Transcript Energy

Energy
A First Look at
Thermodynamics…
Definition…
• Ability to do “work”; capacity of a
system to do “work”
• “Work” is defined for macroscopic
systems as moving an object against a
force. But chemistry deals with many
microscopic systems so the physics
definition of “work” does not always
seem to apply.
Forms…
• Potential – energy due to position
• Kinetic – energy due to motion
• These forms of energy are
interconvertible.
Energy is the capacity to do work.
less stable
change in potential energy
EQUALS
kinetic energy
more stable
A gravitational system. The potential energy gained when a
lifted weight is converted to kinetic energy as the weight
falls.
Energy is the capacity to do work.
less stable
change in potential energy
EQUALS
kinetic energy
more stable
A system of two balls attached by a spring. The potential energy
gained by a stretched spring is converted to kinetic energy when the
moving balls are released.
Energy is the capacity to do work.
less stable
change in potential energy
EQUALS
kinetic energy
more stable
A system of oppositely charged particles. The potential energy
gained when the charges are separated is converted to kinetic energy
as the attraction pulls these charges together.
Energy is the capacity to do work.
less stable
change in potential energy
EQUALS
kinetic energy
more stable
A system of fuel and exhaust. A fuel is higher in chemical potential
energy than the exhaust. As the fuel burns, some of its potential
energy is converted to the kinetic energy of the moving car.
Types of Energy
• There are also different types of energy:
chemical
electrical
gravitational
heat
light (electromagnetic radiation)
magnetic
mechanical
nuclear
These are also interconvertible.
Laws!
• No matter what the conversion the
system must obey the Law of
Conservation of Energy.
• Energy cannot be created or destroyed.
(but may be converted to mass under
certain conditions)
Thermal Energy
• In chemistry we deal primarily with thermal
energy. We also encounter light and electrical
energies.
• Thermal energy = energy due to molecular or
atomic motion
• Heat (q) = transfer of chemical energy due to a
temperature difference
• Thermodynamics = study of heat and its
transformations.
• Thermochemistry = branch of thermodynamics
that deals with the heat involved with chemical
and physical changes.
Temperature
• Measure of the “hotness” or “coldness”
of something
• Measure of the effect of heat on an
object or system
• Measure of particle motion
• Heat always flows from an area of
higher temperature to areas of lower
temperatures.
Temperature Measurement
• Typically measured using a thermometer
which has an expandable liquid (mercury
or alcohol) trapped in a sealed cylinder
• US meteorological and heating unit
temperatures are expressed using the
Fahrenheit scale.
• Water freezes at 32oF and boils at
212oF.
Figure 1.8
The freezing and boiling points of water.
Temperature Measurement
• Other countries and the scientific
community use the Celsius scale.
• Water freezes at 0oC and boils at
100oC.
Figure 1.8
The freezing and boiling points of water.
Silberberg, Principles of Chemistry
Temperature Measurement
• There is a problem using either of these
scales in certain systems. Because there
are negative temperatures certain
mathematical relationships generate
impossible results.
• Lord Kelvin generated a new scale with the
coldest temperature as 0 K (absolute
zero). The unit of temperature is a kelvin
(K). The term degree is not used.
• Water freezes at -273.15 K and boils at
373.15 K.
Figure 1.8
The freezing and boiling points of water.
Silberberg, Principles of Chemistry
Temperature Conversions
• oF  oC
o
F -32 o
= C
1.8
• oC  K
C + 273.15 = K
o
C x 1.8 + 32 = F
o
o
Systems…
• Changes in thermal energy of systems are
determined by transfer of heat (q)
observed by changes in temperature.
• First you need to define what your system
and surroundings are…
• When heat moves out of a system to the
surroundings the process is exothermic for
the system.
• When heat moves into a system from the
surroundings the process is endothermic
for the system.
A system transferring energy as heat only.
Exothermic
Endothermic
Heat Units
Joule (J)
1 J = 1 kg*m2/s2
Calorie (cal)
1 cal = 4.18J
British Thermal Unit
1 Btu = 1055 J
Specific Heat Capacity
• Substances respond differently to the
same quantity of energy.
• Consider a wooden spoon and a metal
spoon in a pot of boiling water. Which
spoon gets hotter?
• Specific heat capacity (c) – energy
required to raise 1 gram of substance by
1oC
J/g. oC (J g-1 oC-1) or kcal/g. oC (kcal g-1 oC-1)
Heat Calculations
q = specific heat capacity x mass x temperature
PROBLEM:
PLAN:
A layer of copper welded to the bottom of a skillet weighs
125 g. How much heat is needed to raise the temperature of
the copper layer from 250C to 300.0C? The specific heat
capacity (c) of Cu is 0.387 J/g*K.
Given the mass, specific heat capacity and change in temperature, we
can use q = c x mass x T to find the answer. T in 0C is the same as
for K.
SOLUTION: q =
0.387 J
g*K
0
x 125 g x (300-25) C
= 1.33x104 J
Silberberg, Principles of Chemistry
Enthalpy
The Meaning of Enthalpy
w = - PV
H = E + PV
where H is enthalpy
H = E + PV
H ≈ E in
1. Reactions that do not involve gases.
2. Reactions in which the number of
moles of gas does not change.
3. Reactions in which the number of
moles of gas does change but q is >>>
PV.
qp = E + PV = H
For most of the processes we will encounter this semester
we can use ΔH as heat exchanged by a system.
Enthalpy
ΔH for a reaction is calculated as:
ΔHproducts – ΔHreactants
If the ΔH is positive, energy of the
products is higher than reactants and
the process is endothermic.
If the ΔH is negative, energy of the
reactants is higher than products and
the process is exothermic.
Enthalpy diagrams for exothermic and endothermic processes.
Enthalpy, H
CH4 + 2O2
H < 0
CO2 + 2H2O
A
CO2(g) + 2H2O(g)
H2O(l)
H2O(g)
Hinitial
heat
out
Enthalpy, H
CH4(g) + 2O2(g)
H > 0
H2O(l)
Hfinal
Exothermic process
B
H2O(g)
Hfinal
heat in
Hinitial
Endothermic process
Determining the
heat exchange
in a chemical
change
A bomb
calorimeter
AMOUNT (mol)
of compound A
Summary of the relationship between
amount (mol) of substance and the heat
(kJ) transferred during a reaction.
AMOUNT (mol)
of compound B
molar ratio from
balanced equation
HEAT (kJ)
Hrxn (kJ/mol)
gained or lost
Using the Heat of Reaction (Hrxn) to Find Amounts
PROBLEM:
The major source of aluminum in the world is bauxite (mostly
aluminum oxide). Its thermal decomposition can be represented by
Al2O3(s)
2Al(s) + 3/2O2(g)
Hrxn = 1676 kJ
If aluminum is produced this way, how many grams of aluminum
can form when 1.000x103 kJ of heat is transferred?
SOLUTION:
1.000x103 kJ x 2 mol Al
1676 kJ
26.98 g Al
1 mol Al
= 32.20 g Al
Reaction Progress
• All reactions
have an energy
barrier that
must be
overcome.
Activation energy (Ea)
This is an exothermic
reaction.
Energy of reactants
Energy of products
Reaction Progress
Activation energy (Ea)
This diagram
illustrates an
endothermic
reaction.
Energy of reactants
Energy of products
Reaction Enthalpy
The difference
between the starting
and ending energies
is ΔH.
ΔH
Catalysts reduce the activation energy of a
reaction.
Reaction energy diagram of a catalyzed and an uncatalyzed process.
Disorder aka Entropy
• Systems tend to greater disorder or
greater energy dispersal.
• Disorder is the preferred state.
• Energy dispersal or system disorder is
called entropy.
• All systems work to achieve lowest
energy and greatest disorder or a
balance between these goals.