4.
Determining the enthalpy
of solution of potassium chloride
with a simple plastic 'coffee cup' calorimeter
(method 1.3a1)
Enthalpy
of solution ΔHsolution(potassium chloride) is defined
as the total heat energy absorbed or released when 1 mole of a
potassium chloride
dissolves in water to form a dilute solution.
The
determination of the enthalpy of solution (dissolution) of
potassium chloride.
KCl(s) + aq ===> K+(aq)
+ Cl-(aq)
Method (for more experimental details
see
method 1.3a1)
Using a
simple polystyrene calorimeter 4.0 g of potassium
chloride was dissolved in 50 ml (50 g) of water.
(Assuming density of water is 1.00 g/cm3,
but you can weigh the water to be more
accurate).
Reading every
30 seconds, measure the temperature for few
minutes before adding the accurately weighed
amount of potassium chloride salt.
When the
potassium chloride is added, stir gently to
make sure all the salt is dissolved and take
temperature readings for another 8-10
minutes.
Typical
temperature-time
graph shape for an calorimetric endothermic experiment
Graphical
analysis (e.g. like above) showed the temperature
fell by 4.40oC.
This is
obtained using the extrapolated purple lines
to correct as much as possible for heat
energy gain as the calorimeter and contents
start warming up as soon as the potassium
chloride dissolves!
Calculation
Calculate
the enthalpy of dissolution (dissolving) of
potassium chloride.
Heat
absorbed by the water = mass of water x SHCwater
x temperature
q = m c
ΔT
(m in g, c = 4.18 Jg-1K-1,
ΔT temperature rise in oC)
Mr(KCl)
= 39 + 35.5 = 74.5; mol KCl = 4.0/74.5 = 0.05369
mol
(i) Using
the value of the heat capacity of pure water and
just the mass of the water.
q = m
c ΔT
q =
50 x 4.18 x 4.40 = 919.6 J, 0.9196 kJ
Scaling up to 1 mole = 0.9196 x (1/0.05369) =
17.1 kJ mol–1
Since
the temperature has fallen, the change is
endothermic
So ΔHdissolution(KCl)
=
+17.1 kJ mol–1
(only accurate to 3 sf)
This
value is not too dissimilar from the
theoretically better calculation (i) below.
(ii)
Using the total mass and the likely SHC of the
potassium chloride solution, assuming it is similar
to sodium chloride at around a 1 molar
concentration.
q = m
c ΔT
q =
54 x 3.90 x 4.40 = 926.6 J, 0.9266 kJ
Scaling up to 1 mole = 0.9266 x (1/0.05369) =
17.3 kJ mol–1
Since
the temperature has fallen, the change is
endothermic
So ΔHdissolution(KCl)
=
+17.3 kJ mol–1
(only accurate to 3 sf)
Both
methods (i) and (ii) give similar results, again
justifying the textbooks/internet pages that ignore
the mass of the salt and using the SHC of pure
water.
Determining the enthalpy
of solution of sodium carbonate
with a simple plastic 'coffee cup' calorimeter
(method 1.3a1)
Enthalpy
of solution ΔHsolution(sodium carbonate) is defined
as the total heat energy absorbed or released when 1 mole of a
sodium carbonate
dissolves in water to form a dilute solution.
The
determination of the enthalpy of solution (dissolution) of
potassium chloride.
Na2CO3(s) + aq ===>
2Na+(aq) + CO32-(aq)
Method
(for experimental details
see
method 1.3a1)
Reading every
30 seconds, measure the temperature for few
minutes before adding the accurately weighed
amount of sodium carbonate.
When the
sodium carbonate is added, stir gently to
make sure all the solid is dissolved and
take temperature readings for another 8-10
minutes.
Typical
temperature-time graph
shape for a typical calorimetric exothermic experiment.
An
example of an exothermic enthalpy of solution
(dissolution) is dissolving anhydrous sodium
carbonate in water.
Na2CO3(s) + aq
===> 2Na+(aq) + CO32-(aq)
On
dissolving 8.5 g of anhydrous sodium carbonate in 50
cm3 of water, the temperature rose by a maximum of
8.1oC e.g. using an extrapolation graph
method illustrated above.
This is
obtained using the extrapolated purple line
to correct as much as possible for heat
energy loss as the calorimeter and contents
start cooling down as soon as the sodium
carbonate dissolves!
Calculation
Calculate
the enthalpy of solution of anhydrous sodium
carbonate.
Formula mass of Na2CO3
= 106
q = m c
ΔT
(m in g, c = 4.18 Jg-1K-1,
ΔT temperature rise in oC)
q = 50 x 4.18 x 8.1 = 1692.9 J, 1.6929 kJ
mol Na2CO3
= 8.5/106 = 0.0802
Scaling
up to 1 mol = 1.6929 x 1.0/0.0802 = 21.108
Since the
temperature rose, exothermic, therefore
So ΔHdissolution(Na2CO3)
= –21.1
kJ mol–1
(only accurate to 3 sf)
Experimental methods of determining enthalpy changes
Energetics–Thermochemistry–Thermodynamics Notes INDEX
thermochemistry of determining the enthalpy of solution of potassium chloride
& sodium carbonate, method, apparatus, data, calculation, plastic cup
calorimeter, thermochemistry for AQA, Edexcel, OCR, Salters, CIE, WJEC Eduqas &
CCEA A-level chemistry exam students, US grades 11-12 K12 AP Honors chemistry
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