1.3a Experimental methods
for determining enthalpy changes and treatment of results
Page introduction
This page
describes experimental
methods for determining enthalpy changes of chemical reactions e.g.
using a simple calorimeter and a bomb calorimeter.
Treatment of
experimental results is covered i.e. how to calculate the molar
enthalpy for the reaction under investigation.
See also
a set of enthalpy problems to
solve, worked out answers given on a separate page!
1.3a Experimental methods
All the methods described here rely on
measuring a temperature change knowing the molar quantities of
reactants and the mass of water used in a calorimeter system
1.3a1
Simple plastic cup calorimeter at room temperature
A simple
polystyrene calorimeter of low heat capacity can be used for any
non–combustion reaction that will happen spontaneously at room
temperature involving solutions or solid reacting/dissolving
with/in a liquid like water and it doesn't matter if the reaction is exothermic or
endothermic.
Reasonably accurate
results can be obtained for a school/college laboratory. The reactants are weighed in if solid
and a known volume of any liquid (usually water or aqueous solution).
The mixture could be a salt and water
(heat change on dissolving) or an acid and an alkali solution (heat
change of neutralisation).
It doesn't matter whether the change is
exothermic (heat energy released or given out, temperature increases) or endothermic (heat
energy absorbed or
taken in, temperature decreases).
Ideally a very accurate mercury thermometer with 0.1 or 0.2
oC graduations should be used or an equally accurate
electronic digital thermometer (can read to 0.01oC via a
thermistor, a solid–state electronic
device).
A double polystyrene cup system provides good thermal
insulation for the system.
Typical apparatus needed for this kind of
calorimetric work with a simple polystyrene cup calorimeter
Safety spectacles, pipette and suction pipette filler
(or a burette, double polystyrene cup calorimeter with insulating lid
plus hole for the thermometer or temperature probe, chemicals of
appropriate mass/volume/concentration, weighing bottle, spatula,
accurate electronic balance, mercury thermometer (preferably 0.1 or 0.2
oC graduations) or electronic thermistor temperature probe,
accurate electronic stop watch or clock.
Note on graphical analysis
To get the best value for the temperature change (ΔT)
you should take multiple readings before and after mixing the reactants
and then plotting a graph versus time.
 Above are
two graphs from simple calorimeter experiments (picture on right).
On the left are typical results from an exothermic
reaction e.g. metal plus acid or metal plus metal salt displacement
reaction.
On the right are typical results from an endothermic
change e.g. when certain salts dissolve in water.
The initial readings
e.g. for 2 minutes before you mixing the final
reactant together give you a baseline, but after
that the
reaction may take a few seconds or a few minutes to completion.
So you cannot get
an immediate true ΔT.
However by drawing a baseline for the initial
temperature and extrapolating back to the start of the reaction (e.g. at
1.5 minutes) you can then estimate the real temperature change.
What happens is quite simple, but it leads to
inaccuracy without doing this extrapolation.
For exothermic reactions the
system will continuously lose heat once the reaction has started, so
the temperature starts to fall once the reaction is complete, so
extrapolating back up gives the true temperature rise.
In this case
ΔT (corrected) = 28.4 -
20.4 = 8.0oC (left graph above)
For endothermic reactions the
system will continuously gain heat once the reaction has started, so
the temperature starts to rise once the reaction is complete, so
extrapolating back down gives the true temperature fall.
In this case
ΔT (corrected) = 19.5 -
11.2 = 8.3oC (right graph above.
Without allowing for these unavoidable
experimental circumstances, you will always measure too low a
temperature change.
Note on the specific heat of water
The specific heat of water is exceptionally high due
to the energy required to break down the hydrogen bonding. When you heat
water, some of the absorbed heat goes into increasing the vibration of
the hydrogen bonds and weakening them and increasing disorder, rather
than being absorbed directly to increase the kinetic energy of the water
molecules i.e. not all the absorbed heat goes to raising the temperature
of liquid water.
When salts (and anything else) are dissolved in
water, the hydrogen bonding is disrupted as water molecules congregate around the ions in
the process called hydration. This has the effect of lowering the specific heat
capacity of water and the more concentrated the salt solution, the lower
the heat capacity of the solution.
Therefore using the specific heat of pure water in
calorimeter data calculations automatically incurs an error!
Examples of
specific heat capacities (SHC) of salt solutions
All of these compounds happen to consist of two ions
and seem to show a similar pattern of specific heat capacity reduction
with increasing salt concentration.
|
molarity
of salt solution, mol dm–3 |
0.0 |
0.1 |
0.5 |
0.9 |
1.0 |
|
SHC
sodium chloride, J g–1 oC–1 |
4.18 |
4.15 |
4.03 |
|
3.90 |
|
approximate
g of NaCl per 100 ml (100 g) water |
0.0 |
0.6 |
2.9 |
|
5.9 |
|
SHC
ammonium chloride, J g–1 oC–1 |
4.18 |
|
4.04 |
3.93 |
|
|
approximate
g of NH4Cl per 100 ml (100 g) water |
0.0 |
|
2.7 |
4.8 |
|
|
SHC
ammonium nitrate, J g–1 oC–1 |
4.18 |
|
4.01 |
3.98 |
|
|
approximate
g of NH4NO3 per 100 ml (100 g) water |
0.0 |
|
4.0 |
7.2 |
|
|
This is all the data I could
find after several hours on the internet!
If anybody finds any useful SHC data for salt
solutions in J g–1 oC–1
please send me the link! |
|
|
|
|
|
Note on
other source of errors
Despite the use of a poorly conducting polystyrene
container and lid, they still absorb/release (exothermic/endothermic)
heat and a small amount of heat will still be lost/gained to/from the
surroundings. The amount of heat involved is small, but not zero!
The glass mercury thermometer or thermistor
temperature probe will absorb some heat (if exothermic)
or release heat (if endothermic), therefore reducing or increasing the
measured temperatures.
If a solid reactant is used or formed e.g. metals in
a displacement reaction, there is a small error from unreacted metal or
metal formed, but the specific heat of metals or other solids is usually
quite small.
The specific heat of aqueous solutions is less than
that of pure water (see examples of data above and the multiple
calculation and discussion in experiment 3.).
A 'low resolution' thermometer, reading to the
nearest 0.5oC is not accurate enough unless the temperature
change is quite big e.g. ΔT >20o.
You can also weigh the
water for a little greater accuracy compared to just
measuring out a volume of water with a measuring
cylinder (not a good idea) or preferable an accurate
pipette.
|
1.3a2 A simple
copper calorimeter for
combustion
A simple system, very inaccurate, but
is specifically
used for determining the heat energy released (given
out) for burning fuels.
The burner is weighed before and after
combustion to get the mass of liquid fuel burned. The thermometer records the
temperature rise of the known mass of water (1g ~ 1cm3 since
density of water is ~1.0 gcm–1).
You can use this system to compare the heat output from burning various
fuels. The bigger the temperature rise, the more heat energy is
released. You can take a series of measurements with time and from the
graph extrapolate the maximum temperature rise.
This is a
very inaccurate method because
of huge losses of heat e.g. radiation from the flame and calorimeter,
conduction through the copper calorimeter, convection from the flame
gases passing by the calorimeter etc. BUT, at least using the same
burner and set–up, you can do a reasonable comparison of the heat output
of different fuels.
Other sources of error: Heat is also absorbed by the
copper calorimeter. The flame may be smoky, indicating combustion is
incomplete.
You can also do a calibration by burning a fixed
amount of material whose enthalpy of combustion is accurately known.
Typical apparatus needed for this kind of
calorimetric work with a copper calorimeter
You need safety spectacles, accurate measuring cylinder for
water (or weigh calorimeter before and after adding the water), copper
vessel calorimeter with insulating lid plus hole for the thermometer or
temperature probe, draught shielding around the burner and calorimeter,
a small burner/lamp with wick, suitable combustible material (alcohols
like ethanol/propanol burn more cleanly than alkanes like hexane),
accurate electronic balance, mercury thermometer (preferably 0.1 or 0.2
oC graduations) or an accurate electronic thermistor
temperature probe, accurate electronic stop watch or clock.
|
1.3a3
The
Adiabatic Bomb Calorimeter
Method 1.3a2 is quite a crude and inaccurate method for determining
enthalpies of combustion. The bomb calorimeter method delivers very
accurate results. The idea is execute the reaction adiabatically –
meaning no heat loss from the system. The compound undergoing combustion is electrically
ignited, and does so under pressure in an atmosphere of pure oxygen gas
and this ensures complete and rapid combustion. There may be several
water baths and thermometers to make sure every joule of energy released
is absorbed by the calorimeter and measured.
Calibration: The heat capacity of a bomb
calorimeter can be very accurately determined by combustion of a
standard substance like benzoic acid whose enthalpy of combustion is
very accurately known from previous experiments.
The bomb calorimeter method involves
measuring the heat released at constant volume and is strictly speaking
called the internal energy change ΔE or
ΔU. Methods 1.3a1 and 1.3a2
involve heat energy changes at constant pressure and directly measure
ΔH.
From 1.3a bomb calorimeter measurements you can calculate the enthalpy change
from the equation
ΔH = ΔE + ΔnRT (Δn = the net change in moles of gas in the
reaction at 298K/101 kPa)
I don't think this equation is needed for any UK pre–university advanced
level chemistry course these days?
If there are no
gaseous reactants or products (i.e. only liquids/solids involved) OR if
moles gaseous reactants = moles gaseous products, then Δn = 0 and ΔH
= ΔE or ΔU).
|
|
1.3b Treatment of
experimental results
(general method for any calorimeter)
In any
calorimeter the heat released or absorbed is given by
energy
transferred = m x
SHCH2O x
ΔT
(sometimes expressed simply as q = m c ΔT)
SHCH2O
= specific heat capacity of water (4.18 J g–1 K–1)
i.e. it takes
4.18 J of heat energy to raise 1g or 1cm3 of water by 1o.
This assumes
the heat capacity of the water is the same as the solution in
method 1.3a1
This
not actually true.
m = mass of
the water absorbing the heat, usually grammes.
This ignores
the mass of the calorimeter, thermometer, insulation etc.
ΔT =
temperature change (Tfinal – Tinitial)
This cannot
take into account heat energy losses, which any experiment
should be designed to minimise, so its only what you can
actually measure directly.
Mass is
converted to moles so you can then relate this heat
energy change (enthalpy) to the mass/molar quantities used to get
the ΔHreaction enthalpy change in kJ mol–1.
Examples of calculations
using data from various calorimetric methods
The calculations are based on
experimental data alone OR a combination of standard data and
experimental data.
Either way, many involve using
Hess's Law, e.g. a 'simple' triangular arrangement (see
Hess's Law
Notes and Hess's Law cycles below, which you must be very
familiar with.
|
|
Application of Hess's Law (i)
ΔHθ1 |
A  B |
|
ΔHθ2 |
 |
|
 |
ΔHθ3 |
|
C |
|
Clearly there are two pathways from A to B
direct and via C following the arrows
direction
∴
from Hess's Law:
ΔHθ1
= ΔHθ2 + ΔHθ3
or
ΔHθ2 = ΔHθ1
– ΔHθ3
or
ΔHθ3 = ΔHθ1
– ΔHθ2 |
|
|
Application of Hess's Law (ii)
ΔHθ1 |
|
A
B |
|
ΔHθ2 |
|
|
 |
ΔHθ3 |
|
C |
|
In this case there are two pathways from A to
C direct and via B following the
arrows direction
∴
from Hess's Law: ΔHθ2
= ΔHθ1 + ΔHθ3
and ΔHθ1
= ΔHθ2 – ΔHθ3
or
ΔHθ3
= ΔHθ2
–
ΔHθ1 |
|
|
Always take care with the ΔH signs
which ever way you set up the triangle to apply Hess's Law |
Although slightly more awkward, (ii) might be
better suited to the way the experiment results are obtained |
1.3c Key revision points
about simple practical calorimetry procedures to determine an
energy change i.e. to determine a
ΔH
Calorimetry is a core practical across A level, IB, CIE, and AP
chemistry courses.
Students must master the
q = mcΔT
equation, understand how to calculate enthalpy changes per mole, and
avoid common pitfalls such as heat loss, incorrect sign conventions,
and misidentifying the limiting reagent.
Exam boards consistently
test both the methodology and the sources
of error in calorimetry experiments.
-
Definition of Enthalpy Change (ΔH):
Heat energy transferred at constant pressure during a chemical
reaction.
-
Calorimetry Equation:
-
q = m
c ΔT
-
Where q = heat energy (J),
m
= mass of solution (g), c = specific heat capacity
(J g⁻¹ K⁻¹), ΔT = temperature change (K).
-
Molar Enthalpy Change:
-
Types of Calorimetry Experiments:
-
Combustion calorimetry
(fuel burned to heat water).
-
Neutralisation calorimetry
(acid + alkali).
-
Solution calorimetry
(dissolving solids in water).
-
Practical Setup: Polystyrene
cup calorimeter, thermometer, lid, stirrer, insulation to reduce
heat loss.
-
Graphical Extrapolation: Plot
temperature vs. time to correct for cooling and determine
maximum temperature change.
Common Misconceptions
about
practical
calorimetry to determine a
ΔH
-
Heat Loss Neglected: Students
often assume all heat goes into the solution. In reality, heat
is lost to the surroundings, thermometer, and calorimeter walls.
-
Wrong Sign Convention:
Forgetting that exothermic reactions have negative ΔH values.
-
Incorrect Limiting Reagent:
Using the wrong reactant when calculating ΔH per mole.
-
Mass Misidentification:
Confusing mass of water with volume (1 cm³ ≠ 1 g unless density
is assumed).
-
Temperature Units: Using °C
instead of K in ΔT. (Numerically the same, but examiners expect
clarity).
-
Specific Heat Capacity Errors:
Forgetting that water has (c = 4.18 , J g-1 K-1
-
Incomplete Combustion: In
combustion calorimetry, assuming all fuel burns completely when
soot formation or incomplete combustion occurs.
Exam Revision Tips
for questions involving
practical
calorimetry to determine a
ΔH
-
Always State Assumptions: e.g.,
“no heat loss to surroundings,” “density of solution = 1 g
cm⁻³.”
-
Draw Diagrams: Label
calorimeter, thermometer, lid, stirrer. Many exam boards award
marks for clear experimental setup sketches.
-
Error Evaluation: Be ready to
discuss improvements (use of bomb calorimeter, insulation, lid,
digital thermometer).
-
Practice Calculations: Convert
between joules and kJ mol⁻¹ carefully.
-
Graph Skills: For
neutralisation experiments, practice extrapolating cooling
curves to find maximum ΔT.
-
Compare Methods: Know
differences between simple polystyrene cup calorimetry and more
accurate bomb calorimetry.
-
Cross-board Consistency:
-
Focus on required practicals and error analysis.
-
BUT, not
neglecting an emphasis on thermodynamics, system versus surroundings, and enthalpy
definitions.
-
Exam Technique:
-
Show all working with units.
-
Clearly identify limiting reagent.
-
State whether ΔH is per mole of reactant or
per mole of reaction.
-
Use negative signs correctly for exothermic
reactions.
Final
tip for
practical
calorimetry to determine a
ΔH:
Examiners love when
students explicitly mention heat loss corrections
and experimental improvements.
Practising past
papers across multiple boards will highlight recurring themes:
calculation accuracy, error discussion, and clear experimental
design.
For the next problem to designed!
If anybody has a good suggestion for a
pre–university practical, I'd be happy to look at it.
Enthalpy calculation problems with worked out answers – based on
enthalpies of reaction,
formation, combustion
Energetics-Thermochemistry-Thermodynamics Notes INDEX
TOP OF PAGE
key phrases: how to
measure enthalpy changes using a calorimeter errors in calorimetric
measurements polystyrene calorimeter what is a bomb calorimeter? how
do you measure enthalpy changes with a bomb calorimeter graphical
analysis of thermochemistry data measurements observations specific
heat of water values for the specific heat of salt solutions how do
you use Hess's Law to calculate enthalpy changes that you cannot
measure directly in a calorimeter how to measure the enthalpy of
combustion of ethanol in a copper calorimeter how to measure the
enthalpy of combustion of an organic compound in a bomb calorimeter
how do you measure the enthalpy of solution dissolution of ammonium
nitrate how do you calculate the enthalpy of solution of ammonium
nitrate how do you measure the enthalpy of solution dissolution of
potassium chloride how do you calculate the enthalpy of solution of
potassium chloride how can you measure the enthalpy of
neutralisation of sodium hydroxide and hydrochloric acid? how can
you measure the enthalpy of reaction for the zinc metal copper
sulfate solution displacement reaction how can you measure the
enthalpy of hydration of anhydrous copper sulfate? how do you
calculate the enthalpy of hydration of anhydrous copper sulfate
using Hess's Law? how can you measure the enthalpy of hydration of
anhydrous magnesium sulfate? how do you calculate the enthalpy of
hydration of anhydrous magnesium sulfate using Hess's Law? How can
you measure the enthalpy of reaction change for the thermal
decomposition of sodium hydrogencarbonate? How can you calculate the
enthalpy of thermal decomposition of sodium hydrogencarbonate using
Hess's Law? using a calorimeter to make enthalpy measurements
calculations
for AQA AS chemistry, using a calorimeter to make enthalpy measurements
calculations
for Edexcel A level AS chemistry, using a calorimeter to make enthalpy
measurements calculations for A level OCR AS chemistry A, using a
calorimeter to make enthalpy measurements calculations for OCR Salters AS chemistry B,
using a calorimeter to make enthalpy measurements
calculations for AQA A level chemistry, using a calorimeter to make
enthalpy measurements calculations for A level Edexcel A level chemistry,
using a calorimeter to make enthalpy measurements
calculations for OCR A level chemistry
A, using a calorimeter to make enthalpy measurements calculations for A level OCR Salters A
level chemistry B using a calorimeter to make enthalpy measurements
calculations for US Honours grade 11 grade 12 using a calorimeter to
make enthalpy measurements calculations for
pre–university chemistry courses pre–university A level revision
notes for using a calorimeter to make enthalpy measurements calculations A level guide
notes on using a calorimeter to make enthalpy measurements calculations for schools colleges academies science course tutors images
pictures diagrams for using a calorimeter to make enthalpy measurements
calculations A level chemistry revision notes on
using a calorimeter to make enthalpy measurements calculations for revising module topics notes to help on understanding of
using a calorimeter to make enthalpy measurements calculations university courses in science
careers in science jobs in the industry laboratory assistant
apprenticeships technical internships USA US grade 11 grade 11 AQA A
level chemistry
notes on using a calorimeter to make enthalpy measurements
calculations Edexcel
A level chemistry notes on using a calorimeter to make
enthalpy measurements calculations for OCR A level chemistry
notes WJEC A level chemistry notes on using a calorimeter to
make enthalpy measurements calculations CCEA/CEA A level
chemistry notes on using a calorimeter to make enthalpy
measurements calculations for university entrance examinations How
to do enthalpy determinations using
a calorimeter system and explaining how to do the calculations, how
to determine the enthalpy of
combustion of an alcohol, how to determine the enthalpy of
combustion of benzoic acid with a bomb calorimeter, how to determine
the enthalpy of solution of ammonium nitrate, how to determine the enthalpy of
dissolution of potassium chloride or sodium carbonate with a simple
plastic
calorimeter, how to determine the enthalpy of
neutralisation of hydrochloric acid and sodium hydroxide, how to
determine the enthalpy of
reaction of zinc displacing copper from copper(II) sulfate solution,
how to use a Hess's Law enthalpy cycle and
experimental data to determine the enthalpy of hydration of
anhydrous copper(II) sulfate, how to use Hess's Law and
experimental data to determine the enthalpy of hydration of
anhydrous magnesium sulfate, how do you use Hess's Law and
experimental data to determine the enthalpy of the decomposition of
sodium hydrogencarbonate to sodium carbonate, water and carbon
dioxide? There are descriptions and discussions of types of calorimeter and procedures
to determine enthalpies of reaction, the principles behind the
enthalpy calculations from calorimeter procedures are described, laboratory equipment needed
for determining enthalpy values, sources of error in enthalpy
experiments, the real (actual) and assumed specific heat of water and salt
solutions is discussed
TOP OF PAGE
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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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