2.3.1 Enthalpy changes
Introduction
It is useful to know how energy changes occur when chemical reactions proceed. This helps us to
understand processes which produce energy; the use of fuels in exothermic reactions is of great
economic importance. As crude oil runs out, there will be a massive incentive to develop other fuels
which are cheap, sustainable and non-polluting. Also, if we are making a product using a chemical
reaction it is useful to know how changes in temperature conditions may affect the reaction.
Measurements of the energy transferred in a chemical reaction must be made under controlled
conditions. A special name is given to the energy exchange with the surroundings when it occurs at
constant pressure and this is enthalpy. It is given the symbol ΔH. All energy changes are measured in
kJmol-1.
Definition: ΔH (enthalpy change) is the heat energy change measured at constant pressure
Student Activity 1
Cross out one of the emboldened words in each sentence.
a. Exothermic reactions can be identified by detecting a temperature decrease/increase.
b. Endothermic reactions can be identified by detecting a temperature decrease/increase.
c. Breaking bonds requires/produces energy.
d. Forming bonds requires/produces energy.
e. Intermolecular and covalent bond breaking is endothermic/exothermic.
f. Forming bonds is endothermic/exothermic.
Reaction pathway diagrams (energy profile diagrams)
Exothermic Reaction: the enthalpy of the products is lower than the enthalpy of the reactants
The enthalpy change ΔH is –ve. The direction of the arrows on the diagram below is important. The
activation energy of a reaction is the minimum energy required for a reaction to occur.
All combustion (burning) reactions are exothermic. You may also have noticed the test tube heating
up when you have added a metal, such as magnesium to an acid, such as HCl(aq).
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Endothermic Reaction: the enthalpy of the products is higher than the enthalpy of the reactants.
The enthalpy change is +ve. These reactions are less obvious in the real world. Examples include:
photosynthesis which involves the combination of carbon dioxide and water to form glucose; the
decomposition of carbonates; cool packs used on the football pitch.
Student Activity 2
Classify the following as either endothermic or exothermic.
Reaction (example)
Oxidation (combustion of CH4)
Thermal decomposition
(CaCO3 → CaO + CO2)
Respiration
Acids reacting with metals
(Li + HCl → LiCl + ½ H2)
Photosynthesis
Endothermic or exothermic
Student Activity 3
a) Complete the following sentences with the words endothermic or exothermic.
A student was monitoring the temperature in three different reactions. The first reaction increased in
temperature by 4.5 °C. This is an ............................. reaction. The second reaction went from 25 °C to
12 °C. This is an .......................... reaction. The third reaction changed in temperature from -5.8 °C to
-3.4 °C. This is an ........................... reaction.
b) Draw enthalpy profile diagrams for the first reaction and the second reaction.
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Standard conditions
The size of the enthalpy change of a particular chemical reaction depends upon: temperature;
pressure; physical state; amount of reactants.
All energy changes therefore must be measured under the same conditions. This is known as
standard conditions. Definition: standard conditions are:-




temperature
= 298 K
pressure
= 1 atm or 100kPa
solution concentration
= 1 moldm-3
all substances should be in their standard sates at 298K and 100kPa
Any enthalpy changes measured under standard conditions are termed a standard enthalpy change
 H 298
The definitions of the different types of standard enthalpy changes
The standard enthalpy of reaction (  H r)
The heat energy change at constant pressure when the amounts of reactants shown in the equation
react under standard conditions to give the products in their standard states.
Example:
The standard enthalpy of formation (  H f)
The enthalpy change when 1 mole of a compound is formed in its standard state from its constituent
elements in their standard states, under standard conditions.
Example:
The standard enthalpy of combustion (  H c)
The enthalpy change when 1 mole of a substance in its standard state reacts completely with oxygen
under standard conditions.
Example:
Bond enthalpy (  H BE)
The enthalpy change when 1 mole of a given type of bond is broken by homolytic fission in the
gaseous state.
Example:
Student Activity 4
a) Write in the names and symbols of the energy changes next to the correct definition.
Energy change
Definition
The enthalpy change when 1 mole of a given type of bond is broken by
homolytic fission in the gaseous state.
The heat energy change at constant pressure when the amounts of reactants
shown in the equation react under standard conditions to give the products in
their standard states.
The enthalpy change when 1 mole of a substance in its standard state reacts
completely with oxygen under standard conditions.
The enthalpy change when 1 mole of a compound is formed in its standard
state from its constituent elements in their standard states, under standard
conditions.
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Symbol
b) Link the reactions to the enthalpy definition that best fits it. One is completed for you.
The standard enthalpy of formation
6C(s) + 6H2(g) → C6H12(l)
2H2(g) + N2(g) → N2H4(l)
½F2O(g) → F(g) + ½O (g)
C(s) + 2H2(g) + ½O2(g) → CH3OH(l)
H2(g) + ½O2(g) → H2O(l)
The standard enthalpy of combustion
¼CH4(g) → ¼C(g) + H(g)
Na(s) + ¼O2(g) → ½Na2O(s)
Bond enthalpy
2Na(s) + ½O2(g) → Na2O(s)
CH4(g) + 2O2(g) → CO2(g) +2H2O(l)
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Note: Mean bond enthalpy is the average energy required to break and separate one mole of bonds
of one type occurring in a range of compounds/molecules. The mean bond enthalpy will never
correspond exactly to the actual bond enthalpy.
Though not required, it is useful to know about the enthalpy of neutralisation (see experiment 13).
Student Activity 5
Label the following statements as either true or false
a) Boiling water is endothermic.
b) The following reaction represents the enthalpy of combustion of ethane.
2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(g)
c) The formation of a molecule of nitrogen molecules from nitrogen atoms will be accompanied by
a temperature decrease.
d) The bond enthalpy of a C-H bond in methane is the same as that of a C-H bond in methanol.
e) The reaction between nitric acid and sodium hydroxide will be expected to have a negative ∆H
sign.
Experimental methods for calculating standard enthalpies of reaction
When experiments have been carried out the enthalpy change can be calculated using:
Where:Q
= Heat Change
m = Mass of substance heated
c
= Specific Heat (for water = 4.18 J g-1 K-1
 T = Change of temperature
Specific heat capacity is defined as the quantity of energy required to heat 1g of substance by 1K.
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Calculating the enthalpy of neutralisation,  Hn
Definition: The standard enthalpy change of neutralisation (  H n) is the enthalpy change when one
mole of water molecules is formed when an acid reacts with an alkali at 298K and 100kPa. The
solutions must have a concentration of 1moldm-3
example: HCl(aq)
+
NaOH(aq)
→
NaCl(aq) + H2O(l)
The apparatus for the experiment is shown in Experiment 13.
Example Calculation
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250cm3 ( 0.4moldm-3 ) of NaOH was added to 250cm ( 0.4moldm-3) of HCl. The temperature rose
from 17.4oC to 20.1oC. Assume that the density of the solution is the same as that of water (=1gcm -3)
Assume the density of the solutions = density of water
Mass of solutions = 500g
c = The specific heat capacity of water is 4.18J g-1 K-1.
Q = m  c  T
Equation: NaOH (aq) +
HCl(aq)
→
NaCl (aq) + H2O (l)
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Calculating the enthalpy of combustion,  HC
The apparatus for the experiment is shown in Experiment 14.
A more accurate way to find this will be demonstrated by your teacher.
Student Activity 6
Past paper question (Jun 2009 Q2b)
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Calculating enthalpy of solution
Definition: The energy change at constant pressure when 1 mole of solid is completely dissolved in
excess water under standard conditions.
In this experiment the temperature is measured over time and a graph is plotted. Extrapolation of the
curve obtained as the solution cools allows correction to be made to the estimated temperature rise.
This corrected value makes an allowance for losses to the surroundings. The temperature of the
liquid is measured at the start of the practical. This is recorded. The solid is added and the
temperature of the liquid is recorded every 30 seconds. There will either be a rise or fall in the
tempearture of the liquid. A typical set of results is given below:
Mass of water = 100g
Mass of solid (sodium hydroxide) = 6.20g
Initial temperature = 15oC
Results Recorded
Time (s)
Temperature (oC)
0
15
30
21
60
28
90
29
120
27
150
26
180
26
210
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In order to find the actual temperature rise a graph of temperature against time is plotted. Plot time
(horizontal axis) against temperature from 15 oC to 30 oC and extrapolate from two straight lines to
get the maximum temperature without cooling.
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Calculation
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Using Hess’s Law to calculate enthalpy changes
Some enthalpies cannot be found by direct experiment. Instead we have to find these enthalpies from
other known enthalpies and the method of drawing energy cycles and using Hess’s law. There are
two different methods which depend on the type of data provided: a) Finding the enthalpy of formation
from enthalpies of combustion and b) Finding the enthalpy of reaction from the enthalpies of
formation.
Definition; Hess's Law (a consequence of the first law of thermodynamics- energy can not be
created or destroyed)
A
B
C
 HAB
=
 HAC
+
 HCB
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a) Finding the enthalpy of formation from enthalpies of combustion
Calculate the enthalpy of formation of methane given the enthalpies of combustion of methane,
carbon and hydrogen.
→ CH4(g)
 Hf = ?
CH4(g) + 2O2(g)
→ CO2(g) + 2H2O (g)
 H1= -890.4 kJ mol
-1
C (s)
→
CO2(g)
 H2= -393.5 kJ mol
-1
→
H2O (g)
 H3= -285.7 kJ mol
-1
C(s)
+
2H2(g)
Combustion Equations
+ O2(g)
H2(g) +
½ O2(g)
We can use Hess' Law to calculate the enthalpy of formation of methane from the enthalpy of
combustion data.
Constructing the Energy Cycle
1. Write the equation for the enthalpy we need to calculate. Make the top arrow long.
 Hf
C(s)
+
2H2(g)
CH4 (g)
2. At the bottom, write the combustion products of carbon dioxide and water which are common to
both sides of the equation and draw in the arrows to make a triangle.
3. Add the values
4. Applying Hess' Law to calculate the missing value on the top equation.
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b) Finding the enthalpy of reaction from the enthalpies of formation
Calculate the enthalpy of reaction for:
NH3 (g) +
HCl (g)
→
NH4Cl(g)
 Hr=?
Given the enthalpies of formation of ammonia and hydrogen chloride.
NH3(g)
HCl(g)
NH4Cl(g)
-46.0kJ mol-1
-92.3kJ mol-1
-315kJ mol-1
Construction of the energy cycle
1. Write the reaction for the enthalpy we need to calculate
 Hr
NH3 (g) +
HCl (g)
NH4Cl(g)
2. Add the elements for formation which are common to both sides of the equation
3. Complete the triangle by adding the values and arrows
4. Applying Hess Law will calculate the value as required
Quick Method
For finding  HӨr when given  HӨf you can use this little shortcut:
 HӨr
=
Σ  HӨf products -
Σ  HӨf reactants
Try it with the examples above and see if it works.
But beware – you cannot use this formula in any other circumstances.
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Finding the enthalpy change in a reaction using bond enthalpies
If enough energy is absorbed by atoms bonded together in a molecule they vibrate so violently that
the bond breaks. The energy required to break these bonds is called the bond enthalpy (or bond
energy). In examinations these energies will always be given.
Bond
H-H
H-Cl
C-C (single)
C-H
C-O
C-Cl
C=C (Double)
C≡C (Triple)
Cl-Cl
Br-Br
H-Br
N≡N (Triple)
H-N
O=O
O-H
C=O
Bond Enthalpy / kJ mol-1
+436
+431
+347
+413
+335
+326
+619
+837
+242
+193
+364
+945
+391
+498
+464
+805
When bonds are made energy is released. The process is exothermic,  H is negative. When bonds
are broken energy is needed. The process is endothermic,  H is positive.
Definition: bond enthalpy is the energy needed to break one mole of bonds in the gaseous state by
homolytic fission.
Definition :mean bond enthalpy is the average energy needed to break one mole of bonds in the
gaseous state by homolytic fission.
Example
Calculate the approximate enthalpy of reaction for the following process:
CH2=CH2
+
H2
CH3-CH3
Bonds Broken
Bonds Formed
Calculate ΔHr as the difference between the energy needed to break the old bonds and the energy
given out when the new bonds are made. This is the sum of bonds broken (which is always positive)
and bonds formed (which is always negative).
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Draw an energy profile for this reaction.
Student Activity 7
Calculate the enthalpy of the combustion of methane using mean bond enthalpies.
Draw an energy profile diagram.
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Why do these values not agree exactly with accurate experimental values?
Remember that most experimental values for finding heats of combustion have big heat losses so
there you always get a smaller value than the book value.
Extension Questions
1. How would the value for the enthalpy of combustion of an alcohol where the steam is allowed to
condense to room temperature compare with that where the steam is kept in the gaseous state?
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2. It is very difficult to determine the standard enthalpy change of formation of hexane directly.
Suggest a reason why.
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3.
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4. Complete the following table by including definitions and symbols
Energy change
Standard enthalpy
of combustion
Definition
Symbol
Standard enthalpy
of formation
Standard enthalpy
of reaction
Standard bond
enthalpy
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