Introduction to Enthalpy (Energy) Changes in
ASSUMED you have studied the GCSE notes on the basics of chemical energy changes
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' 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
introduced and enthalpy changes are clearly distinguished from
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.
of a reaction will be rarely mentioned until the end of Part 3), but all
important ideas from the GCSE page will be restudied in their advanced
level context, but I would still study the GCSE page first!
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
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
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 mol1)
or ΔH = ∑Hproducts
or ΔHθ(reaction) = ∑ΔHθf(products)
The Greek letter delta
implies 'change in' ....
The Greek letter
implies 'sum of' ....'
standard enthalpy of formation
which is explained further down.
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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 =
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
(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
(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.
a convenient point to make the point about the importance of
state symbols via the combustion of hydrogen. eg
1/2O2(g) ==> H2O(l)
ΔH = 285.9 kJ mol1, but for
1/2O2(g) ==> H2O(g)
ΔH = 241.8 kJ mol1
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 mol1
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
(iii) This sort of combustion reaction can
be measured in a calorimeter (see
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
Enthalpy change values are usually quoted
in kJ mol1, but take care in their
interpretation because you must know what equation goes with
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 mol1
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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
==> CaO(s) + CO2(g) ΔH =
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
8001000oC 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
The two diagrams below illustrate
how exothermic (left) and endothermic (right) reactions are specified on
an enthalpy level diagram.
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
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
(K = 273 + oC) and a pressure of 1 atm/101 kPa and a
concentration of 1.00 mol dm3.
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
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
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 whatsoever!
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
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Examples of standard
Standard Enthalpy of Reaction ΔHθr/react/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)
(i) NaOH(aq) + HCl(aq)
==> NaCl(aq) + H2O(aq)
= 57.1 kJ mol1 (can also be described as an
'enthalpy of neutralisation')
CaO(s) + CO2(g) (endothermic)
= +179 kJ mol1 (can also be described as an 'enthalpy
of thermal decomposition')
Standard Enthalpy of Formation ΔHθf/form/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
It may be endothermic or exothermic
(b) Any accompanying equation should
involve the formation of 1 mole of the compound
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
+ 2H2(g) ==> CH4(g) ΔHθf,298(methane)
= 74.9 kJ mol1
+ 2H2(g) ==> C2H4(g) ΔHθf,298(ethene)
= +52.3 kJ mol1
(iii) 2C(s) + 3H2(g)
+ 1/2O2(g) ==> CH3CH2OH(l) ΔHθf,298(ethanol)
= 278 kJ mol1
+ O2(g) ==> NO2(g) ΔHθf,298(nitrogen
dioxide) = +33.9 kJ mol1
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.
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/comb/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).
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.
(i) C3H8(g) + 5O2(g) ==> 3CO2(g) +
= 2219 kJ mol1
(ii) CH3COOH(l) +
2O2(g) ==> 2CO2(g) +
= 876 kJ mol1
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
There is more the enthalpies of
combustion of alkanes and alcohols in section 1.4a
Standard enthalpy of
is the energy released when unit molar quantities of
acids and alkalis completely neutralise each other at 298K (pressure
effects are insignificant for reactions only involving
+ HCl(aq) ==> NaCl(aq) + H2O(l) ΔHθneutralisation
= 57.1 kJ mol1
(ii) Ba(OH)2(aq) + 2HNO3(aq)
==> Ba(NO3)2(aq) + 2H2O(l) ΔHθneutralisation
= 116.4 kJ mol1
(iii) 1/2Ba(OH)2(aq) + HNO3(aq)
==> 1/2Ba(NO3)2(aq) +
= 58.2 kJ mol1
Note! It looks as if the enthalpy
of neutralisation of barium hydroxide is approximately double that
of sodium hydroxide ie ~ twice as exothermic! Well yes it is! and no
Yes ~twice as much
energy is released per mole of soluble base/alkali.
No however, on the basis of heat
released per mole of water formed, they are actually very
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.
enthalpies of neutralisation
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
e.g. for (i) H2(g) ==>
2H(g) ΔH = +436 kJ mol1 for
the HH bond
or for (ii)
CH3CH2Br(g) ==> CH3CH2(g) + Br(g)
ΔH = +276 kJ mol1 for the CBr 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. More
on this in the
bond enthalpy section.
Some examples of points made on this page
with reference to an enthalpy level change diagram
General points: Arrows pointing
downwards represent exothermic changes and arrows pointing upwards
represent endothermic changes
The energy released when 1 mole of aluminium oxide is formed.
value of 1669 kJ mol1
corresponds to the very exothermic enthalpy of formation of Al2O3
or the enthalpy of the complete combustion of two moles of Al.
exothermicity of the reaction suggests, and correctly, that aluminium
oxide is a very stable compound it is thermally stable to at least
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
enthalpy level diagram involving hydrogen, chlorine and hydrogen
The +436 kJmol1 represents the bond enthalpy for
splitting hydrogen molecules into hydrogen atoms.
The +242 kJmol1 is the
bond enthalpy of chlorine molecules i.e splitting 1 mole of chlorine
molecules into chlorine atoms.
The 184 kJmol1 is the enthalpy of
formation of 2 moles of hydrogen chloride gas (-92 kJmol-1).
The very exothermic
862 kJmol1 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 HCl bond enthalpy is +862/2 = 431 kJmol1
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[Use the website search
QUICK INDEX for
GCSE Notes on the basics of chemical energy changes
important to study and know before tackling any of the three Advanced Level
Chemistry pages Parts 13 here
* Part 1ab
ΔH Enthalpy Changes
1.1 Advanced Introduction to enthalpy changes
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 : 1.3ab
for determining enthalpy changes and treatment of results and
Some enthalpy data patterns : 1.4a
The combustion of linear alkanes and linear
1.4b Some patterns in Bond
Enthalpies and Bond Length : 1.4c
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 BornHaber cycle : 2.1ac What happens when a
salt dissolves in water and why? :
cycles involving a salt dissolving : 2.2ac
BornHaber Cycle *** Part 3
ΔS Entropy and ΔG Free Energy Changes
: 3.1ag Introduction to Entropy
entropy values and comments * 3.3a ΔS, Entropy
and change of state : 3.3b ΔS, Entropy changes and the
feasibility of a chemical change : 3.4ad
More on ΔG,
free energy changes, feasibility and
applications : 3.5
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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