Transcript Ch. 6 Section 6.1 Powerpoint

THERMOCHEMISTRY
Courtesy of lab-initio.com
Definitions #1
•Energy: The capacity to do work or produce heat.
•Potential Energy: Energy due to position or
composition.
•Kinetic Energy: Energy due to the motion of the
object and depends on its mass and velocity.
1 2
KE  mv
2
Definitions #2
•Law of Conservation of Energy:
Energy can neither be created nor
destroyed, but can be converted
between forms.
•The First Law of
Thermodynamics: The total
energy content of the universe is
constant.
State Functions refers to a property of the system
that depends only on its present state.
ENERGY IS A STATE
FUNCTION
A person standing at the top of
Mt. Everest has the same
potential energy whether they
got there by hiking up, or by
falling down from a plane 
WORK and HEAT ARE NOT A
STATE FUNCTIONS
WHY NOT???
•Work: defined as force acting over a distance.
•Heat: involves the transfer of energy between
two objects due to a temperature difference.
•Remember heat and temperature are different.
•Temperature is a measure of the hotness or
coldness of something and is proportional to the
average molecular kinetic energy of the atoms,
molecules, or ions present.
•System: The portion of the universe under study.
•Surroundings: Everything else besides the system.
•Interaction: Exchange of energy and or matter between the
system and its surroundings.
•Systems:
•Open System: Exchanges both matter and energy with
its surroundings.
•Closed System: Exchanges only energy with its
surroundings.
•Isolated System: Exchanges neither energy nor matter
with its surroundings.
•Internal Energy (E) of a system = sum of the PE + KE of all
“particles” of the system.
•The joule (J) is the SI unit for energy.
•J = kg•m2 /s2
•Can be changed by a flow of work, heat, or both.
ΔE = q + w
ΔE = change in system’s internal energy
q = heat
w = work
Thermodynamic quantities = two parts: a number, giving
the magnitude of the change, and a sign, indicating the
direction of flow.
q = heat flowing into or out of the system
-q if energy is leaving to the surroundings
+q if energy is entering from the surroundings
w = work done by, or on, the system
-w if work is done by the system on the
surroundings
+w if work is done on the system by the
surroundings
•A common type of work associated with chemical
processes is work done by a gas (through expansion) or
work done to a gas (through compression).
•For example, in an automobile engine, the heat from the
combustion of the gasoline expands the gases in the
cylinder to push back the piston, and this motion is then
translated into the motion of the car.
•The gases are expanding so the volume is increasing
therefore, the change in volume (ΔV) is positive. Work and
volume must have opposite signs.
Work, Pressure, and Volume
w   PV
Expansion
+V (increase)
-w results
Esystem decreases
Work has been done by
the system on the
surroundings
Compression
-V (decrease)
+w results
Esystem increases
Work has been done on
the system by the
surroundings
Exothermic: energy flows out of the system.
Which has lower potential energy? Reactants or products?
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy (heat)
Total energy is conserved = energy gained by the
surroundings must be equal to the energy lost by
the system.
In the case above, the heat flow into the
surroundings results from a lowering of the PE of
the reaction system.
In an exothermic process, the bonds in the products are
stronger (on average) than those of the reactants.
More energy is released by forming the new bonds than is
consumed in breaking the bonds in the reactants.
In any exothermic reaction, some of the PE stored in the
chemical bonds is being converted to thermal energy via
heat.
Endothermic process = reverse situation
Products have higher PE (weaker bonds on average) than the
reactants.
Energy flows into system as heat to increase the PE of the
system.