a electrolyte. • chemical changes will occur to produce electrical energy. • in a simple voltaic cell, the energy transformation which occurs is as follow: chemical energy  electrical energy galvanometer is used to detect the production of electric current. can be replaced by a voltmeter or ammeter.

G Iron plate can be replaced by other conductors like lead, zinc, and carbon

+

iron plate dilute sulphuric acid

Various types of cells and their uses. Dry cell

-

copper plate

dilute sulphuric acid can be replaces by other electrolytes like sodium chloride solution and dilute nitric acid.


• • • • • • •

are the most commonly used electrochemical cells. the zinc casing is the negative terminal. the carbon rod is the positive terminal. the carbon rod is coated with a mixture of carbon powder and manganese(IV) oxide. the carbon powder reduces resistance in the cell. the manganese(IV) oxide absorbs the hydrogen gas released during reaction. the electrolyte is paste of ammonium chloride mixed with zinc chloride.

Lead-acid accumulator

• • • • • • •

the car battery is a type of electrochemical cell called accumulator. the lead-acid accumulator used in cars consists of six cells connected in series. this type of battery supplies 12 volts of electrical energy. the lead electrode is the negative terminal. the lead electrode coated with lead(IV) oxide is the positive terminal. the electrolyte is concentrated sulphuric acid. the accumulator is a type of secondary cell which can be recharged to be used repeatedly.

Alkaline battery


• • • • •

an alkaline battery is similar to a dry cell but it uses a different electrolyte and is lasts long. the zinc casing is the negative terminal. the manganese(IV) oxide is the positive terminal. the electrolyte is potassium hydroxide solution. an alkaline battery is used in watches, torches, radios, electric shavers and toys.

Silver oxide-zinc cell

• • • • •

the shape of this type of battery is like a button. the zinc casing is the negative terminal. the silver oxide is the positive terminal. the electrolyte is potassium hydroxide. this battery is used in watches and electronic toys.

Nickel-cadmium battery • this battery operates on the same principle as the lead-acid accumulator, but it uses different chemical substances. • the cadmium is the negative terminal. • the nickel(IV) oxide is the positive terminal. • the electrolyte is potassium hydroxide. Advantages and Disadvantages of electrochemical cell Type of electrochemical cell Dry cell Lead-acid accumulator Alkaline battery

Advantages -light and can be easily carried along -supplies constant current -rechargeable -supplies high voltage for a long period -long-lasting -supplies higher current

Disadvantage -not long lasting -not rechargeable -heavy and expensive -electrolyte which corrodes spills over easily -not rechargeable -more expensive than


Silver oxide-zinc cell Nickel-cadmium battery

than dry cell although the voltage is same -long-lasting -supplies constant current -long-lasting -rechargeable -concentration of its electrolyte does not change

ordinary dry cell -not rechargeable -expensive

5.7 Chemical Reactions that Occur In The Presence of Light. -

some chemical reactions occur in the presence of light. examples of such reactions are : a. photosynthesis in green plants b. decomposition of certain chemical substances in photography

Photosynthesis. • in photosynthesis, green plants absorb sunlight and convert it into chemical energy (glucose). • water and carbon dioxide are used in photosynthesis to produce glucose and oxygen is released. sunlight

water + carbon dioxide

glucose + oxygen chlorophyll

• • •

the light energy absorbed by the chlorophyll in green plants is used to break water molecules into hydrogen and oxygen. (this process is called photolysis) the hydrogen then combines with carbon and a part of the oxygen in carbon dioxide to produce glucose. the energy transformation that takes place during photosynthesis is : light energy  chemical energy

Effects of light on photosensitive chemicals. • photographic paper is covered with a thin layer of silver bromide or sliver chloride. • when a photographic paper is exposed to light, light energy decomposes the silver bromide to silver atoms. (silver is a dark grey substance) silver bromide

silver + bromine light

by the darker part of the photograph is caused by the formation of the silver atoms.


Storing chemical substances. • chemical substances like chlorine water, sodium hypochlorite solution and silver salt are very sensitive to light. • these chemical substances will decompose to other substance if exposed to sunlight. • as a result, photosensitive chemicals must be stored in dark condition. • chlorine water and sodium hypochlorite solution must be stored in dark bottles. • photographic paper is also stored is a black bag or a black box.


Chapter 5