Energy Changes in Chemistry
Chemical Ideas 4.1 deals with energy changes in chemistry.
Also helpful at this point is:
Chemical Ideas 1.3 Using equations to work out reacting masses
Key terms in this chapter are:
Exothermic.
Endothermic.
Heat change; Enthalpy change; ∆H.
Enthalpy of Combustion, ∆Hc.
Standard conditions.
Standard Enthalpy of Combustion,
o
∆H c, 298
Thermochemical equation.
o
Standard Enthalpy of Formation, ∆H f, 298
Enthalpy level diagram.
Activation enthalpy.
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Joule (J), kJ mol .
Specific heat capacity.
Enthalpy profile diagram.
Hess's Law.
Enthalpy cycle.
Exothermic and Endothermic reactions
In this section you are going to learn:
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What chemists mean by exothermic and endothermic reactions and how to
recognise them.
How to use a sign convention to show a reaction is exothermic or endothermic.
About the units that measure heats of reaction (enthalpy changes, ∆H).
How to define Enthalpy of Combustion.
About standard conditions of temperature and pressure.
That the precise value of ∆H depends on temperature and pressure and that such
values are expressed at standard conditions.
How to define for Standard Enthalpy of Combustion and to write chemical equations
to satisfy the definition.
How to write thermochemical equations to satisfy thermochemical definitions.
How to do simple thermochemical calculations involving ∆H values and balanced
chemical equations.
How to define Standard Enthalpy of Formation and to write chemical equations to
satisfy the definition.
How to draw enthalpy level diagrams.
About the units of energy, and specific heat capacity.
How to work out heat changes using the specific heat capacity of water.
What is meant by Activation enthalpy, Ea.
How to draw enthalpy profile diagrams.
When chemical reactions take place they are often accompanied by heat changes. The
reaction may give out heat, and is called an exothermic reaction. You can tell this has
happened because the reaction vessel feels warmer. Here is an example:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
∆H = -890.3 kJ mol-1
Remember that the balanced chemical equation shows the amounts in moles of reactants and products.
Put simply for now, ∆H means 'heat change of reaction'.
Notice that state symbols are always included for the reactants and products:
(g) gas; (l) liquid; (s) solid; (aq) aqueous solution.
This balanced chemical equation tells us that…
When 1 mole (16 g) of methane burns completely, 809.3 kJ of energy is released.
Since this energy is released by the reaction it is given a negative (-ve) sign. Heat changes
that accompany chemical reactions are also called enthalpy changes.
Note that in this example, 1 mole of methane is being burned. There is an important definition
here.
Enthalpy of Combustion, ∆Hc
The heat evolved when 1 mole of an element or compound completely burns in oxygen.
However, the precise amount of energy released when a given amount of a substance burns is
dependent on both the temperature and pressure at which the reaction is measured. This
means that if we are going to be able to compare enthalpy changes for different reactions,
then they need to be expressed under the same conditions.
Standard conditions of temperature and pressure are used.
298 K (25 °C) and one standard atmosphere pressure (100 00 Nm-2).
This means that the above thermochemical definition can be expressed with regard to
standard conditions. You need to note the differences carefully.
Standard Enthalpy of Combustion, ∆H
o
c, 298
The amount of heat energy given out when 1 mole of an element or compound completely
burns in oxygen, measured under standard conditions of temperature and pressure.
The thermochemical equation now becomes
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
∆Hoc, 298 = -890.3 kJ mol-1
You should note that questions in textbooks and in examinations may or may not refer to standard
conditions, so watch out for this and go along with it.
All the enthalpy values used in this section are standard values and can be found in chemical data books.
You need to be able to write thermochemical equations to satisfy the definition for 'standard
molar enthalpy of combustion'.
Q1.
Write a thermochemical equation for the standard enthalpy of combustion
of hydrogen gas, H2. Use a data book to obtain the enthalpy value and
include all of the correct symbols.
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Q2.
Write a thermochemical equation for the standard enthalpy of combustion
of ethanol, CH3CH2OH. Use a data book to obtain the enthalpy value and
include all of the correct symbols.
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Q3.
Write a thermochemical equation for the standard enthalpy of combustion
of butane, C4H10. Use a data book to obtain the enthalpy value and include
all of the correct symbols.
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Q4.
The following equation indicates that when 1 mole of carbon burns
completely in air or oxygen 393 kJ of heat are evolved.
C(s) + O2(g)
→
CO2(g)
∆Hc = -393 kJ mol-1
What quantity of heat would be given out on complete combustion of
(a) 10 moles of carbon
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(b) 0.25 mole of carbon
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(c) 18 g of carbon ?
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What mass of carbon would have to be burned to produce
(a) 196.5 kJ of heat
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(b) 786 kJ of heat ?
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An endothermic reaction is one that takes in heat, and the reaction vessel feels cooler. Here
is an example:
½H2(g) + ½I2(s) → HI(g)
∆Hof, 298 = + 26.5 kJ mol-1
Note that in this example, 1 mole of hydrogen iodide is being formed. There is an important
definition here.
Standard Enthalpy of Formation, ∆Hof, 298
The amount of heat energy given out or taken in when 1 mole of compound forms from its
elements in their standard states, measured under standard conditions of temperature and
pressure.
Note that the standard enthalpy of formation of an element in its standard state must be zero.
You need to be able to write thermochemical equations to satisfy the definition for 'standard
molar enthalpy of formation'.
Q5.
Write a thermochemical equation for the standard enthalpy of formation of
water, H2O. Use a data book to obtain the enthalpy value and include all of
the correct symbols.
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Q6.
Write a thermochemical equation for the standard enthalpy of formation of
sodium chloride, NaCl. Use a data book to obtain the enthalpy value and
include all of the correct symbols.
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Q7.
Write a thermochemical equation for the standard enthalpy of formation of
ethanol, CH3CH2OH. Use a data book to obtain the enthalpy value and
include all of the correct symbols.
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Q8.
Write a thermochemical equation for the standard enthalpy of formation of
propane, C3H8. Use a data book to obtain the enthalpy value and include
all of the correct symbols.
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Another equation that is useful for calculating the ∆H value for a chemical reaction follows
from the definition of 'standard enthalpy of formation.
∆H = Σ∆Hf (products) - Σ∆Hf (reactants)
Note that the question concerned must provide relevant enthalpy of formation data.
Q9.
By applying the equation above, calculate the standard enthalpy change of
reaction, ∆Hor, 298, for the reaction
CO(g) + Cl2(g) → COCl2(g).
The standard enthalpies of formation for CO(g) and COCl2(g) are -110 kJ
mol-1 and -223 kJmol-1 respectively.
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Enthalpy level diagrams can be used to illustrate exothermic and endothermic changes:
You should note the details that these diagrams include, such as balanced equations and state symbols.
Also, the Enthalpy axis has no numerical values. This is because it is impossible to measure the enthalpy,
H, (heat energy content) of a given amount of a substance. Only differences in total enthalpies (∆H
values) of reactants and products for a chemical reaction can be measured.
A note about heat measurements…
The SI unit of energy is the joule, J. You may have heard of calories. One calorie is,
by definition, equal to 4.184 J (exactly). One calorie raises the temperature of 1 g of
water by 1 ºC.
The specific heat capacity of a substance is the quantity of heat required to warm 1
g of the substance by 1 ºC. The specific heat capacity of water is 4.184 J g-1 ºC-1
(4.184 J g-1 K-1).
Therefore, the quantity of heat transferred (lost or gained), q, is determined by the
temperature change (∆T = Tfinal - Tinitial), the mass, m, of the substance and its
specific heat capacity.
q = specific heat capacity x mass x ∆T
Q10. An exothermic chemical reaction produced enough heat to raise the
temperature of 200 g of water from 17 °C to 34.5 °C. Calculate the
quantity of heat, q, evolved.
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Q11. Using a heat of combustion apparatus, equivalent to 250 g of water and
containing 400 g of water, the enthalpy of combustion of butan-1-ol
(CH3CH2CH2CH2OH) was determined. The initial temperature of the
apparatus was 18 °C. A weighed spirit burner (14.29 g) containing the
alcohol was lit under the apparatus. After a short time the burner was
extinguished and the highest temperature reached by the apparatus was
measured to be 24.4 °C. The spirit burner was re-weighed (13.81 g).
Calculate the molar enthalpy of combustion of butan-1-ol.
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Of course, there will be significant heat losses in the above experiment and insulation is needed in an
attempt to reduce these as much as possible.
Draw a labelled diagram of the calorimeter apparatus use in the above
experiment.
Activation Enthalpy
Methane and oxygen do not react simply when mixed. An input of energy, such as a flame, is
required to get the reaction started, after which its exothermic nature keeps it going. The
amount of energy needed to start the reaction is called the activation enthalpy. The enthalpy
profile diagram below illustrates the activation enthalpy, and in this case is for an exothermic
reaction.
Thermochemical Calculations and Hess's Law
In this section you are going to learn:
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The meaning of Hess's Law and its definition, and why it makes sense.
How to calculate overall enthalpy changes for reactions, ∆H, by drawing enthalpy
cycles and applying Hess's Law.
It is possible for a chemical reaction to take place by more than one route or pathway, that is,
via one lot of chemical intermediates or another. Whichever way a reaction goes, the overall
enthalpy change must be the same.
Imagine a reaction going from reactants to products and taking in 10 kJ of heat energy, and then
going back from products to reactants by another pathway and giving out 20 kJ of heat energy. If
this were possible, which it isn't, 10 kJ of energy would have been created. Energy cannot be
created or destroyed. Hess's Law is a statement of this.
Hess's Law
Hess's Law says that the overall enthalpy change accompanying a chemical reaction is
independent of the route taken in going from reactants to products, provided the initial and
final states are the same.
Some enthalpy changes are impossible to measure by experiment. However, such values can
be calculated by using Hess's Law.
Compare the enthalpy cycles below. They have been drawn to calculate the standard enthalpy
of formation of hexane. They are the same except for one small difference that you should
note. Use which one you prefer.
The enthalpy cycle on the left is now shown on the next page but with the enthalpy values
entered, and the standard enthalpy of formation of hexane calculated.
For the following two questions use the enthalpy values given above.
Q12. Draw an enthalpy cycle to calculate a value for the standard enthalpy
change of formation of butane.
∆Hoc (C4H10) = -2877.0 kJ mol-1.
Work through the following steps:
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Write a chemical equation for the reaction for which ∆H is to be calculated.
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Completely burn the appropriate number of moles of C(s) and H2(g).
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Combine the correct amounts (in moles) of CO2(g) and H2O(l) to form 1 mole of
C4H10(g). (Reverse of the enthalpy of combustion of butane.)
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Add together the total enthalpy changes for each step of the alternative route.
Plan this out on scrap paper first.
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Now apply the above ideas to this slightly different example.
Q13. Draw an enthalpy cycle to calculate the enthalpy of combustion of ethanal,
CH3CHO(l). ∆Hof (CH3CHO) = -192 kJ mol-1.
Plan this out on scrap paper first.
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Q14. Given the following two thermochemical equations:
C(s) + O2(g) → CO2(g)
∆H = -393 kJ mol-1
CO(g) + ½O2(g) → CO2(g)
∆H = -284 kJ mol-1.
Draw an enthalpy cycle to calculate a value for the combustion
C(s) + ½O2(g) → CO(g)
Plan this out on scrap paper first.
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