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MODULE 5
Energy and
Thermodynamics
Thermodynamics & Energy
• Thermodynamics - The science of heat
and work
• Work - A force acting upon an object to
cause a displacement
• Energy - The capacity to do work and
transfer heat
Kinetic Energy
• Kinetic Energy
– KE = 1/2 mv2
– The energy of a moving object,
depending on it’s mass and velocity
– “Energy of motion”
Potential Energy
• Potential Energy
– PE = mgh
– The energy that results in an object’s
position
– “Stored energy”
PE, KE and Work
Internal Energy
• Internal energy () - The sum of the
potential energies and kinetic energies
of the particles within a thermodynamic
system
ETOTAL = PE + KE + 
First Law of Thermodynamics
• Law of Conservation of Energy
– The total energy of the universe is
constant
– Heat, work, and other energy transfers in
an event equal the total energy content
both before and after the event has
occurred
– A battery stores chemical potential energy
Measurement
• 1 cal = 4.184J
• 1000cal = dietary calorie = 1kcal
• 1 calorie = amount of energy required
to raise 1 gram of H2O, 1°C.
Temperature & Heat
• Heat is not the same as temperature
• The more thermal energy, the more kinetic
energy, the more motion the atoms and
molecules will have
• The total thermal energy of an object is the
sum of all the individual energies
• Thermal energy depends on the amount of
substance as well as the temperature
• Temperature changes are measured with a
thermometer by heat transfer
Heat Transfer
• Occurs when 2 objects of different
temperatures are brought into contact
• Heat is transferred from the hotter object to
the colder one
• Transfer will continue until the 2 objects are
at the same temperature, we call the system
at thermal equilibrium
• The amount of heat lost by the hotter one =
the amount of heat gained by the colder one
Heat Transfer
• Exothermic – process where heat is
transferred from a system to it’s
surroundings
• Endothermic – process where heat is
transferred to the system from it’s
surroundings
Specific Heat
The quantity of heat transferred
depends on:
– The amount of material
– The overall change in temperature
– The identity of the material transferring the
energy
Specific Heat Capacity
• Specific heat is the quantity of heat
required to raise the temperature of 1
gram of a substance by one kelvin
• (See Table 6.1 on page 210)
q = C x m x T
Heat Calculations
Calculate the heat absorbed by 15.0g of
water required to raise the temperature
from 20°C to 50°C. Where q=C·m·∆T
Let q = heat = unknown
C = Heat Capacity for H2O = 4.184J/gK
m = mass = 15.0g
∆T = Tf – Ti = 50°C-20°C = 323K-293K = 30K
q = (4.184J/gK)(15.0g)(30K) = 1.88x103J=1.88kJ
Exothermic
•
•
•
•
Heat is given off
Q<0
Enthalpy (H) is negative (-)
Energy of the products is less than the
energy of the reactants
• The balanced equation is written:
4Fe(s) + 3O2(g)  2Fe2O3(s): H = -1648kJ
Endothermic
•
•
•
•
Heat is absorbed
Q>0
Enthalpy (H) is positive (+)
Energy of the products is more than the
energy of the reactants
• The balanced equation is written:
2Fe2O3(s)  4Fe(s) + 3O2(g) : H = +1648kJ
Standard Enthalpies
• See Table 20 on page A-31
Hess’s Law
• Hess’s Law states: If a reaction is the
sum of 2 or more other reactions, the
H for the overall reaction is the sum of
all the H values of those individual
reactions.
• See CD-ROM Screen 6.17
• Calculate H for the lab data.
Enthalpy Change for a Rxn.
• The enthalpy change for a reaction can be
calculated by the sum of the products H
values minus the sum of the reactants H
values.
• See Screen CD-ROM 6.18
• See Table 20, Appendix L starting on page
A.31
• Hrxn = [Hf (products)] -[Hf (reactants)]