Unit 1: Thermochemistry
 Introduction and Some Definitions
 Internal Energy
 The First Law of Thermodynamics
 Enthalpy and Enthalpy Changes
 Calorimetry
 Hess’s Law
 Using Enthalpy’s of Formation
Thermochemistry – Some Definitions
 Most daily activities involve processes that
either use or produce energy:
 Metabolism of food
 Burning fossil fuels
 Photosynthesis
 Pushing a bike up a hill
 Thermodynamics:
 the study of the energy and its
transformations
Thermochemistry – Some Definitions
 Thermochemistry:
 A branch of thermodynamics
 the study of the energy absorbed or
released as heat during a chemical
reaction or process
 Objects (including chemicals) can have two
types of energy:
 Kinetic energy
 Potential energy
Thermochemistry – Some Definitions
 Kinetic energy
 Energy of motion
 Thermal energy
 a type of kinetic energy a substance
possesses because of its
temperature.
 Potential energy
 “stored” energy or energy of position
 Energy that results from attractions and
repulsions between objects
Thermochemistry – Some Definitions
 Chemical energy:
 A type of potential energy stored within a
substance
 Results from electrostatic forces between
charged particles within the substance as
well as from the arrangement of the
atoms (or ions) within the substance
Thermochemistry – Energy Units
 Units of Energy:
 SI unit = joule (J)
 A very small quantity
 ~ the energy required to lift a small
apple one meter into the air
 Kilojoule (kJ)
 1 kJ = 1000 J
Thermochemistry – Energy Units
 Calorie (cal)
 Originally defined as the amount of energy
needed to raise the temperature of 1 g of
water from 14.5oC to 15.5oC
 1 cal = 4.184 J (exactly)
 Kilocalories (kcal)
 1 kcal = 1000 cal = 1 Cal (food calorie)
 British Thermal Unit (BTU)
 1 BTU = 1054.35 J
Thermochemistry – Energy Units
 You are responsible for knowing the
conversion factors shown in red on the
previous slides.
 You must be able to use dimensional analysis
to convert from one energy unit to another.
Thermochemistry – Energy Units
Example: Convert 7.63 kcal to BTU using the
relationships from the previous slides.
Thermochemistry – Energy Units
Example: A particular furnace produces
9.0 x 104 BTU/hr of heat. Use dimensional
analysis to calculate the number of kcal of
heat delivered by the furnace after running
for 2.50 hours.
Thermochemistry – More Definitions
 When studying the amount of heat gained or
lost during a process or reaction, chemists
focus on a limited, well-defined part of the
universe to study:
 System:
 The part of the universe singled out for
study
 Typically the chemicals involved in the reaction
or process
 Surroundings:
 Everything else
Thermochemistry – More Definitions
The system is
usually the
chemicals in the
flask/reactor.
The system
The flask and
everything else
belong to the
surroundings.
Thermochemistry – More Definitions
 A system can be either open or closed.
 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.
Internal Energy
 The internal energy (E) of a system is the
sum of the kinetic and potential energy of all
components of a system.
 For the molecules in a chemical system, the
internal energy includes:
 The motion and interactions of all of the
molecules
 The motion and interactions of the nuclei
and electrons found in the molecules
Internal Energy
 Internal energy (E) is an extensive property.
 Depends on the mass of the system
 Internal energy (E) is a state function.
 A property of the system that is
determined by specifying its condition or
state in terms of T, P, location, etc.
 Depends only on its present condition
and not how it got there
Heat and Work
 The internal energy (E) of a system can
change when the system gains energy
from or loses energy to the
surroundings as either heat (q) or work
(w).
 Work (w):
 Energy used to move an object
against a force
 Lifting your backpack
 Hitting a baseball with a bat
Heat and Work
 Heat (q):
 Energy used to increase the temperature
of an object
 The energy transferred from a hotter
object to a colder one
 Energy:
 The capacity (ability) to do work or
transfer heat
Heat & Work – Sign Conventions
 Energy can be transferred between the
system and the surroundings as either heat
or work.
 Energy gained by the system is always
designated using a positive sign.
 A reaction or process in which the system
gains heat from the surroundings is
endothermic.
system
(+)
(+)
q
w
Heat & Work – Sign Conventions
 Energy lost by the system is always
designated using a negative sign.
system
q
(-)
w
(-)
 A reaction or process in which the system
loses heat to the surroundings is exothermic.
Internal Energy Changes
 The change in the internal energy (DE) that
occurs when energy is gained from or lost to
the surroundings:
D E = Efinal - Einitial
DE = change in internal energy
Efinal = final energy of system
Einitial = initial energy of system
Internal Energy Changes
 A reaction or process that experiences a net
gain of energy from the surroundings is
referred to as endergonic.
 Efinal > Einitial
 DE > 0 (positive)
Internal Energy Changes
 A reaction or process that experiences a net
loss of energy to the surroundings is
referred to as exergonic.
 Einitial > Efinal
 DE < 0 (negative)
Internal Energy Changes
 From a practical perspective, the change in
internal energy (DE) of the system is found
by measuring the amount of heat gained or
lost by the system and the amount of work
done on or by the system:
DE = q + w
Where q = heat
w = work
Be sure to use the correct sign for q and w!
Internal Energy Changes
Example: Calculate the change in internal
energy of the system for a process in which
the system absorbs 197 J of heat from the
surroundings while doing 73 J of work.
Internal Energy Changes
Example: Calculate DE for a system when the
system loses 72 kJ of heat while the
surroundings do 193 kJ of work on the system.
First Law of Thermodynamics
 Energy can be transferred between the
system and the surroundings as heat and/or
work.
 Energy can also be converted from one form
to another.
 Kinetic energy  Potential energy
 First Law of Thermodynamics: Energy
can be converted from one form to
another, but it cannot be created or
destroyed.
 Any energy lost by the system must be
gained by the surroundings and vice versa.
Enthalpy
 Many reactions or chemical processes occur
in open containers (i.e. at constant
pressure).
 The amount of heat gained or lost under
constant pressure conditions (qp) is often
referred to as the enthalpy change (DH)
 Enthalpy (H):
 An extensive property (one that depends
on the amount of substance present) that
is defined by the equation H = E + PV
Calculating the Amount of Work
 Two common types of work done by chemical
systems:
 Electrical work
 Redox reaction incorporated in galvanic
cells (Unit 5)
 Mechanical work (P-V work)
 Work done by expanding gases
 For example, the expanding gases in
a cylinder of a car engine
Calculating the Amount of Work
 The amount of P-V work done at constant
pressure can be found:
 w = -P DV
where P = pressure
DV = change in volume
 The negative sign indicates that work is
being done by the system
DE = DH – P DV
 For a process occurring at constant pressure
in which the only work done is PV work,
DE = qp + w
DH = qp
w = -P DV
DE = DH – P DV
(at constant pressure)
DE = DH – P DV
Example: If the volume of a cylinder increases
from 2.13 L to 3.50 L at a constant pressure
of 1.33 atm while it absorbs 2.515 kJ, what is
the change in internal energy of the system?
(Note: 1 atm.L = 101.3 J)