Topic 5 - Thermodynamics (Part 1) Lecture

GENERAL CHEMISTRY
2
Topic 5 – Thermodynamics (Part 1)

What is thermodynamics?
• Thermodynamics can be defined as
the study of energy.
• Chemical thermodynamics - The
chemistry that deals with energy
exchange, entropy, and the
spontaneity of a chemical process.
Concepts associated with
thermodynamics
• Temperature
• Heat
• Enthalpy
• Endothermic
• Exothermic
• Entropy
• Gibbs free energy

Thermodynamics vs. Kinetics
• Kinetics - Rate of a reaction
depends on the pathway from
reactants to products.
• Thermodynamics tells us whether
a reaction is spontaneous based
only on the properties of reactants
and products.

Objectives
Students should be able to:
• Describe the term ‘energy’ and state forms of energy
• Describe the term ‘state functions’
• Distinguish between the terms ‘temperature’ and ‘heat’
• Distinguish between the terms ‘system’ and ‘surroundings’
• Distinguish between the terms ‘exothermic’ and endothermic’
• State and explain the first law of thermodynamics
• Define internal energy
• Calculate the change in internal energy of a chemical process or reaction.

What is Energy?
• Energy is at the center of our very existence as individuals and as a society
• The food that we eat furnishes the energy to live and work
• Most of our energy have come from carbon-based fossil fuel which was a
readily available and abundant source.
• However we have moved from a period of ample and cheap supplies of
petroleum to one of high prices and uncertain supplies.
• Alternative sources of energy are now being sort and the relationship
between chemistry and energy is being used to find alternatives to fossil fuels

The nature of energy
• Energy can be defined as the ability to do work or produce heat
• Work: Energy used to cause an object that has mass to move.
• Heat: is a measure of the flow of energy due to temperature
differences between objects. Heat always flows from hot to cold.

Topic 5 - Thermodynamics (Part 1) Lecture

Classification
• Energy can be classified as either potential or kinetic energy.
• Potential energy is energy an object possess by virtue of its position or chemical
composition.

• Kinetic energy is the energy an object
possesses by virtue of its motion. It
depends on the mass and velocity of
the object according to the equation:
• KE = ½ mv2
• Law of conservation of energy -
Energy can be converted from one
form to another but can be neither
created or destroyed (first law of
thermodynamics)
• Describe the changes in energy that
takes place as a ball is dropped from
the top of the building.

Intro to energy transformations

Units of energy
• Remember kinetic energy = 1/2mv2
. Energy has units of (mass)
(velocity)2
• The SI unit of energy the joule (J). 1J = 1 Kgm2
/ s2
(1Kgm2
s-2
)
• An older, non SI unit is still in widespread use: The calorie (cal)
‐
• The calorie is defined as the amount of energy (heat) required to
raise the temperature of one gram of water by one Celsius
degree. (1 cal = 4.184 J)

State functions
Look at the following videos
• http://www.youtube.com/watch?v=lL2CXpVti34
• http://www.youtube.com/watch?v=SEapOFaYDlA
Questions:
• Define the term state function
• What are some examples of state functions?
• What are 2 non-state/path functions in thermodynamics?
• What are the 4 state functions that thermodynamics is concerned with measuring?

State Functions
• “State functions are values that depend on the state of the substance, and not on how
that state was reached. For example, density is a state function, because a substance's
density is not affected by how the substance is obtained.
• Consider a quantity of H2O: it does not matter whether that H2O is obtained from the tap,
from a well, or from a bottle, because as long as all three are in the same state, they
have the same density.
• When deciding whether a certain property is a state function or not, keep this rule in
mind: is this property or value affected by the path or way taken to establish it? If the
answer is no, then it is a state function, but if the answer is yes, then it is not a state
function.
• "State" refers to temperature, pressure, and the amount and type of substance present”
Chem.libretexts.org

Analogy
• The main point to remember when trying to identify
a state function is to determine whether the path
taken to reach the function affects the value. The
analogy below illustrates how to tell whether a
certain property is a state function.
• The change in vertical distance (∆y) between the 1st
floor and 2nd
floor in a building stays the same
whether you take the stairs or the elevator. As a
result, ∆y is a state function because its value is
independent of the path taken to establish its value.
• In the same situation, time, or ∆t, is not a state
function. If someone takes the longer way of getting
to the 2nd
floor (climbing the stairs), ∆t would be
greater, whereas ∆t would be smaller if the elevator
is taken. In this analogy, ∆t is not a state function
because its value is dependent on the path.

Temperature vs Heat
What is the difference?
• Temperature is a measure of the random
motions of the components of a
substance. This means that the H2O
molecules in warm water are moving
around more rapidly than the H2O
molecules in cold water.
• Heat can be defined as a measure of the
flow of energy due to a temperature
difference. It is a measure of the total
potential and kinetic energy transferred
between objects.

System vs Surroundings
• The system includes the molecules
we want to study (In this picture the
hydrogen and oxygen molecules)
• The surroundings are everything else
that the system touches and can
exchange energy with
• The piston and cylinder are the
boundary between the system and
surroundings and can be considered
part of the surroundings

Endothermic vs Exothermic
processes
• An exothermic process is one where
the system releases heat to the
surroundings. E.g the combustion of
a match, energy flows out of the
system as heat.
• An endothermic process is one
where the system absorbs energy
from the surroundings. Eg boiling
water to form steam.

First law of thermodynamics
• The first law of thermodynamics, also known as Law of Conservation of
Energy, states that energy can neither be created nor destroyed; energy can
only be transferred or changed from one form to another.
• ΔESystem = Δ
‐ Esurroundings

Internal energy
• The internal energy of a system is the sum of all kinetic and
potential energies of all components of the system; we call it E or U.
• Etotal internal = EKE + EPE + Eelectrons + Enuclei +……
• It is almost impossible to calculate total internal energy
• Instead we always look at the change in energy (ΔE).

Change in Internal energy
• By definition, the change in internal energy, ΔU, is the final energy of the system minus
the initial energy of the system:
• ΔU = Ufinal − Uinital
• Where Δ means change
• ΔU is also used to represent change in internal energy. Change in internal energy is a
state function.

Changes in Internal Energy
• When energy is exchanged between
the system and the surroundings, it
is exchanged as either heat (q)
and/or work (w).
• That is, ΔU = q + w.
• ΔU represents the change in internal
energy of the system.
• W= P ΔV (for gaseous systems)

Thermodynamic Quantities and
Signs

Processes that can occur in a system
Type of Process Characteristics Resulting
Equation
Isovolumetric (constant
volume)
Δ V = 0, this means the volume is constant, hence W = 0 (no
work is being done by or on the system)
Any change in ΔU is due to heat transfer only
ΔU = Q
Isothermal (constant
temperature)
ΔU = 0, this means the internal energy is constant, hence any
heat added to the system is used completely to do work by the
system
Q = -W
Adiabatic (constant
heat)
Q = 0, this means there is no heat transfer hence any change
in internal energy is as a result of work done by or on the
system
ΔU = W
Isolated System Q = 0 and W = 0, this means there is no heat transfer or work
done hence there is no change in internal energy
ΔU = 0

1st
law of thermodynamics Equation
• About half of textbooks, teachers, and professors write the first law of
thermodynamics as ΔU=Q+W and the other half write it as ΔU =Q-W.
• Both equations are correct, and they say the same thing.
• The reason for the difference is that in the formula ΔU=Q+W, we are assuming
that ’W’ represents the work done on the system and it is given a + sign. When
we use ΔU=Q−W, we are assuming that ’W’ represents the work done by the
system and this is also given a + sign.
• For this course, we will be using the equation, ΔU=Q+W
• W = P ΔV (for gaseous systems), where P = pressure and ΔV = change in
volume

1st
law of Thermodynamics
Calculations

Practice Question
A gas in a system has constant pressure. The surroundings around the system
lose 62 J of heat and does 474 J of work onto the system. What is the change
in internal energy of the system?
Solution

Practice Question
A container has a sample of nitrogen gas and a tightly fitting movable piston that
does not allow any of the gas to escape.
During a thermodynamics process, 200 J of heat enter the gas, and the gas
does 300 J of work in the process.
• Calculate the change in internal energy of the system.
• What will happen to the temperature of the nitrogen gas in system?

Practice Question
Four identical containers have equal amounts of helium gas that all start
at the same initial temperature. Containers of gas also have a tightly
fitting movable piston that does not allow any of the gas to escape. Each
sample of gas is taken through a different process as described below:
• Calculate the change in internal energy for each sample of helium
• Rank the samples from highest temperature to lowest temperature after they’ve
gone through the processes given above.

Worked Example #2
A system has constant volume (ΔV=0) and the heat around the system
increases by 45 J.
1.What is the sign for heat (q) for the system?
2.What is ΔU equal to?
3.What is the value of ΔU for the system in Joules?

Practice Question
For a process at constant volume,
A) q = 0, w = 0, and ΔU = 0.
B) w = 0 and ΔU = q.
C) w = 0 and ΔH = q.
D) w = 0 and ΔU = ΔH.

Practice Question
Calculate the work energy, w, gained or lost by the system when a gas
expands from 15 L to 35 L against a constant external pressure of 1.5
atm. [1 L ∙ atm = 101 J]
A) -5.3 kJ
B) -3.0 kJ
C) +3.0 kJ
D) +5.3 kJ

Topic 5 - Thermodynamics (Part 1) Lecture

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