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Appendix
9. Colorimetry – quantitative
analysis and determining the formula of a complex ion
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(a) What is colorimetry?
what is a colorimeter? how does a colorimeter work?
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Colorimetry is
method of determining the concentration of a substance by measuring the
relative absorption of light (usually visible) with respect to a known
concentration of the substance.
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The instrument by which
these measurements are made is called a colorimeter (illustrated
below).
-
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Light from a suitable
source is passed through a light filter to select the most
appropriate wavelength of light, some of which is then absorbed by the
solution held in a special glass cuvet (a sort of 'test tube').
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The amount of light
absorbed is called, and measured as, the absorbance which is a
function of the coloured solute concentration.
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Most expensive
instruments use a double beam system of two cuvets, one is a 'blank'
of water and one the actual coloured solution under test, two photocells and
sophisticated optics of lenses and mirrors which need not concern as at all.
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Cheaper colorimeters
(i.e. in school and illustrated above) allow you to put in a cuvet of
'colourless' water, zero the instrument i.e. set it to read zero absorbance,
replace with a cuvet of the coloured solution and simply read of the
'absorbance'. The 'zeroing' is necessary because even the apparently
'colourless blank' of glass cuvet and water can absorbed a tiny amount of
light. This procedure eliminates this error.
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The filter is chosen
to select the band of wavelengths which are most strongly absorbed by the
coloured solution e.g. this is illustrated on the diagram above, and in
the table below, by using a yellow filter to use in measuring the
concentration of a blue coloured solution like copper(II) sulfate or its
ammine/amine complex.
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|
The wavelength (nm) of the observed transmitted colour of the solution |
The observed transmitted colour of the solution
(* as
in the diagram above) |
The complementary colour of the
solution i.e. the colour of the filter |
|
400–435 |
violet |
yellowish–green |
|
435–480 * |
blue * |
yellow
* |
|
480–490 |
greenish–blue |
orange |
|
490–500 |
bluish–green |
red |
|
500–560 |
green |
purple |
|
560–580 |
yellowish–green |
violet |
|
580–595 |
yellow |
blue |
|
595–610 |
orange |
greenish–blue |
|
610–750 |
red |
bluish–green |
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Although the table
illustrates the 'complementary' colour relationship between the solution and
the filter, in practice it is better to try several filters on a typical
concentration of the solution under test to see which filter gives the
highest absorption value i.e. gives you maximum sensitivity and hence
maximum accuracy in your measurements.
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How do we use
colorimetry to measure the concentration of a transition metal ion?
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If an aqueous transition
metal ion is intensely coloured, its concentration can be measured directly
e.g. manganese concentration can be measured if it is oxidised to the deep
purple manganate(VII) ion, MnO4–.
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However, many ions are
not as intense as the MnO4– ion, BUT if a suitable
ligand or complexing agent is added, then a more intensely coloured complex
may be formed, from which accurate measurements of concentration can be made
e.g.
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The blue
hexaaquacopper(II) ion forms a deeper violet–blue ammine complex with
ammonia.
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Yellow–brown iron(III)
ions form a deep blood–red complex with the thiocyanate ion (SCN–)
by mixing it with ammonium/potassium cyanate.
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Once the method of
producing a more intense colour is established, you then need to derive a
calibration curve.
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This is done by
measuring the absorbance of solutions of known concentrations of the
coloured complex and plotting the calibration curve/graph (see right of
colorimeter diagram).
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The known concentration
range should include any likely absorbance measured from the solutions under
test i.e. those whose concentration is being determined.
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Generally at low
concentrations the calibration curve is linear i.e. it obeys Beer's
Law (Beer–Lambert Law). Without going into the mathematics of Beer's Law
and absorption, it basically states that a solution's absorbance is
directly proportional to the concentration of the coloured solute.
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For various reasons the
calibration may curve upwards (positive deviation from Beer's Law) or curve
over (negative deviation from Beer's Law). However, a linear or otherwise
calibration curve still shows increasing absorption with increasing
concentration and curved calibration graphs are acceptable, if not
advisable, if your methodology is accurate.
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How can we use
colorimetry to deduce the formula of a complex?
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The method depends on
measuring the absorbencies of solutions containing different ratios of
transition metal ion to complexing agent.
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Just for the sake of
argument, imagine that one mole of transition metal (M3+) ion
reacts with two moles of a monodentate ligand (X–). The reaction
equation for the ligand displacement reaction to form the complex would be:
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There are two basic
approaches as to how you vary the transition metal ion – ligand ratio and
the results illustrated in the diagram below.
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Method (1) The mole
ratio method keeping Mn+ constant and gradually increasing
the number of moles of ligand X from zero to a large molar excess.
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Method (2) The
continuous variations method in which you start with zero moles of Mn+
and an excess of the ligand X. In each successive mixture you then increase
the amount of Mn+ and decrease the amount of X– and
keep the total moles of Mn+ and X– constant.
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For the sake of
argument, if you assume both stock solutions are the same molarity,
then the ratio of the volumes automatically gives you the X/M ratio in the
mixture.
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For the fictitious M3+
and F– complex, the peak will occur between the 6X– :
4M3+ mixture (ratio 1.5) and the 7X– : 3M3+
mixture (ratio 2.33). Theoretically the peak should occur at a ratio of 2.0
for X–/M3+ i.e. theoretically, in terms of a total
volume of 10 units, it means a 6.66X– : 3.33M3+
mixture by volume.
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For the
iron(III)–thiocyanate complex in reaction (ii) the peak will occur at 5CNS–
: 5Fe3+
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For the
copper(II)–ammonia complex in reaction (iii) the peak will occur at 8NH3
: 2Cu2+
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In both cases, you work
out the mole ratios present in the mixtures from the known concentrations
and volumes of the stock solutions of Mn+ and X– used
in each experiment.
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OTHER APPLICATIONS of
COLORIMETRY
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The manganate(VII) ion, MnO4–,
e.g. in potassium manganate(VII) solution, is a brilliant purple colour
and its concentration in very dilute solution can be measured by using
colorimetry i.e. by comparing the absorbance of the solution versus a
calibration graph of known concentrations of the manganate(VII) ion.
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This can be applied to kinetic
experiments e.g. measuring the rate of reaction of acidified potassium
manganate(VII) when it oxidises ethanedioic acid/ethanedioate ion. The
reaction is autocatalysed by the manganese(II) ions formed.
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Therefore, theoretically, via
the intensity of the manganate(VII) ion colour, you can investigate the
effects of changing the concentration of potassium manganate(VII),
ethanedioic acid, dilute sulfuric acid and manganese(II) sulfate.
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A simplistic view would be to
write an overall rate expression such as ...
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rate = k[MnO4–]a[(COOH)2]b[H+]c[Mn2+]d,
where a, b, c and d represent the orders of reaction.
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However:
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(i) good results are
not easy to obtain, specifically due to the effect of autocatalysis
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(ii)
it is unlikely the rate expression is as simple as you will encounter in
your pre–university course!
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Periodic
Table - Transition Metal Chemistry - Doc
Brown's Chemistry Revising
Advanced Level Inorganic Chemistry Periodic Table
Revision Notes colorimetric analysis of a 3d block transition metal
complex
