Chapter 9 Chemical Reactions I. Classification of Chemical Reactions

1. Chapter 9 Chemical Reactions I. Classification of Chemical Reactions 1. Combination reactions A + B C heat 2H2 + O2 2H2O I2 + H2 2HI 2Na + Cl2 2NaCl

2. Decomposition Reactions • A  B + C heat H2O + CO2 (gas) H2CO3 light 2H2O2 2H2O + O2 (gas) 3. Single Replacement Reactions Zn + CuSO4 Cu + ZnSO4 2NaI + Br22NaBr + I2 Zn + 2HCl H2 + ZnCl2

4. Double Replacement Reactions – Exchanging partners AgNO3(aq) + NaCl(aq)  NaNO3(aq) + AgCl(s) Solid (precipitation) 2KI(aq) + Pb(NO3)2(aq)  2KNO3(aq) + PbI2(s)

5. 5. Combustion Reactions - Reactions with O2 CH4 + 2O2 CO2 + 2H2O C3H8 + 5O2 3CO2 + 4H2O 2Mg + O2 2MgO

6. 6. Oxidation-reduction (redox) reactions • Consider 2Cu + O2 2CuO reduced oxidized Oxidizing agent reducing agent Oxidation must be accompanied by reduction 2Cu 2Cu2+ + 4e- Oxidized (loss of electrons) (increase of oxidation number) O2 + 4e- 2O2- Reduced (gaining of electrons) (decrease of oxidation number)

7. Oxidation number (Page 226 - 227) • Ox.# of an element in its element state is zero. • Ox.# of monatomic ion = charge of the ion. • Ox.# of H is +1 in most hydrogen containing compounds. • Ox.# of O is –2 in most O containing compounds • Ox.# of F = -1 in F containing compounds. • Ox.# of Cl, Br & I containing compounds are –1 except in • a compound with F or O. • Sum of ox.#’s in a compound = 0 • Sum of ox.#’s in a polyatomic ion = charge of the ion

8. Oxidation number • element =0 • monatomic ion = charge of the ion. • H = +1 in a compounds. • O = –2 • F = -1. • Cl, Br & I = –1 except in a compound with F or O. • Sum of ox.#’s = charge of the species Examples CuCl2 HNO3 Ox# H? Cu O? Cl N?

9. More examples

10. Examples of Redox Reactions Zn + CuSO4 ZnSO4 + Cu

11. Examples of Redox Reactions 2H2 + O2 2H2O 2Na + 2H2O  2NaOH + H2 3H2S + 2HNO3 3S + 2NO + 4H2O 3H2SO3 + 2HNO3 2NO + H2O + 3H2SO4

12. Important of redox reactions 1. Combustion reaction (see examples 1 and 2 above) 2. Resperation Similar to combustion but at low temperature 3. Bleaching 4. Rusting – reaction with oxygen 4Fe + 3O2 2Fe2O3

13. e- Batteries Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Page 151

14. II. Reaction Rates • The kinetic theory • A + B  C + D • A collides with B C + D A + B  (AB) Transition state or Activated complex

15. a) Collision orientation Figure 8.3

16. H+ H H + O + HCl H + Cl- O H H H Cl O H H Incorrect orientation H Cl O H H Correction orientation

17. b). Activation energy (Ea) – energy needed for reaction to occur Energy diagram for an exothermic reaction

18. Energy diagram for an endothermic reaction.

19. c) Exothermic reaction: A reaction in which energy is released A + B  C + D + heat Endothermic reaction: A reaction in which an input of energy is needed (absorbed) A + B + heat  C + D Origin of heat of reaction: Bond breaking – heat is absorbed Bond formation – heat is released

20. Reaction rate depends on 1) collision frequency 2) velocity (kinetic energy) of reactant molecules. Only those with kinetic energy > activation energy can undergo reaction B. Factors that affect reaction rates • Nature of the reactants • Concentration Rate increases as concentration increases – more collisions/s

21. 3) Temperature Reaction rate increases as T increases As T increases, kinetic energy of the reactants increases (more reactants have E > activation energy, Ea) Optimum body temperature = 37oC

22. Catalyst • A catalyst decreases the activation energy • - reaction rate increases A + B + X C + D + X Enzymes are catalysts Figure 8.9

23. 5. Surface area The larger the surface, the faster the reaction.

24. III. Chemical equilibrium • Chemical Equilibrium • Many reactions are reversible After some time, forward rate = reverse rate The system is at a state of chemicalequilibrium (or just equilibrium)

25. B. Equilibrium constant At equilibrium Equilibrium constant At constant T [A] = concentration of A In general Equilibrium (constant) expression

26. Example At equilibrium: [H2] = 0.86 M, [I2] = 0.86 M, [HI] = 0.27 M Large Keq more products than reactants at equilibrium Small Keq more reactants than products at equilibrium

27. PbCl2(s) Pb2+(aq) + 2Cl-(aq)

28. C. Factors that affect the equilibrium constant Le Chatelier’s Principle If a stress is applied to a system in equilibrium, a new equilibrium will be established in which the equilibrium position has been shifted in such a way as to relieve the applied stress. • Equilibrium position – a quantitative indication of the relative amounts of reactants and products at equilibrium a) Effect of concentration

29. Concentration changes that result when H2 is added to an equilibrium mixture involving the system. Equilibrium position shifts to the right

30. In general Adding A or B will cause the equilibrium position to shift to the right. Adding C or D will cause the equilibrium position to shift to the left. Adding A: [A] [B] [C] [D] b) Effect of pressure (for reactions involving gases) A increase of pressure causes the equilibrium to shift to the right

31. c) Effect of Temperature exothermic If T increases, shift left [A] and [B] increase, [C] and [D] decrease, Keq decreases If T decreases, shift right [A] and [B] decrease, [C] and [D] increase, Keq increases endothermic If T increases, shift  If T decreases, shift 

32. d) Catalyst Adding a catalyst speeds up both the forward and reverse reaction rates but does not change the equilibrium position.

33. Example: Equilibrium position Shift Adding NO Removing O2 Shift Adding NO2 shift Increasing T Decreasing T Increasing P by decreasing V Decreasing P by increasing V Adding a catalyst

34. An increase of temperature will cause Keq to An increase of pressure by decreasing volume will cause Keq to Only temperature will can change Keq of a given equilibrium