Unit 6: Thermochemistry
 Introduction
 Heat and Work
 Specific Heat
 Enthalpy (DH)
 Enthalpy of Reaction
 Phase Diagram
Introduction
 Most daily activities involve processes
that either use or produce energy:
 Activities that produce energy
Metabolism of food
Burning fossil fuels
 Activities that use energy:
Photosynthesis
Pushing a bike up a hill
Baking bread
Introduction
 Thermodynamics
 The study of energy and its
transformations
 Thermochemistry:
 A branch of thermodynamics
 The study of the energy (heat)
absorbed or released during chemical
reactions
Introduction
 Objects can have two types of energy:
 Kinetic energy
Energy of motion
Thermal energy
The type of kinetic energy a
substance possesses because of
its temperature
 Potential energy
Energy of position
“stored” energy resulting from the
attractions and repulsions an object
experiences relative to other
objects
Introduction
 Units of Energy
 SI unit = joule (J)
1 J = the kinetic energy of a 2 kg
mass moving at a speed of 1 m/s
A very small quantity
 Kilojoule (kJ)
1 kJ = 1000 J
Introduction
 Units of Energy (cont)
 Calorie (cal)
Originally defined as the amount of
energy needed to raise the
temperature of 1g of water from
14.5oC to 15.5oC.
1 cal = 4.184 J (exactly)
 Kilocalorie (kcal)
1 kcal = 1000 cal
Introduction
Example: Convert 3.02 kJ to J.
Given: 3.02 kJ
Find: J
1 kJ = 1000 J
J = 3.02 kJ x 1000 J = 3020 J
1 kJ
Introduction
Example: Convert 725 cal to kJ.
Given: 725 cal
Find: kJ
1 cal = 4.184 J
1 kJ = 1000 J
J = 725 cal x 4.184 J x 1 k J = 3.03
kJ
1 cal
1000 J
Introduction
 When using thermodynamics to study
energy changes, we generally focus on a
limited, well-defined part of the
universe.
 System:
 The portion of the universe singled out
for study
 Surroundings:
 Everything else
Introduction
The system is
usually the
chemicals in the
flask/reactor.
The system
The flask and
everything else
belong to the
surroundings.
Introduction
 Open system:
 A system that can exchange both
matter and energy with the
surroundings
 Closed system:
 A system that can exchange energy
with the surroundings but not matter
A cylinder with a piston is one
example of a closed system.
Introduction
 In a closed system energy
can be gained from or lost
to the surroundings as:
Work
Heat
 Work:
 Energy used to cause an
object to move against a
force
Lifting an object
Hitting a baseball
Introduction
 Heat:
 The energy used to cause the
temperature of an object to increase
 The energy transferred from a hotter
object to a cooler one
 Energy:
 The capacity to do work or to
transfer heat
Introduction
 The potential energy of a
system can be converted into
kinetic energy and vice versa.
Potential energy
 Energy can be transferred
back and forth between the
system and the surroundings
as work and/or heat.
Kinetic energy
work
The First Law of Thermodynamics
 Although energy can be converted from
one form to another and can be
transferred between the system and the
surroundings:
Energy cannot be created or destroyed.
(First Law of Thermodynamics)
 Any energy lost by the system must be
gained by the surroundings and vice
versa.
The First Law of Thermodynamics
 The First Law of Thermodynamics can
be used to analyze changes in the
Internal Energy (E) of a system.
 The sum of all kinetic and potential
energy of all components of a system
 For molecules in a chemical system, the
internal energy would include:
 the motion and interactions of the
molecules
 the motion and interactions of the
nuclei and electrons found in the
molecules
The First Law of Thermodynamics
 Internal Energy:
 Extensive property
depends on mass of system
 Influenced by temperature and
pressure
 Has a fixed value for a given set of
conditions
 State function
The First Law of Thermodynamics
 The internal energy of a system is a
state function.
 A property of the system that is
determined by specifying its condition
or its state in terms of T, P, location,
etc
 Depends only on its present condition
 Does not depend on how the system
got to that state/condition
The First Law of Thermodynamics
 The internal energy of a system can
change when:
 heat is gained from or lost to the
surroundings
 work is done on or by the system.
 The change in the internal energy
D E = Efinal - Einitial
DE
= change in internal energy
Efinal = final energy of system
Einitial = initial energy of system
The First Law of Thermodynamics
 If Efinal > Einitial,
 DE >0 (positive)
 the system has gained energy
from the surroundings.
endergonic
The First Law of Thermodynamics
 The decomposition of water is endergonic
(DE > 0):
2 H2O (l)
2 H2 (g) + O2 (g)
H2 (g), O2 (g)
final
E
H2O (l)
initial
Energy must be
gained from
the
surroundings.
The First Law of Thermodynamics
 If Efinal < Einitial,
 DE < 0 (negative)
 the system has lost energy to the
surroundings.
exergonic
The First Law of Thermodynamics
 The synthesis of water is exergonic (DE <
0)
2 H2 (g) + O2 (g)
2 H2O (l)
H2 (g), O2 (g)
E
initial
H2O (l)
final
Energy is lost
to the
surroundings in
this reaction.
The First Law of Thermodynamics
 The internal energy of a system can change
when energy is exchanged between the
system and the surroundings
 Heat
 Work
 The change in internal energy that occurs
can be found:
DE=q+w
Where q = heat
w = work
The First Law of Thermodynamics
 By convention:
 q = positive
Heat added to the
system
 w = positive
Work done on the system
by the surroundings
 q = negative
Heat lost by the system
 w = negative
Work done by the
system on the
surroundings
The First Law of Thermodynamics
Example: Calculate the change in internal
energy of the system for a process in
which the system absorbs 140. J of heat
from the surroundings and does 85 J of
work on the surroundings.
Given: system absorbs 140. J heat = + 140.J
system does 85 J work =
- 85 J
Find: D E
The First Law of Thermodynamics
DE=q+w
D E = +140 J + (-85 J)
D E = +55 J