Download Thermochemistry

Thermochemistry
Thermodynamics
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Study of energy transformations
Thermochemistry is a branch of
thermodynamics which describes energy
relationships in chemical reactions
Energy
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Capacity to do work or to transfer heat
Mechanical work (w) is the product of force (F)
operating on an object and the distance (d)
through which it moves
W=Fxd
Energy is required to do work
Heat (Q)
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Heat is the energy transferred from one object
to another due to a difference in temperature
Forms of Energy
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Kinetic Energy – energy of motion
- magnitude depends on the mass of the
object and its velocity
- EK = ½ m v2
- both mass and speed determine how
work it can do
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Potential Energy – stored energy
Other forms of energy are simply types of
kinetic or potential on an atomic or molecular
level
Energy Units
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Joule (J)
1J = 1 kg m2/ s2
A calorie (cal) is the amount of energy required
to raise the temp of 1 g of water 1 ºC
1 cal = 4.184 J
Example
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A 145 g baseball is thrown with a speed of
25 m/s. Calculate the kinetic energy in Joules.
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What is the kinetic energy in calories?
Systems
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Portion we single out for study
Surroundings is everything else outside the
system
When studying energy changes in a chemical
reaction, the reactants and products are the
system and everything else is the surroundings
Law of Conservation of Energy
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Energy can be converted from one form to
another but cannot be created or destroyed
Also called “First Law of Thermodynamics”
Internal Energy
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Total energy of a system – sum of kinetic and
potential energies
Cannot determine exact internal energy
Can only determine a change in internal
energy
ΔE = Efinal – Einitial
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If ΔE is positive there is a gain in internal energy in the
system
If ΔE is negative the system lost energy to its
surroundings
Higher energy systems tend to lose energy and are
therefore less stable
Heat and Work
Any system can exchange energy with
surroundings in two ways – as heat or work
 Internal energy increases as heat is added to
or work is done on a system
ΔE = Q + w
Q is positive if heat is added to system
w is positive if work is done on the system
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Heat Changes
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Exothermic Reactions – when heat is given off
by the reaction (-Q)
Endothermic Reactions – when heat is used by
the reaction (+Q)
Example
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As a combustion reaction occurs the system
loses 550 J of heat to its surroundings and it
does 240 J of work in moving a piston. What is
the change in its internal energy?
State Function
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These are systems for whom the value of ΔE
does not depend on the previous history of the
sample, only on the present condition
Energy is a state function
Work and heat are not state functions