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• Define the terms exothermic reaction,
endothermic reaction and standard enthalpy
change of reaction (∆Hѳ)
• State that combustion and neutralization are
exothermic processes.
• Apply the relationship between temperature
change and the classification of a reaction as
endothermic and exothermic.
• Deduce, from an enthalpy level diagram, the
relative stabilities of reactants and products,
and the sign of the enthalpy change for the
reaction.
Exothermic and Endothermic
reactions
 Chemical reactions involve the ________ and
__________of bonds.
 It requires energy to break bonds and energy is given
out when new bonds are formed.
 If more energy is given out than needs to be put in,
then this will be released as heat to the surroundings,
and the reaction is said to be___________.
 If more energy is required to break the reactant
bonds than is given out when the products are formed,
then extra heat will be need to be added, and the
reaction is said to be____________.
making
endothermic
breaking exothermic
Enthalpy
 The energy contained in chemical bonds that can
be converted into heat is known as enthalpy and is
given the symbol H.
 It is impossible to measure the actual heat
content or enthalpy of a particular substance, but
what can be readily measured is the enthalpy
change for a reaction, ΔH.
Exothermic reactions
 For exothermic reactions the heat content (enthalpy) of
the products is less than that of the reactants.
Reactants
Enthalpy
- ΔH
Products
Extent of the reaction
 By convention the enthalpy change for the reaction, ΔH,
is said to be negative (-).
Examples of exothermic
reactions
 Examples of exothermic reactions include:
 the combustion of fuels, such as hydrogen or gasoline
(petrol)
C8H18(l) + 12½O2(g) → 8CO2(g) + 9H2O(l) ΔH = -5512kJ mol-1
 the neutralization of an acid by a base:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ΔH -57.3kJ mol-1
Endothermic reactions
 For endothermic reactions, where the enthalpy of the
products is more than the enthalpy of the reactants, the
enthalpy change, ΔH, will have a positive value (+).
Enthalpy
Products
+ ΔH
Reactants
Extent of the reaction
 An example of an endothermic reaction is the combination
of nitrogen with oxygen to form nitrogen dioxide.
N2(g) + O2(g) → NO2(g) ΔH = +33.9 kJ mol-1
 Another example of an endothermic reaction is the
reaction between solid hydrated barium hydroxide and
solid ammonium chloride:
Ba(OH)2.8H2O(s) + 2NH4Cl(s) → BaCl2.2H2O(s) + 2NH3(g/aq) + 8H2O(l)
ΔH = +96.0 kJ mol-1
 This is interesting not only because it is a reaction
between two solids, but also because the temperature
decrease is such that the flask will often stick to the
bench after the reaction.
 When there is a change of state from solid to a liquid,
or a liquid to a gas at a constant temperature, the
process is also endothermic.
 For example:
Na (s) → Na (g) ΔH = +108 kJmol-1
Standard Enthalpy Change of
Reaction
 The actual amount of heat given out or taken in during a
chemical reaction will depend on several factors:
 The nature of the reactants and products; different




reactions have different values of enthalpy changes, ΔH, as
different bonds with different strengths are being broken
and formed.
The amount (or concentration) of reactants; the greater
the amount that reacts , the greater the heat change.
The states of the reactants and products – changing state
involves an enthalpy change, and so will affect the total
amount of heat given out or taken in.
The temperature of the reaction
The pressure at which the reaction is carried out
 All these factors are taken into account by referring
to standard enthalpy changes.
 The standard enthalpy change of the reaction, ΔH θ,
is defined as:
the enthalpy change when molar quantities of reactants in
their normal states react to form products in their
normal states under standard conditions of temperature
and pressure
 Standard pressure is 101.3 k Pa (1 atmosphere) and
standard temperature is any specified temperature,
but usually this is taken as 298 K (25oC).
 The symbol θ signifies that the enthalpy change has
been measured under standard conditions.
 It is worth noting that enthalpy changes are
normally carried out under constant pressure.
 If the volume is kept constant then the value for
the heat change will be slightly different from
that at constant pressure, because some work is
involved in expanding or contracting a gas.
 Strictly speaking the enthalpy change, ΔH, is the
heat change if the pressure is kept constant; if
the volume is kept constant then the total change
is designated by ΔU.
Activity
 Carry out the experiment: Measuring some energy
changes.
Contents
 Calculate the heat energy change when the
temperature of a pure substance is changed.
 Design the suitable experimental procedures for
measuring the heat energy changes of reactions.
 Calculate the enthalpy change for a reaction using
experimental data on temperature changes,
quantities of reactants and mass of water
 Evaluate the results of experiments to determine
enthalpy changes.
Measuring enthalpy changes
 The general name for any reaction is the standard
molar enthalpy change of reaction, ΔH θ.
 It is measured in kilojoules per mole, kJ mol-1.
 We write a balanced symbol equation for the
reaction and then find the heat change for the
quantities in moles given by the reaction.
 For example, ΔH for: 2NaOH + H2SO4 → Na2SO4 + 2H2O
is the enthalpy change when 2 moles of sodium
hydroxide react with one mole of H2SO4
Heat and Temperature
 The words ‘heat’ and ‘temperature’ are often used to
mean the same thing in daily conversation but in
science they are quite distinct and you must be clear
about the difference.
 Temperature is a measure of the average kinetic
energy of the particles in a system.
 As the particles move faster, their average kinetic
energy increases and the temperature goes up.
 But it does not matter how many particles there are –
temperature is independent of the number present.
We measure temperature with a thermometer.
 Heat is a measure of the total energy of all the




particles present in a given amount of substance.
It does depend on how much of the substance is
present.
The energy of every particle is included. There is no
instrument that can measure heat directly.
So a bath of lukewarm water has much more heat than
a sparkler even though the temperature is less.
Heat always flows from high to low temperature, so
heat will flow from the sparkler into the bath water,
even though the water has much more heat than the
sparkler.
Calculating the enthalpy
change of reaction
 The enthalpy change of a reaction is the heat
given out or taken in as the reaction heats up or
cools down.
 There is no instrument that measures heat
directly.
 To measure the enthalpy change we need to know
three things:
 The mass of the substance that is being heated up
or cooled down.
 The temperature change
 The specific heat capacity of the substance
Specific Heat Capacity
 The specific heat capacity, c, is the amount of heat





needed to raise the temperature of 1g of substance
by 1K.
Its units are joules per gram per kelvin, or J g-1 K-1
Water has a high specific heat capacity and is able to
absorb considerable quantities of heat without its
temperature increasing significantly.
One gram of water requires 4.18 J of heat energy for
the temperature of the water to increase by 1K.
This is known as the specific heat capacity of water,
and is often quoted in SI units as; 4.180 J g-1 K-1
It is important to make sure that the units agree with
each other when carrying out enthalpy calculations.
Enthalpy change
 To calculate enthalpy change we use the equation:
Enthalpy
change
=
ΔH
=
Mass of
substance
m
x
x
Specific heat
x
capacity
c
x
Temperature
change
ΔT
 Mass and temperature change are found out from
experiment.
 The specific heat capacity of water (or dilute
aqueous solution) is 4.18 J g-1 K-1
Reactions carried out in
 It is relatively easy to measure heat
solution
changes for reactions that take




place in solution.
The heat is generated in the
solutions themselves and only has to
be kept in the calorimeter.
We often use expanded polystyrene
beakers for the calorimeters.
These are good insulators (this
reduces heat loss through their
sides) and they have a low heat
capacity so the beaker itself
absorbs very little heat.
We usually take the specific heat
capacity of dilute solutions to be the
same as water, 4.18J g-1 K-1
Exothermic or endothermic?
A
B
Watch the video clips. Which reaction is exothermic
and which is endothermic? Calculate the enthalpy
change for each reaction, show your working.
The Simple Calorimeter
 We can use this apparatus to find the




enthalpy change when a fuel burns.
The enthalpy change when one mole of
substance in its standard state is
completely burned in oxygen is called
the standard molar enthalpy of
combustion, ΔH θc
We burn the fuel to heat a known mass
of water and then measure the
temperature rise of the water.
We assume all the heat from the fuel
goes into the water.
The apparatus used is called a
calorimeter.
An example
 A calorimeter was used to measure the enthalpy
change of combustion of methanol:
CH3OH(l)
+
1½ O2(g)
CO2(g)
→
+
2H2O(l)
 0.32g (0.010mol) of methanol was burned and heated
200g of water, the water temperature rose by 4K.
Heat change
=
ΔH
=
m
x
=
200
x
c
4.18
x
ΔT
x
4
=
3344J
 0.01 mol gives 3344 J so 1 mol would give 334400J or
334 kJ
 ΔHc = -334 kJ mol-1 (negative because heat is given
out)
 This value is less than half the accepted value
since not all the heat reaches the water, can you
suggest reasons why?:
 There may be draughts that take heat away from
the calorimeter
 Heat is lost from the top and the sides of the
calorimeter.
 The beaker and the rest of the apparatus itself
gets heated rather than just the water
 Not all the methanol burns completely to carbon
dioxide. This is incomplete combustion. If there is
not enough oxygen available, then we get some
carbon monoxide and carbon formed. Soot (carbon)
on the bottom of the beaker will show if this has
happened.
Flame calorimeter
 Watch the video and record
the results in a table similar
to the one below;
Name
of
Alcohol
Mass
(g)
Initial
temp
(oC)
End
temp
(oC)
Propan1-ol
Octan1-ol
 Calculate the enthalpy of
combustion for each of these
reactions
Contents
 Determine the enthalpy change of a reaction that
is the sum of two or three reactions with known
enthalpy changes.
Hess’s Law
 Consider a chemical reaction in which reactant A is
converted into product B.
A→B
 Let us suppose it can either be converted directly into
B with an enthalpy change of ΔHx, or can be
converted directly into B via an intermediate product
C with a total enthalpy change of ΔHy + ΔHz
A
B
C
 In other words, the reaction can be proceed by two
different pathways.
 Hess’s Law states that the total enthalpy change
of a reaction depends only on the difference
between the enthalpy of the products and the
enthalpy of the reactants. It is independent of
the reaction pathway.
 This is a useful law in chemistry, because it
enables us to work out some enthalpy change
values indirectly.
 Consider the standard enthalpy of formation of
methane. The equation for this reaction, in which
methane is formed from its elements in their
normal states under standard conditions, is
ΔHf θ
C(s) + 2H2(g) → CH4(g)
 Superficially this looks a simple enough reaction.
 We could react carbon with hydrogen and
determine the heat change.
 However, it is difficult to get carbon to react
with hydrogen, and when it does there are many
possible compounds of carbon and hydrogen
(called hydrocarbons) that could be formed, so
the reaction is actually impossible to achieve
experimentally.
 However, what we can do easily is burn carbon,
hydrogen and methane separately and determine
the standard enthalpy change of combustion,
ΔHcθ, directly for the three substances
ΔHf θ
C (s)
+
2H2
1.
CH4
(g)
3.
2.
CO2
(g)
(g)
H2O
(l)
According to Hess Law the total enthalpy of the
reaction is independent of the reaction pathway
1. C (s) + O2 (g) → CO2 (g) ΔHcθ = -393 kJ mol-1
2. H2 (g) + ½O2 (g) → H2O (l) ΔHcθ = -286 kJ mol-1
3. CH4(s) + O2
(g)
→ CO2 (g) + H2O(l) ΔHcθ = -890 kJ mol-1
This reaction is
impossible to
achieve
experimentally
But all three
components in the
reaction can be
reacted with
oxygen easily –
this is called the
enthalpy of
combustion
Contents
 Define the term average bond enthalpy.
 Explain in terms of average bond enthalpies, why
some reactions are exothermic and others are
endothermic.