Outline

Understanding Heats of Reaction

Thermodynamics is the science of the relationships between heat and
other forms of energy.

Thermochemistry is one area of thermodynamics. It concerns the study

of the quantity of heat absorbed or evolved (given off) by chemical
reactions.

Understanding Heats of Reaction

Practical reasons

 you could calculate the cost of the fuel per unit of heat energy
produced;
 you could calculate the quantity of heat obtained per unit mass of
fuel;
 you are able to calculate the amount of energy needed to break
a particular kind of chemical bond and so learn something about
the strength of that bond
A reaction that absorbs heat: Two crystalline substances, barium hydroxide
octahydrate and an ammonium salt, are mixed thoroughly in a flask. Then the
flask, which feels quite cold to the touch, is set in a puddle of water on a board. In
a couple of minutes, the flask and board are frozen solidly together. The board can
then be inverted with the flask frozen to it.
 Energy and Its Units
Energy can be defined as the potential or capacity to move matter. Energy
exists in different forms that can be interconverted.
A fuel is burned to heat water and generate steam. The steam expands against a
piston (or turbine), which is connected to a drive shaft that turns an electrical coil
in a magnetic field. Electricity is generated in the coil. The fuel contains chemical
energy, which is converted to heat. Part of the heat is then converted to motion
of the drive shaft, and this motion is converted to electrical energy. The
electrical energy could be used to run a motor, transforming the electrical energy
back to the energy of motion.

Conversion of light energy to kinetic energy. Solar-powered vehicles use

panels of photovoltaic cells.
 Energy and Its Units
Kinetic energy is the energy associated with an object by virtue of its
motion. The SI unit of energy, kg.m2/s2, is given the name joule (J)

The joule is an extremely small unit. Watt is a measure of the quantity of energy
used per unit time and equals 1 joule per second.

A 100-watt bulb, for example, uses 100 joules of energy every second. A kilowatt-
hour, the unit by which electric energy is sold, equals 3600 kilowatt-seconds, or 3.6
million joules

The calorie (cal) is a non-SI unit of energy commonly used by chemists, originally
defined as the amount of energy required to raise the temperature of one gram of
water by one degree Celsius.
 Energy and Its Units
Potential energy is the energy an object has by virtue of its position in a
field of force.

For example, water at the top of a dam. As a quantity of water falls over the
dam, its potential energy decreases from mgh at the top of the dam to zero at the
earth’s surface. The potential energy decreases and the kinetic energy increases.

Water at the top of the dam has

potential energy. As the water falls
over the dam, this potential energy is
converted to kinetic energy.
 Internal Energy
The sum of the kinetic and potential energies of the particles making up a
substance is referred to as the internal energy, U, of the substance.
Therefore, the total energy, Etot, of a quantity of water equals the sum of
its kinetic and potential energies as a whole (Ek + Ep) plus its internal
energy.

Law of Conservation of Energy

Energy may be converted from one form to another, but the total quantity
of energy remains constant.
System is the specific part of the universe that is of interest to us. For chemists,
systems usually include substances involved in chemical and physical changes.
For example, in an acid-base neutralization experiment, the system may be a
beaker containing 50 mL of HCl to which 50 mL of NaOH is added.

The surroundings are the rest of the universe outside the system.

There are three types of systems.

An open system can exchange mass and energy, usually in the form of heat
with its surroundings
a closed system, which allows the transfer of energy (heat) but not mass.
an isolated system, which does not allow the transfer of either mass or energy.
 Heat
Heat is defined as the energy that flows into or out of a system because of
a difference in temperature between the thermodynamic system and its
surroundings.
As long as a system and its surroundings are in thermal contact, energy
(heat) flows between them to establish temperature equality, or thermal
equilibrium. Heat flows from a region of higher temperature to one of
lower temperature; once the temperatures become equal, heat flow stops.
once heat flows into a system, it appears in the system as an increase in its
internal energy.
Heat and temperature are sometimes confused. When you add heat to a
gas, you increase its internal energy—and therefore its total kinetic energy.
This increase in kinetic energy will be distributed over the molecules in the
sample. Therefore, the increase in average kinetic energy per molecule
and thus the increase in temperature) depends on the size of the gas
sample. A given quantity of heat will raise the temperature of a sample
more if the sample is small.
Heat (q) is positive if heat is absorbed by the system and negative if heat is
evolved
 Heat of a reaction
 The heat of reaction (at a given temperature) is the value of q required to
returna system to the given temperature at the completion of the reaction.

 Chemical reactions or physical changes are classified as exothermic or

endothermic.
 An exothermic process is a chemical reaction or a physical change in which
heat is evolved (q is negative);
 An endothermic process is a chemical reaction or a physical change in which
heat is absorbed (q is positive).
 In the exothermic reaction, the reaction flask initially warms; in the
endothermic reaction, the reaction flask initially cools.
 Safety Match

Safety matches have a head containing mostly an oxidizing agent and

require a striking surface containing nonpoisonous red phosphorus.
 Enthalpy and Enthalpy Change
 Enthalpy (H) is an extensive property of a substance that can be used to
obtain the heat absorbed or evolved in a chemical reaction. (An extensive
property is a property that depends on the amount of substance. Other
examples of extensive properties are mass and volume.)
 Enthalpy is a state function. A state function is a property of a system
that depends only on its present state, which is determined by variables
such as temperature and pressure, and is independent of any previous
history of the system. This means that a change in enthalpy does not
depend on how the change was made, but only on the initial state and
final state of the system.

Musa Ibrahim: Tibet

Muhit/Nishat: Nepal

Altitude here is analogous to a

thermodynamic state function.
 Enthalpy and Enthalpy Change

The change in enthalpy for a reaction at a given temperature and pressure

is called the enthalpy of reaction.

The enthalpy of reaction equals the heat of reaction at constant pressure.

 Enthalpy and Internal Energy

Consider a reaction system at constant pressure P.

The last term in this equation (-P∆V) is the energy required by the system
to change volume against the constant pressure of the atmosphere. (The
minus sign means that energy is required to increase the volume of the
system.) This required energy is called the pressure–volume work.
 Enthalpy and Internal Energy

When hydrogen gas is released during the reaction, it pushes upward on

the piston and raises the weight. It requires energy to lift a weight
upward in a gravitational field. If you calculate this pressure–volume
work at 25oC and 1.00 atm pressure, you find that it is –P∆V=-2.5 kJ.
Problem: A certain gas expands in volume from 2.0 L to 6.0 L at constant
temperature. Calculate the work done by the gas if it expands (a) against a
vacuum and (b) against a constant pressure of 1.2 atm.
 Thermochemical Equations
A thermochemical equation is the chemical equation for a reaction
(including phase labels) in which the equation is given a molar
interpretation, and the enthalpy of reaction for these molar amounts
is written directly after the equation.

This equation says that 2 mol of sodium reacts with 2 mol of water to
produce 2 mol of sodium hydroxide and 1 mol of hydrogen gas, and
368.6 kJ of heat evolves.

Enthalpy change, ∆H, depends on the phase of the substances

 Thermochemical Equations

1.

2.
 Thermochemical Equations
Aqueous sodium hydrogen carbonate solution (baking soda solution) reacts with
hydrochloric acid to produce aqueous sodium chloride, water, and carbon dioxide
gas. The reaction absorbs 12.7 kJ of heat at constant pressure for each mole of
sodium hydrogen carbonate. Write the thermochemical equation for the reaction.
 Thermochemical Equations
When 2 mol H2(g) and 1 mol O2(g) react to give liquid water, 572 kJ of
heat evolves.
2H2(g) + O2(g) 2H2O(l ); ∆H = -572 kJ
Write this equation for 1 mol of liquid water. Give the reverse equation,
in which 1 mol of liquid water dissociates into hydrogen and oxygen.
Calculating the Heat of Reaction from the Stoichiometry
Calculating the Heat of Reaction from the Stoichiometry

How much heat is evolved when 9.07 ×105 g of ammonia is produced

according to the following equation? (Assume that the reaction occurs at
constant pressure.)
N2(g) + 3H2(g) 2NH3(g); ∆H = -91.8 kJ
 Heat Capacity

The heat capacity (C) of a sample of substance is the quantity of heat

needed to raise the temperature of the sample of substance one
degree celsius (or one kelvin). Changing the temperature of the
sample from an initial temperature ti to a final temperature tf requires
heat equal to

where ∆t is the change of temperature and equals tf - ti. The heat

capacity will depend on whether the process is constant pressure or
constant-volume.
The molar heat capacity of a substance is its heat capacity for one
mole of substance.
 Specific heat capacity
The specific heat capacity (or simply specific heat) is the quantity of
heat required to raise the temperature of one gram of a substance by
one degree Celsius (or one kelvin) at constant pressure.

The heat q required to raise the temperature of a sample, multiply the

specific heat of the substance, s, by the mass in grams, m, and the
temperature change, ∆t.
Problem: Iron metal has a specific heat of 0.449 J/(g.oC). How much
heat is transferred to a 5.00-g piece of iron, initially at 20.0oC, when it
is placed in a pot of boiling water?
 Measurement of Heat of Reaction

Calorimeter, a device used to measure the

heat absorbed or evolved during a physical or
chemical change.

The coffee-cup calorimeter is a constant-

pressure calorimeter. The heat of the
reaction is calculated from the temperature
change caused by the reaction, and since this
is a constant-pressure process, the heat can
be directly related to the enthalpy change,
∆H. Research versions of a constant-pressure
calorimeter are available, and these are used
when gases are not involved.
 Measurement of Heat of Reaction
For reactions involving gases, a bomb calorimeter is generally used. Consider
the heat of combustion of graphite. To measure the heat released when
graphite burns in oxygen, The graphite is surrounded by oxygen, and the
graphite and oxygen are sealed in a steel vessel, or bomb. An electrical circuit
is activated to start the burning of the graphite. The bomb is surrounded by
water in an insulated container, and the heat of reaction is calculated from
the temperature change of the calorimeter caused by the reaction.

Because the reaction in a bomb

calorimeter occurs in a closed vessel, the
pressure does not generally remain
constant. Rather the volume remains
constant, and under these conditions the
heat of reaction does not in general equal
∆H; a small correction is usually needed.
A lead (Pb) pellet having a mass of 26.47 g at 89.98°C was placed in a
constant-pressure calorimeter of negligible heat capacity containing
100.0 mL of water. The water temperature rose from 22.50°C to 23.17°C.
What is the specific heat of the lead pellet?
Suppose 0.562 g of graphite is placed in a calorimeter with an excess of oxygen at
25.00 oC and 1 atm pressure. Excess O2 ensures that all carbon burns to form CO2.
The graphite is ignited, and it burns according to the equation
C(graphite)+O2(g) CO2(g)
On reaction, the calorimeter temperature rises from 25.00oC to 25.89oC. The heat
capacity of the calorimeter is 20.7 kJ/oC. What is the heat of reaction at 25.00oC
and 1 atm pressure? Express the answer as a thermochemical equation.
Suppose 33 mL of 1.20 M HCl is added to 42 mL of a solution containing excess
sodium hydroxide, NaOH, in a coffee-cup calorimeter. The solution
temperature, originally 25.0oC, rises to 31.8oC. Give the enthalpy change, ∆H,
for the reaction HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
Express the answer as a thermochemical equation.
 Hess’s Law
In 1840, the Russian chemist Germain Henri Hess, a professor at the
University of St. Petersburg, discovered this result by experiment.
Hess’s law of heat summation states that for a chemical equation that
can be written as the sum of two or more steps, the enthalpy change
for the overall equation equals the sum of the enthalpy changes for
the individual steps.

excess of oxygen form only carbon dioxide,

limited quantity of oxygen form a mixture of carbon monoxide and
carbon dioxide
How can you obtain the enthalpy change for the preparation of pure
carbon monoxide from graphite and oxygen?
 Hess’s Law

or

Apply Hess’s law to obtain the enthalpy change for a reaction that is
difficult to determine by direct experiment.
 Enthalpy diagram illustrating Hess’s Law

The diagram shows two different ways to go from graphite and oxygen
(reactants) to carbon monoxide (products). Going by way of reactions 1 and
2 is equivalent to the direct reaction 3.
Problem: What is the enthalpy of reaction, ∆H, for the formation of
tungsten carbide, WC, from the elements?
The enthalpy change for this reaction is difficult to measure directly,
because the reaction occurs at 1400oC. However, the heats of combustion
of the elements and of tungsten carbide can be measured easily:

Solution
 Problem: Manganese metal can be obtained by reaction of manganese
dioxide with aluminum.
What is ∆H for this reaction? Use the following data:

Solution:
 Standard Enthalpies of Reaction

The term standard state refers to the standard

thermodynamic conditions chosen for substances when
listing or comparing thermodynamic data: 1 atm pressure
and the specified temperature (usually 25oC).

The enthalpy change for a reaction in which reactants in

their standard states yield products in their standard states
is denoted ∆Ho. The quantity ∆Ho is called the standard
enthalpy of reaction.
 Standard Enthalpies of Formation

An allotrope is one of two or more distinct forms of an element in the same

physical state
rhombic sulfur monoclinic sulfur

Left: An evaporating dish contains rhombic sulfur, the stable form of the
element at room temperature. Right: When this sulfur is melted, then
cooled, it forms long needles of monoclinic sulfur, another allotrope. At
room temperature, monoclinic sulfur will slowly change back to rhombic
sulfur. Both forms contain the molecule S8, depicted by the model.
O2, O3
Graphite and diamond
 Standard Enthalpies of Formation

The reference form of an element for the purpose of

specifying the formation reaction is usually the stablest form
(physical state and allotrope) of the element under standard
thermodynamic conditions. The reference form of oxygen at
25oC is O2(g); the reference form of carbon at 25oC is
graphite.

The standard enthalpy of formation (also called the standard

heat of formation) of a substance, denoted ∆Hfo, is the
enthalpy change for the formation of one mole of the
substance in its standard state from its elements in their
reference form and in their standard states.
Standard Enthalpies of Formation
Fuels—Foods, Commercial Fuels, and Rocket Fuels
A fuel is any substance that is burned or similarly reacted to provide heat
and other forms of energy.

The earliest use of fuels for heat came with the control of fire, which was
achieved about 750,000 years ago.
 Foods as Fuels

Foods fill three needs of the body:

they supply substances for the growth and repair of
tissue,
they supply substances for the synthesis of compounds
used in the regulation of body processes,
they supply energy.

About 80% of the energy we need is for heat. The rest is

used for muscular action, chemical processes, and other
body processes. The human body requires about as much
energy in a day as does a 100-watt light bulb.
 Foods as Fuels

One gram of glucose yields 15.6 kJ (3.73 kcal) of heat when

burned.

glyceryl trimyristate (fat)

One gram of this fat yields 38.5 kJ (9.20 kcal) of heat when
burned.
 Fossil Fuel

 All of the fossil fuels in existence today were created millions of years
ago when aquatic plants and animals were buried and compressed by
layers of sediment at the bottoms of swamps and seas.
 Over time this organic matter was converted by bacterial decay and
pressure to petroleum (oil), gas, and coal.

 Anthracite, or hard coal, the oldest variety of coal, was laid down as
long as 250 million years ago and may contain over 80% carbon.

 Bituminous coal, a younger variety of coal, has between 45% and 65%
carbon.
 Fossil Fuel
Fuel values of coals are rated in Btu’s (British thermal units) per pound, which are
essentially heats of combustion per pound of coal.

A typical value is 13,200 Btu/lb. A Btu equals 1054 J, so 1 Btu/lb equals 2.32 J/g.

32.8 kJ/g

This value of ∆Ho is equivalent to 50.1 kJ per gram of fuel.

octane

This value of ∆Ho is equivalent to 44.4 kJ/g. These combustion values

indicate another reason why the fluid fossil fuels are popular: they release
more heat per gram than coal does.
Fossil Fuels
 Coal Gasification and Liquefaction
 It has been estimated that petroleum supplies will be 80% depleted by
about the year 2030.

 Natural-gas supplies may be depleted even sooner.

 Coal supplies are sufficient to last several more centuries.

 This abundance has spurred much research into developing commercial

methods for converting coal to the more easily handled liquid and
gaseous fuels.

Different catalysts and different reaction conditions result in liquid fuels. An

added advantage of coal gasification and coal liquefaction is that sulfur, normally
present in fossil fuels, can be removed during the process. The burning of sulfur-
containing coal is a major source of air pollution and acid rain.
 Rocket Fuel
Rockets are self-contained missiles propelled by the ejection of gases from an
orifice. Usually these are hot gases propelled from the rocket by the reaction of
a fuel with an oxidizer.

Hydrogen is the element of lowest density, and at the same

time it reacts exothermically with oxygen to give water. 120
kJ/g of fuel (H2)

The first stage of liftoff used kerosene and oxygen, and an

unbelievable 550 metric tons (550x103 kg) of kerosene were
burned in 2.5 minutes!

The landing module for the Apollo mission used a fuel made
of hydrazine, N2H4, and a derivative of hydrazine. The oxidizer
was dinitrogen tetroxide, N2O4.
Solid propellants are also used as rocket fuels