6$03/(Saartjie Spiesdocscientia.co.za/images/stories/pdf/Gr11_Chemistry_WB...2.3 Application of exothermic and endothermic reactions 180 Experiment 12 182 Experiment 13 184 2.4 Bond

Physical SciencesChemistry Grade 11

Textbook and Workbook

Elma de Vos • Santie du Plessis Carlien Fanoy • Elize Jones

Patricia Lees-Rolfe • Judy McDougall Janetta Nel • Karen Reynecke

Saartjie SpiesSAMPLE

Doc ScientiaPosbus 7011

Ansfrere 1711

www.docscientia.co.za

For any queries and feedback: [email protected]

Jacques Fanoy or Stephan FanoyOffice: 011 472 8728

Fax: 086 546 1423

ISBN: 978-1-920537-17-3

First edition December 2009Revised edition December 2010; 2011

Second edition December 2012Revised edition December 2013; 2014; 2015

Graphic design: Helene Jonck

All rights reserved. No part of this publication may be reproduced in any form or by any means –

mechanical or electronic, including recordings or tape recordings and photocopying – without the prior permission of the publisher.SAMPLE

Unit Page

KNOWLEDGE AREA MATTER AND MATERIAL 13

Unit 1 ATOMIC BONDS: MOLECULAR STRUCTURE 13

1.1 Chemical bonds 14

1.2 Bonding models 14

1.3 Valence electrons 16

1.4 Valency 16

1.5 Lewis structure 16

Activity 1 17

1.6 Formation of bonds 17

1.7 Single covalent bonds 18

1.8 Molecules with multiple bonds 19

1.9 Dative covalent bond 22

Exercise 1 23

1.10 Valence Shell Electron Pair Repulsion theory (VSEPR) and molecular shapes 24

Activity 2 25

Practical activity 1 28

Exercise 2 29

1.11 Electronegativity 30

Activity 3 34

Exercise 3 36

1.12 Forces between atoms 37

1.13 Electrical potential energy changes in the formation of a molecule 38

1.14 Bond energy and length 39

Exercise 4 41

Summary of Unit 1 43

Mind maps of Unit 1 45

Unit 2 INTERMOLECULAR FORCES (BETWEEN MOLECULES) 49

2.1 Types of intermolecular forces 50

2.1.1 Ion-dipole forces 51

2.1.2 Ion-induced dipole forces 52

2.1.3 Dipole-dipole forces 53

2.1.4 Dipole-induced dipole forces 54

2.1.5 Induced dipole force (dispersion or London forces) 54

2.1.6 Hydrogen bonds 55

2.2 Difference between intermolecular forces and interatomic forces 56

Exercise 5 57

2.3 The influence of intermolecular forces 59

2.3.1 State of molecules 59

2.3.2 Size of molecules 59

2.3.3 Density 60

2.3.4 Boiling and melting points 60

2.3.5 Viscosity 61

Activity 4 61

2.3.6 Thermal expansion 64

INDEX

SAMPLE

2.3.7 Thermal conductors 64

Experiment 1 64

Experiment 2 66

Experiment 3 68

Experiment 4 69

Experiment 5 70

Exercise 6 71

2.4 Chemistry of water 73

2.4.1 Microscopic structure of water 73

Activity 5 75

2.4.2 Properties of water 75

Experiment 6 77

Experiment 7 78

Experiment 8 79

Exercise 7 80



Unit 3 IDEAL GASES AND THERMAL PROPERTIES 89

3.1 Movement of particles 89

3.2 Average kinetic energy of gas molecules and temperature 90

3.3 The ideal gas model 90

Exercise 8 92

3.4 The relationship between volume and pressure of a gas 95

Experiment 9 95

Exercise 9 100

3.5 Relationship between volume and temperature of a gas 103

Experiment 10 103

3.6 Relationship between temperature and pressure of a gas (constant volume of gas) 106

3.7 Kelvin temperature scale 107

Exercise 10 109

3.8 Relationship between pressure, volume and temperature of a gas

110

Exercise 11 111

3.9 The ideal gas equation 113

Exercise 12 114



Question paper 119

KNOWLEDGE AREA CHEMICAL CHANGE 125

Unit 1 QUANTITATIVE ASPECTS OF CHEMICAL CHANGE: STOICHIOMETRY 125

1.1 Mole concept 126

1.2 Molecular atomic mass (A), molar molecular mass and molar formula mass (M) 127

Exercise 13 128

1.3 Relationship between moles, mass and molar mass 130

Exercise 14 131

1.4 Molar gas volume 133

1.4.1 Avogadro’s law 133

SAMPLE

1.4.2 Molar volume in balanced chemical equations 134

Exercise 15 135

1.5 Concentration of solutions 136

1.5.1 Calculating the concentration of a solution 136

1.5.2 Preparation of a standard solution 139

Practical demonstration 139

1.5.3 Diluting solutions 140

Exercise 16 141

1.6 Percentage composition of compounds 143

1.6.1 Calculations 143

1.6.2 Determination by practical quantitative methods 145

Exercise 17 147

1.7 Empirical formula 148

Exercise 18 149

1.8 Molecular formula 150

Exercise 19 151

1.9 Stoichiometric calculations 152

1.9.1 Calculations based on balanced equations 152

1.9.2 Limiting reactants 158

1.9.3 Mass of precipitate formed during precipitation reactions 161

1.9.4 Percentage yield 161

1.10 Applications of exothermic reactions that produce large amounts of gas 162

Experiment 11 163

Exercise 20 167



Unit 2 ENERGY AND CHEMICAL CHANGE 177

2.1 Introduction 177

2.2 Breaking and forming of bonds during a chemical reaction 178

2.3 Application of exothermic and endothermic reactions 180

Experiment 12 182

Experiment 13 184

2.4 Bond energy 186

Exercise 21 187

2.5 Activation energy and activated complex 190

2.5.1 Energy profiles for exothermic and endothermic reactions 192

2.5.2 Catalysts 192

Experiment 14 194

Exercise 22 195



Unit 3 ACID-BASE REACTIONS 201

3.1 Acids 201

3.2 Bases 203

3.3 Strong and weak acids and bases 204

3.4 Concentrated and dilute acids and bases 205

3.5 Acid-base reactions 205

Exercise 23 207

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3.6 Ampholytes 210

3.7 Reactions of acids 210

3.7.1 Acid-metal carbonate reactions 211

3.7.2 Acid-metal oxide reactions 211

3.7.3 Acid-metal hydroxide reactions 211

3.7.4 Acid-base reactions: applications 214

3.8 Indicators 214

Practical activity 2 215

Experiment 15 215

Experiment 16 216

Experiment 17 218

Exercise 24 219



Unit 4 REDOX REACTIONS 225

4.1 Oxidation numbers 226

4.2 Reactions 228

4.2.1 Displacement reactions 228

Experiment 18 228

4.2.2 Synthesis reactions 230

Experiment 19 230

4.2.3 Decomposition reactions 233

Experiment 20 233

4.2.4 Reducing operation 234

Experiment 21 234

4.2.5 Oxidising effect 236

Experiment 22 236

Exercise 25 237



Question paper 243

KNOWLEDGE AREA CHEMICAL SYSTEMS 251

Unit 1 THE DEVELOPMENT OF MINING 251

1 Introduction 251

1.1 People through the ages 251

1.2 Mineral sources in South Africa 253

Exercise 26 256

Summary: of Unit 1 257

Unit 2 MINING AND MINERAL PROCESSING 258

A GOLD 259

2.1 Gold in South Africa 259

2.1.1 Reasons for mining gold 259

2.1.2 Locations 260

2.1.3 Gold mining and processing 260

2.1.4 Gold mining and the environment 261

Exercise 27 262

2.1.5 Ancient mining methods 265

Case study: Mapungubwe: South Africa’s lost city of gold 265

B IRON 268

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2.2 Iron in South Africa 268

2.2.1 Reasons for mining iron 268

2.2.2 Why South Africa exports iron ore 269

2.2.3 Locations 269

2.2.4 Iron ore mining and processing 269

Experiment 23 271

Experiment 24 272

2.2.5 Iron mining and the environment 273

2.2.6 Processing of iron 274

Experiment 25 274

Exercise 28 276

C PHOSPHATES 279

2.3 Phosphates in South Africa 279

2.3.1 Reasons for mining phosphates 279

2.3.2 The most important mining activities in South Africa

279

2.3.3 Mining and processing of phosphate rock 279

2.3.4 Uses 281

2.3.5 The impact of phosphate mining on the environment

281

Exercise 29 282

Project 283



Unit 3 GLOBAL WARMING 293

3.1 Increasing human population and global warming 293

3.2 Burning of fossil fuels 294

3.3 What is global warming? 295

3.4 Consequences of global warming 295

Exercise 30 295



Question paper 298

Information sheets 309

Work cited 311

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Doc Scientia CHEMISTRY textbook and workbook - Grade 11 177

KNOWLEDGE AREA: CHEMICAL CHANGE

UNIT 2 ENERGY AND CHEMICAL CHANGE

2.1 IntroductionThermodynamics is the study of energy or heat transfer during chemical reactions.When you eat an orange, the sugar in the orange reacts with the oxygen in your body to form CO2 and H2O. During this chemical process, energy is released that is used by your body to make your muscles work. Some reactions release energy and others use energy.

When a mixture of methane gas and air is ignited, enough heat is released to cook a piece of beef on a gas grill:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + ENERGY

The formation of glucose, C2H12O6, by the process of photosynthesis, requires the absorption of light energy:

6CO2(g) + 6H2O(ℓ) + ENERGY → C6H12O6(s) + 6O2(g)

Chemical reactions are classified as exothermic or endothermic on the grounds of energy changes that occur during the reaction.

Breaking and forming of bonds

Application of exothermic and endothermic reactions

Bond energy

Energy profiles

Catalysts

Exothermic reaction: A reaction that releases energy/heat – heat flows from the system to the surroundings.

A test tube in which an exothermic reaction occurs feels warm, since heat is released by the reaction. Combustion reactions are examples of exothermic reactions.The tragic explosion of the spacecraft Challenger was the result of an uncontrolled exothermic reaction between hydrogen and oxygen. Hydrogen, stored as a liquid, is the fuel used to drive spaceships out of the earth’s gravitational field.

Energy and chemical change

Activation energy and activated complex

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178 Doc Scientia CHEMISTRY textbook and workbook - Grade 11

Endothermic reaction: A reaction that absorbs energy. Heat flows from the surroundings into a system.

A beaker in which KNO3(s) is dissolved in water will feel cold, as the reaction absorbs heat from the surroundings. Another example of an endothermic reaction is the melting of ice. (This is not a chemical reaction.)

Enthalpy: The total amount of chemical potential energy in a chemical system.

Because the energy during a chemical reaction is either absorbed or released, the enthalpy of the system changes during the reaction. This module focuses on the change in enthalpy or the heat of reaction, in other words, on the energy transfer during a chemical reaction.

Enthalpy is represented by the symbol H and the change in enthalpy (heat of reaction) for a reaction system is indicated by ΔH:

ΔH = Eproducts - Ereactants

If the products in a reaction system have a greater enthalpy than the reactants, then ΔH will be positive. The system therefore absorbs heat and the reaction is endothermic.

Endothermic reaction: ΔH > 0

If the enthalpy of the products in a reaction system is less than that of the reactants, then ΔH will be negative. The system releases heat and the reaction is exothermic.

Exothermic reaction: ΔH < 0

2.2 Breaking and forming of bonds during a chemical reactionChemical bonds form as a result of the forces of attraction between the atoms or ions in a compound. During chemical reactions, existing bonds in reactants are broken, and new bonds are formed to create products. To break bonds, the attractive forces between particles must be overcome, and energy is absorbed. The energy absorbed can be in the form of heat, light or electrical energy. When new bonds are formed, energy is released.

Quick factsThe word enthalpy comes from the Greek word “enthalpein”, which means “to heat”.

SAMPLE

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Bond energy: The energy absorbed when bonds break, or released when new bonds form.

A study of bond energies helps us to understand, at microscopic level, why some reactions are exothermic and others are endothermic.

Energy change = Σenergyabsorbed (breaking bonds) - Σenergyreleased (forming bonds)

If the energy that is absorbed to break the existing bonds is less than the energy released when new bonds form, the reaction will be exothermic: ΔH < 0.

ExamplesThe single bonds between the H atoms in each of the two H2 molecules, as well as the double bond between the two O atoms in the O2 molecule, break and energy is absorbed:

Total energy absorbed (breaking bonds) = 2(436) + 498 = 1 370 kJ

2H2 + O2 → 2H2Obonds involved

and another

436 kJ

436 kJ

494 kJ

breaking

breaking

breaking

H – H

H – H

O = O

H

H

H

H

O

O

Quick factsΣ means “the sum of” and is read as ‘sigma’.

SAMPLE

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The loose atoms are now called an activated complex, and move around so that collisions occur between them. Some of the collisions are effective and lead to bond formation. Effective collisions occur when the colliding atoms are correctly orientated and when they have more than enough energy to overcome the repulsive forces between them.

Total energy released(forming bonds) = 2 × 2(463) = 1 852 kJΔH = ΣEabsorbed + (-ΣEreleased) = 1 366 + (-1 852) = -486 kJ for 2 mol of water = -243 kJ·mol-1 for 1 mol of water

If the energy that is absorbed to break the existing bonds, is more than the energy released when new bonds form, then the reaction is endothermic: ΔH > 0

ExamplesStudy the energy change in the reaction below:

2HI(g) → H2(g) + I2(g)H─I bonds need 299 kJ·mol-1 energy to break. According to the balanced equation there are 2 HI molecules.Total energy absorbed(breaking bonds) = 2 × 299 = 598 kJ·mol-1

One H─H bond forms: 436 kJ·mol-1 is released.One I─I bond forms: 149 kJ·mol-1 is released.Total energy released(forming bonds) = 436 + 149 = 585 kJ·mol-1

ΔH = ΣEabsorbed + (-ΣEreleased) = 598 + (-585) = +13 kJ·mol-1

ΔH is positive, the reaction is endothermic.

2.3 Application of exothermic and endothermic reactions

• Combustion Fuel such as petrol for the propulsion of carsCharcoal for power generationCharcoal, wood or gas (mainly butane) for cooking and heating Natural gas (mainly methane, CH4) for industrial use

Exothermic reaction

463 kJ (Energy is released when a new bond forms.)

463 kJ (Energy is released when a new bond forms.)

During the bond formation step, energy is released:

SAMPLE

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Interesting factsRespiration is the reverse process of photosynthesis.Photosynthesis Respiration

• Occurs only in plants.• Oxygen and sugar are formed during the process.• Light is necessary for photosynthesis.• Food is produced.• Energy is stored.• Only takes place in cells that contain chlorophyll.

• Occurs in animals and plants.• Water and carbon dioxide are formed

during the process.• Light is not necessary for respiration.• Food is digested.• Energy is released.• Takes place in all cells.

• Neutralisationacid + base → salt + water + energy

• Hydration Reaction of some anhydrous salts with water to form hydrated salts, e.g. CuSO4(s) + 5 H2O(ℓ) → CuSO4·5H2O(s) + energy

• Respiration Oxidation of glucose to form CO2, water and energy in plant and animal cells, provides a supply of energy for growth and bodily functions.

Endothermic reactions:• Photosynthesis

Production of sugar (glucose) and oxygen from CO2(g) by plants in the presence of light.6CO2(g) + 6H2O(ℓ) + light energy → C6H12O6(aq) + 6O2(g)

• Electrolysis Electroplating and removal of metals from ores.

• Thermal decompositionMany compounds decompose in the presence of heat energy.

CaCO3(s) + heat → CaO(s) + CO2(g) • Reaction between steam and coal to prepare H2(g).

C(s) + H2O(g) → CO(g) + H2(g) • Ice packs used for sports injuries.SAMPLE

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The enthalpy changes for exothermic and endothermic reactions can be represented as follows:

Experiment 12 Date:Aim: Investigate the reaction of a number of elements and compounds with water, and classify the reactions as exothermic or endothermic.

Hypothesis:

________________________________________________________________________________

________________________________________________________________________________

Variables:

Independent variable(Which is changed.)

Dependent variable(Which is measured.)

Controlled variable(s)(Which remain(s) the same.)

Exothermic reactions Endothermic reactions

ΔH < 0

E absorbed during breaking of bonds

E released during bond formation

Eactivated complex Eactivated complex

E absorbed during breaking of bonds

E released during bond formation

ΔH > 0

EreactantsEproducts

Ereactants Eproducts

SAMPLE

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Apparatus: Chemicals:• Test tube rack ● NH4NO3(s)• Twelve test tubes ● KNO3(s)• Thermometer ● MgSO4(s)• Spatula ● KBr(s)• Distilled water ● NH4OH(s)• Stirring rod ● BaCℓ2(s) ● Vinegar (aq) ● Na2CO3(s) ● NaHCO3(s) ● Na2S2O3(s) ● Citric acid ● Cal-C-Vita tablet Method:1. Half-fill twelve test tubes with distilled water.2. Measure the temperature of the water.3. Add a spatula tip of a salt to the water in the test tube and stir to dissolve the salt.4. Feel the outside of the test tube and measure the temperature of the solution.5. Add another spatula tip of the salt to the water and stir again to dissolve the salt.6. Measure the temperature of the solution again.7. Make a note of the observations.8. Repeat steps 2 to 7 for each of the other salts.

Safety measures: Do not handle the chemicals with your hands. Be especially careful with the NaOH and the H2SO4. Concentrated acids must always be added to water and not the other way round. Always use SMALL amounts. Do not inhale any fumes. Use a fume cupboard or a well-ventilated work space.

Observations:

Reactants Ti Tf ΔT ObservationExothermic or

endothermic reaction

NH4NO3 + H2O

KNO3 + H2O

MgSO4 + H2O

KBr + H2O

NH4OH + H2OSAMPLE

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BaCℓ2 + H2O

Vinegar + H2O

Na2CO3 + H2O

NaHCO3 + H2O

Na2S2O3 + H2O

Citric acid + H2O

Cal-C-Vita + H2O

NaOH + H2O

Conclusions:What can be deduced from the temperature change during the different reactions?

_______________________________________________________________________________

_______________________________________________________________________________

Experiment 13 Date:Aim: Investigate a number of chemical reactions and classify the reactions as exothermic or endothermic.

Hypothesis:

_______________________________________________________________________________

_______________________________________________________________________________

Variables:

Independent variable(Which is changed.)

Dependent variable(Which is measured.)

Controlled variable(s)(Which remain(s) the same.)SA

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Apparatus: Chemicals:• Test tube rack ● CaCℓ2(s)• Eight test tube ● CuSO4·5H2O(s)• Thermometer ● Li• Spatula ● Mg ribbon and Mg powder• Distilled water ● H2SO4(d)• Stirring rod ● NH4Cℓ • Glass beaker ● Glycerine • Glass bowl

Safety measures: Some of the reactions are fierce and large amounts of heat are produced. These experiments should only be done by the teacher as a demonstration, and the safety of the learners should be considered at all times.

Method:1. Half-fill eight test tubes with distilled water.2. Measure the temperature of the water.3. Add a spatula tip of CaCℓ2(s) to the water in a test tube and stir to dissolve the salt.4. Feel the outside of the test tube and measure the temperature of the solution.5. Add another spatula tip of CaCℓ2(s) and stir to dissolve the salt.6. Measure the temperature of the solution.7. Write down your observations.8. Repeat steps 2 to 7 with CuSO4·5H2O(s), Li, Mg ribbon, Mg powder, H2SO4 and CaCℓ2(s).9. Half-fill the glass bowl with water.10. Measure the temperature of the water.11. Add a small piece of lithium (approximately the size of a pea) to the water.12. Observe what happens, and take the temperature of the water again.

Observations:

Reactants Ti Tf ΔT ObservationExothermic or

endothermic reaction

CuSO4·5H2O + H2O

Mg ribbon + H2O

Mg powder + H2O

H2SO4(d) + H2O

CaCℓ2 + H2O

Li + H2O

Glycerine + H2O

NH4Cℓ + H2OSAMPLE

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Conclusions:What can be deduced from the temperature changes during the different reactions?

________________________________________________________________________________

________________________________________________________________________________

Interesting factsGlycerine (glycerol) is very hygroscopic.Which means it absorbs water from the air.It is often used for hair products.

2.4 Bond energy

When bonds break, bond energy is absorbed, and it therefore has a positive value. When bonds form, the energy value is negative, as energy is released during bond formation. The table below shows the estimated bond energy values of specific compounds, in kJ·mol-1.

Bond Bond energy (kJ·mol-1) Bond Bond energy

(kJ·mol-1)H–H 436 C=C 619

H–O 463 Cℓ–Cℓ 243

H–C 413 I–I 149

H–N 389 Br–Br 190

H–F 569 O=O 494

H–S 338 C–O 358

H–Cℓ 431 C–C 348

H–I 299 C=O 799

H–Br 366 C≡C 839F–F 159 N≡N 946

Bond energy: The energy that is absorbed to break an existing bond, or the energy that is released when a bond forms.

SAMPLE

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ExamplesUse the table of bond energies to calculate ΔH for the reaction below and say if the reaction is exothermic or endothermic.

2HF(g) + Cℓ2(g) → 2HCℓ(g) + F2(g)

SolutionEnergy to break bonds:Cℓ–Cℓ: 243

H–F: 569

H–F: 569

+1 381 kJ

Energy to form bonds:

H–Cℓ: 431

H–Cℓ: 431

F–F: 159

-1 021 kJ

ΔH = ΣEabsorbed + ΣEreleased = (1 381) + (-1 021) = +360 kJ

ΔH > 0 The reaction is endothermic.

Exercise 21 Date:1 Classify the following reactions as exothermic or endothermic.

1.1 Fe + S → FeS ΔH negative ____________________

1.2 2KNO3 → 2KNO3 + O2 ΔH positive ____________________

1.3 Mg + I2 → MgI2 ΔH < 0 ____________________

1.4 2MgO → 2Mg + O2 ΔH > 0 ____________________

1.5 2HI → H2 + I2 ΔH = +40 kJ·mol-1 ____________________

1.6 CO + NO2 → CO2 + NO ΔH = -226 kJ·mol-1 ____________________SAMPLE

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2 Determine with the aid of the table of bond energies, what the net energy change will be for the following reactions. Also indicate whether the reactions are exothermic or endothermic.2.1 C(s) + H2O(g) → CO(g) + H2(g)

_____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________ _____________________________________________________________________________

_____________________________________________________________________________

2.2 N2 + 3H2 → 2NH3

_____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________ _____________________________________________________________________________

_____________________________________________________________________________

2.3 Ethanol forms ethene and water. It is represented by the following reaction:

C2H5OH C2H4 + H2O

_____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________ _____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________

H HH H

H O HC CC OC +

H HH H

H

H

SAMPLE

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3. The reaction between two gases can be represented as follows:

X2 + 3Y2 2XY3

Use the information in the table below to calculate the heat of reaction (ΔH) for this reaction.

Symbolic representation Diagrammatic representation

Energy to break one bond. Energy unit

X2 → X + X

5

Y2 → Y + Y 2

XY3 → XY2 + Y 3

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_____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________

_____________________________________________________________________________

4. The reaction between two substances can be represented as follows:

CH4 + 2O2 CO2 + 2H2O

O O+ +

H H

H H

H

H C H

H

O

OO

O

OC OSAMPLE

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Use the table of bond energies to calculate the heat of reaction (ΔH). Also indicate if the reaction is exothermic or endothermic.

______________________________________________________________________________

______________________________________________________________________________

______________________________________________________________________________

______________________________________________________________________________

______________________________________________________________________________

______________________________________________________________________________

______________________________________________________________________________

______________________________________________________________________________

Interesting factsSprains and injuries are often treated with ice packs.An ice pack usually consists of:- a liquid (water).- solids (ammonium nitrate and ammonium chloride).When an ice pack is used, the seal between the substances is broken. The result is an endothermic reaction. The ice pack turns cold.

2.5 Activation energy and activated complex

In 1888 the Swedish chemist, Svante Arrhenius, suggested that molecules must have a certain minimum amount of kinetic energy before they can react with each other to form a new bond.

Study the following diagram:Initially the ball is at position A. At position B the ball will be at a lower (more stable) energy state.

To get the ball from position A to position B though, it needs to have enough energy to get over the hump, which is blocking its way. The ball must be moved over the energy barrier before it can come to rest at the lower-energy position B – it needs “activation energy”.

A BSAMPLE

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For a reaction to occur between stable molecules, a certain minimum amount of energy must be absorbed to break the bonds keeping the particles in the reactant compounds together. This amount of energy is necessary to separate atoms or ions in the reactants so that they can be rearranged to form new bonds, and thereby form the products.

The energy of the activated complex is the barrier that prevents some reactions from occurring. The reactant particles must therefore gain activation energy to overcome the energy barrier between the reactant state and the final product state.

It is therefore the energy that is absorbed by the reactant:• to break existing interatomic bonds.• to change the particles into an activated complex.

Fast reactions have low activation energies and slow reactions require more activation energy.

Stable molecules, before converting to products, must go through an unstable, high-energy intermediate phase. This specific rearrangement of atoms, which has maximum chemical potential energy, is referred to as the activated complex.

The steps according to which the reaction occurs, are called the reaction mechanism.

We can visualise the reaction mechanism of CO with NO2 as follows:

CO + NO2 CO2 + NO reactants activated complex products

According to this representation, we can interpret the activation energy as the difference between the energy of the activated complex and the energy of the reactant molecules.

Spontaneous reactions occur when the reactants have enough energy for the reaction to start by itself spontaneously.

The activation energy is supplied by the reactants themselves. Non-spontaneous reactions occur when the reactants themselves do not have enough energy to start the reaction, and additional energy is needed in the form of light or heat energy.

Activation energy (EA) is the minimum energy needed to allow a chemical reaction to occur.

The activated complex is a temporary, high energy, unstable transition state between reactants and products.

activated complex

reactantsCO + NO2

[C; 3 O; N]

activation energy

productsO=C=O + N=OSA

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2.5.1 Energy profiles for exothermic and endothermic reactions

There are different ways of writing an equation to show whether it is exothermic or endothermic.Exothermic Endothermic

2.5.2 Catalysts

A catalyst can be explained with the following analogy:Welcome to the chemistry class. I am the teacher and my role is that of a catalyst – to increase your understanding without being used up in the process!

Reactants → products + energyReactants → products ΔH < 0Reactants → products ΔH = -20 kJ·mol-1Reactants → products + 20 kJ·mol-1

activated complexEP

reactants

course of reactionproducts

EA

ΔHenthalpy change

Enthalpy change: ΔH = Hproducts - Hreactants

ΔH < 0 (negative)

Reactants + energy → productsReactants → products ΔH > 0Reactants → products ΔH = 20 kJ·mol-1Reactants +20 kJ·mol-1 → products

activated complex

EP

reactants

course of reaction

productsEA

ΔHenthalpy change

Enthalpy change: ΔH = Hproducts - Hreactants

ΔH > 0 (positive)

Catalyst: A substance that changes/increases the rate of a reaction without itself undergoing a permanent change or being used up.

Quick factsSpontaneous reactions take place without the addition of external energy or stimulus, e.g. phosphorus spontaneously starts burning in air.

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The effectiveness of a catalyst in increasing the rate of a reaction, is explained in terms of the change in the reaction system’s activation energy. If the activation energy for a reaction is low, then a greater fraction of the reactant molecules will have greater energy than the activation energy, and the reaction will take place quickly. When the activation energy is high, then only a few of the reactant molecules have enough energy to change to an activated complex, and the reaction will be slow.

A positive catalyst:• lowers the activation energy of a reaction.• does not permanently change during the reaction.• does not change the energy of the reactants or the products, therefore ΔH stays the same.

Heterogeneous catalysts are not in the same state as the reactants. A solid catalyst, like platinum (Pt) for example, is used during the oxidation of ammonia (NH3) to increase the reaction rate:

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(ℓ)

The reactant is in the gaseous state but the catalyst is a solid. The catalyst temporarily bonds with the reactant molecules, and in so doing weakens the existing interatomic forces. Less energy is therefore needed to break the bonds in the reactant molecules. The activation energy decreases and the reaction rate increases.

Homogeneous catalysts are in the same state as the reactants.Exhaust fumes from motor cars include, amongst others, the poisonous gas carbon monoxide (CO). Modern cars have catalytic converters in their exhaust systems, which usually contain platinum (Pt) and rhodium (Rh) as catalysts. The CO(g) and NO(g) in the gas mixture are then quickly converted to harmless CO2(g) and N2(g):

2CO(g) + 2NO(g) → 2CO2(g) + N2 (g)

Pt

ΔH

enthalpy change

activation energy

EP

reactants

course of reaction

activated complex

EA

with catalyst

products

SAMPLE

chemical change


Experiment 14 Date:Aim: To investigate the concept of activation energy, by making use of magnesium ribbon that burns in air and oxygen.

Hypothesis:

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Apparatus:• Two deflagration spoons• Magnesium ribbon• Two glass cylinders• Oxygen• Two glass covers• A stopwatch• Bunsen burner

Method:1. Fill one of the glass cylinders with oxygen and cover the opening with a glass cover.2. Fill the other cylinder with air and cover the opening with the other glass cover.3. Scrape two pieces of magnesium ribbon, equal in length, until clean.4. Roll both pieces of ribbon and place on separate deflagraton spoons.5. Set one on fire and lower it into the cylinder with air.6. Use the stopwatch to measure the time it takes to burn out from the moment it was lowered into

the cylinder.7. Carefully observe the brightness and the rate of combustion.8. Repeat by lowering the other deflagrating spoon with burning magnesium ribbon into the

cylinder with oxygen.

Observations:

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_______________________________________________________________________________Results:Use the same axes and draw a sketch graph for the potential energy versus time.

Conclusions:

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chemical change


Exercise 22 Date:1. Draw an energy diagram for each of an exothermic and an endothermic reaction and indicate the

following on the diagram: 1. Reactants 2. Products 3. Activation energy 4. ΔH

2 Are the following reactions exothermic or endothermic?

2.1 S + O2 → SO2 (ΔH is negative.) ____________________

2.2 A + B + energy → AB ____________________

2.3 Food is digested. ____________________

2.4 NH4Cℓ(s) → NH4+(aq) + Cℓ-(aq) (ΔH > 0) ____________________

2.5 A2(g) + B2 → 2AB(g) (ΔH = + 80 kJ·mol-1) ____________________

2.6 CO2 + H2O + energy → glucose and oxygen ____________________

2.7 A2 + B → A2B (ΔH > 0) ____________________

2.8 Acid is diluted in water. ____________________

3 The energy diagram below indicates the energy change during a chemical reaction.

3.1 Name the energies 1 – 3.

1 _______________________________

2 _______________________________

3 _______________________________

course of reaction

pote

ntia

l ene

rgy

2

3

1SAMPLE

Change in energy Endothermic reaction:

Exothermic reaction Catalyst

ENERGY AND CHEMICAL CHANGE

SAMPLE

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6$03/(Saartjie Spiesdocscientia.co.za/images/stories/pdf/Gr11_Chemistry_WB...2.3 Application of exothermic and endothermic reactions 180 Experiment 12 182 Experiment 13 184 2.4 Bond