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1.1 Advanced
Introduction to Enthalpy (Energy) Changes in
Chemical Reactions
I have
ASSUMED you have studied the GCSE notes on the basics of chemical energy changes
1.1a
Introduction to enthalpy changes
This page is an introduction to advanced
level ideas on exothermic or endothermic energy changes in chemical
reactions – referred to as 'enthalpy changes' at constant
pressure e.g. enthalpy of reaction, enthalpy of
formation, enthalpy of combustion and also bond enthalpies ('bond
energies' needed to break bonds).
The concept of enthalpy level diagrams is described
and
introduced and enthalpy changes are clearly distinguished from
activation energies.
The notion of 'standard conditions' is described
and why it is necessary to have data based on a standard defined
temperature, pressure and concentration.
You need to be completely
ok in interpreting enthalpy level and activation energy diagrams such as
...
 
... and can clearly distinguish between enthalpy change and
activation energy and the latter energy does change with a catalyst
BUT a catalyst does NOT change the
enthalpy value of the reaction.
The
activation energy
of a reaction will be rarely mentioned until the end of Part 3), but all
important ideas from the GCSE page will be re–studied in their advanced
level context, but I would still study the GCSE page first!
REMEMBER –
all chemical changes are accompanied by energy changes or energy
transfers, many of which can be directly measured, or, theoretically
calculated from known enthalpy change values.
1.1b Enthalpy Changes and Thermochemistry
Some important initial definitions and
examples:
The system: The reactants and products of the reaction being
studied i.e. the contents of the calorimeter.
The surroundings: The means
the rest of the 'world' including the i.e. a copper calorimeter, the
surrounding air etc. etc.
 Enthalpy H: The
heat energy
content of a substance. This cannot be determined absolutely but
enthalpy changes for a chemical reaction can be measured directly or
indirectly from theoretical
calculations using known enthalpy values.
Enthalpy change ΔH: The net
heat energy transferred to a system from the surroundings or from the surroundings
to a system at constant pressure. The Greek letter delta Δ
in maths implies a change, in this case a net heat energy change.
ΔH = Hfinal
– Hinitial (the units of delta H are kJ mol–1)
or ΔH = ∑Hproducts
– ∑Hreactants
 or
ΔHθ(reaction) = ∑ΔHθf(products)
– ∑ΔHθf(reactants)
The Greek letter delta
Δ
implies 'change in' ....
The Greek letter
∑
implies 'sum of' ....'
ΔHθf
denotes a
standard enthalpy of formation –
which is explained further down.
TOP OF PAGE
An exothermic reaction
A
reaction in which heat energy is given out from the system to the surroundings
i.e. the enthalpy of the reacting system decreases and the temperature
of the system and surroundings rises.
This means Hreactants
> Hproducts so that
ΔH is negative (–ve).
The enthalpy of the reaction
system is decreasing.
Example: All combustion
reactions are exothermic
e.g.
CH4(g) + 2O2(g)
==> CO2(g) + 2H2O(l) ΔH =
–890 kJmol–1
i.e. the figure
of 890 kJ released refers to the complete combustion of 1 mole
of gaseous methane (24 dm3), using exactly 2 moles of gaseous oxygen
(48 dm3) to
form exactly 1 mole of gaseous carbon dioxide (24 dm3) and 2 moles of
liquid water. These values refer to 298K (25oC
and 1 atm/101 kPa)
Note some general points (which apply to all
exothermic or endothermic changes, physical or chemical
changes):
(i) All enthalpy values must be quoted with
referenced to the ambient/assumed temperature and pressure
of the system undergoing the physical or chemical change.
The usual
standard reference conditions are
298K
(25oC and 1 atm/101 kPa), and other criteria
may apply e.g. 1 molar solution if applicable.
(ii) Not only the molar quantities must
clearly indicated BUT the physical states of all the substances must
be clearly stated too.
This is a convenient point to make the point about the
importance of state symbols via the combustion of hydrogen.
e.g.
H2(g) +
1/2O2(g) ==> H2O(l)
ΔH = –285.9 kJ mol–1, but for
H2(g) +
1/2O2(g) ==> H2O(g)
ΔH = –241.8 kJ mol–1
If the water forms
remains as steam/vapour/gas, then 44.1 kJ less heat energy
is released to the surroundings, because condensation is an
exothermic process (g ==> l) and forming liquid water
releases an extra 44.1 kJ. The –285.9 (~–286) kJ mol–1
is the usual value for the enthalpy of combustion of
hydrogen you will encounter in your studies because at the
standard temperature of 298K water is a liquid in its normal
stable state.
(iii) This sort of combustion reaction can
be measured in a calorimeter (see
section 1.3).
BUT, however the enthalpy change is measured, all equations should be read in molar terms when dealing with
enthalpy values i.e. a delta H value goes with a specific
equation.
(iv)
Enthalpy change values are usually quoted
in kJ mol–1, but take care in their
interpretation because you must know what equation goes with
the
ΔH
value!
e.g. the
enthalpy of combustion usually refers to the complete
combustion of one mole of the combustible material as
for water above, BUT if you double the equation you must
also double the enthalpy values for that equation
2H2(g) + O2(g) ==> 2H2O(l)
ΔH = 2 x –285.9 =
–571.8 kJ mol–1
TOP OF PAGE
An endothermic reaction
A
reaction in which the system takes in or absorbs heat energy from the
surroundings i.e. the enthalpy of the system increases and the
temperature of the system and surroundings falls OR the system must be
heated to initiate the reaction and provide the heat absorbed.
This means Hproducts
> Hreactants so that ΔH is positive (+ve).
The enthalpy of the reaction
system is increasing.
Example: The thermal
decomposition of calcium carbonate
CaCO3(s)
==> CaO(s) + CO2(g) ΔH =
+179 kJmol–1
i.e. 179 kJ
of heat energy must be absorbed to decompose 1 mole of solid
calcium carbonate into 1 mole of solid calcium oxide and 1 mole
of gaseous carbon dioxide. Mr(CaCO3) =
100, so 17.9 kJ of heat energy is absorbed in decomposing 10g of
limestone. This reaction requires an experimental temperature of
800–1000oC to achieve an appreciable rate of reaction
and cannot be studied quantitatively in the laboratory. However
it can be theoretically calculated from known enthalpy change
values by means of a
Hess's Law cycle
calculation.
Reminder of energy profile differences between exothermic and
endothermic reactions
The sets of diagrams below illustrate
how exothermic (left) and endothermic (right) reactions are specified on
an enthalpy level diagram.
 
Simple enthalpy level diagrams for
exothermic and endothermic reactions
More complex enthalpy level diagrams for
exothermic and endothermic reactions showing the activation energy
'hump' and reduction of activation energies by employing a catalyst
1.1c. Standard conditions
defined for thermochemistry and thermodynamics in general
Standard conditions for referencing
enthalpy values are essential for communicating accurate data
throughout the scientific community.
It means values
measured/calculated in one laboratory/research team can be used
in another scientific establishment anywhere!, OR checked for
accuracy by any other scientists.
In this way
accurate enthalpy data can be built up and through time
validated and perhaps more accurately measured with
technological developments and theoretical calculations
become more reliable.
A superscript
θ
means a standard enthalpy value (see further down the
page)
This
means the reactants/products start/finish at a specified temperature,
pressure and concentration whatever the 'temporary' temperature change
in the reaction – which is required to calculate the enthalpy change.
The net energy change is based on the products returning to the same
temperature and pressure that the reactants started at.
The most
frequently used standard conditions are:
a
temperature of 298 K/25oC
Kelvin:
K = 273 +
oC, is used in any thermodynamic calculation.
In calorimetric calculations, the
temperature change can be measured in oC.
a pressure of 1 atm/101 kPa
Units
such as mmHg are never used in calculations, and atm are being
increasingly less used in favour of the SI unit Pascal)
a concentration of 1.00 mol dm–3.
The use of standard
conditions enables a database of delta H change to be assembled from
which you can do theoretical calculations (see section 1.2 using
Hess's Law).
Strictly speaking the standard conditions should be
indicated in terms of the standard temperature and the reactants
involved and standard delta H values are denoted with the Greek letter
theta (θ).
By using data based on standard. agreed
and defined conditions, then the data can be used universally by any
laboratory around the world and also allows scientists to check each
others experimental results.
Its pertinent here to consider the question – how can you have an standard enthalpy of combustion at 25oC
when the flame temperature is perhaps peaking at over 1000oC
!!!
The answer applies
to all enthalpy changes what–so–ever!
The enthalpy change represents the
heat energy change needed to restore the products to the
temperature of the reactants at the start e.g. room
temperature/25oC.
TOP OF PAGE
1.1d. Examples of standard
enthalpies
Standard Enthalpy of Reaction ΔHθr
or reaction
is the enthalpy change (heat absorbed/released, endothermic/exothermic)
when molar quantities of reactants as stated in an equation react under
standard conditions (i.e. 298K/25oC, 1 atm/101kPa)
Examples
(i)
NaOH(aq) + HCl(aq)
==> NaCl(aq) + H2O(aq)
(exothermic)
ΔHθr,298
= –57.1 kJ mol–1
(can also be described as an
'enthalpy of neutralisation')
(ii)
CaCO3(s) ==>
CaO(s) + CO2(g) (endothermic)
ΔHθr,298
= +179 kJ mol–1
(can also be described as an 'enthalpy
of thermal decomposition')
Standard Enthalpy of Formation ΔHθf
or formation
is the enthalpy change when 1 mole of compound is formed from its
constituent elements with both the compound and elements in their
standard states ('normal stable states) i.e. at 298K/25oC, 1 atm/101kPa
(a)
It may be endothermic or exothermic
(b) Any accompanying equation should
involve the formation of 1 mole of the compound
The
standard state is the most stable state
of a substance at the standard
temperature and pressure e.g. the physical state at 298K/25oC and 1 atm/101kPa
pressure e.g.
H2(g)
H2O(l) C(s)
O2(g), C3H8(g)
C8H18(l) C24H50(s)
CO2(g) CH3CH2OH(l)
etc.
Examples -
some are just theoretical and some can actually happen
(i)
C(s)
+ 2H2(g) ==> CH4(g) ΔHθf,298(methane)
= –74.9 kJ mol–1
(ii)
2C(s)
+ 2H2(g) ==> C2H4(g) ΔHθf,298(ethene)
= +52.3 kJ mol–1
(iii)
2C(s) + 3H2(g)
+ 1/2O2(g) ==> CH3CH2OH(l) ΔHθf,298(ethanol)
= –278 kJ mol–1
(iv)
1/2N2(g)
+ O2(g)
==> NO2(g) ΔHθf,298(nitrogen
dioxide) = +33.9 kJ mol–1
Note (a) The values can be
positive/endothermic or negative/exothermic.
(b)
The enthalpy of formation of elements in their standard
stable states is arbitrarily assigned a value of zero.
This
definition, together with experimental values of enthalpy
changes allows a body of enthalpy change data to be accumulated
and extended via theoretical calculations.
Standard Enthalpy of Combustion ΔHθc
or combustion
is the enthalpy change when 1 mole of a fuel (or any combustible
material) is completely burned in oxygen (or air containing oxygen)
equated to standard conditions (298K/25oC, 1 atm/101kPa).
You
should ensure just 1 mole of fuel appears in the equation to accompany
the delta H value which is always negative i.e. always exothermic.
Examples
(i)
C3H8(g) + 5O2(g) ==> 3CO2(g) +
4H2O(l) ΔHθc,298K(propane)
= –2219 kJ mol–1
(ii)
CH3COOH(l) +
2O2(g) ==> 2CO2(g) +
2H2O(l) ΔHθc,298K(ethanoic
acid)
= –876 kJ mol–1
In the calculations explained below
just the subscripted letters r/f/c will be used for brevity and a
temperature of 298K and a constant pressure 1atm assumed unless
otherwise stated.
There is more the enthalpies of
combustion of alkanes and alcohols in section 1.4a
and specifically a detailed
discussion on the
combustion of alcohols
Standard enthalpy of
neutralisation
ΔHθneutralisation
is the energy released when unit molar quantities of
acid and alkali solutions completely neutralise each other at 298K/101
kPa (though pressure
effects are insignificant for reactions only involving
liquids, solutions or solids)
(i)
NaOH(aq)
+ HCl(aq) ==> NaCl(aq) + H2O(l) ΔHθneutralisation
= –57.1 kJ mol–1
(ii)
Ba(OH)2(aq) + 2HNO3(aq)
==> Ba(NO3)2(aq) + 2H2O(l) ΔHθneutralisation
= –116.4 kJ mol–1
(iii)
½Ba(OH)2(aq) + HNO3(aq)
==>
½Ba(NO3)2(aq) +
H2O(l) ΔHθneutralisation
= –58.2 kJ mol–1
Note! It looks as if the enthalpy
of neutralisation of barium hydroxide is approximately double that
of sodium hydroxide i.e. ~ twice as exothermic! Well yes it is! and no
it isn't!
Yes – ~twice as much
energy is released per mole of the soluble base (alkali).
No – however, on the basis of heat
released per mole of water formed, they are actually very
similar to the other enthalpies of neutralisation in which one
mole of water is formed.
In
other words, which value you quote, depends on which point you
want to make via a 'molar' equation.
Yet another example of carefully qualifying enthalpy values with
respect to the context.
More on
enthalpies of neutralisation
1.1e.
Bond enthalpy, enthalpy level diagrams and enthalpies of formation
Bond Enthalpy ('bond energy')
This is the
average energy absorbed to break 1 mole of a specified bond when all species involved are in the
gaseous state.
e.g. for (i)
H2(g) ==>
2H(g) ΔH = +436 kJ mol–1 for
the H–H bond
or for (ii)
CH3CH2Br(g) ==> CH3CH2(g) + Br(g)
ΔH = +276 kJ mol–1 for the C–Br bond
It is always endothermic and the
reverse process – bond formation, is always exothermic. In many cases
the values are averaged from a variety of 'molecular' situations because
a specific bond enthalpy can vary slightly because of small
differences in the molecular environment of the bond.
More
on this in the
bond enthalpy section.
Some examples of points made on this page
with reference to an enthalpy level change diagram
shown below which enthalpies of formation and reference to bond
enthalpies as well.
General points: Arrows pointing
downwards represent exothermic changes and arrows pointing upwards
represent endothermic changes
1.
The energy released when 1 mole of aluminium oxide is formed.
The ΔH
value of –1669 kJ mol–1
corresponds to the very exothermic enthalpy of formation of Al2O3
or the enthalpy of the complete combustion of two moles of Al.
The very
exothermicity of the reaction suggests, and correctly, that aluminium
oxide is a very stable compound – it is thermally stable to at least
2500oC.
2.
The endothermic enthalpy of
formation of gold(III) oxide
It is a compound not readily formed and it
decomposes on heating at ~150oC, so contrast this thermal
instability with that of aluminium oxide.
3. This is a much
more complex
enthalpy level diagram involving hydrogen, chlorine and hydrogen
chloride.
The +436 kJmol–1 represents the bond enthalpy for
splitting one mole of hydrogen molecules into two moles of hydrogen atoms.
The +242 kJmol–1 is the
bond enthalpy of chlorine molecules i.e. splitting 1 mole of chlorine
molecules into two moles of chlorine atoms.
The –184 kJmol–1 is the enthalpy of
formation of 2 moles of hydrogen chloride gas, which can be expressed as -92 kJmol-1.
Take care in matching
ΔH
values with the number of moles it refers too!
The very exothermic
–862 kJmol–1 is the
energy released theoretically when two moles of hydrogen chloride are
formed directly from 2 moles of hydrogen atoms and 2 moles of chlorine atoms.
The latter indicates
that the H–Cl bond enthalpy is +862/2 = 431 kJmol–1
TOP OF PAGE
1.1f
Key revision points for
this page on introducing enthalpy changes
Enthalpy change (ΔH))
is the heat energy transferred at constant pressure during a chemical
reaction.
It can be exothermic (negative
ΔH)
or endothermic (positive
ΔH.
Students must master
definitions of standard enthalpy changes e.g. formation, combustion,
neutralisation and bond enthalpy and
apply Hess’s Law, and avoid common
misconceptions about signs, conditions, and bond enthalpy averages.
-
Definition of
Enthalpy Change
-
Enthalpy (H) = heat energy
content of a system at constant pressure.
-
ΔH
= enthalpy of products – enthalpy of reactants.
-
Exothermic: (ΔH
< 0) (heat energy released).
-
Endothermic: (ΔH
> 0) (heat energy absorbed).
-
Standard Conditions
-
298 K (25 °C), 100 kPa,
substances in standard states.
-
Always quote these when
defining standard enthalpy changes.
-
Types of Enthalpy
Changes
-
Reaction:
Energy change for a given balanced equation.
-
Formation (ΔHf):
1 mole of compound formed from elements in standard states.
-
Combustion (ΔHc):
Energy released when 1 mole of substance completely burned in oxygen.
-
Neutralisation (ΔH(neut):
Acid + base → 1 mole of water formed (per mole H2O).
-
Bond Enthalpy:
Energy to break 1 mole of bonds in gaseous molecules (average values).
-
Hess’s Law
(not dealt with on this page)
Common
Misconceptions about enthalpy changes
Exam
Revision Tips for enthalpy questions
-
Precise Definitions:
Learn word-perfect definitions for each enthalpy change type. Examiners
penalise vague answers.
-
Units:
Always use kJ mol⁻¹. Forgetting “per mole” loses marks.
-
Equations:
Write balanced chemical equations alongside enthalpy definitions.
-
Diagrams:
Practice energy profile diagrams (exothermic versus endothermic).
-
Hess’s Law:
Draw clear enthalpy cycles; label arrows with (\Delta H).
-
Bond Enthalpy
Calculations (not on this
page)
-
Break bonds (positive
values).
-
Make bonds (negative
values).
-
ΔH
= energy released in forming bonds - energy absorbed in breaking bonds
-
See
Bond Enthalpy Calculations
-
Neutralisation:
Remember strong acid–strong base values are consistent; weak acid/base
values are smaller (less exothermic).
-
Combustion:
Always “complete combustion” in oxygen, not partial.
-
Formation:
Elements must be in their standard states (e.g., O2 gas,
graphite for carbon).
-
Cross-board alignment:
All exam boards (AQA, Edexcel, OCR, Salters, WJEC, CCEA, CIE, IB, AP)
require:
-
Definitions under standard
conditions.
-
Application of Hess’s Law.
-
Bond enthalpy
calculations.
-
Recognition of exothermic
versus endothermic processes.
Summary Table
|
Enthalpy Change |
Definition |
Typical Sign |
Key Exam Tip |
|
Reaction |
Energy change for balanced
equation |
± |
Always specify reaction
conditions |
|
Formation |
1 mole compound from
elements |
± |
Elements in standard
states |
|
Combustion |
1 mole substance burned in
O2 |
Negative |
Must be complete
combustion |
|
Neutralisation |
Acid + base → 1 mole H2O |
Negative (~–57 kJ mol⁻¹ of
H2O) for strong acid/base) |
Weak acids/bases differ |
|
Bond Enthalpy |
Energy to break 1 mole
bonds (average) |
Positive (breaking bonds) |
Use averages; approximate
values |
Final Tip:
Examiners love clarity. Always
define ..., clear diagrams, calculate, and check signs.
Mislabelled signs and missing
“per mole” are the most common lost marks.
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QUICK INDEX for
Energetics:
GCSE Notes on the basics of chemical energy changes
– important to study and know before tackling any of the three Advanced Level
Chemistry pages
INDEX of ALL
advanced level pages on thermochemistry and thermodynamics
Parts 1–3 here
* Part 1a–b
ΔH Enthalpy Changes
1.1 Advanced Introduction to enthalpy changes
of reaction,
formation, combustion etc. : 1.2a & 1.2b(i)–(iii)
Thermochemistry – Hess's Law and Enthalpy
Calculations – reaction, combustion, formation etc. : 1.2b(iv)
Enthalpy of reaction from bond enthalpy
calculations : 1.3a–b
Experimental methods
for determining enthalpy changes and treatment of results and
calculations :
1.4
Some enthalpy data patterns : 1.4a
The combustion of linear alkanes and linear
aliphatic alcohols
:
1.4b Some patterns in Bond
Enthalpies and Bond Length : 1.4c
Enthalpies of
Neutralisation : 1.4d Enthalpies of
Hydrogenation of unsaturated hydrocarbons and evidence of aromatic
ring structure in benzene
:
Extra Q page
A set of practice enthalpy
calculations with worked out answers **
Part 2 ΔH Enthalpies of
ion hydration, solution, atomisation, lattice energy, electron affinity
and the Born–Haber cycle : 2.1a–c What happens when a
salt dissolves in water and why? :
2.1d–e Enthalpy
cycles involving a salt dissolving : 2.2a–c
The
Born–Haber Cycle *** Part 3
ΔS Entropy and ΔG Free Energy Changes
: 3.1a–g Introduction to Entropy
: 3.2
Examples of
entropy values and comments * 3.3a ΔS, Entropy
and change of state : 3.3b ΔS, Entropy changes and the
feasibility of a chemical change : 3.4a–d
More on ΔG,
free energy changes, feasibility and
applications : 3.5
Calculating Equilibrium
Constants from ΔG the free energy change : 3.6
Kinetic stability versus thermodynamic
feasibility - can a chemical reaction happen? and will it happen?
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