Energetics Answers

Enthalpy Changes

1. A reaction in which heat energy is released into the surroundings. Chemical potential energy (enthalpy) is converted into heat energy.
2. The enthalpy increases before it decreases because energy is required to break the existing bonds before the energy is released when the new bonds are made
3. +890 kJmol^-1
4. a) 5560 kJ  b) 9 kJ c) 899 g
5. A reaction in which heat energy is absorbed from the surroundings. Heat energy is converted into chemical potential energy (enthalpy).
6. -2802 kJmol^-1
7. a) 15600 kJ  b) 43 kJ   c) 4 g

Bond Breaking, Bond Making

1. The energy required to break one mole of a covalent bond
   1. Homolytically
   2. In the gas phase
   3. Averaged over a range of different environments

1. a) -40 kJmol^-1 b) -818 kJmol^-1 c) -537 kJmol^-1 d) -96 kJmol^-1
2. 298 kJmol^-1
3. Bond energies are average values and the average value may be different from the bond energy in that particular environment

 

Calorimetry

1. -26.3 kJmol^-1
2. -193.5  kJmol^-1
3. +11.1 kJmol^-1
4. -118 kJmol^-1
5. -56.8 kJmol^-1
6. -52.0 kJmol^-1
7. -930 kJmol^-1
8. -1530 kJmol^-1
9. They do not take into account heat loss to the surroundings or the heat capacity of the calorimeter.

 

Formation and Combustion Equations

1. Enthalpy change when one mole of a compound is formed from its elements with all reactants and products in their standard states under standard conditions

2. a) Mg(s) + 1/2O₂(g) MgO(s)
3. b) C(s) + O₂(g) CO₂(g)
4. c) 4C(s) + 5H₂(g) C₄H₁₀(g)
5. d) 2C(s) + 3H₂(g) + 1/2O₂(g) C₂H₆O(l)
6. e) 2Al(s) + 3/2O₂(g) Al₂O₃(s)
7. By definition – because they are already elements in their standard states
8. Enthalpy change when one mole of a substance is completely burned in excess oxygen with all reactants and products in their standard states under standard conditions

5. Write equations which represent the standard enthalpy of combustion of the following substances:
6. a) CH₄(g) + 2O₂(g) CO₂(g) + 2H₂O(l)
7. b) C₆H₆(g) + 7.5O₂(g)  6CO₂(g) + 3H₂O(l)
8. c) C₂H₆O(g) + 3O₂(g) 2CO₂(g) + 3H₂O(l)
9. d) H₂(g) + 1/2O₂(g)   H₂O(l)
10. e) Al(s) + 3/4O₂(g) 1/2Al₂O₃(s)
11. O₂, CO₂, H₂O

Hess’s Cycle

1. a) i) C₂H₆(g) + 3½O₂(g)  2CO₂(g) + 3H₂O(l)
2. ii) C₂H₄(g) + 3O₂(g) 2CO₂(g) + 2H₂O(l)
3. b) i) -1558.9 kJmol^-1       ii) -1410 kJmol^-1
4. -152 kJmol^-1
5. -1532 kJmol^-1
6. diborane: -2027.4 kJmol^-1 benzene: -3167.9 kJmol^-1
7. -265.1 kJmol^-1
8. -126.8 kJmol^-1
9. -126 kJmol^-1
10. a) -75 kJmol^-1 b) -606 kJmol^-1

Born-Haber Cycles

1. -362 kJmol^-1
2. -775 kJmol^-1
3. a) -386.5 kJmol^-1
4. b) lattice energy of CaCl is -245.5 kJmol^-1. This is less exothermic than the lattice energy of CaCl₂ the reaction: 2CaCl(s) Ca(s) + CaCl₂(s) is exothermic and so should be spontaneous

 

Ionic Compounds in Solutions

1. +11 kJmol^-1
2. Ba(OH)₂: -45 kJmol^-1, Ca(OH)₂: +80 kJmol^-1, Mg(OH)₂: +155 kJmol^-1 the more exothermic a reaction, the more likely it is to be spontaneous so Ba(OH)₂ is the most soluble, followed by Ca(OH)₂ and then Mg(OH)₂

3. +77 kJmol^-1 this is more endothermic than the enthalpy of solution of NaCl so dissolving AgCl is less spontaneous than dissolving NaCl.

 

Gibb’s Free Energy

1. ∆H = +135 kJmol^-1   ∆S = +334 JK^-1mol^-1

Reaction feasible above 404 K (131 ^oC)

2. ∆S = -99.4 JK^-1mol^-1

Reaction feasible below 462 K (189 ^oC); the lower the temperature, the higher the yield

3. ∆S = +305.3 JK^-1mol^-1 Reaction feasible above 102 K

(feasible at all temperatures for which water is liquid)