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Chapter 6 –
Thermochemistry
A. The chemistry related to heat change in
chemical reactions
1. Energy –
ability to do work or produce heat
work = force x distance
w = F x d
a. Kinetic – motion
b. Potential – stored
c. Radiant – travels through space
What is Energy?
• the resources for producing usable power
• that which is needed to oppose natural attractions
• the capacity to do work or to produce heat
Forms of Energy
Potential – Energy due to
position or composition
Kinetic – Energy due to the
motion of an object
EK = ½ mv2
Energy is the capacity to do work
•
Thermal energy is the energy associated with
the random motion of atoms and molecules
•
Chemical energy is the energy stored within the
bonds of chemical substances
•
Nuclear energy is the energy stored within the
collection of neutrons and protons in the atom
•
Electrical energy is the energy associated with
the flow of electrons
•
Potential energy is the energy available by virtue
of an object’s position
2. Units – Joule – J
EK = ½
mv2
= ½ (2kg) (1 m/s)2
= 1 kg • m2
s2
= 1 J (joule)
4.184 J = 1 calories
1 CAL = 1000 cal = 1 kcal
4.184 kJ = 1 CAL = 1 kcal
mass = 2 kg
height = 1 meter
3. System vs. Surroundings
System : That part of the Universe whose change we are
going to measure.
Surroundings : Every thing else that is relevant to the change
is defined as the “surroundings”.
Heat
• energy transferred between two objects as a
result of the temperature difference between
them.
Temperature
• A measure of kinetic energy
Energy Changes in Chemical Reactions
Heat is the transfer of thermal energy between two bodies that
are at different temperatures.
Temperature is a measure of the thermal energy.
Temperature = Thermal Energy
900C
400C
greater thermal energy
B. 1st Law of Thermodynamics
• The energy of the universe is constant.
• Law of Conservation of Energy
1. Internal Energy – Energy of a system is the
sum of all the kinetic and
potential parts
E = Efinal  Einitial
 E if energy leaves system
+ E if energy enters system
2. State Function
• A function or property whose value depends
only on the present state (condition) of the
system, not on the path used to arrive at that
condition.
• Note the E of a system doesn’t depend on
how system got there
Thermodynamics
State functions are properties that are determined by the state
of the system, regardless of how that condition was achieved.
energy, pressure, volume, temperature
Potential energy of hiker 1 and hiker 2
is the same even though they took
different paths.
E is a state function, q and w are not.
3. Heat and Work
E=q+w
heat gain or loss
REMEMBER:
Ein is + (endo)
Eout is – (exo)
work done = -PV
Example #1 : Heat and Work
A system performs 50 kJ of work on its
surroundings and absorbs 20 kJ of heat from
its surroundings. What is the change in
internal energy of the system?
E=q+w
= 20 kJ + (-50 kJ)
= -30 kJ
Determining Energy Change in a System
Problem: In the internal combustion engine, the heat produced by the
combustion of the fuel causes the carbon dioxide and water that is
produced during the combustion to expand, pushing the pistons. Excess
heat is removed by the cooling system. Determine the change in energy
( E) in J, kJ, and kcal if the expanding gases do 515 J of work on the
pistons, and and the system loses 407 J of heat to the cooling system.
q = - 407 J
J
w = -515
 E = q + w = - 407 J + ( - 515 J) = - 922 J
Thermodynamics
E = q + w
E is the change in internal energy of a system
q is the heat exchange between the system and the surroundings
w is the work done on (or by) the system
w = -PV when a gas expands against a constant external pressure