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Thermochemistry
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Thermochemistry- heat changes that
occur during chemical reactions
 Introductory Objectives
1. Explain the relationship between energy
and heat.
2. Distinguish between heat capacity and
specific heat.
TIP : Do not confuse standard conditions with
STP used in gas law calculations.
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Energy
 Energy - the capacity for doing work or supplying heat
 Chemical potential energy – the energy stored within
the structural units of chemical substances
 Different substances store different amounts of energy.
The kinds of atoms and their arrangement in the
substance determine the amount of energy stored in
the substance.
 All energy in a process can be accounted for as work,
stored energy, or heat
 Law of Conservation of Energy – In any physical or
chemical process, energy is neither created not
destroyed
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Heat (q)
 Heat – energy that transfers from one object
to another because of a temperature
difference between them
 Cannot be detected by the senses or
instruments
 Only changes caused by heat can be
detected
 Always flows from a warmer object to a cooler
object
 If two objects remain in contact, eventually
the temperature of both objects will be the
same
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Systems
 System- part of the universe on which you
focus your attention
 Surroundings – include everything else in the
universe
 Universe – the system and the surroundings
 Example: Chemicals and water are in a
beaker. (Universe) Your system includes the
chemicals and water. The beaker is the
surrounding.
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Endothermic and Exothermic
Reactions
 Endothermic reaction - heat, q, flows into a
system (heat absorbed), >0 (positive number)
Examples: melting of ice, evaporation of a
puddle, sublimation of a mothball, heat used to
cook food
During endothermic phase changes, energy
absorbed does not increase the temperature
because the energy is being used to overcome
attractions between particles.
Bond-breaking
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What value of q is endothermic?
 Exothermic reaction – heat, q, flows out of the
system (heat is given off), <0 (negative
number)
Examples: combustion of fossil fuels, cooling
of skin as perspiration evaporates, freezing of
water
Bond-formation
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Heat vs. Temperature
 Temperature – a measurement of the
average kinetic energy of the particles
Can be detected with a thermometer
 Heat cannot be measured with a
thermometer
 Heat can also increase the potential (rather
than kinetic) energy. This occurs during
phase changes: solid to liquid AND liquid to
gas
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Think!
 Suppose two identical candles are used to heat two
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samples of water. One sample is a cup of water; the
other is 10 gallons of water in a drum.
1. How will the change in temperature of the samples
compare?
Practically no change in the drum; a large increase in
the cup
2. How will the amount of heat received by each
container compare?
Both containers receive the same amount of heat
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 Joule – the SI unit of heat and energy
A joule of heat raises the temperature of 1 g
of pure water 0.2390 °C
 1 J = 0.2390 cal
4.184 J = 1 cal
 Heat capacity – amount of heat needed to
increase the temperature of an object exactly
1°C
 Besides varying with mass, the heat capacity
of an object also depends on its chemical
composition
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Heat Capacity and Specific Heat
 calorie- the quantity of heat needed to raise
the temperature of 1g of pure water 1°C
 Calorie = 1000 calories (refers to energy in
food)
 1 Calorie = 1kilocalorie = 1000 calories
 “10g of sugar has 41 Calories” means that
10g of sugar releases 41 kilocalories of heat
when completely burned to produce carbon
dioxide and water
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Questions
 What is the relationship between a joule and
a calorie? Calorie and a dietary calorie?
 What is the difference between specific heat
capacity and heat capacity? Give examples.
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 Note that the temperature of water changes
less than the temperature of iron because the
specific heat capacity of water is larger.
 Specific heat capacity (specific heat) – the
amount of heat it takes to raise the
temperature of 1 g of the substance 1°C
 Specific heat (C) is a measure of a substance
to store heat. The specific heats of
substances can be compared because the
quantity (1 g) of matter involved is specified.
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Specific Heat
 C = q ÷ (m x
T)
C or cp = specific heat at a given pressure
q = energy lost or gained (heat)
m = mass of the sample
T = difference between initial
temperature and final temperature
 q = cp x m x T
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Objectives
 1. Construct equations that show the heat
changes for chemical and physical processes
 2. Calculate heat changes in chemical and
physical processes
 Think! A match won’t ignite unless you strike
it and add the heat produced from friction. Is
the burning of a match an endothermic
reaction?
 Is there a way to measure how much heat is
released from a burning match?
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Answer to Think!
 No; the reaction releases more energy in the
form of heat and light than the amount of
energy it absorbs to start.
 Yes, but only indirectly. If the reaction were
confined, then any temperature changes in
the surroundings could be attributed to heat
transfer from the reaction.
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Calorimetry
 The accurate and precise measurement of heat
change for chemical and physical processes
 Need insulated container
1. Constant pressure calorimeter
2. Bomb calorimeter – constant volume
Measures the heat released from burning a
compound; closed system: the mass of the system is
constant
 The heat released by the system is equal to the heat
absorbed by its surroundings
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Enthalpy (H)
 Heat changes for reactions carried out at constant
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pressure
Because the reactions presented in the textbook
occur at constant pressure, the text uses heat and
enthalpy interchangeably
Heat change for a chemical reaction carried out in
aqueous solution:
q= H=mxCx T
Reacting chemicals = system
Known volumes of water = surroundings
Exothermic – negative number
Endothermic – positive number
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Thermochemical Equations
 An equation that includes the heat change
 Heat of reaction – the heat change for the
reaction exactly as it is written (Usually heat
change at constant pressure)
 The physical state of the reactants and
products must be given
 Standard conditions = 101.3 kPa (1atm) and
25 °C
 Amount of heat absorbed or released
depends on the number of moles
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Heat of Combustion
 Heat of reaction for the complete burning of
one mole of a substance
 Like other heats of reaction, heats of
combustion are reported as the enthalpy
changes when the reactions are carried out at
101.3 kPa of pressure and the reactants and
products are in their physical states at 25 °C
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Objectives
 Classify, by type, the heat changes that occur
during melting, freezing, boiling, and
condensing
 Calculate heat changes that occur during
melting, freezing, boiling, and condensing
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Review
 Specific heat capacity (specific heat) – the amount of heat it
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takes to raise the temperature of 1 g of the substance 1°C
C = q ÷ (m x T)
Enthalpy (H) - Heat changes for reactions carried out at constant
pressure
Heat change for a chemical reaction carried out in aqueous
solution:
q= H=mxCx T
Like other heats of reaction, heats of combustion are reported
as the enthalpy changes when the reactions are carried out at
101.3 kPa of pressure and the reactants and products are in
their physical states at 25 °C
H is enthalpy or heat content
H represents a change in the heat content
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Key Terms and Concepts
 Solid ----------------Liquid -------------------Vapor
+ Fusion
+Vaporization
Low enthalpy-----------------------High enthalpy
 Molar heat of fusion ( Hfus) – heat absorbed
by one mole of a substance in melting from a
solid to a liquid at a constant temperature
 Molar heat of vaporization ( Hvap) – the
amount of heat necessary to vaporize one
mole of a given liquid
 Endothermic reactions
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 Vapor------------------Liquid------------------Solid
-Condensation
-Solidification
High enthalpy-----------------------Low enthalpy
 Molar heat of condensation ( Hcond) –
amount of heat released when one mole of
vapor condenses
 Molar heat of solidification ( Hsolid) – the heat
lost when one mole of a liquid solidifies at a
constant temperature
 Exothermic reactions
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 Solid ----------------Liquid -------------------Vapor
+Fusion
+Vaporization
-Solidification
-Condensation <
 The molar heat of fusion is the heat absorbed
by one mole of a substance in melting from a
solid to a liquid at a constant temperature.
The heat lost when one mole of a liquid
solidifies at a constant temperature is the
molar heat of solidification. Because energy
is conserved in all chemical and physical
changes, the quantity of heat absorbed by the
melting solid must equal the quantity of heat
lost when the liquid solidifies.
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Hfus = - Hsolid
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Hvap = - Hcond
Values are numerically the same, but the
values have different signs
 Fusion—endothermic—Vaporization (+)
 Solidification – exothermic—Condensation (-)
 The melting of one mole of ice at 0°C to one
mole of water at 0°C requires the absorption
of 6.01 kJ of heat. What is the heat of fusion?
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 The heat of fusion is 6.01 kJ/mol.
 The heat of solidification is -6.01 kJ/mol.
 Ice is commonly used to refrigerate
perishable foods. What happens to the
temperature of the ice as it begins to melt?
 The ice and the water are both at 0°C. The
temperature will not rise above 0°C until all of
the ice has melted.
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Problem Solving
 How many grams of ice at 0°C and 101.3 kPa
could be melted by the addition of 2.25 kJ of
heat?
 Standard conditions for ice exist. Use heat of
fusion for water.
 Grams for one mole (18g)/6.01 kJ = x/2.25 kJ
 Answer = 6.74 g ice
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Heat of Solution
Hsoln – heat change caused by the
dissolution of one mole of a substance
 Examples:
1. Exothermic molar heat of solution:
sodium hydroxide dissolved in water, hot
pack that mixes calcium chloride and water
2. Endothermic molar heat of solution: cold
pack that allows water and ammonium nitrate
to mix (Heat is released from the water and
the temperature of the solution decreases.)
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Problem Solving
 How much heat (in kJ) is absorbed when
24.8 g of H2O(l) at 100°C is converted to
steam at 100°C ?
 Use heat of vaporization for water.
 24.8g/x = 18g/40.7 kJ
 Answer = 56.1 kJ
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Objectives
 Apply Hess’s law of heat summation to find
heat changes for chemical and physical
processes
 Calculate heat changes using standard heats
of formation
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Hess’s Law of Heat Summation
 Hess’s law of heat summation – If you add
two or more thermochemical equations to
give a final equation, then you can also add
the heats of reaction to give the final heat of
reaction
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Standard Heats of Formation
 Standard heat of formation ( Hf°) – the
change in enthalpy that accompanies the
formation of one mole of a compound from its
elements with all substances at their standard
states at 25°C
 The standard heat of formation of a free
element in its standard state is arbitrarily set
at 0. (Includes diatomic molecules and
graphite form of carbon)
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Standard Heat of Reaction
 The standard heat of reaction ( H°) is the
difference between the standard heats of
formation of all the reactants and products.
 H° = Hf° (products) - Hf° (reactants)
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