CP Chapter 17 Thermochemistry
Thermochemistry
 Thermochemistry is the study of energy __________________ that occur during chemical reactions and
phase changes (changes of state)
The Nature of Energy
 Energy is the ability to do work or produce __________________
 Heat is energy transferred from a sample of ____________________ temperature to
____________________ temperature. Represented by the letter Q or q.
Kinetic vs. Potential Energy
 Two main types of energy – kinetic and potential
o Kinetic – energy of _________________
o Potential – energy due to position or energy ____________________ in chemical bonds
 Chemical potential energy - the energy stored in a substance because of its composition
 Example: gasoline
Law of Conservation of Energy
 Law of conservation of energy states that in any chemical reaction or physical process, energy can be
________________________ from one form to another, but it is neither created nor destroyed.
Temperature vs. Heat
 Temperature is a measure of the average kinetic energy of the molecular _____________________ in a
sample. The Kelvin temperature and kinetic energy are ____________________ proportional.
Temperature is a _______________________________ of energy.
 Heat is the total energy of molecular motion, dependent upon amount, size, and type of particles. Heat is
_______________________.
System and Surroundings
 System – The part of the universe on which contains the ____________________ or process you wish to
study.
 Surroundings – Everything in the universe _________________ than the system.
 When heat is transferred it can flow in or out of the system
Endothermic vs. Exothermic
 An Exothermic process is one that _____________________ (evolved) heat to its surroundings (feels
warm)
o Energyproducts < Energyreactants
 An Endothermic process is one that _____________________ heat from the surroundings (feels cold)
o Energyproducts > Energyreactants
Exothermic Process
Endothermic Process
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Measuring Heat
 calorie - the amount of heat required to raise the temperature of one gram of pure _________________by
one degree Celsius
 Calorie – nutritional calorie;
o 1 Calorie = 1000 calories = 1 kilocalories (kcal)
 Joule – SI unit of heat
o 1 calorie = ________________________J
Specific Heat = Cp
 Specific heat of a substance is the amount of heat required to raise the temperature of one gram of that
________________________ by one degree Celsius.
 Unit for specific heat is J/goC
Specific Heat cont . . .
 Each substance has a _________________________ specific heat
 Water = 4.184 J/goC
 Gold = 0.129 J/goC
 Copper = 0.386 J/goC
 The lower the specific heat the lower the amount energy is required to raise its temperature.
Calculating Heat Released and Absorbed
 Q = m(T)Cp
o Q = Heat
o m is mass
o T is temperature change Tfinal -Tinitial
o Cp is specific heat at a constant pressure
 The unit for heat (Q) is Joules (J) or kilojoules (kJ)
 If reaction is ______________________________ heat (Q) is positive
 If reaction is exothermic heat (Q) is ___________________________
Converting Energy Units
Calorie/calorie/ kilocalorie
calorie/Joule
A cereal has 155 nutritional Calories per serving. How many calories, kilocalories and Joules is this?
A person on a diet consumed 1350 Calories in one day. How many calories, kilocalories and Joules is this?
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Calculating Heat
If the temperature of 56.6 g of ethanol increases from 45.0 oC to 80.0oC, how much heat has been absorbed by
the ethanol? Specific heat of ethanol = 2.44 J/g oC
A 4.00 g sample of a substance was heated from 274 K to 314 K and absorbed 32 J of heat, what is the specific
heat of the substance?
If 98,000 J of energy are added to 6200 g of water at 14.00oC, what will the change in temperature of the water
be? Specific heat of water = 4.184 J/g oC
Calorimetry and Enthalpy
Calorimetry
 Calorimetry is the _________________________ of the heat flow into or out of a system for chemical or
physical processes
o Based on the law of conservation of ___________________________
Calorimetry Example
 Heat __________________ by the system is equal to the heat absorbed by the surroundings
 Heat absorbed by the system is equal to the heat released by the surroundings
 An ice cube is added to a warm cup of water.
o The amount of heat used to melt the ice cube is the same amount of heat lost by the warm water.
Measuring Heat
 Heat changes that occur during chemical and physical processes can be measured by a
________________________.
 A calorimeter is an _________________________ device for measuring the amount of heat absorbed or
released during chemical or physical processes
Enthalpy and Enthalpy Changes
 Enthalpy (H) is the ____________________ content of a system at constant
____________________________
 Thermochemistry uses the _____________________ in enthalpy(H) to study heat changes
 At constant pressure, Q = H
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Endothermic and Exothermic Processes
 All chemical and physical changes release or absorb heat
 Exothermic process heat is released and goes from the system to the surroundings
o Feels ________________ - Eproducts < Ereactants
 Endothermic process heat is absorbed and goes from the surrounding to the system
o Feels ________________ - Eproducts > Ereactants
Heat of Reaction
 Heat of reaction (Hrxn) – The change of enthalpy in a _____________________ reaction
 Exothermic reaction Hrxn is negative
o Products have a ____________________ enthalpy than the reactants (Energy released) H products <
Hreactants
 Endothermic reaction Hrxn is positive
o Products have a ____________________ enthalpy than the reactants (Energy absorbed) Hproducts >
Hreactants
Thermochemical Equation
 A thermochemical equation is a balanced chemical equation that includes the physical states of all
reactants and products (1 atm, 25 oC) and the energy _________________________, expressed as a
change in enthalpy, H
 NH4NO3(s)  NH4+(aq) + NO3-(aq) H = 27 kJ
 Endothermic
 CaO(s) + H2O(l)  Ca(OH)2(s) H = -65.2 kJ
 Exothermic
Thermochemical Equations for Endothermic Reactions
 Endothermic : H is positive; heat is on the _______________________ side of the equation (Heat
absorbed)
 NH4NO3(s)  NH4+(aq) + NO3-(aq) H = 27 kJ
 NH4NO3(s) + Heat  NH4+(aq) + NO3-(aq)
 NH4NO3(s) + 27kJ  NH4+(aq) + NO3-(aq)
Thermochemical Equations for Exothermic Reactions
 Exothermic : H is negative; heat is on the _________________________ side of the equation (Heat
released)
 CaO(s) + H2O(l)  Ca(OH)2(s) H = -65.2 kJ
 CaO(s) + H2O(l)  Ca(OH)2(s) + Heat
 CaO(s) + H2O(l)  Ca(OH)2(s) + 65.2 kJ
Heat of Combustion
 Heat of combustion is the enthalpy change for the complete burning of one mole of the substance
 CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -891kJ
 CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) + Heat
 CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) + 891kJ
A piece of metal with a mass of 4.68 g absorbs 2556 J of heat when its temperature increases by 182oC.
What is the specific heat of metal?
How much heat is released by the combustion of 250.0 g of octane, C 8H18?
Octane ΔHcomb = -5471 kJ/mol
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Changes of State (Phase Changes)
Phase Change
 A phase change is going from one _____________________ of matter to another (Physical change)
o Gas to liquid
o Liquid to solid
 Changes of state either ____________________ or ________________________ (evolve/liberate)
energy
Phase Changes cont. .
 Endothermic phase changes (Absorb energy)
o Melting (solid to liquid)
o Evaporation/___________________________ (liquid to gas)
 Exothermic phase changes (Release energy)
o ____________________________ (gas to liquid)
o Freezing (liquid to solid)
Temperature and Phase Change
 During a phase change, as energy is added or removed, there is _________ temperature change.
Latent Heat
 Since there is no __________________________ change, we cannot use Q = m(T)Cp to calculate
enthalpy change.
 Instead we use latent (hidden) heat, which is the quantity of heat released or absorbed during a phase
change.
 Units are usually Joule/gram or kilojoule/mole
Changes of State
 Latent Heat of Vaporization is the heat required to ________________________ one mole of a liquid
o Going from liquid to gas
o For water, Hvap = 40.7 kJ/mol
 Latent Heat of Condensation is the heat required to _______________________ one mole of a gas
o Going from gas to liquid
o For water, Hcond = -40.7 kJ/mol
Thermochemical Equations for Vaporization and Condensation
 The reverse process of vaporization is condensation
o The _________________________ process has an opposite sign
o Hvap = - Hcond
o H2O(l)  H2O(g) Hvap = 40.7 kJ/mol
o H2O(g)  H2O(l) Hcond = -40.7 kJ/mol
Changes of State
 Latent Heat of Fusion is the heat required to _______________________ one mole of a solid substance.
o Going from solid to liquid
o For water, Hfus = 6.01 kJ/mol (Endothermic)

Latent Heat of Solidification is the heat required to _____________________ one mole of a liquid
substance.
o Going from liquid to solid
o For water, Hsolid = -6.01 kJ/mol (Exothermic)
Thermochemical Equations for Fusion and Solidification
 The reverse process of fusion is solidification.
o The reverse process has an ________________________ sign
o Hfus = - Hsolid
o H2O(s)  H2O(l) Hfus = 6.01 kJ/mol
o H2O(l)  H2O(s) Hsolid = -6.01 kJ/mol
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Heating Curve
Explanation of Heating Curve
 When there is a temperature change use Q=m(T)Cp to calculate the amount of energy used.
 When the heating curve is flat (no temperature change), there is a phase change. Use the latent heat to
calculate the amount of energy used.
Calculate the heat required to melt 45.6 g of water at its melting point .
Latent heat of fusion (∆Hfus) = 6.01 kJ/mol
Calculate the heat evolved to condense 75.4 g of water at its boiling point
Latent heat of vaporization (∆Hvap) = 40.7 kJ/mol
What mass of ammonia (NH3) must be vaporizes that absorbs 345 kJ of heat?
Latent heat of vaporization (∆Hvap) = 23.3 kJ/mol
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Hess’s Law - Calculating Enthalpy Change (Heat of Reaction)
Hess’s Law states that if you can ___________ two or more __________________________ equations to
produce a __________________ equation for a reaction, then the sum of the enthalpy changes for the individual
reactions is the ____________________ change for the____________________ reaction.
General principles for combining thermochemical equations.
1. If a reaction is reversed, the sign of H is reversed
2. If the coefficients are multiplied, then the H is multiplied by the same factor.
Calculate the enthalpy change for:
2S(s) + 3O2(g)  2SO3(g) ∆H = ?
S(s) + O2(g)  SO2(g) ∆H = -297 kJ
2SO3(g)  2SO2(g) + O2(g) ∆H = 198 kJ
2NO(g) + O2(g)  2NO2(g) ∆H = ?
N2(g) + O2(g) - 2NO(g) ∆H = 180.6 kJ
N2(g) + 2O2(g)  2NO2(g) ∆H = 66.4 kJ
2C (s) + O2(g)  2CO(g) ∆H = ?
C(s) + O2(g)  CO2(g) ∆H = 393.5 kJ
2CO + O2  2CO2 ∆H = -566kJ
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C(s) + 2H2(g)  CH4(g) ∆H = ?
C(s) + O2(g)  CO2(g)
2H2(g) + O2(g)  2H2O(l)
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
∆H =-393.5 kJ
∆H =-571.6 kJ
∆H = -890.8 kJ
Standard Heat of Formation



Standard heat of formation is defined as the _______________ in ______________ that
accompanies the ___________________ of one mole of the compound at its ________________
state from its ____________________ elements in their standard state.
Standard state is _________________________________________
Every free element in its standard state is assigned a ∆Hf of _____________. Free elements include
___________________ elements and any ___________________ molecule (H 2O2N2Cl2Br2I2F2)
Hrxn =
___________________________________________________________________________________________________
Use the standard enthalpies of formation to calculate ∆Hrxn
CH4(g) +2O2(g)  CO2(g) + 2H2O(l)
∆Hf(CO2(g)) = -394kJ
∆Hf(H2O(l)) = -286 kJ
∆Hf(CH4(g)) = -75kJ
∆Hf(O2(g)) = 0 kJ
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Use Table C-13 to calculate ∆Hrxn for the following reactions
2HF(g)  H2(g) + F2(g)
2H2S(g)
+ 3O2(g)  2H2O(l)
+ 2SO2(g)
4Fe(s) + 3O2(g)  2Fe2O3 (s)
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2NO(g)
+ O2(g)
 2NO2(g)
2CO(g)
+ O2(g)
 2CO2(g)
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