Chemical thermodynamics

Introduction to chemical thermodynamics

Chemical thermodynamics is the study of the heat associated

with the reaction.

Therefore, chemical thermodynamics refers to the interconversion of various

kinds of energy and the changes in physical properties that are involved.
 Introduction to chemical thermodynamics ct’d
Thermodynamic is the study of the flow of heat or any other form of energy into
or out of a system as it undergoes a physical or chemical transformation

These energies can be:

∆H (Enthalpy) ∆S ( Entropy) ∆G (Gibb’s free energy)

Heat of a reaction direction of reaction spontaneity of a reaction

Introduction to chemical thermodynamics ct’d
In some situations, the energy produced by chemical reactions is actually of
greater interest to chemists than the material products of the reaction.

For example, the controlled combustion of organic molecules, primarily sugars

and fats, within our cells provides the energy for physical activity, thought, and
other complex chemical transformations that occur in our bodies.
Theminologies: System and Surroundings

System: part of the universe we are interested in.

Surrounding: the rest of the universe.
Exothermic: energy released by system to surrounding.
Endothermic: energy absorbed by system from surroundings.
 Types of thermodynamic system

•A closed system can exchange energy but not matter with the
surrounding , like a closed balloon or beaker without insulation.

•An open system can exchange matter and energy with the surrounding,
like a pot of boiling water.

•An isolated system cannot exchange both energy or matter with the
surrounding, such as an isolated bomb calorimeter.
Thermodynamic system ct’d
Thermodynamic processes ct’d
 Thermodynamic processes ct’d
Exothermic process is any process that gives off heat – transfers thermal energy
from the system to the surroundings.
2H2 (g) + O2 (g) 2H2O (l) + energy

H2O (g) H2O (l) + energy

Endothermic process is any process in which heat has to be supplied (heat is

absorbed) to the system from the surroundings.

energy + 2HgO (s) 2Hg (l) + O2 (g)

energy + H2O (s) H2O (l)

Spontaneous Processes

• Spontaneous processes are those that can proceed

without any outside intervention.
• The gas in vessel B will spontaneously effuse into vessel
A, but once the gas is in both vessels, it will not
spontaneously effuse into one of the vessel
Spontaneous Processes

Processes that are spontaneous in one

direction are nonspontaneous in the
reverse direction.
Spontaneous Processes
• Processes that are spontaneous at one temperature may be
nonspontaneous at other temperatures.
• Above 0C it is spontaneous for ice to melt.
• Below 0C the reverse process is spontaneous.
Reversible Processes

In a reversible process the system changes in such a way that the

system and surroundings can be put back in their original
states by exactly reversing the process

• Simply, it is a process in which the system and environment can

be restored to the same initial states by a backward process
Irreversible Processes

• Irreversible processes cannot be undone (restored) by exactly reversing the

change to the system.

All Spontaneous processes are irreversible.

First law of thermodynamic

Energy

1st Law of Thermodynamics

Enthalpy / Calorimetry

Hess' Law Enthalpy of Formation

First Law of Thermodynamics: Conservation of
energy
• Energy is neither created nor destroyed. Thus, the total energy is a constant.
• In other words, Energy can, however, be converted from one form to another or
transferred from a system to the surroundings or vice versa.

“The total energy of a system is equal to the sum of potential

energy and kinetic energy.”
 Energy
What does energy mean ? What are the types of energy do you know ?
• Energy is the ability to do work or transfer heat.
• Energy used to cause an object that has mass to move is called work.
• Energy used to cause the temperature of an object to rise is called heat.

Potential Energy
Potential energy is energy an object possesses because of its position relative
to other objects or chemical composition.

Kinetic Energy
Kinetic energy is energy an object has because of its motion. © 2009, Prentice-Hall, Inc.
 Internal Energy of the system
The change in the internal energy of a system is: E = q + w
q = Amount of heat energy that enters of leaves the system (heat in
and out of the system)
w = work (work done by the system or done on the system)

© 2009, Prentice-Hall, Inc.

 Heat (q)
Heat flows from warmer objects to cooler objects.

∆E = q + w

Heat (q) is:

(+) heat absorbed by the system (heat flows from
surrounding into the system): Endothermic

(-) heat released by the system (heat flows out of the system
into the surrounding): Exothermic
© 2009, Prentice-Hall, Inc.
 Heat (q)
Heat (q) can be calculated by:
q = m C ∆T
m = mass (g)
∆E = q + w C= specific heat capacity (j/g. oC). For H2O (C =
4.184 j/g. oC)
∆T= temperature change ( in kelvin or celsius)

Example:
Calculate the amount of energy required to heat up 50 g of water from 25 to 75
oC? The specific heat capacity of water is 4.184 j/g oC.
Heat (q): Example

Solution: q = m C∆T
m=50 g
C = 4.184 j/g oC.
∆T= T2-T1 = 75 -25 = 50
q = 50 g x 4.184 j/g oC x 50 oC

q = 10,460 j
 Heat absorbed or released: ∆H
If you need to calculate the heat energy that is absorbed or released
whenever there is a change, we use the equation (s):
q = m ∆H or q = n ∆H

∆H can be enthalpy of vaporization (∆Hvap) or fusion (∆Hfus)

m and n will depend on the units of ∆H

Example 1:
How much heat energy is required to melt 54 g of ice at 0 degree into liquid
water at the same temperature? Heat of fusion of for water is 6 Kj per mole
 Heat (q): Example
Solution: q = n ∆H
m = 54 g
n= m/Mm =54g/18g mol-1 = 3 mol

q = 3 mol x 6 kj/mol = 18 kj

Example 2:
Consider the combustion of propane:
C3H8 + 5O2 → 3CO2 + 4H2O + 1200kj
(a)If 64 g of oxygen has reacted, how much heat energy is released.
(b) If 3600 kj of heat was released, how many grams of CO2 was produced
 Work (w)
• Energy/force (F) used to move an object over
some distance (d) is work (w = F  d)

∆E = q + w

Work (w) is:

(+) Work is done on the system

(-) Work is done by the system

© 2009, Prentice-Hall, Inc.

 Work (w)
For a system ( chemistry point of view)
w = - p∆v; p (pressure) and ∆v (change in volume)

When a gas expands (increase in volume), the work is negative because the
system can exert a force to move a piston.
Now, to compress a gas you need to apply a force (work is positive). Whenever
gas is compressed the change in volume is negative (∆v = v2-v1 and v2< v1)
 Work (w)

When the work is done by the system (w = (-), then ∆E = q – w (∆E

goes down/decreases/lose energy to do work) and when work is done
on the system, (w = (+), then ∆E= q + w (∆E increases/gain energy)
 Internal energy, work (w): Example
If 300 joules of heat energy was absorbed by the system causes a gas to expand
from 2 L to 3 L at a pressure of 5 atm. Calculate the change in the internal
energy of the system. Given 101.3 j = 1L . atm

Solution: ∆E= q + w
W= -p∆v; -5 atm x (3 L-2 L) = -5 atm. L (convert this into joule)
101.3 j = 1L.atm
-5 L.atm = -5 x 101.3 = -506.5 j (because there is gas expansion, the system has
done work, so (w) it is negative)
∆E= q + w
∆E= 300 j + (-506.5 j) = -206.5 j
Limitations of the First law of thermodynamic
Enthalpy of Reaction
This quantity, H, is called the enthalpy of reaction, or the heat of reaction.

© 2009, Prentice-Hall, Inc.

∆Hro = Standard bond enthalpy for a reaction
Negative (-)∆Hro = Exothermic reaction
Positive (+)∆Hro = Endothermic reaction
Energy in Foods
Most of the fuel in the food we eat comes
from carbohydrates and fats.

© 2009, Prentice-Hall, Inc.

EXERCISES
 Exercise 2
3.6 g ethanol is burned in a bomb calorimeter with a heat capacity of
2.3 kj/c. The temperature of the calorimeter increases from 12.1C to
55.5 C. Calculate the energy of combustion per mole.

Exercise 3
How much energy is required to heat 80 grams of water from 26 C to
48 C? The specific heat capacity of water is 4.184 j/gC
 Exercise 4
How much energy is required to melt 75 g of ice? The heat of fusion
of ice is 334 j/g
Answer: Use q = m∆Hfusion
 Exercise 5
100g of iron metal (c = 0.45 j/g C) was placed in 200g of water at 25
oC. What is the final temperature of the mixture?