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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 chemical changes
happen?
• Substances spontaneously
move towards greater
stability – in energy terms
(see later how this is defined)
• High energy state is unstable
with respect to lower energy
state
• Simple (but incomplete)
analogy is ball rolling
downhill
Energy
• Is capacity to do 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
F x d
Energy: forms
• Kinetic energy is energy of motion
• Potential energy is energy stored – by
position, in spring, in chemical bond, in
nucleus
2
2
1mvEK
mghEP
Energy: units
• From definition of kinetic energy (1/2mv2), we
get units of energy:
kg m2/s2
• S.I. unit for energy is joule (J) = 1 Nm
• Another common unit is calorie (cal):
Energy required to raise temperature of 1 g
water 1ºC
1 cal = 4.184 J
• Note the food calorie (Cal) = 1 000 cal
Interchange and conservation
• Energy can be changed from one form to another – Stationary ball on 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 (all molecules at same T have
same E)
• 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 SYSTEM
contained within SURROUNDINGS
– When energy changes are measured in chemical
reaction:
– system is reaction mixture
– surroundings are flask + room + rest of universe
Internal energy (U or E)
• Internal energy is sum of all types of energy
(kinetic and potential) of system. It
measures capacity to do work
• Typically we don’t know absolute value of U
for system
– Internal energy usually has symbol U. Other
sources use E (I know, that is confusing)
• We measure change to internal energy
initialfinal UUU
Work and internal energy
• Work done on system increases its
internal energy
• Work done by system decreases its
internal energy
• ΔU and w have same sign
ΔU = w
Workin’ for a livin’
• Mechanical work is
force applied over a
distance
W = F x d
• In chemical process
release of gas
allows work to be
done (PΔV)
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)
VPw
Expansion work
• Work done by gas expanding:
w = -PexΔV • In expansion the ΔV > 0; w < 0
Internal energy decreases ΔU < 0
• In contraction, ΔV < 0; w > 0
Internal energy increases ΔU > 0
Deposits and withdrawals
• Process is always viewed from perspective of system
• Energy leaving system has negative sign – (decreases internal energy – lowers energy bank balance)
• Energy entering system has positive sign – (increases internal energy – increases energy bank balance)
• Change in internal energy due to heat and/or work
wqU
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.)
Significance of state functions
• 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
• Any function made from other state
functions is also a state function
• Change in state function between two
states is independent of pathway
• Given two states of system:
– ΔU is always the same
– q and w depend on type of change
Heat and work
• Any chemical process may have heat
and work terms
• Total internal energy change = sum of
contributions from each
• In closed system ΔV = 0, so q = ΔU
VPqwqU
VPUq
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
• Enthalpy (H) is defined as:
VPUqP
PVUH
Heats of reaction and enthalpy
• Absolute enthalpy of system is not known
• Enthalpy change is measured
• Enthalpy change is known as heat of reaction
– 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 by system to expand
• If no work done, all internal energy change is enthalpy
VPUqH P
Comparing ΔH and ΔU
• In reactions involving volume
change at constant finite P, ΔH and
ΔU are different. How big is it?
• Consider reaction:
• 1 additional mole of gas is produced
• Work done by system, w < 0
ΔU = - 2045 kJ, ΔH = - 2043 kJ
PΔV = + 2kJ (P = 1 atm, ΔV = 20 L)
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