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1/25/2010
Thermochemistry
Reading: Chapter 5 (omit 5.8)
As you read ask yourself …
What is meant by the terms system and surroundings? How are they related
to each other? How does energy get transferred between them?
What makes up the internal energy of matter?
What kind of work do chemical system do?
What impact does heat and work transfer have on the internal energy of a
system?
Does how we effect the transfer of heat and work matter to the value of the
internal energy change? Why or why not?
How can we measure heat flow?
How is temperature change related to the transfer of heat?
Does the composition of the substance that undergoes a change in
temperature matter,
matter if so why?
Is the same thermodynamic quantity measured in all calorimetry
experiments? Why or why not?
Since we can’t measure every reactions enthalpy directly, what methods can
we use to figure out the enthalpy of any given reaction?
What fundamental property of enthalpy makes Hess’s Law work?
Why don’t we have a value for the heat of formation of O2(g) and C(graphite)?
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Thermochemistry
Why do chemical reactions occur?
energy is an important factor
in determining stability
stability!
products are more stable
than reactants
energy
less stable
Nature of energy:
capacity to do work or transfer heat
energy when a
force causes a
mass to move
reaction
more stable
energy to cause
an increase in
temperature
objects possess energy
work and heat are ways to transfer energy
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Energy can be classified as either kinetic or potential energy
kinetic energy
EK = ½ m
energy of motion
v2
thermal energy – energy associated
with temperature of object
thermal energy depends on T and quantity
potential energy
stored energy
arises from position or composition
chemical energy – energy associated
with electrons and nuclei
electrostatic energy – interaction
between charges is the force
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Units
Joules
1 Joule = 1 kg m2 s-2
e.g. a 4 kg mass moving at 1 m/s
EK = ½ m v2 = ½ (4 kg) (1 m s-1)2 = 2 kg m2 s-2
Calories
= 2 Joules
defined as 1 cal = 4.184 J exactly
old definition: energy required to raise temp. of 1 g of water from
14.5 ° to 15.5 °C
Calorie = 1000 cal = 1 kcal
Nutritional calorie
Kilowatt-hour
1 watt = 1 J
1 kW h = 3.6 x
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SI unit of power (rate of energy conversion)
s-1
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100 watt incandescent bulb
uses 100 J s-1 = 0.1 kW h
J
CFL 23 watts
0.023 kW h
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First Law of Thermodynamics: energy
is neither created or
destroyed therefore the total energy of the universe is
constant (energy is conserved).
energy can be transferred
energy can be
b converted
t d ffrom
one form to another
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System and surroundings
to understand energy transfers and transformations
we have to define the part of the universe
we are studying
system:
part of universe chosen
for study
surroundings:
part of universe outside of the system
with which the system interacts
isolated
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most observations are made in the surroundings
example: chemical reaction
Zn(s) = HCl(aq) J Zn2+(aq) + H2(g) + 2Cl-(aq)
once Zn is added system
y
is closed
mass is moved so work is done = w
heat is given off = q
energy is transferred
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change in
temperature
energy transferred from hot to
cold until equilibrium (both at
same T)
move an object against a
force
w=Fxd
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size of force and distance
moved determine
quantity of work
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Internal energy, E
internal energy is the sum of all the
kinetic and potential energy of all
components of the system
hard to obtain absolute value of E
normally determine the change in internal
energy ΔE = E
–E
final
initial
for a chemical reaction:
ΔE = Eproducts – Ereactants
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How is ΔE changed?
depends on heat transfer (q) and work (w)
heat added to
or liberated by system
ΔE = q + w
work done by
or done on system
the magnitude of ΔE depends on the size of q and w and their
relative signs
note that the
values of ΔE,
ΔE q
and w refer to
the SYSTEM
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system loses heat, ΔE L so q is
negative
system uses its energy & does work,
ΔE L so work is negative
in this case Efinal < Einitial, ΔE is < 0
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exothermic
system gains heat, ΔE K so q is +
work done on system, ΔE K so w is +
in this case Efinal > Einitial, ΔE is > 0
endothermic
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example: A balloon is heated by adding 900 J of heat, it expands
and does 422 J of work.
What is the change in internal energy?
If balloon is heated by adding 1500J it expands and does 800J of
work.
Suppose the balloon is then cooled by removing 520 J of heat and
compressed by doing 298 J of work on the balloon
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how the energy transfer is divided between work and heat depends on
the process
the total energy transferred does not depend on the
p
pathway
y
in the example:
analogy to altitude:
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internal energy is a state function
A function or property whose value depends only on the
present state or condition of the system, not on the path
used to arrive at that state
a state function
Examples of state functions:
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ΔE in chemical reactions
ƒ Most chemical changes occur at constant
atmospheric pressure
ƒ Heat and work are exchanged with surroundings
ƒWork is typically mechanical (change in volume
of gases) or electrical
expansion
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Expansion work (P-V work)
work – the force that moves an object through a distance
pressure = force / area
sign convention, when ΔV is positive,
work is negative
Units:
w = − pressure x volume = atm x L
from Data Sheet: 1 J = 0.009869 L atm
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Example: If 0.225 mol of N2 at constant T is compressed by
15.1 L at P = 0.750 atm, what is the work involved?
What does this mean for our definition of internal energy?
Substitute PV work into expression for ΔE
If reaction
ti takes
t k place
l
in
i a sealed
l d
container, ΔV = 0
then no PV work is done
What if external pressure is zero (expansion against a vacuum)?
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chemical changes typically take place at constant pressure
define a new thermodynamic quantity for
the heat change at constant pressure,
enthalpy H
defined as H = E + PV
because E, P and V are state functions
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Example. ΔE = -186.9 kJ/mol, what is sign of PΔV? What is sign
and approximate magnitude of ΔH? Assume that T is constant.
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Enthalpy and chemical reactions
ΔH = Hfinal − Hinitial
so in chemical reactions
(ΔH is often written
chemical reaction and ΔH together are thermochemical equation
relationship between amounts of chemicals and heat involved in
1 provides
reaction
2H2O2(ℓ)
2H2O2(ℓ) 2H2O(ℓ) + O2(g) ΔH = −196 kJ
2
3
−196 kJ
+196 kJ
2H2O2(ℓ) 2H2O(g) + O2(g) ΔH = −108 kJ
Chem 102
2H2O(ℓ) + O2(g)
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Example: Sulfuric acid is produced by reacting sulfur trioxide with water
according to the equation:
SO3 (g) + H2O (ℓ) → H2SO4 (ℓ) ΔH° = −131.8 kJ/mol
How much heat is evolved when 75.0 g of SO3 reacts with a
stoichiometric amount of H2O(ℓ)?
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Heat transfer
define system and surroundings
Effectiveness of heat transfer depends on nature of substances
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Heat capacity and specific heat
how much the temperature changes per
amount of heat added depends on the
nature of the substance
Heat capacity
p
y – amount of heat required
q
to
raise the temperature of
an object by 1 K (or 1°C)
heat capacity is proportionality constant, C
heat capacity depends on
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Heat capacity for pure substances
intrinsic property
of substance
Specific heat capacity – defined for 1 gram of substance and a
temperature increase of 1°C or 1 degree Kelvin
specific heat =
amount of heat transferred
(grams of substance) (change in temperatur e)
Molar heat capacity – defined for 1 mole of substance and a
temperature increase of 1°C or 1 degree Kelvin
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Specific heat capacity (molar heat capacity) depends on
bonding
complexity
physical state
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d t
determining
i i specific
ifi h
heat:
t iin step
t ((a)) the
th lead
l d is
i heated
h t d tto 100°C
100°C, th
then
added to water in step(b) and the final temperature is measured in step(c)
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Measuring heat transfer
measure the magnitude of the
temperature change as heat flows
calorimetry:
main idea:
heat
change in
water in
device
heat change in the
reaction
Chem 102
measure change in T
calorimeter
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Constant pressure calorimetry
usually reactions are in solution
what is the system?
what are the surroundings?
When 10.0 mL of a 1.00 M AgNO3 solution is added to 10.0 mL of 1.00 M
NaCl solution at 25.0 °C in a constant pressure calorimeter, a white
precipitate of AgCl is formed and the temperature of the aqueous mixture
increases to 32.6 °C. Assuming that the specific heat of the aqueous
mixture
i t
is
i 4
4.18
18 J g-11 °C-11, that
th t th
the d
density
it off the
th mixture
i t
is
i 1.00
1 00 g mL
L-11, and
d
that the calorimeter absorbs no heat, calculate ΔH in kJ for the reaction.
Ag+(aq) + Cl−(aq) J AgCl(s)
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Constant volume calorimetry
device often used for combustion reactions
e.g. C6H6(ℓ) + 15/2 O2(g) J 6 CO2(g) + 3 H2O(g)
heat flows
out of reaction chamber
into water and heats
calorimeter
what is the system?
what are the surroundings?
qrxn = ΔE … why?
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Example: Combustion of a liquid rocket fuel, methylhydrazine CH6N2(s), produces
CO2(g), N2(g) and H2O(ℓ). When 4.00g of methyhydrazine is burned in a
bomb calorimeter, the temperature increased from 25.0 to 39.5°C. The
heat capacity of the calorimeter is 7.794 kJ/°C. What is the heat of
reaction for combustion of 1.0 mole of CH6N2 in the calorimeter?
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Reaction enthalpies
the enthalpy change for every reaction can
not be easily measured
1
2
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Enthalpy is a state function, so can use known enthalpies for
stepwise processes that have the initial and final states of interest
N2(g) + 3H2(g) J 2 NH3(g)
ΔHrxn = ?
N2(g) + 3H2(g) J N2H4(g) ΔH = 95.4 kJ
N2H4(g) + H2(g) J 2 NH3(g) ΔH = − 187.6 kJ
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Hess’s Law
The enthalpy change of an overall process is the sum of the
enthalpy changes of its individual steps
P
Procedure:
d
combine
bi the
th individual
i di id l reactions
ti
so their
th i sums give
i the
th
desired reaction
™Arrange reactions so all reactants appear on the left and all products
appear on the right
™All intermediates must occur on both the right and the left so they
cancel
™Any reaction that is reversed must have the sign of its ∆H changed
™A reaction can be multiplied by a coefficient as necessary, but ∆H for
that reaction must be multiplied by the same coefficient.
Chem 102
Example: Calculate ΔH for the reaction
S(s) + 3/2 O2(g) J SO3(g)
From the data:
S(s) + O2(g) J SO2(g)
2 SO2(g) + O2(g) J 2 SO3(g)
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ΔH1 = − 296.8 kJ
ΔH2 = − 198.4 kJ
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Hess’s law requires tabulated data about the enthalpy change of a
reaction
many types of physical and chemical changes are tabulated
enthalpy of vapourization, ΔHvap &, enthalpy of fusion, ΔHfus
enthalpy of combustion, ΔHcomb
enthalpy of formation, ΔHf
Must define the physical state , temperature and
pressure to use tabulated values
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Thermodynamic Standard state
pure substance in its most stable form
1 atm pressure
25°C (chosen as reference state)
1 M concentration for all substances in solution
Standard Enthalpy change, ΔH°
a standard enthalpy of a reaction is the enthalpy change for a
reaction with reactants and products in their standard states
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Standard enthalpy of formation:
the enthalpy change ΔHf° for the (hypothetical) formation of 1 mol
of substance in its standard state from its constituent elements in
their standard states
example:
l for
f ethanol:
th
l C2H5OH(ℓ)
by definition, ΔHf° = 0 for
The ΔHf° values can be used in Hess’s law
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Example:
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What is the enthalpy change for this rxn?
NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(ℓ)
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Hess’s law using standard heats of formation can be generalized
for the example:
The value of the ΔHf° for a substance can be determined in a calorimeter by
measuring the heat evolved in a combustion experiment and using the known
ΔHf° for CO2 and H2O
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Example:
Combustion of 1 gram of 2,3,4-trimethylpentane (C8H18,
mol. mass = 114 g/mol) in a bomb calorimeter raises the temperature of the
calorimeter plus contents by 3.8 °C. The calorimeter’s total heat capacity is
11.66 kJ/°C. What is the heat of formation of 2,3,4-trimethylpentane?
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