Page 1

Unit 6 Chemical Reactions Chapter 7


Objective 1 • Interpret chemical equations in terms of reactants, products, and conservation of mass.


Chemical Reactions • When a substance undergoes a chemical change, a chemical reaction is said to take place. • In order to understand chemical reactions, you first must be able to describe them.


Chemical Reactions • A useful description of a chemical reaction tells you the substances present before and after the reaction. • Reactants – the substances that undergo change in a chemical reaction. • Products - the new substances formed as a result of that change.


Chemical Reactions • During a chemical reaction, the reactants change into products. • You can summarize this process with a word equation. Reactants  Products Always on the left always on the right


Chemical Reactions

• Burning is an example of a chemical reaction.


Chemical Reactions • To describe the burning of charcoal, you can substitute the reactants and products of the reaction into the word equation that follows. Carbon + Oxygen  Carbon dioxide


Chemical Reactions • You can then simplify the word equation by writing the reactants and products as chemical formulas. C + O2  CO2


Chemical Reactions • A chemical equation is a representation of a chemical reaction in which the reactants and products are expressed as formulas. Carbon + oxygen → carbon dioxide C

+

O2

CO2


Conservation of Mass

• As charcoal burns, it gets smaller and smaller until it is finally reduced to a pile of ash. • Although the charcoal seems to disappear as it burns, it is actually being converted to carbon dioxide gas. • If you measured the mass of the carbon dioxide produced, it would equal the mass of the charcoal and oxygen that reacted.


Law Of Conservation of Mass • Antoine Lavoisier demonstrated that mass is conserved in a chemical reaction. • He formulated the Law of Conservation of Mass-mass is neither created or destroyed in a chemical reaction.


Law of Conservation of Mass • By demonstrating that mass is conserved in various reactions, Lavoisier laid the foundation for modern chemistry.

• The mass of the reactants will always equal the mass of the product s.


Objective 2 • Balance chemical equations by manipulating coefficients.


Balancing Chemical Equations • This picture shows workers preparing hydrazine to be loaded into spacecraft. • When hydrazine burns in the presence of oxygen, the reaction produces nitrogen, water vapor, and heat.


Unbalanced Equation Hydrazine + Oxygen → Nitrogen + water + heat N2H4 + O2 → N2 + H2O + heat • The number of atoms on the left side does not equal the number of atoms on the right. • The equation is NOT balanced.


Balancing Chemical Equations • In order to show that mass is conserved during a reaction, a chemical equation must be balanced. • You can balance a chemical equation by changing the coefficients, the numbers that appear before the formulas.


N2H4 + O2 → N2 + H2O + heat • In the unbalanced equation above, the coefficients are understood to be 1. • When you change a coefficient, you change the amount of that reactant or product represented in the chemical equation. • As you balance equations, you should never change the subscripts in a formula. • Changing the formula changes the identity of that reactant or product.


N2H4 + O2 → N2 + H2O + heat • The first step in balancing an equation is to count the number of atoms of each element on each side of the equation.

Reactants

Products

N

2

N

2

H

4

H

2

O

2

O

1


N2H4 + O2 → N2 +

H2O + heat

• The next step is to change one or more coefficients until the equation is balanced. • Try changing the coefficient of water to 2. Reactants

Products

N

2

N

2

H

4

H

2

O

2

O

1


N2H4 + O2 → N2 + 2H2O + heat • Multiply each subscript in water by 2. • Change the number of atoms of hydrogen and oxygen in the Products column. Reactants

Products

N

2

N

2

H

4

H

2 4

O

2

O

21

• The equation in now balanced.


N2H4 + O2 → N2 + 2 H2O + heat • Notice that the oxygen and nitrogen molecules are written as O2 and N2. • These are diatomic molecules. • You need to memorize the diatomic molecules. • They always occur as two atoms combined.


Diatomic Molecules • The diatomic molecules are oxygen, iodine, hydrogen, bromine, nitrogen, fluorine, and chlorine. • An easy way to remember: “Oh, I had better not fail chemistry!”

• O2 • I2 • H2 • Br2 • N2 • F2 • Cl2


Diatomic Molecules • Copy and memorize the next slide.


Molecule Oxygen

Formula O2

Saying Oh

Iodine

I2

I

Hydrogen

H2

Had

Bromine

Br2

Better

Nitrogen

N2

Not

Fluorine

F2

Fail

Chlorine

Cl2

Chemistry


Zn + HCl ďƒ¨ H2 + ZnCl2 Balance the above equation. Reactants

Products

Zn

1

Zn

1

H

1

H

2

Cl

1

Cl

2


Zn + HCl  H2 + ZnCl2 Reactants

Products

Zn

1

Zn

1

H

1

H

2

Cl

1

Cl

2


Zn + 2 HCl  H2 + ZnCl2 Reactants

Products

Zn

1

Zn

1

H

2

H

2

Cl

2

Cl

2


Fe + S  FeS Reactants

Products

Fe

1

Fe

1

S

1

S

1

This equation is balanced with no coefficients!


3. CH4 + O2 → CO2 + H2O

Reactants

Products

C

1

C

1

H

4

H

2

O

2

O

3


CH4 + 2 O2 → CO2 + 2 H2O

Reactants

Products

C

1

C

1

H

4

H

4

O

4

O

4


Your turn 4. KClO3 → KCl + 3O2 5. Al + HCl → AlCl3 + H2 6. Sb + O2 → Sb4O6 7. Al2O3 + C + Cl2 → AlCl3 + CO 8. Mg + HCl  H2 + MgCl2


Answers 4. 2KClO3 → 2KCl + 3O2 5. 2Al + 6HCl → 2AlCl3 + 3H2 6. 4Sb + 3O2 → Sb4O6 7. Al2O3 + 3C + 3Cl2 → 2AlCl3 + 3CO 8. Mg + 2HCl  H2 + MgCl2


CsNO3 + Mg(OH)2 CsOH + Mg(NO3)2 Cs

1

Cs

1

NO3

1

NO3

2

Mg

1

Mg

1

OH

2

OH

1


Balanced Equation • 2CsNO3 + Mg(OH)2 2CsOH + Mg(NO3)2


Balancing Chemical Equations • The following website gives you chemical equations to practice. http://funbasedlearning.com/ • Click on Classic Chembalancer or when you are ready click on Brain Boggle Chembalancer.


Warm-Up Question Balance the following: 1.

C2H4 + O2  CO2 + H2O

2.

Fe + Cl2  FeCl3


Key to Warm-Up Question Balance each of the following: 1.

C2H4 + 3O2 2 CO2 + 2 H2O

2.

2 Fe + 3 Cl2 2 FeCl3


Lab: Balancing Chemical Equations • What element is represented by the letter “H”?

5 H2

• What number represents the coefficient? • What number represents the subscript? • How many “Hs” do you have?


Example No. 1

Reactants

Products

H 2 + O2 → H 2 O Reactants

Products

H

2

H 2

O

2

O

1


Balanced Equation (write this on the right)

2 H2 + O2 → 2 H2O Reactants Final

Products Final H

4

H

4

O

O 2 2 Equation is balanced!


Example No. 2

Reactants

Products

H 2 O 2 → H 2 O + O2 Reactants H

2

O

2

Products H 2 O

3


Balanced Equation

2 H2O2 → 2 H2O + O2 Reactants

Products

H

4

H

4

O

4

O

4

Equation is balanced!


Objective 3 • Convert between moles and mass of a substance using molar mass.


Please get out the following: • Calculator • Periodic table • Pencil • Notes Packet


Counting Things • Chemists need practical units for counting things. • Although you can describe a reaction in terms of atoms and molecules, these units are too small to be practical. • Because chemical reactions often involve large numbers of small particles, chemists use a counting unit called a mole.


What is a Mole?

• A mole is an amount of a substance that contains approximately 6.02 x 1023 particles of that substance. • This number is known as Avogadro’s number. • In chemistry, a mole of a substance generally contains 6.02 x 1023 atoms, molecules, or ions of that substance. • For instance, a mole of iron is 6.02 x 1023 atoms of iron.


Molar Mass • A dozen eggs has a different mass than a dozen oranges. • Similarly, a mole of carbon has a different mass than a mole of sulfur. • The mass of one mole of a substance is called a molar mass. • For an element, the molar mass is the same as its atomic mass expressed in grams. • For example, the atomic mass of carbon is 12.011 amu, so the molar mass of carbon is 12.011 grams.


Atomic Mass- use PT

Element

Symbol

Carbon

C

Phosphorus

P

Sodium

Na

Atomic Mass, g/mol


Molar Mass

• For a compound, you can calculate the molar mass by adding up the atomic masses of its component atoms, and then expressing this sum in grams. • A carbon dioxide molecule is composed of one carbon atom (12.011 amu) + two oxygen atoms (2 x 16. ) = 44.01 grams.


Molar Mass • Round calculated molar masses to

• 0.01 g/mole


Molar Mass

Compound

Formula

Carbon dioxide

CO2

Ammonia

NH3

Aluminum chloride

AlCl3

Sulfuric Acid

H2SO4

Molar mass g/mol


Mole-Mass Conversions • Once you know the molar mass of a substance, you can convert moles of that substance into mass, or a mass of that substance into moles. • For either calculation, you need to express the molar mass as a conversion factor.


Sample Problem

1. Convert 55.0 g of CO2 to moles.

2. 1.25 moles of CO2 to grams.


Converting from Moles to Grams Moles X atomic mass = grams Moles X molar mass = grams


Converting from Moles to Grams-do together 3. Convert 10.2 moles of carbon to grams. 4. Convert 3.85 moles of CO to grams. 5. Convert 15.6 moles of NH3 to grams.


You try: 6. Convert 25.0 moles of P to grams. 7. Convert 3.22 moles of AlCl3 to grams.


Answers 6. 774 g 7. 429 g


Converting Grams to Moles • Mass divided by atomic mass = moles • Mass divided by molar mass

= moles


How many moles in the following?do together 8. 25.0 g of MgO 9. 75.0 g of silver 10. 110 . g of Ca(NO3)2


You try: 11. 225 g of sugar, C12H22O11 12. 45 g of Zn(C2H3O2)2


Mole Conversions • Label each of the following as atomic, molecular or a formula unit.

• CO • Ar • N2 • AlCl3 • PbBr2 • CH3COOH • C4H6 • Cu


Mole Diagram


Mole Map

Liters

Moles Atoms

Grams


Mole Conversions • Use your mole map to 1. Convert 3.6 moles of Ar to • Grams • Liters • atoms

2. Convert 5.2 X 1018 atoms of Ar to moles. 3. 125 g Ar to atoms


Mole Conversions 4. Convert 325 g of C4H6 to molecules. How many carbon atoms?


Mole Conversions 5. 7.2 X 1024 formula units of AlCl3 is how many moles? 6. 215 g of AlCl3 is how many formula units?


Objective 4 Calculate amounts of reactants or products by using molar mass, mole ratios, and balanced chemical equations.


Mole/Mass ratios a. Mg + S MgS This means one mole of magnesium reacts with one mole sulfur to produce one mole of magnesium sulfide. Let’s see how the mass adds up. Mg

+

S

MgS

6.02 X 1023 atoms + 6.02 X 1023 atoms 6.02 X 1023 formula units 24.32 g

+ 32.066 g

56.386 g


Mole/Mass Ratio Fe + Cl2 FeCl2

b.

This means one mole of iron reacts with one mole of chlorine to produce one mole of iron (II) chloride.

Fe

+

Cl2

FeCl2

6.02 X 1023 atoms + 6.02 X 1023 molecules  6.02 X 1023 formula units

55.85 g

+

70.91 g

126.76 g


Mole/Mass Ratio c. S + O2  SO2 This means one mole of sulfur reacts with one mole of oxygen to produce one mole of sulfur dioxide. S

+

O2

SO2

6.02 X 1023 atoms + 6.02 X 1023 molecules  6.02 X 1023 molecules

32.066 g

+

32 g

64.066g


Mole/Mass Ratio • You can see each equation supports the Law of Conservation of Mass. • Use atoms when an element is listed. • Use molecules when two or more nonmetals are combined.


Mole/Mass Ratio • Let’s see when we would use ions. • Let’s dissolve the chemical magnesium sulfate in water. • The solid would break apart into the ions that formed it. • The ions would then be floating in solution.


Mole/Mass Ratio d. MgSO4(s)  Mg+2(aq) + SO4-2(aq) One mole of magnesium sulfate produces one mole of magnesium ions and one mole of sulfate ions.

MgSO4(s)

Mg+2(aq)

+ SO4-2(aq)

6.02 X 1023 formula units  6.02 X 1023 ions + 6.02 X 1023 ions 120.39 g

 24.32 g

+

96.07 g


Summary • A mole is 6.02 x 1023 particles. • The particles can be • Atoms – when the particle is an element – except • Molecules – when the particle is a diatomic element (O2, I2, H2, Br2, N2, F2, Cl2)


Summary, continued • The particles can be • Molecules – when the compound is covalent – usually two nonmetals • Formula units – when the compound is ionic – when a metal and nonmetal (or polyatomic ion) combine • Ions – when ions are represented –

Fe+2, K+, (PO4)-3, etc.


Mole Ratios & Chemical Equations • Balance the following equation C2H6 + O2  CO2 + H2O


Mole Ratios & Chemical Equations 2C2H6 + 7O2 ďƒ 4CO2 + 6H2O 2 moles of C2H6 would produce how many moles of CO2? how many moles of water?


Mole Ratios & Chemical Equations 2C2H6 + 7O2 ďƒ 4CO2 + 6H2O a. 3.8 moles of C2H6 will produce how many moles of water? b. 7.9 g of oxygen would produce how many liters of carbon dioxide?


Lab: What is a Mole? • Part One: Counting Particles • Teacher Demo/Need 2-3 volunteers • 1 mole = 6.02 x 1023 atoms • I will obtain 1 mole of carbon atoms • Volunteers will count out 1 mole of rice grains

• Discuss • Efficiency of using molar mass • Mass, mole, and atoms of C in briquette


Lab: What is a Mole? • Part Two: Modeling a Mole • Measure mass of bolt, nut, washer • Molecule of BN2W2 • Predict mass of BN2W2 • Find mass • Construct three different models


• Homework • Finish lab • Complete Review Sheet, Part I • Read pages 199-204. Study for QUIZ.


Today • • • • • •

Turn in RS.1 and Lab What is a Mole? Complete warm up questions Quiz: Balancing Equations Notes: Objective 5 – Types of Chemical Reactions Lab Double Replacement Homework • Finish lab • Complete WS II


Warm-Up Question 1. One mole of which of the following would contain 6.02 x 1023 atoms? A.

Fluorine

B.

Carbon

C.

CH4

D.

NaCl


Warm-Up Question 2. One mole of which of the following would contain 6.02 x 1023 molecules? A.

Magnesium

B.

Oxygen

C.

CuCN

D.

Fe2(PO4)3


Warm-Up Question 3. One mole of which of the following would contain 6.02 x 1023 formula units? A.

Carbon dioxide

B.

Sodium chloride

C.

Hydrogen

D.

Diphosphorus pentoxide


Warm-Up Question 4. One mole of which of the following would contain 6.02 x 1023 ions? A.

Cl2

B.

Na2SO4

C.

Ag+

D.

CH4


Warm-Up Question Determine the formula mass of the following: 1. Ag2SO4 2. CsC2H3O2 3. Mg(OH)2 4. (NH4)2C4H4O6 5. Aspirin's chemical formula is C9H8O4. What is its formula mass?


Key to Warm-UP 1.

Ag2SO4 – 311.83 u

2.

CsC2H3O2 – 191.96u

3.

Mg(OH)2 – 58.34 u

4.

(NH4)2C4H4O6 – 184.16 u

5.

C9H8O4 – 180.16 u


• Use the factor-label method to work each problem below. 1. What is the mass in grams of 0.452 mole of Ag2SO4? 2. Determine the number of magnesium ions in 3.05 mole of Mg(OH)2. 3. How many molecules are in 39.0 grams of C9H8O4? 4. Calculate the mass in grams of 2.23 moles of (NH4)2C4H4O6. 5. How many atoms are in 6.4 grams of copper?


Objective 5 •

Classify chemical reactions as synthesis, decomposition, single replacement, double replacement, or combustion reactions.


Synthesis Reaction • Synthesis Reaction - two or more substances react to form a single substance. • The reactants may be either elements or compounds. • The product synthesized is always a compound. • The general equation is: A + B  AB • 2Na + Cl2  2NaCl • Zn + S  ZnS


Synthesis Reaction

Zn + S  ZnS


Decomposition Reaction • Decomposition reaction - a compound breaks down into two or more simpler substances. • It is the opposite of a synthesis reaction. • The reactant must be a compound. • The products may be elements or compounds. • The general equation is: AB  A + B • CaCO3  CaO + CO2 • 2NH3  3H2 + N2 • H2CO3  H2O + CO2


Mercury (II) oxide

• Decomposes with heat into oxygen and mercury


Single Replacement Reactions •

Single Replacement - is a reaction in which one element takes the place of another in a compound.

General Form: A + BC B + AC

ZnSO4

Al + 3CuCl2  3Cu + 2AlCl3

K + 2HOH  H2 + 2KOH

+

Mg

 MgSO4 + Zn


Most Reactive potassium calcium

Activity Series

sodium magnesium aluminum zinc iron tin lead hydrogen copper mercury silver gold Least Reactive

•The table to the left lists metals in order of reactivity, or tendency to react. •A metal on the table will usually displace any below it on the table.


Single-Replacement Reactions • • 1. 2. 3. 4. 5.

Use the activity series to predict which of these single-replacement reactions can and cannot occur. Complete and balance the reactions that can occur. Zn + CaCl2  Na + Fe2O3  Sn + Al2O3  Mg + HCl  AgNO3 + Cu 


Single Replacement- Halogen • Usually a metal is replacing a metal Zn CaCl2 • Or a metal is replacing hydrogen Mg

+ +

HCl

• A special case is when a halogen replaces a halogen.


Single Replacement- Halogen • Use an activity series. • Complete the following: • F2 + NaI • Cl2 + KBr • Br2 + CaF2 


Double Replacement Reactions • Double Replacement - is one in which two different compounds exchange positive ions and form two new compounds. • General Form: AB + CD AD + CB • Notice that two replacements take place. • A is replacing C; D is replacing B.


Double Replacement • Complete and balance the following reactions: a. Pb(NO3)2 + KI  b. AgNO3 + NaCl  c. FeS + HCl 


Double Replacement d. CaCO3 + 2HCl  CaCl2 + H2CO3 • The H2CO3 that forms breaks down into water and carbon dioxide. So this equation becomes: • CaCO3 + 2HCl  CaCl2 + H2O +

CO2


• Lab Double Replacement • Six solutions will be reacted together and carefully observed. • Two unknowns will be reacted with the six solutions. • Write double replacement reactions and balance.


Today • Turn in Lab Double Replacement & WS II • Quiz: Moles • Lab Single Replacement • Homework • Finish lab • Complete WS III • Read pages 206-209


Today

• Notes: Objectives 5-6 • Combustion Reactions and Activity Series


Warm-Up Question •

Classify each of the following reactions by type:

1.

S8 + 8 O2 → 8 SO2 + heat

2.

Pb(NO3)2 + K2CrO4 → PbCrO4 + 2 KNO3

3.

2 NaHCO3 → Na2CO3 + H2O + CO2

4.

Zn + 2 HCl → ZnCl2 + H2


Warm-Up Question • Hydrogen peroxide, H2O2, decomposes to produce water and oxygen gas. The balanced equation for this reaction is ___. • 2 H2O2 → 2 H2O + O2


Combustion Reactions • Combustion- is one in which a substance reacts rapidly with oxygen, often producing heat and light. • The products of the reaction are carbon dioxide and water.


Combustion a. The combustion of an alcohol, ethanol. Write the reactants on the the left, the products on the right. C2H5OH + O2 → CO2 + Balance the equation.

H2O


Answer

a. C2H5OH +

3O2 ягн 2CO2 +

3H2O


Complete the following b. The combustion of propane gas, C3H8. c. The combustion of butane gas, C4H10.


Objective 7 • Describe the energy changes that take place during chemical reactions.


Energy Changes • The heat produced by a propane grill is a form of energy. • When you write the chemical equation for the combustion of propane, you include “heat” on the right side of the equation. • C3H8 + 5 O2 3 CO2 + 4 H2O + Heat • This equation states that the heat released in the reaction came from the reactants.


Energy Changes • Chemical energy - is the energy stored in the chemical bonds of a substance. • The chemical energy of propane is the energy stored in the bonds between the C and H of this molecule. • Likewise, oxygen, carbon dioxide, and water molecules all have energy stored in their chemical bonds.


Energy Changes • Energy changes in all chemical reactions are determined by changes that occur in chemical bonding. • Chemical reactions involve the breaking of chemical bonds in the reactants and the formation of new chemical bonds in the products. • In the combustion of propane, the bonds in propane and oxygen molecules are broken, while the bonds in carbon dioxide and water molecules are formed.


Breaking Bonds • Each propane molecule reacts with five oxygen molecules. • In order for the reaction to occur, eight C – H single bonds, two C – C single bonds, and five O  O double bonds must be broken. • Breaking chemical bonds requires energy. • This is why propane grills have an igniter, a device that produces a spark. • The spark provides enough energy to break the bonds of reacting molecules and get the reaction started.


Forming Bonds

• For each mole of propane burned, three moles of carbon dioxide and four moles of water are formed. • This means six C  O double bonds and eight O – H single bonds are formed in the reaction. • The formation of chemical bonds releases energy. • The heat and light given off by a propane stove result from the formation of new chemical bonds. • The bonds form as the carbon, hydrogen, and oxygen atoms in the propane and oxygen molecules are rearranged into molecules of carbon dioxide and water.


Objective 8 • Classify chemical reactions as exothermic or endothermic.


Exothermic Reactions • An exothermic reaction releases energy to its surroundings. • The energy released by the bonds of the reactants is divided. Some goes into the bonds of the products and the excess is released. • Combustion is an example of an extremely exothermic reaction. • When 1 mole of propane reacts with 5 moles of oxygen, 2220 kJ of heat is released.


Endothermic Reactions • An endothermic reaction absorbs energy from its surroundings. • The products end up with more energy than the reactants. • The decomposition of mercury (II) oxide into its elements is endothermic.


2 HgO + 181.7 kJ -> 2 Hg + O2 • When mercury (II) oxide is heated to a temperature of about 450°C, it breaks down into mercury and oxygen. • The decomposition of mercury (II) oxide is an endothermic reaction that can be described by the equation above. • Because heat is absorbed, the energy term appears on the left side of the equation. • For every 2 moles of HgO that decomposes, 181.7 kJ of heat must be absorbed.


Exothermic or Endothermic? a. When solid KBr is dissolved in water, the solution gets colder. b. Natural gas is burned in a furnace. c. When concentrated sulfuric acid is added to water, the solution gets very hot. d. Water is boiled in a teakettle.


Objective 9 • Explain how energy is conserved during chemical reactions.


Conservation of Energy • In an exothermic reaction, the chemical energy of the reactants is converted into heat plus the chemical energy of the products. • In an endothermic reaction, heat plus the chemical energy of the reactants is converted into the chemical energy of the products.


Law of Conservation of Energy

• In both cases, the total amount of the energy before the reaction is the same. • This principle is known as the Law of Conservation of Energy - energy cannot be created or destroyed.


Objective 10 • Explain what a reaction rate is.


Reaction Rate • Any change that happens over a period of time can be expressed as a rate. • For example, speed is the rate that distance changes over time. • Reaction rate - is the rate at which reactants change into products over time.


Reaction Rates • Reaction rates tell you how fast a reaction is going. • That is, • how fast the reactants are being consumed, • how fast the products are being formed, or • how fast energy is being absorbed or released.


Objective 11 • Describe the factors affecting chemical reaction rates.


Chemical Reactants • Recall that chemical reactions involve collisions between particles of reactants. • The reaction rate depends on how often these particles collide. • If the collisions occur more frequently, then the reaction rate increases. • If the collisions occur less frequently, then the reaction rate decreases. • Almost any reaction rate can be changed by varying the conditions under which the reaction takes place.


Factors that affect Reaction Rate •

Factors that affect reaction rate include 

temperature

surface area

stirring

concentration

catalysts


Temperature

• Increasing the temperature of a substance causes its particles to move faster, on average. • Particles that move faster are both more likely to collide and more likely to react. • If the number of collisions that produce reactions increases, then the reaction rate increases.


Surface Area • The smaller the particle size of a given mass, the larger its surface area. • An increase in surface area increases the exposure, the more collisions there are that involve reacting particles. • With more collisions, more particles will react. • This is why increasing the surface area of a reactant tends to increase the reaction rate.


Stirring • You can increase the exposure of reactants to each other by stirring them. • For example, when you wash your clothes in a washing machine, particles of detergent react with particles of the stains on your clothes. • This reaction would go slowly if you just left your clothes soaking in a tub of water and detergent.


Stirring, continued • A washing machine speeds up the reaction by stirring the contents back and forth. • Collisions between particles of the reactants are more likely to happen. • Stirring the reactants will generally increase the reaction rate.


Concentration • Concentration refers to the number of particles in a given volume. • The more reacting particles that are present in a given volume, the more opportunities there are for collisions involving those particles. • The reaction rate is faster.


Concentration, continued • Increasing the partial pressure of a gas, increases the rate of reaction because that increases the concentration of the gas.


Catalyst • A catalyst is a substance that affects the reaction rate without being used up in the reaction. • Chemists often use catalysts to speed up a reaction or enable a reaction to occur at a lower temperature. • Because a catalyst is not consumed, it can be used to speed up the same reaction over and over again.


Catalysts, continued • Recall that in order for a reaction to take place, the reacting particles must collide with enough energy to break the chemical bonds of those particles. • A catalyst lowers this energy barrier. • One way that a catalyst can do this is by providing a surface on which the reacting particles can come together. • This causes the reacting particles to be more likely to react.


Demo Rates of Reaction • The factors of temperature, surface area, concentration and catalyst will be investigated.

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