1. Energetics 6.1 What is Energetics?
6.2 Ideal Enthalpy Changes Related to Breaking and Formation of Bonds
6.3 Standard Enthalpy Changes

2. 6.4 Experimental Determination of Enthalpy Changes by Calorimetry
6.5 Hess’s Law
6.6 Calculations involving Enthalpy Changes of Reactions

3. 6.1
What is Energetics?
Energetics is the study of energy changes associated with chemical reactions.
Thermochemistry is the study of heat changes associated with chemical reactions.
Some terms
Enthalpy(H) = heat content in a substance
Enthalpy change(H) = heat content of products - heat content of reactants = Hp - Hr

4. Internal Energy and Enthalpy
6.1
Internal Energy and Enthalpy
e.g Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Heat change at constant pressure =
Change in internal energy +
Work done on the surroundings
(Heat change at constant volume)
Enthalpy change

5. Exothermic and Endothermic Reactions
6.1
Exothermic and Endothermic Reactions
An exothermic reaction is a reaction that releases heat energy to the surroundings. (H = -ve)

6. Exothermic and Endothermic Reactions
6.1
Exothermic and Endothermic Reactions
An endothermic reaction is a reaction that absorbs heat energy from the surroundings. (H = +ve)

7. Enthalpy Changes Related to Breaking and Forming of Bonds
6.2
Enthalpy Changes Related to Breaking and Forming of Bonds
e.g. CH4 + 2O2 CO2 + 2H2O
In an exothermic reaction, the energy required in breaking the bonds in the reactants is less than the energy released in forming the bonds in the products (products contain stronger bonds).

8. Enthalpy Changes Related to Breaking and Forming of Bonds
6.2
Enthalpy Changes Related to Breaking and Forming of Bonds
In an endothermic reaction, the energy required in breaking the bonds in the reactants is more than the energy released in forming the bonds in the products (reactants contain stronger bonds).

9. Mean Bond Enthalpy (kJ mol-1)
6.2
Bond Enthalpies
Bond
Mean Bond Enthalpy (kJ mol-1)
H – H
C – C
C≡C
C≡ C
N – N
N ═ N
N≡ N
436
348
612
837
163
409
944
To be discussed in later chapters.

10. Standard Enthalpy Changes
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ mol-1
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ mol-1
As enthalpy changes depend on temperature and pressure, chemists find it convenient to report enthalpy changes based on an internationally agreed set standard conditions:
1. elements or compounds in their normal physical states; 2. a pressure of 1 atm ( Nm-2); and 3. a temperature of 250C (298 K)
Enthalpy change under standard conditions denoted by symbol: H ø

11. Standard Enthalpy Changes of Reactions
Standard Enthalpy Change of Neutralization
The standard enthalpy change of neutralization (Hneut) is the enthalpy change when one mole of water is formed from the neutralization of an acid by an alkali under standard conditions. ø
e.g. H+(aq) + OH-(aq)  H2O(l) Hneut = kJ mol-1 ø
e.g. The standard enthalpy change of neutralization between HNO3 and NaOH is kJ mol-1
e.g. The standard enthalpy change of neutralization between HCl and NaOH is kJ mol-1

12. Standard Enthalpy Changes of neutralization
-57.1
-57.2
-52.2
-68.6
NaOH
KOH
NH3
HCl HF
Hneu
Alkali
Acid ø
H+(aq) + OH-(aq)  H2O(l)

13. Standard Enthalpy Changes of Reactions
6.3
Standard Enthalpy Changes of Reactions
Standard Enthalpy Change of Solution
The standard enthalpy change of solution (Hsoln) is the enthalpy change when one mole of a solute is dissolved to form an infinitely dilute solution under standard conditions. ø
e.g. NaCl(s) + water  Na+(aq)+Cl-(aq) Hsoln=+3.9 kJ mol-1 ø
e.g. LiCl(s) + water  Li+(aq) + Cl-(aq) Hsoln=-37.2 kJ mol-1 ø
Note that enthalpy changes of solution associate with physical changes.

14. Standard Enthalpy Changes of solution
6.3
Standard Enthalpy Changes of solution
-44.7
+3.9
-57.8
+20.0
NaOH
NaCl
KOH
KBr
Hsoln(kJ mol-1)
Salt ø

15. Standard Enthalpy Changes of Reactions
6.3
Standard Enthalpy Changes of Reactions
Standard Enthalpy Change of Formation
The standard enthalpy change of formation (Hf) is the enthalpy change of the reaction when one mole of the compound in its standard state is formed from its constituent elements under standard conditions. ø
e.g. 2Na(s) + Cl2(g)  2NaCl(s) H = -822 kJ mol-1 ø
Standard enthalpy change of formation of NaCl is -411 kJ mol-1.
Na(s) + ½Cl2(g)  NaCl(s) Hf = -411 kJ mol mole ø

16. Standard Enthalpy Changes of Reactions
6.3
Standard Enthalpy Changes of Reactions
What is Hf [N2(g)] ? ø
N2(g)  N2(g)
Hf [N2(g)] = 0 ø
The enthalpy change of formation of an element is always zero.

17. Standard Enthalpy Changes of Reactions
6.3
Standard Enthalpy Changes of Reactions
Standard Enthalpy Change of Combustion
e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) H1 = kJ
2C3H8(g) + 10O2(g) 6CO2(g) + 8H2O(l) H2 = ?
H2 = kJ
 It is more convenient to report enthalpy changes per mole of the main reactant reacted/product formed.

18. Standard Enthalpy Changes of Reactions
6.3
Standard Enthalpy Changes of Reactions
Standard Enthalpy Change of Combustion
The standard enthalpy change of combustion (Hc) of a substance is the enthalpy change when one mole of the substance burns completely under standard conditions. ø
e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l)
Hc = kJ mol mole ø
The standard enthalpy change of combustion of propane is kJ mol-1

19. Standard Enthalpy Changes of Reactions
6.3
Standard Enthalpy Changes of Reactions
-285.8
-395.4
-393.5
-283.0
-890.4
H2(g)
C (diamond)
C (graphite)
CO(g)
CH4(g)
Hc (kJ mol-1)
Substance ø

20. Experimental Determination of Enthalpy Changes by Calorimetry
6.4 Experimental Determination of Enthalpy Changes by Calorimetry (p. 151)
Experimental Determination of Enthalpy Changes by Calorimetry
Calorimeter = a container used for measuring the temperature change of solution
Determination of Enthalpy Change of Neutralization

21. Heat evolved = (m1c1 + m2c2) ΔT Where m1 is the mass of the solution,
6.4 Experimental Determination of Enthalpy Changes by Calorimetry (p. 151)
Heat evolved = (m1c1 + m2c2) ΔT
Where m1 is the mass of the solution,
m2 is the mass of calorimeter,
c1 is the specific heat capacity of the solution,
c2 is the specific heat capacity of calorimeter,
And Δ T is the temperature change of the reaction.

22. Determination of Enthalpy Change of Combustion
6.4 Experimental Determination of Enthalpy Changes by Calorimetry (p. 153)
Determination of Enthalpy Change of Combustion

23. Heat evolved = (m1c1 + m2c2) ΔT
6.4 Experimental Determination of Enthalpy Changes by Calorimetry (p. 153)
Heat evolved = (m1c1 + m2c2) ΔT
Where m1 is the mass of water in the calorimeter,
m2 is the mass of the calorimeter,
c1 is the specific heat capacity of the water,
c2 is the specific heat capacity of calorimeter,
And Δ T is the temperature change of the reaction.

24. Hess’s Law A + B C + D E Route 1 H1 H3 Route 2 H2 H1 = H2 + H3
6.5 Hess’s Law (p. 157)
Hess’s Law
Route 1
H1
A + B
C + D
H3
Route 2 E
H2
H1 = H2 + H3
Hess’s Law states that the total enthalpy change accompanying a chemical reaction is independent of the route by which the chemical reaction takes place.
Why?

25. Importance of Hess’s Law
6.5 Hess’s Law (p. 158)
Importance of Hess’s Law
The enthalpy change of some chemical reactions cannot be determined directly because:
the reactions cannot be performed in the laboratory
the reaction rates are too slow
the reactions may involve the formation of side products
But the enthalpy change of such reactions can be determined indirectly by applying Hess’s Law.

26. Enthalpy Change of Formation of CO(g)
6.5 Hess’s Law (p. 158)
Enthalpy Change of Formation of CO(g)
Given: Hf [CO2(g)] = kJ mol-1; Hc [CO(g)] = kJ mol-1 ø
C(graphite) + ½O2(g)
CO(g)
Hf [CO(g)] ø
H2
+ ½O2(g)
CO2(g)
H1
+ ½O2(g)
Hf [CO(g)] + H2 = H1 ø
Hf [CO(g)] = H1 - H2 ø
= ( )
= kJ mol-1

27. Enthalpy Change of Formation of CO(g)
6.5
Enthalpy Change of Formation of CO(g)
Hf CO(g)] + H2 = H1 ø

28. Enthalpy Change of Hydration of MgSO4(s)
6.5
Enthalpy Change of Hydration of MgSO4(s)
MgSO4(s) + 7H2O(l)
MgSO4·7H2O(s)
Mg2+(aq) + SO42-(aq) + 7H2O(l) ΔH
ΔH2 aq
ΔH1 ø aq
ΔH = enthalpy of hydration of MgSO4(s)
ΔH1 = molar enthalpy change of solution of anhydrous magnesium
sulphate(VI)
ΔH2 = molar enthalpy change of solution of magnesium
sulphate(VI)-7-water
ΔH = ΔH1 - ΔH2 ø

29. Calculations involving standard enthalpy changes of reactions
6.6 Calculations involving Enthalphy Changes of Reactions (p. 161)
Calculations involving standard enthalpy changes of reactions
reactants
products
Hreaction
 Hf [products] ø
elements
-  Hf [reactants] ø
Hreaction =  Hf [products] -  Hf [reactants] ø

30. The END