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Ch. 5: Thermochemistry
OpenStax
Chemistry
Joseph DePasquale
Outline
5.1: Energy Basics
5.2: Calorimetry
5.3: Enthalpy
5.1: Energy Basics
 The energy changes that accompany chemical
reactions and physical changes are important to our
everyday lives.
 The reactions involved in metabolism of food.
 The reactions involved in generation of electricity.
 The reactions involved in the powering of our automobiles.

Sliding a match head along a rough surface initiates a combustion reaction that produces energy in the
form of heat and light. (credit: modification of work by Laszlo Ilyes)
Energy
•
Energy is the capacity to supply heat or do work.
•
All forms of energy are either kinetic or potential.
• Potential energy is the energy possessed by an object
by virtue of its position, composition, or condition.
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•
Kinetic energy (Ek) is the energy that an object posses
because of its motion.
•
One form of kinetic energy of particular interest to chemists is
thermal energy, which is the energy associated with the
random motion of atoms and molecules.
Thermal Energy
Potential Energy
• Potential energy is the energy
possessed by an object by virtue of its
position, composition, or condition.
• One important form of potential
energy of interest to chemists is
Chemical energy.
• Chemical energy is energy stored
within the structural units or chemical
bonds in a substance.
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If the water flows through generators at the bottom
of a dam, such as the Hoover Dam shown here, its
kinetic energy is converted into electrical energy.
(credit a: modification of work by Steve Jurvetson;
credit b: modification of work by
“curimedia”/Wikimedia commons)
Energy and Energy Changes
• Energy is interconvertible.
One form can be converted to
another form.
Water that is higher in elevation, for example, at the top of Victoria Falls, has a
higher potential energy than water at a lower elevation. As the water falls,
some of its potential energy is converted into kinetic energy. (credit a:
modification of work by Steve Jurvetson; credit b: modification of work by
“curimedia”/Wikimedia commons)
Law of conservation of energy – During a chemical or
physical change, energy can be neither created nor destroyed,
although its form can change.
Although energy can assume many forms, the total energy of
the universe is constant.
Thermal Energy and Temperature
• Thermal energy – Kinetic energy associated with the random
motion of atoms and molecules.
• Temperature – A quantitative measure of “Hot” and “Cold”.
• Fast moving molecules  High thermal energy  “Hot”
• Slow moving molecules  Low thermal energy  “Cold”
• Heat (q) – The transfer of thermal energy between two
bodies at different temperature.
• Thermochemistry – The study of the heat flow during a chemical
or physical change.
What do we know about heat?
 Heat will flow from a hotter object to a colder object.
(a)
Substances H and L are initially at different temperatures, and their atoms have different average kinetic energies.
(b)
When they are put into contact with each other, collisions between the molecules result in the transfer of kinetic
(thermal)energy from the hotter to the cooler matter.
(c)
The two objects reach “thermal equilibrium” when both substances are at the same temperature, and their molecules have the
same average kinetic energy.
Principles of Heat Flow
 Definitions:
 The System: The substance or substances undergoing the
chemical or physical change.
 The Surroundings: The rest of the universe.
 We are usually only concerned with the surroundings
that are in direct contact with the system.
 We will discuss the flow of heat between the system and
surroundings.
 Direction of Heat flow
 Matter undergoing a chemical or physical process can
absorb or release heat.
 Endothermic process – A reaction or change that
absorbs heat.
 Heat flows from the surroundings to the system.
 The sign on q is positive for an endothermic process.
 Exothermic process – A reaction or change that
releases heat.
 Heat flows from the the system to the surroundings.
 The sign on q is negative for an exothermic process.
Combustion of Methane (CH4)
CH4(g) + O2(g)  CO2(g) + H2O(g)
Experiment
 Mix solids samples of
barium hydroxide
and ammonium
nitrate then place on
a wet board.
Ba(OH)2(s) + NH4NO3  Ba(NO3)2(s) + NH3(g) + H2O(l)
After Mixing
 What has happened?
Units of Heat (q)
 Heat (q) is expressed in joules (or kilojoules)
 Named after James Joule (1818-1889), whom came up
with the technique known as calorimetry.
 Alternate unit of heat is the calorie
 1 calorie = 4.184 J
 1 kilocalorie = 4.184 kJ
 Nutritional calories on your food labels are
kilocalories (kcal).
Heat Capacity and Specific Heat Capacity
 Heat capacity (C) – Amount of heat either absorbed or
released by a body of matter when it experiences a
temperature change (DT) of 1 °C.
 Used with objects of fixed mass, that are typically
composed of more than one pure substance.
 Units: (J/ °C)
 Specific heat capacity (c) - Amount of heat needed to
raise the temperature of one gram of a substance by 1 °C.
 Also called “specific heat”
 Used with pure substances of variable mass.
 Units: (J/g . °C)
Thought Experiment
 Consider an iron pan and water. Each is placed in a 400
ºF oven for 5 minutes.
 Which one will feel hotter if you touch with
your bare hand?
Specific heat (c) - Amount of
heat needed to raise the
temperature of one gram of a
substance by one degree Celsius.
Some substances are easier to heat than others.
Climate
 Only two states in the United States have never
reached 100 °F. Which ones are they?
Magnitude of Heat Flow
 In a thermo chemical process, we are interested in both the
direction of heat flow and in the magnitude of heat flow.
 q = c × m × ΔT
q = C × ΔT
 q = heat
c = Specific heat (J/g . °C)
 m = mass (g)
C = Heat capacity (J/ °C)
 ΔT = Change in Temperature (° C)
5.2: Calorimetry
 A calorimeter is a device used to measure the amount of heat
and direction of heat flow in a chemical or physical process.
 The calorimeter contains water and/or other materials of known
heat capacity.
 The walls of the calorimeter are insulated so all of the heat flow is
between the process and the calorimeter components.
 System = Substance(s) undergoing the chemical or physical change.
 Surroundings = Calorimeter components (water, etc.)
 When thermal equilibrium is reached the system and
surroundings will be at the same final temperature.
 The heat flow for the reaction system is equal in magnitude and
opposite in sign of the heat flow of the calorimeter.
qreaction = - qcalorimeter
In a calorimetric determination, either (a) an exothermic process occurs and heat, q, is negative, indicating
that thermal energy is transferred from the system to its surroundings, or (b) an endothermic process
occurs and heat, q, is positive, indicating that thermal energy is transferred from the surroundings to the
system.
Types of Calorimeters
Coffee – cup calorimeter
• Made out of two styrofoam
coffee cups.
• Filled partially with water.
• Works well for reactions in
aqueous solution.
• You will use this type of
calorimeter in lab.
• Not the most accurate.
A simple calorimeter can be constructed from two
polystyrene cups. A thermometer and stirrer extend
through the cover into the reaction mixture.
Coffee-Cup Calorimeter
 Assumptions:
1) Reaction is done at constant pressure (atmospheric
pressure) because the walls of the cup are flexible.
2) All of the heat evolved by the reaction is absorbed by
the water.
Bomb Calorimeter
• The reaction is run in a “bomb” - a container with
heavy, sturdy walls.
• The bomb is submerged in water and kept in an
insulated outer container.
• The bomb calorimeter typically yields more
accurate results and is essential when the reaction
either
1) Produces large amounts of heat
2) Involves gases
Bomb Calorimeter
Bomb Calorimeter
 Assumptions:
1) Reaction is done at constant volume.
2) All of the heat evolved by the reaction is absorbed by
the water and the steel bomb.
 Cbomb is a value unique to every calorimeter, and can be
measured experimentally.
8.3: Enthalpy
 Most chemical and physical process happen at constant
pressure.
 Enthalpy (H) - The heat flow at constant pressure.
 It is most valuable to know the change in heat flow at
constant pressure - the change in Enthalpy (ΔH).
 Is the heat flow in a bomb calorimeter equal to ΔH?
 How about the heat flow in a coffee cup calorimeter?
State Variable (or Function)
• State Variable – A variable with a value that only
depends on the state of the system, not on the path
that the system took to reach that state.
• Altitude or Elevation is a state variable.
• Enthalpy is also a state variable
ΔH = Hproducts – Hreactants
• The sign of ΔH is the same as for heat flow:
ΔH > 0 for endothermic reactions
ΔH < 0 for exothermic reactions
Paths X and Y represent two different routes to the summit of Mt. Kilimanjaro. Both have the same change
in elevation (altitude or elevation on a mountain is a state function; it does not depend on path), but they
have very different distances traveled (distance walked is not a state function; it depends on the path).
(credit: modification of work by Paul Shaffner)
DH of Exothermic and Endothermic Reactions
Thermochemical Equations
 A thermochemical equation represents the changes in both
matter and energy in a chemical or physical process.
 ΔH is written on the right side of the chemical equation.
 Example:
H2 (g) + Cl2 (g)  2HCl (g)
ΔH = -185 kJ
 Conversion factors involving enthalpy can be be obtained
from a balanced thermochemical equation.
 -185 kJ/1 mole H2
 -185 kJ/1 mole Cl2
 -185 kJ/2 mol HCl
Conventions for Thermochemical Equations
1. The sign of ΔH indicates whether the reaction is
endothermic or exothermic at constant pressure.
2. The coefficients of the thermochemical equation
represent the number of moles of reactants and
products and can be used to derive useful conversion
factors.
3. The phases of all reactants and products must be
specified.
4. The value of ΔH applies when products and reactants
are at the same temperature, usually 25 °C, unless
stated otherwise.
ΔH Notes
• The magnitude of ΔH is directly proportional to the
amount of reactant or product.
• ΔH is an extensive property
 Consider the dissolving of CaCl2
CaCl2 (s)  Ca+2 (aq) + 2Cl- (aq)
ΔH = -81.4 kJ
What is ΔH when 5.00 g of calcium chloride dissolves in water?
ΔH Notes
•
ΔH for a reaction is equal in
magnitude but opposite in sign
of ΔH for the reverse reaction.

Consider the precipitation of CaCl2
Ca+2 (aq) + 2Cl- (aq)  CaCl2 (s)
Standard State
 ΔH is usually tabulated for reactions in which both the
reactants and products are at the same conditions.
 Standard state conditions:
 Pressure = 1 bar
 Concentration of dissolved species = 1 M
 There is no standard state temperature, but usually at 25
°C (298 K).
 ΔH° represents the change in enthalpy of a process at
standard conditions.
Standard Enthalpy of Combustion (ΔH°C)
 Standard Enthalpy of Combustion (ΔH°C) –
The enthalpy change when 1 mole of a
substance burns (undergoes a combustion
reaction) under standard conditions.
 Sometimes called heat of combustion or standard
molar enthalpy of combustion.
Standard Enthalpy of Combustion (ΔH°C)
Standard Enthalpy of Formation (ΔHf° )
 The standard molar enthalpy of formation, ΔHf°,
of a compound is the enthalpy change for when exactly
one mole of the compound is formed from its
elements in their most stable states at standard
conditions.
 ΔHf° of some common substances are recorded in
Appendix G in the back of the textbook.
Notes on Enthalpy of Formation Values (ΔHf°)
 ΔHf° for most compounds are negative.
 Exothermic process
 But some are endothermic.
Example: ΔHf° C2H2 = +226.7 kJ/mol.
 ΔHf° of an element in its most stable state at
standard conditions is equal to zero.
 ΔHf° Br2(ℓ) = 0
 ΔHf° Fe(s) = 0
 ΔHf° S8(s) = 0
Hess’s Law
•
•
•
For some reactions it is very difficult to measure ΔH.
Luckily there are ways to calculate ΔH.
Because ΔH is a state variable, its value is the same whether
the reaction occurs in one step or as a series of steps.
Hess’s Law: If a process can be written as the sum of
several stepwise processes, the ΔH of the total process
equals the sum of the enthalpy changes of the various steps.
• If,
equation 1 + equation 2 = equation 3
• Then,
ΔH1 + ΔH2 = ΔH3
Hess’s Law
Calculation of ΔH°
 The ΔH° of a reaction can also be calculated from the ΔHf°
values of the reactants and products by the following equation.
DH = å DHf products - å DHf reactants
 The symbol Σ refers to “the sum of”
 Elements in their standard states may be omitted, as their
enthalpies of formation are zero.
 The coefficients of reactants and products in the balanced
thermochemical equation must be accounted for.