Introduction to
thermochemistry
Heat, work, energy and the First
Law
Learning objectives
• Define energy and identify types of energy
• Compare and contrast heat and work
• Describe internal energy and how it changes
during a process
• Describe basic properties of state functions
• Apply first law of thermodynamics to
determine heat flow and work
• Define enthalpy
Behind it all
• Why do things happen in chemistry?
• Substances spontaneously move towards a
position of greater stability – in energy terms
• A high energy state is unstable with respect
to a state of lower energy
• A simple (but incomplete) analogy is a ball
rolling downhill
Energy
• Is capacity to perform work
• Mechanical work is application of force
over distance
• Heat is energy transferred by virtue of
temperature gradient – associated with
molecular motion
• Joule demonstrated experimentally that
heat and work are interchangeable
forms of energy
Energy: forms
• Kinetic energy is the energy of motion
1 2
E K  mv
2
• Potential energy is energy stored – by
position, within a spring, within a chemical
bond, within the particles of a nucleus
EP  mgh
Energy: units
• From the definition of kinetic energy
(1/2mv2), we get the units of energy:
kg m2/s2
• S.I. unit for energy is the joule (J) = 1Nm
• Another common unit is the calorie (cal): the
energy required to raise the temperature of 1
g of water by 1ºC
1 cal = 4.184 J
• Note the food calorie (Cal) = 1 000 cal
Interchange and conservation
• Energy in its many forms can be
changed from one to another
– A stationary ball on a hill has potential
energy (P.E.) by virtue of position but no
kinetic energy (K.E.). As it rolls down, it
gains K.E. at the expense of P.E.
Energy conservation
• There is no gain or loss:
Energy cannot be created or destroyed; it can
only be changed from one form to another
– Chemical processes involve conversion of
chemical potential energy into other forms and
vice versa
– Energy never goes away, but in some forms it is
more useful than others
– Efficient energy use means maximizing the useful
part and minimizing the useless part
Some like it hot
• Thermal energy is the kinetic energy of
molecular motion
– Temperature measures the magnitude of the
thermal energy
• Heat is the transfer of thermal energy from a
hotter to a cooler body
– Temperature gradient provides the “pressure” for
heat to flow
• Chemical energy is the potential energy
stored in chemical bonds
System and surroundings
• Any process can be divided into the SYSTEM
contained within the SURROUNDINGS
– When energy changes are measured in a chemical
reaction, the system is the reaction mixture and
the surroundings are the flask, the room, and the
rest of the universe.
Internal energy
• Internal energy is the sum of all of the types
of energy (kinetic and potential) of the
system. It is the capacity of the system to do
work
• Typically we don’t know the absolute value of
U for the system
– (Internal energy usually has symbol U. Other
sources use E)
• We can measure the change to the internal
energy
ΔU = Ufinal - Uinitial
Work and internal energy
• Work done on system increases its
internal energy
• Work done by system decreases its
internal energy
ΔU = w
Workin’ for a livin’
• Mechanical work is force applied over a
distance
W=Fxd
• In chemical process release of gas
allows work to be done by system
Work done at constant pressure
• Gas generated in reaction pushes against the
piston with force: P x A
• At constant P, volume increases by ΔV and
work done by system is:
w = -PΔV (ΔV = A x d)
– Work done by system is –ve in expansion (ΔV > 0)
• ΔU < 0 (ΔV > 0, -PΔV < 0)
– Work done by system is +ve in contraction (ΔV <
0)
• ΔU > 0 (ΔV < 0, -PΔV > 0)
Expansion work
• Work done by gas expanding:
w = -PexΔV
• In expansion the ΔV > 0; w < 0
ΔU < 0
• In contraction, ΔV < 0; w > 0
ΔU > 0
Heat and internal energy
• Heat is transfer of energy by virtue of
temperature gradient
ΔU = q
• If system is cooler than surroundings
q>0
• If system is hotter than surroundings
q<0
Deposits and withdrawals
• Process is always viewed from perspective of
system
• Energy leaving system has negative sign
– (decreases internal energy – lowers the chemical bank
balance)
• Energy entering system has positive sign
– (increases internal energy – increases chemical bank
balance)
• Useful process is one where change is negative
• Energy is in the form of heat or work
–
ΔU = q + w
First Law of Thermodynamics
Total internal energy of isolated
system is constant
– Energy change is difference between final
and initial states (ΔU = Ufinal – Uinitial)
– Energy that flows from system to
surroundings has negative sign (Ufinal <
Uinitial,)
– Energy that flows into system from
surroundings has positive sign (Ufinal >
Uinitial.)
Functions of state
• State Function
A property that depends only on present
state of the system and is independent
of pathway to that state
• Internal energy is a state function, as
are pressure, volume and temperature
Significance of state functions
• Change in state function between two
states is independent of pathway
• Given two states of a system:
– ΔU is always the same
– q and w depend on type of change
Heat and work
• Any chemical process may have
associated with it heat and work terms
• The total internal energy change will be
the sum of the contributions from each
ΔU = q + w = q - P ΔV
q = ΔU + P ΔV
• In a sealed system ΔV = 0, so q = ΔU
Cracked pots and enthalpy
• Most reactions are conducted in open
vessels where P is constant and ΔV ≠ 0
• The heat change at constant pressure is
qP = ΔU + P ΔV
• Enthalpy (H) is defined as:
H = U + PV
Heats of reaction and enthalpy
• Absolute enthalpy of system is not known
• Enthalpy change is measured
• Enthalpy change is known as heat of reaction
ΔH = qP = ΔU + P ΔV
– If reaction is exothermic and involves
expansion:
• ΔU < 0, ΔV > 0 ΔH less negative than ΔU
• Enthalpy change is portion of internal energy
available as heat after work is done
• If no work done, all the internal energy change is
enthalpy
Comparing ΔH and ΔU
• In reactions involving volume change
at constant P, ΔH and ΔU are different.
How big is it?
• Consider reaction:
C3 H 8 ( g )  5O2 ( g )  3CO2 ( g )  4H 2O( g )
• 1 additional mole of gas is produced
ΔU = - 2045 kJ, ΔH = - 2043 kJ
PΔV = + 2kJ