Slide 1

Introduction to
Thermodynamics
and Kinetics

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Slide 2

Topics
Types

of energy and units of energy
exothermic vs. endothermic reactions
heat capacity and specific heat
energy transfers in changes of state
calorimetry
enthalpy and entropy
Kinetics and catalysts
equilibrium and LeChatelier’s Principle
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Slide 3

Energy is defined as the
capacity to do work.
There

are two catagories of energy:
Kinetic energy.
Potential energy
Kinetic Energy is the energy of motion,
like climbing steps, or gears moving,
etc.
Potential energy is energy of position,
like a rock sitting at the edge of a cliff, or
a tightly wound up spring. It can
convert into Kinetic energy.

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Slide 4

Chemical energy is a form of
energy stored in the structure of a
chemical substance.

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Slide 5

Consider

the potential energy of the
rock up on the cliff. If one rock is up
on a 1000 ft. cliff and another is up
on a 10 ft. hill, the rock that is 1000ft.
up will hit the ground with much more
energy than the other if it falls.

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Slide 6

Chemical

energy has similar
differences in that the structure
of one substance may contain
much more potential energy
than another substance.

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Slide 7

Thermochemistry is the study
of heat transfers occurring in
chemical and physical changes
of substances.
 Thermal

energy is the energy
associated with the random motion of
atoms and molecules.
 Heat is the transfer of thermal energy
between two bodies that are different
temperatures.
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Slide 8

The Universe is made up of
the system and the
surroundings
The system is the part of the universe
that is of interest to us. This could be, for
example, a beaker in which a chemical
reaction is taking place.
The surroundings is everything outside of
the system.
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Slide 9

Exothermic vs. Endothermic
Processes
Any

process that gives off
heat to the surroundings is
an Exothermic process.
When bonds are made
between particles, the
process is exothermic.

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Slide 10

When

a process absorbs
heat from the surroundings
it is an Endothermic
process. When ice melts it
absorbs heat from the
surroundings, thus it is an
endothermic process.
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Slide 11

The symbol DH is used to
represent the change in heat
into or out of the system. It is
defined as the change in
enthalpy.

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Slide 12

The enthalpy of reaction is the
difference between the enthalpies
of the products and reactants.
DH(rxn) = DH(products) - DH(reactants)

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Slide 13

The diagram illustrates energy given
off when bonds are made between H2
and O2 , (a). Energy is absorbed when
bonds are broken in HgO , (b).

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Slide 14

The units of energy are
Kilojoules and Calories.
Enthalpy of reaction will be
expressed as Kilojoules (Kj).
A conversion factor to go
between these two units is
1 cal = 4.184 J
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Slide 15

It is important to note the following
when calculating the enthalpy of
reaction:
The enthalpy of a substance in its
standard state is equal to zero.
Enthalpy is dependent on the
quantity of the substance therefore the
enthalpy of a substance must be
multiplied by its coefficient in the
chemical reaction.
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Slide 16

Enthalpy values of compounds are
found in tables of thermodynamic data.
Ex. 1

Determine the enthalpy for the following
chemical reaction:
CO2 (g) + 2 H2O (l) -----> 2 O2 (g) + CH4 (g)
continued...
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Slide 17

Use table 7.2 in your text to find the
individual enthalpy values.
CO2 (g) = -393.509, H2O (l) = -285.83
CH4 (g) = -74.81 ,

O2 (g)= 0

CO2 + 2H2O ----> 2 O2 + CH4
DH = DH(products) – DH(reactants)
DH = (0 x -74.81)-(-393.509 + 2(- 285.83))
= -74.81 + 965.17 = 890.36 Kj
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Slide 18

Specific heat and heat Capacity
Heat capacity is the amount of heat
required to raise the temperature of a
given quantity of a substance by 1o C.

Specific Heat is the amount of heat
required to raise the temperature of 1
gram of a substance by 1o C.
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Slide 19

These can be thought of as a substance’s
ability to absorb heat and to store heat.
For example, a metal does not require
much energy to heat it up and it does not
hold the heat for a long time.
A liquid like water (which contains strong
intermolecular forces, hydrogen bonds),
requires much more energy to heat it up
and it does hold the heat for a longer
period of time.
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Slide 20

The metal has a low heat capacity and
the water has a high heat capacity.
The specific heat capacity of water is
an important and easy number to
remember ; 1 cal / g oC ,
(or 4.184 J / goC

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Slide 21

Calorimetry is a technique used
in the lab to measure the enthalpy
of a reaction. One apparatus
used is the Bomb Calorimeter. It
is a Heavy walled, steel container
which has a known specific heat.
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Slide 22

The heat of the reaction (q) is
equal to the negative of the heat
absorbed by the (bomb plus the
water) surrounding the bomb.

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Slide 23

q is used to represent the
quantity and direction of heat
transferred, using a calorimeter.

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Slide 24

Necessary Equations:
q = (specific heat)(mass)(change in Temp.)

q = (J/gK)(g)(DT)
q(rxn) = - (q of the water + q of the bomb)

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Slide 25

Subliminal message.....

Wake up !!!
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Slide 26

Ex. 2

Calorimetry

A 466g sample of water is heated from
8.50oC to 74.60oC. Calculate the amount
of heat absorbed by the water.
q = (sp. heat)(mass)(DT)
q = (4.184 J/goC)(466g)(74.60-8.50oC)
q = 1.29 x 105 J = 129 KJ
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Slide 27

Ex. 3

1.435 g of Naphthalene (molar mass
=128.2) was burned in a bomb
calorimeter. The temp. rose from
20.17oC to 25.84 oC. The mass of
the water surrounding the
calorimeter was 2000.g and the heat
capacity of the bomb was 1.80
KJ/oC. Calculate the heat of
combustion of Naphthalene.
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Slide 28

q(rxn) = -(q water + q bomb)

q(water) = (2000g)(4.184 J/goC)(5.67oC)
= 4.74 x 104

q(bomb) = (1.80 x 103J / oC)(5.67oC)
= 1.02 x 104 J

q(rxn)

= -(4.74 x 104 J + 1.02 x 104 J)
= -5.76 x 104 J
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