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Thermodynamics and Thermochemistry
Thermodynamics and Thermochemistry
Thermodynamics -- the science of energy and its transformations
Thermochemistry -- the branch of thermodynamics specifically
focused on the changes in energy and transfer of heat related to
chemical reactions
On a basic level, you can think of thermodynamics as an
accounting system for keeping track of changes in energy
associated with physical and chemical processes
System and Surroundings
System and Surroundings
In thermodynamics, the system is a limited, well-defined portion
of the universe on which we focus in order to keep track of the
energy changes that occur
Example: A heated block of copper is placed in a beaker of water
System: The block of copper
surroundings
system
surroundings
Surroundings: The water, the
beaker, and everything outside
Everything else in the universe outside the boundaries of the
system is called the surroundings
System and Surroundings
Example: Hydrogen and oxygen gases confined in a cylindrical
container with a moveable piston for a lid
System and Surroundings
Example: Water in a sealed, insulated vessel (calorimeter)
equipped with a heating element
System: The H2 and O2 molecules
inside the container
System: The water
Surroundings: The walls of the
c o n t a i n e r, t h e p i s t o n , a n d
everything outside
Surroundings: The heating element,
t h e r m o m e t e r, c a l o r i m e t e r, a n d
everything outside
Different types of systems
open system -- a system in which energy and matter can be
exchanged with the surroundings
System and Surroundings
Example: A beaker of water boiling on a hot plate
closed system -- a system in which energy can be exchanged
with the surroundings, but matter can not be exchanged with
the surroundings
System: The water
isolated system -- a system in which neither energy nor
matter can be exchanged with the surroundings
Surroundings: The beaker, the
hot plate, and everything outside
System and Surroundings
Example: Hydrogen and oxygen gases confined in a cylindrical
container with a moveable piston for a lid
OPEN SYSTEM
System and Surroundings
Example: Water in a sealed, insulated vessel (calorimeter)
equipped with a heating element
System: The H2 and O2 molecules
inside the container
System: The water
CLOSED SYSTEM
ISOLATED SYSTEM
Surroundings: The heating element,
t h e r m o m e t e r, c a l o r i m e t e r, a n d
everything outside
Surroundings: The walls of the
c o n t a i n e r, t h e p i s t o n , a n d
everything outside
Internal energy
Internal energy (E): the total energy of a thermodynamic system
-- i.e., the sum of all the different forms of energy contained by
all components of the system
Internal energy comprises two broad categories:
kinetic energy (molecular) -- translational, rotational, and
vibrational motions of the atoms and molecules making up the
system
• i.e., thermal energy
Internal energy
It is usually not possible or practical to calculate an exact value
for the internal energy of a system (E)
But it usually is feasible (and more interesting) to calculate the
change in the internal energy of a system ( E)
Initial system conditions
Final system conditions
E initial
E final
potential energy -- binding energy of the atoms and molecules
that make up the system
• primarily chemical bonds (covalent, ionic)
• also intermolecular forces, binding energy of nuclei, etc.
E = E final – E initial
thermodynamics focuses on changes in the energy of a system
Changes in internal energy
When the internal energy of a system increases (E final > E initial),
the change in energy ( E) is positive
Changes in internal energy
Law of Conservation of Energy: Energy can be neither created
nor destroyed (i.e., energy is conserved)
If the internal energy of a system increases, the system must
gain energy from the surroundings
E = E final – E initial
positive value
Energy is transferred to the system
from the surroundings
E is
positive
When the internal energy of a system decreases (E final < E initial),
the change in energy ( E) is negative
E = E final – E initial
negative value
If the internal energy of a system decreases, the system must
lose energy to the surroundings
Energy is transferred from the system
to the surroundings
E is
negative
Methods of energy transfer: Heat and work
Energy can be transferred to/from a system by two methods:
Energy transfer via work
Example: Compressing a gas in an insulated container
(i.e., no heat transfer)
work (w) -- movement of an object against a force
• mechanical work (primarily related to expansion of gases)
• electrical work
heat (q) -- random motion of atoms/molecules
• heat transfer occurs when a temperature difference exists
work (w)
E
heat (q)
system
Energy transfer via work
Example: A pressurized gas in an insulated container is
allowed to expand (i.e., no heat transfer)
Work is performed on the system (work is positive)
• Energy is transferred to the system
• Energy of the system increases ( E is positive)
Energy transfer via heat
Example: A bunsen burner is used to heat water in a
sealed container with fixed walls (i.e., constant volume)
T = 20 °C
Work is performed by the system (work is negative)
• Energy is transferred from the system
• Energy of the system decreases ( E is negative)
T = 50 °C
Heat is transferred to the system (heat is positive)
• Energy is transferred to the system
• Energy of the system increases ( E is positive)
Energy transfer via heat
First law of thermodynamics
Example: Propane is combusted in a sealed, uninsulated
container with fixed walls (i.e., constant volume)
C3H8(g)
The first law of thermodynamics is a consequence of the law
of conservation of energy
When a system undergoes any chemical or physical change:
The change in internal energy is equal to the heat transferred
to/from the system plus the work performed on/by the system
O2(g)
CO2(g)
E = q+w
H2O(g)
heat
work
• internal
Heat is transferred from the system (heat is negative)
• Energy is transferred from the system
• Energy of the system decreases ( E is negative)
energy increases when heat is added to a system or
work is done on a system
• internal energy decreases when heat is removed from a system
or work is done by a system
Banking analogy
Summary of sign conventions
The system is like a bank where energy constitutes the assets
• deposits and withdrawals can be made in the form of heat or work
positive (+)
negative (–)
Deposits
Withdrawals
SYSTE
M
Heat (q)
system gains
heat
system loses
heat
heat (+)
heat (–)
Work (w)
work done on
system
work done by
system
work (+)
work (–)
Change in
Internal energy Internal energy
internal energy
of system
of system
( E)
increases
decreases
For any transaction:
Net change in assets ( E) =
Propane and oxygen are burned in a cylindrical container with a freefloating piston for a lid.
The expansion of the product gasses
performs 560 J of work on the surroundings. The container loses
1830 J of heat to the surroundings. What is the change in the
internal energy of the chemical substances in the container?
deposits /
withdrawals
of heat (q)
+
When a sample of nitrogen is heated, it absorbs 125 J of heat and
performs 32 J of work on its surroundings due to its expansion What
is the change in the internal energy of the nitrogen?
E = q+w
E = q+w
q = 125 J
q = –1830 J
w = –32 J
w = –560 J
E = (125 J) + (-32 J)
E = (–1830 J) + (-560 J)
E = 93 J
E = –2390 J
deposits /
withdrawals
of work (w)
A solid block of copper is transferred from an oven to a beaker of cold
water. As it cools, the block of copper loses 785 J of heat to the
water. What is the change in the internal energy of the copper?
A balloon filled with helium is heated and expands, performing 370 J
of work on the surrounding gasses in the atmosphere? The internal
energy of the helium increases by 1220 J. How much heat did the
helium absorb?
E = q+w
E = q+w
q = –785 J
w=0J
w = –370 J
E = (–785 J) + (0 J)
E = 1220 J
1220 J = q + (–370 J)
E = –785 J
1220 J + 370 J = q
1590 J = q
Thermodynamics definitions
Thermodynamics definitions
System: A limited, well-defined portion of the universe
Surroundings: Everything else in the universe outside the
boundaries of the system
State: The condition of a system as defined by a specific set
of measurable, macroscopic properties (e.g., temperature,
pressure, etc.)
• the properties of a state are uniquely determined (i.e., they
have one uniform value throughout the entire system)
• the properties of a state remain constant over time (i.e., the
In thermodynamics, we keep track of:
system is in equilibrium)
• changes in the energy of the system
• the transfer of energy (as work and heat)
Process: A thermodynamic process leads to a change in the
thermodynamic state of a system
between the system and the surroundings
• physical process (e.g., compressing a gaseous system)
• chemical process (e.g., reaction between two components
of a system)
Thermodynamics states and processes
Thermodynamics states and processes
Example: Isothermal compression of 0.25 mole of nitrogen gas
(assuming ideal gas behavior)
Example: Combustion of propane and oxygen to form carbon
dioxide and water vapor under constant pressure
System: The N2 gas molecules in the container
System: The C, H, and O atoms that make up the
reactant and product molecules
Initial state
Final state
Initial state
Final state
P1 = 1.0 atm
P2 = 8.2 atm
P1 = 1.0 atm
P2 = 1.0 atm
V1 = 6.1 L
V2 = 3.0 L
V1 = 2.0 L
V2 = 5.0 L
T1 = 25 °C
T2 = 25 °C
T1 = 25 °C
T2 = 40 °C
Thermodynamics definitions
State function: A property of a system that depends only on
the current state of the system, not on the path that the
system took to reach that state
• state functions are also referred to as state variables
• state functions are path-independent
• changes in state functions depend only on the initial and
State function analogy:
Latitude, longitude, and altitude
Whenever you are in Denver, your “state” is defined by Lat = 39°43’N,
Lon = 105°01’W, and Alt = 5280 ft.
It doesn’t matter how you got there -- if the “state” of your “system” is
being in Denver, the latitude, longitude, and altitude will have those values
final states of the system, not on how the change occurred
(i.e., changes in state functions are also path-independent)
Denver
Latitude, longitude
and altitude are
path-independent
Denver
Latitude: 39°43’N
Longitude: 105°01’W
Altitude: 5280 ft
State function analogy:
Latitude, longitude, and altitude
State function analogy:
Latitude, longitude, and altitude
Consider a trip from Chicago to Denver:
Consider a trip from Chicago to Denver:
Regardless of your route, the difference in latitude, longitude and altitude
will be the same (i.e., path-independent)
But total distance travelled will depend on the route you take
Lat = 39°43’ – 41°51’ = –2°18’
Lon = 105°01’ – 87°38’ = 17°23’
Alt = 5280 ft – 656 ft = 4684 ft
Chicago
Latitude: 41°51’N
Longitude: 87°38’W
Altitude: 596 ft
Chicago
Denver
Denver
Latitude: 39°43’N
Longitude: 105°01’W
Altitude: 5280 ft
Albuquerque
Memphis
Changes in
latitude, longitude
and altitude are
path-independent
State functions of thermodynamic systems
Example: For a system consisting of an ideal gas, pressure, volume
and temperature are state functions (i.e., path-independent)
Chicago
Latitude: 41°51’N
Longitude: 87°38’W
Altitude: 596 ft
Chicago
Distance travelled is
not path-independent
Denver
Denver
Latitude: 39°43’N
Longitude: 105°01’W
Altitude: 5280 ft
Albuquerque
Memphis
State functions of thermodynamic systems
For any thermodynamic system:
Internal energy is a state function
compression
heat loss
heating
expansion
Distance travelled (Direct route) = 996 miles
Distance travelled (Scenic route) = 1957 miles
• E and
E are path-independent
Work performed and heat transferred are not state functions
• w and q are not path-independent
Initial state
Final state
Initial state
P = 1.0 atm
P = 1.0 atm
P = 0.80 atm
V = 2.0 L
V = 3.0 L
V = 6.0 L
T = –75 °C
T = 25 °C
T = 225 °C
Internal energy, work and heat
Example: Expansion of an ideal gas following two different paths
Path A
Internal energy, work and heat
Example: Expansion of an ideal gas following two different paths
Path B
heating
under
constant
pressure
cooling
while keeping
volume
constant
cooling
while keeping
volume
constant
heating
under
constant
pressure
q = 10130 J
w = – 4050 J
q = – 4560 J
w=0J
q = – 1515 J
w=0J
q = 5060 J
w = – 2025 J
Initial state
P = 2.00 atm
V = 10.0 L
T = –29.4 °C
Intermediate state
P = 2.00 atm
V = 30.0 L
T = 458 °C
Final state
P = 1.00 atm
V = 30.0 L
T = 92.4 °C
Internal energy, work and heat
Initial state
P = 2.00 atm
V = 10.0 L
T = –29.4 °C
Intermediate state
P = 1.00 atm
V = 10.0 L
T = –151 °C
Final state
P = 1.00 atm
V = 30.0 L
T = 92.4 °C
State functions of thermodynamic systems
Example: Expansion of an ideal gas following two different paths
SUMMARY OF MAIN POINTS
Path A
q = (10130 J) + (– 4560 J) = 5570 J
State functions are path-independent
w = (– 4050 J) + (0 J) = – 4050 J
E = (5570 J) + (– 4050 J) = 1520 J
Changes in state functions are also path-independent
• the change in a state function only depends on the
initial and final states of the system
Path B
q = (– 1515 J) + (5060 J) = 3545 J
w = (0 J) + (– 2025 J) = – 2025 J
E = (3545 J) + (– 2025 J) = 1520 J
E is path-independent
q and w are not path-independent
• the change in a state function does not depend on
the process by which the system moved from the
initial state to the final state