Thermochemistry is the study of heat and energy changes in chemical and physical processes. It is concerned with energy changes that accompany physical and chemical processes at constant volume or constant pressure. The first law of thermochemistry states that the quality of heat evolved or absorbed during a chemical reaction is the same, regardless of the pathway between the initial and final states. Hess's law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for the individual steps of the reaction. Enthalpy change (ΔH) of a reaction is the heat absorbed or released when one mole of the reaction occurs at constant pressure. It can be measured using a bomb calorimeter. Common standard

1. Introduction to chemical
thermodynamics/ thermochemistry
Heat, work, energy and the First
Law

2. Introduction
• Thermochemistry is part of broader subject
called thermodynamics, which is the study of
relationship between heat and other forms of
energy involved in a chemical or physical
processes.
• It is the part of chemistry concerned with energy
changes that is accompanied with physical and
chemical processes either at constant volume or
at constant pressure

3. Energy
• Is capacity to perform work
• Mechanical work is application of force
over distance
• Heat is energy transferred by virtue of
temperature gradient – associated with
molecular motion
• Joule demonstrated experimentally that
heat and work are interchangeable
forms of energy

4. Energy: forms
• Kinetic energy is the energy of motion
• Potential energy is energy stored – by
position, within a spring, within a chemical
bond, within the particles of a nucleus
2
2
1
mv
EK 
mgh
EP 

5. Energy: units
• From the definition of kinetic energy (1/2mv2),
we get the units of energy:
kg m2/s2
• S.I. unit for energy is the joule (J) = 1Nm
• Another common unit is the calorie (cal): the
energy required to raise the temperature of 1 g
of water by 1ºC
1 cal = 4.184 J

6. System and surroundings
• Any process can be divided into the SYSTEM
contained within the SURROUNDINGS
– When energy changes are measured in a chemical
reaction, the system is the reaction mixture and
the surroundings are the flask, the room, and the
rest of the universe.

7. Types of systems
• Open system: this can exchange mass and
energy with the surroundings. E.g, a quantity
of water in an open container.
• Closed system: this system allows the
transfer of energy and not mass. E.g, a
closed flask of water.
• Isolated system: this does not allow the
exchange either of mass or energy. E.g,
flask of water inside an insulator.

8. Internal energy
• Internal energy is the sum of all the types of
energy (kinetic and potential) of the system.
It is the capacity of the system to do work
• Typically we don’t know the absolute value
of U for the system
– (Internal energy usually has symbol U. Other
sources use E)
• We can measure the change to the internal
energy
ΔU = Ufinal - Uinitial

9. Work done at constant pressure
• Gas generated in reaction pushes against the
piston with force: P x A
• At constant P, volume increases by ΔV and
work done by system is:
w = -PΔV (ΔV = A x d)
– Work done by system is –ve in expansion (ΔV > 0)
• ΔU < 0 (ΔV > 0, -PΔV < 0)
– Work done by system is +ve in contraction (ΔV <
0)
• ΔU > 0 (ΔV < 0, -PΔV > 0)

10. Expansion work
• Work done by gas expanding:
w = -PexΔV
• In expansion the ΔV > 0; w < 0
ΔU < 0
• In contraction, ΔV < 0; w > 0
ΔU > 0

11. Heat and internal energy
• Heat is transfer of energy by virtue of
temperature gradient.
• If system is cooler than surroundings
q > 0
• If system is hotter than surroundings
q < 0

12. Deposits and withdrawals
• Process is always viewed from perspective of
system
• Energy leaving system has negative sign
– (decreases internal energy – lowers the chemical bank
balance)
• Energy entering system has positive sign
– (increases internal energy – increases chemical bank
balance)
• Useful process is one where change is negative
• Energy is in the form of heat or work
ΔU = q + w
q = heat (released or absorbed by the system);
w=work (do on the system or by the system)

13. First Law of Thermodynamics
“ The total energy of a system can be converted from one
form to another but cannot be created or destroyed”
…Total internal energy of isolated system is
constant.
– Energy change is difference between final and initial
states (ΔU = Ufinal – Uinitial)
– Energy that flows from system to surroundings has
negative sign (Ufinal < Uinitial)
– Energy that flows into system from surroundings has
positive sign (Ufinal > Uinitial)

14. Functions of state
• State Function
A property that depends only on present state
of the system and is independent of pathway
to that state.
• It is defined by values of all relevant
macroscopic properties, such as composition,
energy, temperature, pressure, and volume.
• When state of a system changes, the
magnitude of change in any state depends
only on initial and final states of the system

15. Heat and work
• Any chemical process may have
associated with it heat and work terms
• The total internal energy change will be
the sum of the contributions from each
ΔU = q + w = q - P ΔV
q = ΔU + P ΔV
• In a sealed system ΔV = 0, so q = ΔU

16. Open systems and enthalpy
• Most reactions are conducted in open
vessels where P is constant and ΔV ≠ 0
• The heat change at constant pressure is
qP = ΔU + P ΔV
• Enthalpy (H) is defined as:
H = U + PV

17. Exothermic and Endothermic
Reactions
Exothermic reaction is a reaction in which heat is given out to the
surroundings. Thus ΔH is negative i.e. less than zero.
Endothemic reaction is a reaction in which heat is absorbed from
the surroundings. Thus, ΔH is positive i.e. greater than zero.

18. Enthalpy Changes
Enthalpy of a reaction can be defined as the heat given out or absorbed in
a chemical reaction at constant pressure and can be measured using
bomb calorimeter. Enthalpy changes is the difference between the heat
energy content of the product and that of the reactant. It is also called the
heat of a reaction.
Standard Enthalpies Change
1.Standard enthalpy of combustion
2.Standard enthalpy of formation
3.Standard enthalpy of neutralization
4.Standard enthalpy of solution
5.Standard enthalpy of dilution
Standard enthalpy of combustion
This is the heat changes i.e. heat given out when one mole of a substance
burnt completely in oxygen under standard conditions of temperature and
pressure. It is always negative i.e. exothermic. It is denoted by ΔHθ
c
C(s) + O2(g) CO2(g) ΔHθ
c = - 393 kJmol-1

19. Standard enthalpy of formation
This is the heat given out or absorbed when one mole of a substance is
formed from its constituent element under standard conditions of
temperature and pressure. Its symbol is ΔHθ
f. It can negative or positive.
Standard enthalpy of neutralization
This is the heat changes when one mole of H+ from an acid is just
neutralized by one mole of hydroxyl ion (OH-) from a base to form one
mole of water at standard conditions of temperature and pressure.
OH- + H+ H2O(l) ΔHθ = -57.4kJmol-1
Standard enthalpy of solution and dilution
This is the amount of heat evolved when one mole of a solute is
dissolved in a specific amount of solvent at standard conditions of
temperature and pressure.
Heat of solution is the total heat changes per mole of a solute when the
solution is completely formed at constant temperature and pressure.

20. Hess’s Law
• The change in enthalpy that occur when reactants are converted to
products in a reaction is the same whether the takes place in one
step or in series of steps.
• Hess’s law ague that ∆H for the net reaction is the sum of the ∆H
values for the individual reactions. Thus ∆H is a state function
• i.e. ∆H = ∑i∆H (kJmol-1)
• ∆Hθ
Reaction = ∑m∆Hθ
Product − ∑n∆Hθ
Reactant
• For example, in the combustion of graphite (carbon) to carbon
monoxide.
2C(graphite) + O2(g) ------ 2CO(g)
• To obtain the enthalpy change for the preparation of pure CO, the
Hess’s law is applied.

21. • Imagine that the reaction occur in two separate steps
2C(graphite) + 2O2(g) ------ 2CO2(g) (1st step)
2CO2(g) ------ 2CO(g) + O2(g) (2nd step)
(According to Lavoiser and Laplace laws, if a chemical reaction is
reversed, the sign of enthalpy changes but the magnitude remain the
same)
2C(graphite) + 2O2(g) ------ 2CO2(g ∆H1= (-393.5kJ)X(2)
2CO2(g) ------ 2CO(g) + O2(g) ∆H2= (-566.0kJ)X(-1)
2C(graphite) + O2(g) ------ 2CO(g) ∆H3= (-221.0kJ)
Example 2: Consider the following data/equations:
S(s) + O2 ----- SO2(g) ∆H = (-297kJ)
2SO3(g) ------ 2SO2 + O2 ∆H1= (198kJ)
How would you use these data to obtain the enthalpy change for the following
equation
2S(s) + 3O2 ---- 2SO3(g)

22. S(s) + O2 ----- SO2(g) ∆H = (-297kJ)
2SO3(g) ------ 2SO2 + O2 ∆H = (198kJ)
2S(s) + 2O2 ----- 2SO2(g) ∆H = (-297kJ)X(2)
2SO2 + O2 ----- 2SO3(g) ∆H = (198kJ)X(-1)
2S(s) + 3O2 ---- 2SO3(g) ∆H = -792kJ
Exercise: Consider the formation of tungsten carbide, WC, from W and C
with this reaction
W(s) + C(graphite) ---- WC(s)
The three steps reactions for this reaction is given as
W(s) + 3O2(g) ---- 2WO3(s); ∆H = (-1685kJ) (1)
C(graphite) + O2(g) ---- CO2(g); ∆H = (-393.5kJ) (2)
2WC(s) + 5O2 --- 2WO3(s) + 2CO2(g) ∆H = (-2391.8kJ) (3)
Calculate the enthalpy of formation of WC.

23. Examples
i. Calculate the standard heat of reduction of ferric oxide by aluminum which
proceeds according to the following reaction:
2Al(s) + Fe2O3(s) 2Fe(s) +Al2O3(s) ΔH = x kJmol-1
Given that the standard heat of formation of Fe2O3 and Al2O3 are -820.06kJmol-1
and -1669.83kJmol-1 respectively.
= (0 – 1669.83) – (0 – 820.06) = - 849.77kJmol-1
ii. Calculate the heat of the following reaction given that the heat of formation of
H2O(l), CO(g), and CO2(g) are -285.77, -108.78 and -393.30kJmol-1 respectively
at 20oC.
CO(g) + H2O(l) CO2(g) + H2(g)
= [(0 + (-393.30)) – (-108.78 + (-285.77))] = 1.25kJmol-1
Measurement of Heat of Reaction
This is usually measured accurately using the calorimeter. A calorimeter
containing the aqueous solution is often used to measure the heat absorbed or
released by a chemical reaction.
Thus; Q = MCΔT
Where, M = mass of water or solution; C = specific heat capacity of water or
solution, ΔT = Tempt rise
The specific heat capacity of water is 4.2jkg-1k-1.
Specific heat capacity = Heat capacity / mass

24. Example: When 2g of ethane C2H6 is burnt, the heat
liberated raises the temperature of 50g of water from 35oC
to 50.5oC. What is the heat of combustion of ethane? Given
that the specific heat capacity of water is 4.2Jg-1oC-1.
Q = MCΔT = 50 x 4.2 x 15.5 = -3255J or -3.255kJ
Molar mass of ethane = 30gmol-1
2g of ethane liberates -3255J
30g of ethane liberates = -3255/2 x 30 = -48825kJmol-1
Practice Question
The amount of energy required to change 15g of water to
steam is +45kJ. Calculate the standard heat of
vaporization.
Answer = +54kJmol-1

25. The laws of Thermochemistry
There are two laws of Thermochemistry which are based on the
principle of conservation of energy. The first law was given by
Lavoiser and Laplace in 1780 while the second law was given by
Hess’s in 1940.
First law of thermochemistry (Lavoiser and Laplace laws): It
states that the quality of heat which must be supplied to decompose
compound into its elements is equal to the heat evolved when the
compound is formed from its element.
e.g. H2(g) + 1/2O2(g) H2O H = - 16.32kJmol-1
H2O H2(g) + 1/2O2(g) H = + 16.32kJmol-1
Second law of thermochemistry (Hess’s law of heat summation):
It states that the heat given out or taken in during a chemical
reaction is always constant and it is independent of the route, time,
and any other intermediate changes involved. This simply means
that total changes accompanying a chemical reaction is
independent of the path taken.

26. Example 1: Given that the heat of formation of water and gaseous
carbon(iv)oxide are -285kJmol-1 and -393kJmol-1 respectively and the
heat of combustion of gaseous butane is -2874kJmol-1. Calculate the
standard enthalpy of formation of gaseous butane.
C(s) + O2(g) CO2(g) ΔH = -393kJmol-1
H2(g) + 1/2O2(g) H2O(l) ΔH = -285kJmol-1
C4H10(g) + 13/2O2(g) 4CO2(g) + 5H2O(l) ΔH = -2874kJmol-1
-2874 = [(4 x -393) + (5 x -285)] – [hfC4H10 + (13/2 x 0)]
Hf(C4H10) = -123kJmol-1
Method 2: Using Hess’s law; ΔH1 = ΔH2 + ΔH3 + ΔH4 = (4 x -393) + (5 x -
285) + (+2874) = -123kJmol-1
Example 2: Find the enthalpy change for the reaction
C2H4(g) + H2(g) C2H6(g)
Given that the enthalpy of formation of C2H6 and C2H4 are -84.6kJmol-1
and 52.3kJmol-1 respectively.
Hreactant = -84.6 – [52.3 + 0] = -136.9kJmol-1
Method 2: Using Hess’s law; ΔH3 + ΔH1 = ΔH2
H3 = H2 – H1 = -846 – (+52.3) = -136.9kJmol-1

27. Class Test 3
1. Write three different units of energy_______,
________, and ________.
2. Write the mathematical relationship between energy,
heat and work.
3. ________of a reaction can be defined as the heat given
out or absorbed in a chemical reaction at constant
pressure.
4. Heat of reaction can be measured using a__________.
5. Second law of thermochemistry is also known
as__________.
6. Mention any three standard enthalpy
changes_________,_________, and ___________.

28. Spontaneity of Reaction
Three factors determine the spontaneity of reaction:
- Enthalpy;
- Entropy
- Free energy
i. Entropy: This is a measure of the degree of disorder or randomness of
a system. It is represented by letter S and a change in it by ΔSθ.
ΔSθ = ΔSθ
2 - ΔSθ
1
The SI unit of entropy is JK-1
Example: Nitrogen gas reacts with hydrogen gas to form ammonia gas.
Given that the standard entropies of nitrogen, hydrogen and ammonia
gases are: 192, 131, and 193Jk-1mol-1 respectively. Calculate the
standard state entropy change for the reaction.
N2(g) + 3H2(g) 2NH3(g) ΔSθ = ?
= [2(193) – (192 + 393)] = - 199JK-1mol-1
Since the standard entropy change is negative, there is a decrease in
randomness or disorder of the system.

29. ii. Gibb’s Free energy (ΔG)
This is the measure of the useful work a system is capable of doing (apart
from volume changes). A change in Gibb’s free energy is represented by
ΔGθ at standard temperature and pressure. ΔGθ = ΔGθ
2 - ΔGθ
1
ΔGθ = ΔHθ – TΔSθ
Spontaneous reaction, ΔGθ = - ve;
Non Spontaneous reaction, ΔGθ = +ve
Equilibrium reaction, ΔGθ = 0
Example: Calculate the standard free energy at 25oC for the thermal
decomposition of CaCO3 from the following data.
CaCO3(s) CaO(s) + CO2(g)
ΔHθ = +178.3KJ(178300J); ΔSθ = +160.6JK-1, T = 25oC = 25 + 273 = 298K;
ΔGθ = + 130441.2J= +130.4KJ
Compound ΔHθf (KJ) ΔSθ (JK-1mol-1)
CaCO3 - 1206.9 92.9
CaO - 635.1 39.8
CO2 - 393.5 213.7

30. CLASS TEST 4
1. _________ is a measure of the degree of disorder in a
system.
2. _________,_________,and________ determines the
spontaneity of a reaction.
3. Free energy is also known as ____________.
4. The value of free energy is_________when a reaction is
spontaneous.
5. The industrial process by which ammonia is manufactured
is called __________.
6. When ΔG is zero, what will be the mathematical
relationship between ΔH, T and ΔS?
7. The value of standard temperature in Kelvin is ___________.
8. For an endothermic reaction ΔH is ___________.