Chapter 33 Energy changes in chemical reactions 33.1 Conservation of energy 33.2 Exothermic and endothermic reactions Key terms Progress check Summary Concept map P. 1 / 49
33.1 Conservation of energy System and surroundings When a chemical reaction takes place, there is an energy change. The energy change usually results in an energy transfer between the system (the substances that are involved in the reaction) and its surroundings (the system s environment). Learning tip Physical and chemical changes are accompanied by energy changes. The study of such energy changes is called energetics. P. 2 / 49
Precipitation of silver chloride AgNO 3 (aq) + NaCl(aq) AgCl(s) + NaNO 3 (aq) System: the reactants (AgNO 3 (aq) + NaCl(aq)) the products (AgCl(s) + NaNO 3 (aq)) Surroundings: the container (e.g. test tube) the atmosphere over the solution everything else in the universe 33.1 Conservation of energy P. 3 / 49
add NaCl(aq) heat heat AgNO 3 (aq) AgCl(s) + NaNO 3 (aq) heat Figure 33.1 Heat is given out during the precipitation of silver chloride. 33.1 Conservation of energy P. 4 / 49
Law of conservation of energy The Law of conservation of energy states that energy can neither be created nor destroyed. Energy can be changed from one form to another. The total amount of energy remains constant. According to the Law of conservation of energy, the total amount of energy of the system and its surroundings before and after the reaction remains constant. 33.1 Conservation of energy P. 5 / 49
During the precipitation of silver chloride, the energy lost by the system must be equal to that gained by the surroundings. surroundings surroundings system Ag + (aq), NO 3 (aq) Na + (aq), Cl (aq) energy system AgCl(s), Na + (aq), NO 3 (aq) energy initial Total energy (initial) = energy of the system + energy of the surroundings final Total energy (final) = (energy of the system energy transferred) + (energy of the surroundings + energy transferred) Figure 33.2 The total amount of energy of the system and the surroundings during precipitation of silver chloride remains constant. 33.1 Conservation of energy P. 6 / 49
When sodium chloride is dissolved in water, energy is taken in from the surroundings. surroundings surroundings system NaCl(s), H 2 O(l) energy system Na + (aq), Cl (aq) energy initial Total energy (initial) = energy of the system + energy of the surroundings final Total energy (final) = (energy of the system + energy transferred) + (energy of the surroundings energy transferred) Total energy (initial) = Total energy (final) Figure 33.3 The total amount of energy of the system and the surroundings during dissolution of sodium chloride remains constant. 33.1 Conservation of energy Class practice 33.1 P. 7 / 49
Change in internal energy and enthalpy change If a chemical reaction is carried out at constant volume, the heat change measured is called the change in internal energy of the reaction. If a chemical reaction is carried out at constant pressure, the heat change measured is called the enthalpy change of the reaction. Learning tip The internal energy of a substance is the sum of the total kinetic energy and potential energy of all particles in that substance. 33.1 Conservation of energy P. 8 / 49
Reaction between Mg and dilute HCl Mg(s) + 2HCl(aq) MgCl 2 (aq) + H 2 (g) Reaction in a closed system (e.g. stoppered flask) The volume remains constant throughout the reaction. 433.8 kj of heat is given out for every mole of magnesium reacted. Key point Heat change at constant volume Change in = = internal energy 433.8 kj 33.1 Conservation of energy P. 9 / 49
Learning tip A negative sign indicates that heat is given out to the surroundings. Closed system (at constant volume) hydrogen magnesium ribbon dilute hydrochloric acid magnesium chloride solution Figure 33.4 The reaction between magnesium and dilute hydrochloric acid takes place at constant volume. In this case, heat change at constant volume is equal to the change in internal energy. 33.1 Conservation of energy P. 10 / 49
Reaction in an open system The pressure remains constant throughout the reaction. Only 430.8 kj of heat is given out. hydrogen gas produced from the reaction has to push back the air in the atmosphere, thus doing work and consuming energy (3.0 kj). 33.1 Conservation of energy P. 11 / 49
Key point Heat change at constant pressure = Change in internal energy (i.e. 433.8 kj) = 430.8 kj = Enthalpy change of the reaction + Work done on the surroundings (i.e. 3.0 kj) Learning tip This relationship is important particularly for reactions involving changes in gas volumes. 33.1 Conservation of energy P. 12 / 49
Open system (at constant pressure) atmospheric pressure (can be imagined as weightless piston) work done against atmospheric pressure magnesium ribbon dilute hydrochloric acid magnesium chloride solution Figure 33.5 The reaction between magnesium and dilute hydrochloric acid takes place at constant pressure. In this case, heat change at constant pressure is equal to the enthalpy change of reaction. 33.1 Conservation of energy P. 13 / 49 hydrogen
Enthalpy and enthalpy change All substances have a certain enthalpy. Enthalpy is an indication of the energy content; cannot be measured directly. The enthalpy change between reactants and products can be measured experimentally. The enthalpy, sometimes called heat content, is denoted by the symbol H. The symbol for enthalpy change is H. 33.1 Conservation of energy P. 14 / 49
Key point Enthalpy change of a reaction ( H) = Enthalpy of products (H p ) Enthalpy of reactants (H r ) i.e. H = H p H r A positive value of H indicates that heat is taken in (or absorbed) from the surroundings. A negative value of H indicates that heat is given out (or released) to the surroundings. Class practice 33.2 33.1 Conservation of energy P. 15 / 49
33.2 Exothermic and endothermic reactions Nature of exothermic and endothermic reactions in terms of enthalpy change Enthalpy change of exothermic reactions When zinc is added to copper(ii) sulphate solution, a displacement reaction takes place. Zn(s) + CuSO 4 (aq) Cu(s) + ZnSO 4 (aq) The reaction mixture becomes warm as heat is given out (released) from the reacting system to the surroundings. This reaction is an exothermic reaction. P. 16 / 49
copper(ii) sulphate solution (a) before zinc is added (b) after zinc is added Figure 33.6 When zinc displaces copper from its aqueous solution, the reaction mixture becomes warm. 33.2 Exothermic and endothermic reactions P. 17 / 49
An exothermic reaction Heat is given out to the surroundings. The total enthalpy of the products (H p ) is less than that of the reactants (H r ). The enthalpy of the system decreases and the enthalpy change ( H) is negative. Key point An exothermic reaction is a reaction that gives out heat. The temperature of the reaction mixture will increase. 33.2 Exothermic and endothermic reactions P. 18 / 49
heat H r reactants heat system hotter than surroundings heat Enthalpy H p products H is negative (a) heat (b) H = H p H r Since H p < H r, H is negative Figure 33.7 In an exothermic reaction, (a) heat is given out to the surroundings and (b) the enthalpy change, H, is negative. Learning tip The diagram shown in Figure 33.7(b) is called enthalpy level diagram. 33.2 Exothermic and endothermic reactions P. 19 / 49
Examples of exothermic reactions 1. All combustion reactions. E.g., the combustion of town gas. 2H 2 (g) + O 2 (g) 2H 2 O(l) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) Figure 33.8 Hydrogen and methane are the main components of Hong Kong town gas (about 50% and 30% respectively). Burning of town gas in a gas burner gives out a large amount of heat energy. 2. All acid-alkali neutralization reactions. E.g., the reaction between sodium hydroxide solution and hydrochloric acid. NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) 33.2 Exothermic and endothermic reactions P. 20 / 49
3. All precipitation reactions. For example, the reaction between silver ions and chloride ions. Ag + (aq) + Cl (aq) AgCl(s) 4. The reaction between calcium oxide (quicklime) and water. CaO(s) + H 2 O(l) Ca(OH) 2 (s) 33.2 Exothermic and endothermic reactions P. 21 / 49
Enthalpy change of endothermic reactions When hydrated barium hydroxide (Ba(OH) 2 8H 2 O) is mixed with ammonium chloride (NH 4 Cl) at room conditions, the reaction mixture becomes cold. Heat is taken in (absorbed) from the surroundings. This reaction is an endothermic reaction. 33.2 Exothermic and endothermic reactions P. 22 / 49
hydrated barium hydroxide ammonium chloride water 1. Hydrated barium hydroxide is mixed with ammonium chloride in a beaker. 2. A large drop of water is placed on a block of wood. 3. The bottom of the beaker becomes cold enough to freeze the water drop and the beaker sticks to the wood. Figure 33.9 The reaction between hydrated barium hydroxide and ammonium chloride absorbs heat from the surroundings. This can be demonstrated by the above simple experiment, in which a large water drop is frozen due to the drop in temperature. 33.2 Exothermic and endothermic reactions P. 23 / 49
An endothermic reaction Heat is taken in from the surroundings. The total enthalpy of the products (H p ) is greater than that of the reactants (H r ). The enthalpy of the system increases and the enthalpy change ( H) is positive. 33.2 Exothermic and endothermic reactions P. 24 / 49
heat H p products heat system colder than surroundings heat Enthalpy H r reactants H is positive (a) heat (b) H = H p H r Since H p > H r, H is positive Figure 33.10 In an endothermic reaction, (a) heat is taken in from the surroundings and (b) the enthalpy change, H, is positive. Key point An endothermic reaction is a reaction that takes in heat. The temperature of the reaction mixture will decrease. 33.2 Exothermic and endothermic reactions P. 25 / 49
Examples of endothermic reactions (1) Cracking of petroleum fractions. For example, the cracking of decane. C 10 H 22 (l) C 4 H 8 (g) + C 6 H 14 (l) (2) Thermal decomposition of calcium carbonate CaCO 3 (s) CaO(s) + CO 2 (g) 33.2 Exothermic and endothermic reactions P. 26 / 49
Exothermic reaction Endothermic reaction Heat transfer to/from the surroundings Heat is given out to the surroundings. Heat is taken in from the surroundings. Relationship between H r and H p H p < H r H p > H r Sign convention for H (i.e. H p H r ) negative ( ) positive (+) H r reactants H p products Diagrammatic representation Enthalpy H p products H is negative H r Enthalpy reactants H is positive Examples Combustion reactions Precipitation reactions Acid-alkali neutralization reactions Cracking of petroleum fractions Thermal decomposition of calcium carbonate Table 33.1 Comparison between exothermic and endothermic reactions. 33.2 Exothermic and endothermic reactions Class practice 33.3 P. 27 / 49
Some applications of exothermic and endothermic processes A number of salts are employed as chemical reagents in instant heat packs and instant cold packs. Figure 33.11 An instant heat pack Figure 33.12 An instant cold pack 33.2 Exothermic and endothermic reactions P. 28 / 49
Learning tip Dissolving a substance in water without hydrolysis (reaction between water and substance(s)) is commonly regarded as a physical change, not a chemical change. 33.2 Exothermic and endothermic reactions P. 29 / 49
Instant heat packs The salts inside release heat when dissolved in the solvent (e.g. calcium chloride dissolved in water). Instant heat packs can be used for keeping warm and treating muscular fatigue. Instant cold packs The salts inside absorb heat when dissolved in the solvent (e.g. ammonium nitrate dissolved in water). Instant cold packs can be used for treating sports injuries and reducing swelling or pain. Experiment 33.1 Experiment 33.1 STSE connections 33.1 33.2 Exothermic and endothermic reactions P. 30 / 49
Nature of exothermic and endothermic reactions in terms of breaking and forming of chemical bonds Chemical reactions involve the breaking and forming of bonds. Breaking of bonds takes in energy. Forming of bonds gives out energy. Key point Breaking of bonds is an endothermic process; forming of bonds is an exothermic process. Think about 33.2 Exothermic and endothermic reactions P. 31 / 49
Exothermic reactions involving a net release of energy Combustion of hydrogen H atom 2H 2 (g) + O 2 (g) 2H 2 O(l) Animation (Exothermic reaction and endothermic reaction) O atom Bond-breaking 2 H H and 1 O=O (a) endothermic process Bond-forming 4 O H (b) exothermic process strong covalent bonds strong covalent bonds Figure 33.13 (a) Bond-breaking processes and (b) bond-forming processes during the combustion of hydrogen. 33.2 Exothermic and endothermic reactions P. 32 / 49
The formation of four O H bonds gives out more energy than that is needed for the breaking of two H H and one O=O bonds. heat is given out to the surroundings. Combustion of hydrogen is an exothermic reaction. 4H + 2O Enthalpy bond-breaking 2H 2 + O 2 bond-forming H = ve 2H 2 O Figure 33.14 A diagram showing the relative enthalpies of different species in the combustion of hydrogen, which is an exothermic reaction. 33.2 Exothermic and endothermic reactions P. 33 / 49
Endothermic reactions involving a net absorption of energy Decomposition of ammonia H atom N atom 2NH 3 (g) N 2 (g) + 3H 2 (g) Bond-breaking 6 N H (a) endothermic process Bond-forming 1 N N and 3 H H (b) exothermic process Figure 33.15 (a) Bond-breaking processes and (b) bond-forming processes during the decomposition of ammonia into nitrogen and hydrogen. 33.2 Exothermic and endothermic reactions P. 34 / 49
The bond-forming processes give out less energy than that is needed for the bond-breaking processes. heat is taken in from the surroundings. Decomposition of ammonia into nitrogen and hydrogen is an endothermic reaction. Enthalpy 2N + 6H bond-breaking 2NH 3 bond-forming N 2 + 3H 2 H = +ve Figure 33.16 A diagram showing the relative enthalpies of different species in the decomposition of ammonia, which is an endothermic reaction. 33.2 Exothermic and endothermic reactions P. 35 / 49
Key point Enthalpy change of a reaction Energy change in = bond-breaking processes Energy change in bond-forming processes 33.2 Exothermic and endothermic reactions P. 36 / 49
Processes involved and their relative energy changes Exothermic reaction Bond-breaking < Bond-forming Endothermic reaction Bond-breaking > Bond-forming atoms or ions atoms or ions Diagrammatical representation Enthalpy bond-breaking reactants bondforming Enthalpy bond-breaking bond-forming products H = ve products reactants H = +ve Heat transfer to/from the surroundings Heat is given out to the surroundings. Heat is taken in from the surroundings. Sign convention for H negative ( ) positive (+) Table 33.2 Comparison between exothermic and endothermic reactions in terms of forming and breaking of chemical bonds. 33.2 Exothermic and endothermic reactions Class practice 33.4 P. 37 / 49
Key terms 1. change in internal energy 內能變化 2. endothermic reaction 吸熱反應 3. enthalpy 焓 4. enthalpy change 焓變 5. exothermic reaction 放熱反應 6. Law of conservation of energy 能量守恆定律 7. surroundings 環境 8. system 體系 P. 38 / 49
Progress check 1. What are the system and surroundings in a chemical reaction? 2. What does the Law of conservation of energy state? 3. In what way does the enthalpy change of a reaction relate to the change in internal energy? 4. What is an exothermic reaction? 5. What are the common examples of exothermic reactions? 6. What is an endothermic reaction? 7. What are the common examples of endothermic reactions? P. 39 / 49
8. Is bond-breaking an endothermic or exothermic process? 9. Is bond-forming an endothermic or exothermic process? 10.How can we represent an exothermic or endothermic reaction diagrammatically? Progress check P. 40 / 49
Summary 33.1 Conservation of energy 1. The Law of conservation of energy states that energy can neither be created nor destroyed but can be changed from one form to another. The total amount of energy of a system and its surroundings remains constant. 2. The heat change of a reaction measured at constant volume is the change in internal energy of that reaction. P. 41 / 49
3. The heat change of a reaction measured at constant pressure is the enthalpy change of that reaction. 4. Enthalpy change of a reaction ( H) = Enthalpy of products (H p ) Enthalpy of reactants (H r ) 33.2 Exothermic and endothermic reactions 5. An exothermic reaction is a reaction that gives out heat. The enthalpy change of an exothermic reaction is negative. Summary P. 42 / 49
6. Combustion, precipitation, acid-alkali neutralization and the reaction between calcium oxide and water are examples of exothermic reactions. 7. An endothermic reaction is a reaction that takes in heat. The enthalpy change of an endothermic reaction is positive. 8. Cracking of petroleum fractions and thermal decomposition of calcium carbonate are examples of endothermic reactions. 9. Breaking of bonds takes in energy and is an endothermic process. Summary P. 43 / 49
10. Forming of bonds gives out energy and is an exothermic process. 11. The table on the next page compares the nature of exothermic and endothermic reactions Summary P. 44 / 49
Heat transfer to/from the surroundings Temperature change of the reaction mixture Exothermic reaction Heat is given out to the surroundings. increases Endothermic reaction Heat is taken in from the surroundings. decreases Relationship between H r and H p H p < H r H p > H r Sign convention for H (i.e. H p H r ) Processes involved and their relative energy changes negative ( ) positive (+) Bond-breaking < Bond-forming Bond-breaking > Bond-forming Diagrammatic representation H r Enthalpy H p reactants products H is negative H p Enthalpy H r products reactants H is positive Summary P. 45 / 49
Concept map Law of conservation of energy are governed by under constant volume ENERGY CHANGES IN CHEMICAL REACTIONS under constant pressure Change in internal energy Enthalpy change P. 46 / 49
Enthalpy change depends on Energy change in bondbreaking processes Energy change in bondforming processes is equal to Enthalpy of products (H p ) Enthalpy of reactants (H r ) Concept map P. 47 / 49
Energy change in bondbreaking processes Energy change in bondforming processes positive negative Endothermic reaction Exothermic reaction Concept map P. 48 / 49
Enthalpy of products (H p ) Enthalpy of reactants (H r ) positive negative Endothermic reaction Exothermic reaction Concept map P. 49 / 49