Energetics
6.1 What is Energetics?
6.2 Ideal Enthalpy Changes Related to
Breaking and Formation of Bonds
6.3 Standard Enthalpy Changes
6.4 Experimental Determination of
Enthalpy Changes by Calorimetry
6.5 Hess’s Law
6.6 Calculations involving Enthalpy
Changes of Reactions
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
6.1
Internal Energy and Enthalpy
e.g.
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Heat change at
constant pressure
Enthalpy change
=
Change in internal
Work done on the
+
energy
surroundings
(Heat change at
constant volume)
6.1
Exothermic and Endothermic Reactions
An exothermic reaction is a reaction that releases
heat energy to the surroundings. (H = -ve)
6.1
Exothermic and Endothermic Reactions
An endothermic reaction is a reaction that absorbs
heat energy from the surroundings. (H = +ve)
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).
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).
6.2
Bond Enthalpies
Bond
H–H
C–C
C≡C
C≡ C
N–N
N═N
N≡ N
Mean Bond Enthalpy (kJ mol-1)
436
348
612
837
163
409
944
To be discussed in later chapters.
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 (101325 Nm-2); and
3. a temperature of 250C (298 K)
Enthalpy change under standard conditions denoted by symbol: H
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.
e.g. H+(aq) + OH-(aq)  H2O(l)
Hneut = -57.3 kJ mol-1
e.g. The standard enthalpy change of neutralization between
HNO3 and NaOH is -57.3 kJ mol-1
e.g. The standard enthalpy change of neutralization between
HCl and NaOH is -57.1 kJ mol-1
Standard Enthalpy Changes of neutralization
Acid
HCl
HCl
HCl
HF
Alkali
NaOH
KOH
NH3
NaOH
Hneu
-57.1
-57.2
-52.2
-68.6
H+(aq) + OH-(aq)  H2O(l)
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.
e.g. NaCl(s) + water  Na+(aq)+Cl-(aq) Hsoln=+3.9 kJ mol-1
e.g. LiCl(s) + water  Li+(aq) + Cl-(aq) Hsoln=-37.2 kJ mol-1
Note that enthalpy changes of solution associate
with physical changes.
6.3
Standard Enthalpy Changes of solution
Salt
NaOH
NaCl
KOH
KBr
Hsoln(kJ mol-1)
-44.7
+3.9
-57.8
+20.0
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.
e.g.
2Na(s) + Cl2(g)  2NaCl(s)
Na(s) + ½Cl2(g)  NaCl(s)
H = -822 kJ mol-1
Hf = -411 kJ mol-1
1 mole
Standard enthalpy change of formation of NaCl is -411 kJ mol-1.
6.3
Standard Enthalpy Changes of Reactions
What is Hf [N2(g)] ?
N2(g)  N2(g)
Hf [N2(g)] = 0
The enthalpy change of formation of an element
is always zero.
6.3
Standard Enthalpy Changes of Reactions
Standard Enthalpy Change of Combustion
e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) H1 = -2220 kJ
2C3H8(g) + 10O2(g) 6CO2(g) + 8H2O(l) H2 = ?
H2 = -4440 kJ
 It is more convenient to report enthalpy changes
per mole of the main reactant reacted/product formed.
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.
e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l)
Hc = -2220 kJ mol-1
1 mole
The standard enthalpy change of combustion
of propane is -2220 kJ mol-1
6.3
Standard Enthalpy Changes of Reactions
Substance
Hc (kJ mol-1)
H2(g)
C (diamond)
C (graphite)
CO(g)
CH4(g)
-285.8
-395.4
-393.5
-283.0
-890.4
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
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.
6.4 Experimental Determination of Enthalpy Changes by Calorimetry (p. 153)
Determination of Enthalpy Change of Combustion
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.
6.5 Hess’s Law (p. 157)
Hess’s Law
Route 1
A+B
H1
H2
C+D
H3
E
Route 2
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?
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.
6.5 Hess’s Law (p. 158)
Enthalpy Change of Formation of CO(g)
Given: Hf [CO2(g)] = -393.5 kJ mol-1; Hc [CO(g)] = -283.0 kJ mol-1
C(graphite) + ½O2(g)
+ ½O2(g)
Hf [CO(g)]
CO(g)
+ ½O2(g)
H1
CO2(g)
H2
Hf [CO(g)] + H2 = H1
Hf [CO(g)] = H1 - H2
= -393.5 - (-283.0 )
= -110.5 kJ mol-1
6.5
Enthalpy Change of Formation of CO(g)
Hf CO(g)] + H2 = H1
6.5
Enthalpy Change of Hydration of MgSO4(s)
ΔH
MgSO4(s) + 7H2O(l)
MgSO4·7H2O(s)
aq
ΔH1
aq
ΔH2
Mg2+(aq) + SO42-(aq) + 7H2O(l)
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
6.6 Calculations involving Enthalphy Changes of Reactions (p. 161)
Calculations involving standard enthalpy change
of reactions
reactants
Hreaction
products
 Hf [products]
-  Hf [reactants]
elements
Hreaction =  Hf [products] -  Hf [reactants]
The END