Transcript Chapter 6 PowerPoint

Chapter 6
Energy
Thermodynamics
Energy is...
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The ability to do work.
Conserved.
made of heat and work.
a state function.
independent of the path, or how you get
from point A to B.
• Work is a force acting over a distance.
• Heat is energy transferred between objects
because of temperature difference.
Energy
Literally means “work within,” however no
object contains work
Energy refers to the capacity to do work
– that is, to move or displace matter
2 basic types of energy:
– Potential (possibility of
doing work because of
composition or position)
– Kinetic (moving objects
doing work)
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Energy
Potential Energy – in
a gravitational field
(= position)
PE = mgh
m = mass (kg)
g = gravity
constant (m s2)
h = height (m)
units are kg m2 s–2  J
Kinetic Energy –
energy of motion
KE = 1/2mv2
m = mass (kg)
v = velocity (m s–
1)
v2 = (m2 s–2)
units are kg m2 s–2
=J
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Potential Energy
Potential energy is energy an object possesses by
virtue of its position or chemical composition.
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Kinetic Energy
Kinetic energy is energy an object possesses by
virtue of its motion.
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KE =  mv2
2
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Conversion of Energy
• Energy can be converted from one type to
another.
• For example, the cyclist above has potential
energy as she sits on top of the hill.
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Conversion of Energy
• As she coasts down the hill, her potential energy is
converted to kinetic energy.
• At the bottom, all the potential energy she had at
the top of the hill is now kinetic energy.
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Units of Energy
• The SI unit of energy is the joule (J).
kg m2
1 J = 1 
s2
• An older, non-SI unit is still in widespread
use: the calorie (cal).
1 cal = 4.184 J
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The universe
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is divided into two halves.
the system and the surroundings.
The system is the part you are concerned with.
The surroundings are the rest.
Exothermic reactions release energy to the
surroundings.
• Endo thermic reactions absorb energy from the
surroundings.
Thermochemistry
System
Surroundings
Universe
Surroundings
Thermochemistry is the study of energy changes
that occur during chemical reactions
Surroundings
Chapter 6: Thermochemistry
Focus is on heat
and matter
transfer
between the
system ...
and the
surrounding
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Definitions:
System and Surroundings
• The system includes the
molecules we want to study
(here, the hydrogen and
oxygen molecules).
• The surroundings are
everything else (here, the
cylinder and piston).
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Thermochemistry
Types of systems one can study:
Matter
Matter
Energy
OPEN
Energy
Matter Matter
Energy
Energy
CLOSED
Matter Matter
Energy Energy
ISOLATED
EOS
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Definitions: Work
• Energy used to move
an object over some
distance is work.
• w=Fd
where w is work, F is
the force, and d is the
distance over which
the force is exerted.
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Internal Energy
Internal Energy (U) is the total energy
contained within the system, partly as kinetic
energy and partly as potential energy
Kinetic involves three types of molecular
motion ...
EOS
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Internal Energy
Internal Energy (U) is the total energy
contained within the system, partly as kinetic
energy and partly as potential energy
Potential energy involves
intramolecular interactions ...
and
intermolecular
interactions ... EOS
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Heat
• Energy can also be
transferred as heat.
• Heat flows from
warmer objects to
cooler objects.
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Heat (q)
Heat is energy transfer resulting from thermal
differences between the system and
surroundings
“flows”
spontaneously from
higher T  lower T
“flow” ceases at
thermal equilibrium
EOS
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Potential energy
CH 4 + 2O 2  CO 2 + 2H 2 O + Heat
CH 4 + 2O 2
Heat
CO 2 + 2 H 2 O
N 2 + O 2 + heat  2NO
Potential energy
2NO
Heat
N2 + O2
Same rules for heat and work
• Heat given off is negative.
• Heat absorbed is positive.
• Work done by system on surroundings is
positive.
• Work done on system by surroundings is
negative.
• Thermodynamics- The study of energy and the
changes it undergoes.
Work (w)
Work is an energy transfer between a
system and its surroundings
Recall from gas laws … the product PV =
energy
Pressure–
volume work is
the work of
compression (or
expansion) of a
gas
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Calculating Work (w)
PV work is calculated as follows:
w = –PDV
Sign conventions: think
SYSTEM WORK
from the perspective of
the system
If work is done by the
system, the system
loses energy equal to –
w
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Calculating Work (w)
SYSTEM WORK
Expansion is an example of
work done by the system—
the weight above the gas is
lifted
compression (or
expansion) of a
gas
ExpansionWork
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What is work?
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Work is a force acting over a distance.
w= F x Dd
P = F/ area
d = V/area
w= (P x area) x D (V/area)= PDV
Work can be calculated by multiplying
pressure by the change in volume at constant
pressure.
• units of liter - atm L-atm
Work needs a sign
• If the volume of a gas increases, the system
has done work on the surroundings.
• work is negative
• w = - PDV
• Expanding work is negative.
• Contracting, surroundings do work on the
system w is positive.
• 1 L atm = 101.3 J
Examples
• What amount of work is done when 15 L of
gas is expanded to 25 L at 2.4 atm pressure?
• If 2.36 J of heat are absorbed by the gas
above. what is the change in energy?
• How much heat would it take to change the
gas without changing the internal energy of
the gas?
Surroundings
System
Energy
DE <0
Surroundings
System
Energy
DE >0
Direction
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Every energy measurement has three parts.
A unit ( Joules of calories).
A number how many.
and a sign to tell direction.
negative - exothermic
positive- endothermic
First Law of Thermodynamics
• Energy is neither created nor destroyed.
• In other words, the total energy of the universe is a
constant; if the system loses energy, it must be gained
by the surroundings, and vice versa.
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First Law of Thermodynamics
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The energy of the universe is constant.
Law of conservation of energy.
q = heat
w = work
DE = q + w
Take the systems point of view to decide signs.
Internal Energy
The internal energy of a system is the sum of all
kinetic and potential energies of all components of
the system; we call it E.
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Internal Energy
By definition, the change in internal energy, DE, is the
final energy of the system minus the initial energy of
the system:
DE = Efinal − Einitial
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Changes in Internal Energy
• When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
• That is, DE = q + w.
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DE, q, w, and Their Signs
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Exchange of Heat between System
and Surroundings
• When heat is absorbed by the system from the
surroundings, the process is endothermic.
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Exchange of Heat between System
and Surroundings
• When heat is absorbed by the system from the
surroundings, the process is endothermic.
• When heat is released by the system into the
surroundings, the process is exothermic.
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States of a System
The state of a system refers to its exact
condition, determined by the kinds and
amounts of matter present, the structure of
this matter at the molecular level, and the
prevailing pressure and temperature
Example: internal energy (U) is a
function of the state of the
system ...
EOS
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State Functions
A state function is a
property that has a
unique value that
depends only on the
present state of a system
and not on how the state
was reached, nor on the
history of the system
DU = Uf – Ui
EOS
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State Functions
• However, we do know that the internal energy of a
system is independent of the path by which the
system achieved that state.
– In the system below, the water could have reached room
temperature from either direction.
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State Functions
• Therefore, internal energy is a state function.
• It depends only on the present state of the system,
not on the path by which the system arrived at that
state.
• And so, DE depends only on Einitial and Efinal.
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State Functions
• However, q and w are not
state functions.
• Whether the battery is
shorted out or is
discharged by running
the fan, its DE is the
same.
– But q and w are different
in the two cases.
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Enthalpy
Symbol is H
Change in enthalpy is DH
delta H
If heat is released the heat content of the
products is lower
• DH is negative (exothermic)
• If heat is absorbed the heat content of the
products is higher
• DH is positive (endothermic)
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Enthalpy
• If a process takes place at constant pressure
(as the majority of processes we study do)
and the only work done is this pressurevolume work, we can account for heat flow
during the process by measuring the enthalpy
of the system.
• Enthalpy is the internal energy plus the
product of pressure and volume:
H = E + PV
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Enthalpy
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abbreviated H
H = E + PV (that’s the definition)
at constant pressure.
DH = DE + PDV
• the heat at constant pressure qp can be
calculated from
• DE = qp + w = qp - PDV
• qp = DE + P DV = DH
Enthalpy
• Since DE = q + w and w = -PDV, we can
substitute these into the enthalpy
expression:
DH = DE + PDV
DH = (q+w) − w
DH = q
• So, at constant pressure, the change in
enthalpy is the heat gained or lost.
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Inc.
Endothermicity and Exothermicity
• A process is
endothermic when
DH is positive.
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Endothermicity and Exothermicity
• A process is
endothermic when
DH is positive.
• A process is
exothermic when
DH is negative.
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Enthalpy of Reaction
The change in enthalpy,
DH, is the enthalpy of
the products minus the
enthalpy of the
reactants:
DH = Hproducts − Hreactants
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Enthalpy of Reaction
This quantity, DH, is called the enthalpy of
reaction, or the heat of reaction.
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The Truth about Enthalpy
1. Enthalpy is an extensive property.
2. DH for a reaction in the forward direction is
equal in size, but opposite in sign, to DH for
the reverse reaction.
3. DH for a reaction depends on the state of
the products and the state of the reactants.
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Inc.
Calorimetry
Calorimetry is a technique used to measure
heat exchange in chemical reactions
A calorimeter is the device
used to make heat
measurements
Calorimetry is based on the law
of conservation of energy
EOS
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Calorimetry Relationships
The heat capacity (C) of a system is the quantity
of heat required to change the temperature of the
system by 1 oC
calculated from C = q/DT
units of J oC–1
or J K–1
Specific heat is the heat capacity of a one-gram
sample
Specific heat = C/m = q/mDT
units of J g–1 oC–1 or J g–1 K–1
EOS
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Specific Heats
Molar heat capacity is the product of specific
heat times the molar mass of a substance
units are J mol–1 K–1
A useful form of the specific heat equation is:
q = m CDT
If DT > 0, then q > 0 and heat is gained by the
system
If DT < 0, then q < 0 and heat is lost by the
system
EOS
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Calorimetry
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Measuring heat.
Use a calorimeter.
Two kinds
Constant pressure calorimeter (called a coffee
cup calorimeter)
• heat capacity for a material, C is calculated
• C= heat absorbed/ DT = DH/ DT
• specific heat capacity = C/mass
Calorimetry
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molar heat capacity = C/moles
heat = specific heat x m x DT
heat = molar heat x moles x DT
Make the units work and you’ve done the
problem right.
A coffee cup calorimeter measures DH.
An insulated cup, full of water.
The specific heat of water is 1 cal/gºC
Heat of reaction= DH = sh x mass x DT
Heat Capacity and Specific Heat
The amount of energy required to raise the
temperature of a substance by 1 K (1C) is its heat
capacity.
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Heat Capacity and Specific Heat
We define specific heat capacity (or simply
specific heat) as the amount of energy required to
raise the temperature of 1 g of a substance by 1 K.
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Heat Capacity and Specific Heat
Specific heat, then, is
Specific heat =
s=
heat transferred
mass  temperature change
q
m  DT
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Examples
• The specific heat of graphite is 0.71 J/gºC.
Calculate the energy needed to raise the
temperature of 75 kg of graphite from 294 K
to 348 K.
• A 46.2 g sample of copper is heated to 95.4ºC
and then placed in a calorimeter containing
75.0 g of water at 19.6ºC. The final
temperature of both the water and the copper
is 21.8ºC. What is the specific heat of copper?
Calorimetry
• Constant volume calorimeter is called a bomb
calorimeter.
• Material is put in a container with pure
oxygen. Wires are used to start the
combustion. The container is put into a
container of water.
• The heat capacity of the calorimeter is known
and tested.
• Since DV = 0, PDV = 0, DE = q
Constant Pressure Calorimetry
By carrying out a reaction
in aqueous solution in a
simple calorimeter such as
this one, one can indirectly
measure the heat change
for the system by
measuring the heat
change for the water in
the calorimeter.
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Constant Pressure Calorimetry
Because the specific heat
for water is well known
(4.184 J/g-K), we can
measure DH for the
reaction with this
equation:
q = m  s  DT
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Bomb Calorimeter
• thermometer
• stirrer
• full of water
• ignition wire
• Steel bomb
• sample
Properties
• intensive properties not related to the amount
of substance.
• density, specific heat, temperature.
• Extensive property - does depend on the
amount of stuff.
• Heat capacity, mass, heat from a reaction.
Hess’s Law
• Enthalpy is a state function.
• It is independent of the path.
• We can add equations to to come up with the
desired final product, and add the DH
• Two rules
• If the reaction is reversed the sign of DH is
changed
• If the reaction is multiplied, so is DH
Hess’s Law
Hess’s law states that
“[i]f a reaction is carried
out in a series of steps,
DH for the overall
reaction will be equal to
the sum of the enthalpy
changes for the
individual steps.”
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Hess’s Law
Because DH is a state
function, the total
enthalpy change depends
only on the initial state of
the reactants and the final
state of the products.
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H (kJ)
O2 NO2
-112 kJ
180 kJ
N2 2O2
NO2
68 kJ
Standard Enthalpies
The standard enthalpy of reaction (DHo) is the
enthalpy change for a reaction in which the
reactants in their standard states yield products
in their standard states
The standard enthalpy of formation (DHof) of a
substance is the enthalpy change that occurs in
the formation of 1 mol of the substance from its
elements when both products and reactants are
in their standard states
EOS
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Standard Enthalpy
• The enthalpy change for a reaction at
standard conditions (25ºC, 1 atm , 1 M
solutions)
• Symbol DHº
• When using Hess’s Law, work by adding the
equations up to make it look like the answer.
• The other parts will cancel out.
Standard Enthalpies of Formation
Standard enthalpies of formation, DHf°, are
measured under standard conditions (25 °C and
1.00 atm pressure).
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Calculation of DH
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
DH = [3(-393.5 kJ) + 4(-285.8 kJ)] – [1(-103.85 kJ) + 5(0 kJ)]
= [(-1180.5 kJ) + (-1143.2 kJ)] – [(-103.85 kJ) + (0 kJ)]
= (-2323.7 kJ) – (-103.85 kJ) = -2219.9 kJ
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Example
5
• Given
C 2 H 2 (g) + O 2 (g)  2CO 2 (g) + H 2 O( l)
2
DHº= -1300. kJ
C(s) + O 2 (g)  CO 2 (g)
DHº= -394 kJ
1
H 2 (g) + O 2 (g)  H 2 O(l)
2
DHº= -286 kJ
calculate
DHº for
this
reaction
2C(s) + H (g)  C H (g)
2
2
2
Example
Given
O 2 (g) + H 2 (g)  2OH(g) DHº= +77.9kJ
O 2 (g)  2O(g) DHº= +495 kJ
H 2 (g)  2H(g) DHº= +435.9kJ
Calculate DHº for this reaction
O(g) + H(g)  OH(g)
Standard Enthalpies of Formation
• Hess’s Law is much more useful if you know
lots of reactions.
• Made a table of standard heats of formation.
The amount of heat needed to for 1 mole of a
compound from its elements in their standard
states.
• Standard states are 1 atm, 1M and 25ºC
• For an element it is 0
• There is a table in Appendix 4 (pg A22)
Standard Enthalpies of Formation
• Need to be able to write the equations.
• What is the equation for the formation of NO2
?
• ½N2 (g) + O2 (g)  NO2 (g)
• Have to make one mole to meet the
definition.
• Write the equation for the formation of
methanol CH3OH.
Since we can manipulate the
equations
• We can use heats of formation to figure out
the heat of reaction.
• Lets do it with this equation.
• C2H5OH +3O2(g)  2CO2 + 3H2O
• which leads us to this rule.
Since we can manipulate the
equations
• We can use heats of formation to figure out
the heat of reaction.
• Lets do it with this equation.
• C2H5OH +3O2(g)  2CO2 + 3H2O
• which leads us to this rule.
( DH of products) - ( DH of reactants) = DH o
Calculations Based on
Standard Enthalpies of Formation
General Expression:
DHo = Snp × DHof (products) – Snr × DHof (reactants)
Each coefficient is multiplied by the standard
enthalpy of formation for that substance
The sum of numbers for the reactants is
subtracted from the sum of numbers for the
products
With organic compounds, the measured DHo
f
is often the standard enthalpy of combustion
DHocomb
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