UNIT 5 ESTIMATION OF COPPER

Structure
5.1 Introduction TAP&^
Objectives
5.2 Iodimetry and I
Indicator

Solution by Iodometric Method

Principle
Requirements
Procedure
Observations
Calculations "
Result
5.4 Colorimetry
Beer-Lambert Law
Principle of Colorimeter
5.5 Experiment 8: Determination of the Percentage of Copper in the Given
Solution by Colorimetric Method
Principle
Colorimeter
Calibration of Colorimeter for ~olori&tric Measurement
Requirements
Procedure
Observations
Calibration Curve
Calculations
Result
5.6 Answers to SAQs

5.1 INTRODUCTION
In the previous unit, you have estimated the mount of ferrous iron, Fe2+,in a
sample of iron filings by using two redox titrations, namely, permanganatometry and
chromatometry. In this unit, we would estimate the amount of copper in a given
sample. Here too, you would perform two experiments, one of which is based on a
redo3 reaction, iodometry, while the other is based on colorimetric determination.
Iodomeaic titrations make use of 12/1- redox reaction and the end point is detected
by using starch as an indicator. Colorimetry, on the other hand, is a method of
analysis based on comparing the colour intensity of an unknown with that of a
standard solution, i.e., the solution of a definite known concentration. The theory
behind iodometric and colorimetric determination of cupric ions, Cu2+is given along
with the procedural details of the experiments.

Objectives
After studying this unit and performing the experiments, you will be able to:
dehne and differentiate between iodometry and iodirnetry,
explain the redox reactions involved in iodometry,
explain the use of indicator in iodometry and standardise the given sodium
thiosulphate solution,
use the iodometric method in estimating Cu2+ions,
state Beer-Lambert law,
explain the principle of colorimetry,
d d b e the calorimeter and its calibration, and
use the colorimetric method in estimating Cu2+ions,

5.2 IODMEIXY AND IODOMEX'RY

lodine is a mild oxidising agent and in the presence of a suitable reducing agent, gets
reduced to iodide ions, I-, according to the following equation:
12+2e+ 21, I?--0.54V ....(5.1)
Quantitative Analysis-[I On the other hand, a variety of oxidising agents can oxidise I- ions into I,. In-fact,
both these reactions are made use of in analytical chemistry. T i t r a w involving the
use of I, as a titrant to estimate the reducing agents are tenned as k W d r k
lodimetric titrations are used for
. .
trtrrrtnnrs. Iodine, being a much weaker oxidising agent than potassium p-
estimating reducing agents while and potassium dichromate, has limited applicability. Moreover, it is very volatile iq
iodometric titrations are used for nature and also has poor solubility.
oxidising agents.
In certain cases, the oxidising agent to be determined is mixed with an excess of
potassium iodide, KI, and kept for some time. The iodine, liberated during the
reaction, is titrated against a standard solution of a reduciug agent,e.g., sodium
lodimetry: Titration with iodine
lodometry: Titration of iodine thiosulphate, Na,S,O,. These titrations are referred to as iodometric tig.tiolls. Since
produced by a Cu2+ions, can behave as an oxidising agent by getting reduced to Cu+ ions, we can
chemical reaction use iodomehic method for their determination. ,-

Ideally an iodometric titration should be a titration using KI as a titrant to titrate the

oxidising agent. In such a reaction, more and more of iodine is liberated from iodide
ions as the titration proceeds. The end point of such a titration would be a stage
where the liberation of iodine ceases. It is impossible to detect this end point with the
help of an indicator. Starch can be used to detect the 'just appeamnce' or the 'just
disappearance' of iodine but not the cessation of I, formation.

An indirect method of end point determination becomes essential in such cases. A

known amount of the solution of the oxidising agent (to be determined) is meamred
and mixed with an excess of a solution of KI and acid. The solution is then left for
about five minutes in the dark for the reaction to complete and the libawd iodine is
titrated with a standardised solution of sodium thiosulphate using starch as the
indicator. The following reaction takes place:

An excess of KI is used because iodine has got very poor solubility in water. Iodine
forms an unstable complex, KI,, with KI which is readily soluble in H20.

In fact, iodine in an aqueous solution containing KI exists mainly as the triiodide ion,
I; and there is an equilibrium between I; ion and I,. In the course of the titration, as
I, is consumed, more and more of I; ions dissociate to give I, which reacts with
thiosulphate. Further, such a titration should be carried out in d d , as I, is volatile
and also the indicator, starch, loses its sensitivity at high temperatures.
SAQ 1
Give two limitations of I, as a titraot.

In principle, iodine can be used as a self indicator like KMnO,, as a drop of iodine
can impart a pale yellow colour to a solution. Since the cdm imparted by iodine is
quite faint, in practice, it becomes diBcult to use this as an indication of the end
point. Iodine is known to form a blue cdoured admpt~oncomplex with stuch. This
property of starch is exploited in using it as an indicator for titratioas invoking
iodine.

In an iodomehic determination, we titrate I, with s,o:- ions and at the end point,
The use of starch enhances the
addition of one drop of s,o:- ions should just decolourise the blue d o u r of
-tivity of the determination of starch-iodine complex. In such titrations, starch should be added just before the end
the end point. point, when a very little amount of I, remains and the solution being titrated has a
faint straw yellow d o u r . If starch is added earlier, i.e., when a large amount of
iodine is present, a large amount of starch-iodine complex is formed. This complex
20 reacts quite slowly with s,o:- and it is likely that the solution is over titrated.
 Estimation of Copper
5.2.2 Standardisation of Thiosulphate
As said above, in iodometry we titrate the liberated iodine with a standardised
solution of sodium thiosulphate. Though sodium thiosulphate, Na2S20,.5H,0, can be
obtained chemically pure, a standard solutio~of thiosulphate cannot be made by
exact weighing. This is because thiosulphate reacts with atmospheric 0,and also the
CO, &solved in water. More so, even some microorganisms decompose
thiosulphate.

A of oxidising agents are available for the standardisation of Na,S,03.

Potassium dichromate is normally used for the purpose.

In acidic medium c~,o:-

ion gets reduced to Cr (111) as shown in the following
equation :
+ +
c~,o;- 14H+ 6e G===+ 2Cr3+ 7H20 + ...(5.3)
and iodide ions from KI get oxidised to I,:
21- d 2; 2e +
To maintain electron balance, multiplying the above equation by 3, :ve get,
61- F=--' 31, + 6e ...(5.4)

The overall ionic equation for the tiuation can be obtained by adding Eq.5.3 and
Eq. 5.4,
cr20;- + 14H++ 61- -----* + + 7H20
2 ~ ? 31, ...(5.5)
We see from Eq. 5.5 that one mole of potassium dichromate reacts with 6 moles of
potassium iodide liberating 3 moles of iodine.

The liberated iodine, in turn, reacts with sodium thiosulphate solution as,
2~~0:- S,O;- 2e + ...(5.6)
...(5.1)
I, + 2e- 21-

Since three moles of I, are liberated by c~,o;-, ,

+
31, 6e+ 61- ...(5,7)

The overall ionic equation for the titration of liberated I, with sodim thiosulphate
can be obtained by adding Eq. 5.6 and Eq. 5.7,

The net chemical reaFon involving a titration of potassium dichromate and sodium
thiosulphate in the presence of excess potassium iodide can be written by combining
Eq. 5.5 and Eq. 5.8,

We see from Eq. 5.9 that one mole of potassium dichromate is equivalent to 6 moles
of sodium thiosulphate. Therefore, substituting the values of p and q in Eq. 1.8, the
molarities are related by the following relationship.
Factor '6' here signifies that one
mole of K2Cr,0, Liberates 3
mdes of I, which is equivalent to
6 moles of sodium thiosulphate.
or MI V, = 6M2V2 . ----
where M, and M, molarities o f , s o d ~thi&%=&&otassium
dichromate V, re-nt the volumes of e r n thiosblp a e
respectively. -"3
Quantitative Analysis-I1

PERCENTAGE OF COPPER IN THE GIVEN

SOLUTION BY IODO-C MEIMOD
Many a time, an analytical chemist - h n f r o n t e d with the problem of finding out the
amounts of some metals, e.g., Fe, Cu, etc. in a given sample. The sample may be of
an ore or an alloy. Let us see how we carry out such an estimation for Cu in a given
sample. We can do this by iodometric titrations. As said before, like
permanganatometry and chromatometry, this titration is also based on a redox
reaction.

To determine the amount of Cu in a given sample, a known mass of it is dissolved by

suitable chemical treatment giving a solution of Cu2+ions. This solution is titrated
against a standard solution of sodium thiosulphate in the presence of an excess of
KI. The Cu2+ions on reacting with KI get reduced to Cu+ ions and liberate an
equivalent amount of I, by oxidking I- ions. This liberated iodine then reacts
quahtitatively with S,O$- ions, and in turn, gets reduced to I- ions. The principle and
the equations involved are given in the next suh&on.

5.3.1 Principle

-
The reaction between Cu2+and N%S203in acidic medium, in the presence of excess
of KI, involves ~xidationof S,Oj- to S,O&, tetrathionate ion, and reduction of Cu2+
to Cu'. The reaction between Cu2+and KI is given as,
21- +
I, 2e ...(5.10)
Cu2++ e-===+Cu+ ...(5.11)

Balancing the reaction between Cu2+and potassium iodide by combining Eq. 5.10
and Eq. 5.11, we get,

We see that two moles of Cu2+react with two moles of potassium iodide and tie
liberated iodine reacts with sodium thiosulphate, shown earlier also, in the following
manner:
2S20!- 4S,O$- + 2e ...(5.6)
+
I, 2e= 21- ...(5.7)
I, + 2S2@- - 21- + S40g- ...(5.8)
The net chemical reaction involving a titration of copper (I][) and sadium thiosulphate
in the presence of excess ~ t a s s i u miodide can be written by combining Eq. 5.12 and
Eq.5.8.

We see from Eq. 5, lFr t b t two moles of copper (II) are equivalent to two mdes of
sodium thiosulphate. In o&e~'wo& one mole of copper (11) is equivalent to one
mole of sodium thiosulphate.

Therefore, substituting the values of p and q in Eq. 1.8, the molarities are related by
the following r e l a t i d p :
- - ----
4- -
M , < C 1

of sodium thiosulphate and c6pper @)

solutions, and V, and V, the volumes of sodium thiosulphate and copper @)
solutions, respectively.
According to the above discussion, the iodometric determination of Cu2+ions is Estimation of Copper
based on the following reaction :
2Cu2+ KI + -2CuZ 4K+ I2
+ + + ...(5.14)
(ex-)
where cupric ions are reduced to cuprous ions and iodide ions are o x i M to iodine.
A look at the standard reduction potentials of Cu2+/Cu+and 12/1- couples :
Cu?++e----"Cu+ l?'=0.17V ...(5.15)
+
1, 2e d 21- i? = 0.54 V ...(5.1)
suggests that the r d o n represented by Eq. 5.14 should proceed in the revme
direction, i.e. iodine should oxidise Cu+to Cu2+,but actually the reaction occurs as
given in &. 5.14. The Cd formed during the reaction has a very low solubility in
water, therefore, the amcatration of the reduced form, Cu+, is greatly reduced and
the potential of Cu2+/Cu+-couplebecomes greatei than that of 12/21-. This explains
the achal course of reaction.
SAQ 2
Write the chemical equations involving a titration of copper (11) with thiosulphate in
presence of excess KI.
[Hint : It involves two steps]

MP-h chenicrls
Burette (50 cm3) 1 - Potassium dichromate
Pipette (20 an3) 1 - Dilute sulphuric acid
Gmical~(250an3)-1 Potassium iodide
Beaker (250 an3) - 2 Glacial acetic acid
Fuanel (small) - 1 Potassium thioqamte
Volumetric flask (250 an3) 1 - Distilled water
MecuWing cylinder (10 m3) 1 -
Test Tube - 1
W a s h b o t t l e f o r ~ e d w a t e r -1
aFeigbinglwttle-1
v~floalr.(ioooan3)- 1
Burettb stand 1 -
Solutions Provided
PrGidures f& the prepation of these solutions are given for the sake of
informatioa These dutions would be prepared for you by the counsellor.

Preparation of solution of cu2+ ions from Cu wire

-me clean copper m e a taken. If tamidmi, it is cleaned first with fine emery cloth,
a rinsed with dilute sulphuric acid, washed well with water, and dried before
w-g. An mount of 1.5 g of the wire is weighed and placed in a 250 an3conical
flask. Then 5-10 an3of 6Mnitric acid is added. If the reaction is slow to start, a few
drops of concentrated nitric acid are added. If the readion goes too fast, a small
watch glaas is put over the top of the flaskto catch the spray. The copper wire is
dissolved by warming the solution on a water bath over a low flame. When all the
aqps bss dissolved, tbe solution is diluted with about 50 an3water and boiled
gmtly for 10 minutes to remove oxides of nitrogen. Then 1 g of urea is added and
the solution boiled f a five minutes. The solution is coded to room temperature and
neutralised with 1:3 ammonia solution adding ammonia carefully, mixing well, until a
faint permanent li@t blue precipitate of Cu(OH), appears. In case the solution
becomes deep blue on addition of ammonia, the latter is boiled off. Then glacial
&c add is added a drop at a time, until the praipitate is dissolved and ~esolution
ir dear. The sdutioll is transferred to a 1000 an3 vdumehic fl8sk, made up to the
d with distilltd water, and shaken well to get a homogeneous soluticm.
Sodium thiosulphate solution (approx. M50)
About 5g of sodium thiosulphate crystals are dissolved in I& of water that has
Quantitative Analysis-11 been recently boiled and coded. An amount of 0.2 g of sodium bicarbom& is added
as a preservative and the solution stored in a clean bottle. Sodium thiosulphate
solutions are somewhat unstable. Apart from oxygen and dissolved CO, they are
easily attacked by air-borne bacteria with the liberation of sulphur. In case any
turbidity is observed, the solution should be discarded.

Starch Sdutioo: About 150 cm3distilled water is heated to boiling in a beaker.

While this is being heated, 0.5 g to 1 g of soluble starch is stimd with about 10 cm3
of distilled water to give a paste. The paste is stirred into the boiiing water and boiled
gently for a few minutes and cooled. The solution should be almost clear. It is kept in
a stoppered bottle. (Starch sdution should be freshly prepared before use).

Potassium -1 SOklfio.: Prepared by dissolviDg 5.0 g KT in 100 cm3of distiUed

water.

Preparation of gto.dud potadtun clichnwrrte sdmtbm Prepare this sdution using

the same procedure as given in Experiment 6,

S-tion of sodipla tLIosrdphte ddibn

Pipette 20 cm3of m u m didwarnate duhn in a 250 cm3oonicPl flPlsL, add 10
an3of dilute sulphuric acid and 1g sodium hydrogen carbonate with gentle mhhg
to liberate carbon dioxide. Sodium hydrogen carbonate maintains an atmosphere of
CO, in the solution which displaces the air and prevents the oxidation of iodide from
+ + +
air. The reaction 41- 0, 4H* =F====+ 21, W , O is catalysed by light, heat and
aci&.ThensddO.Sg~\~1iodi&or10cm~of5%K1sdutioa,swk.larvathe
flask with a watch glpss and allow the solution to stand for 5 minutes in a dark place.
Titrate against sodium thiosulphate solution from the burette until a light pale ydow
dourofiodinerrppeus.lbaradd2cm3sturhduhandcoatiauetoti~ until
~ be green doured becaw of the
titrant. 'Ihe h a l s d u t i o ~will d
the blue colour of starch-iodine complex disappears on addition of a drop

chromium (III) ions. Record the burette readings before and afta the titration i~
obsen&on Table I. Repeat the same exercise to get at least two amaxdnrrt
readings.

lltmtloa of Copper (11) souiom

M t m standardising sodium thi-te solution, you can titrate fhesolution
containing Cu2+ioas. For this,pipette out 20 cm3aliquot in@ a 250 cm3conical
flask, and add 0.5 g solid +um iodide or 104fm3of 5% KI solution, swirl it to
dissolve; then titrate with the standat- miibm t h i d p b a t e which is taken in a
burette.Whenthebro~11cdourofiodinebecomespbydiou,add2cm~of~
starch solution. The colour of the sdution at this s w e is dkep blue. Swirl the flask
forabout 15secondsandwInpletethetitrrtionoddiDgsodiumthi~soluh
dropwise. During the titEation, as Cul is f d , it absorbs I; on the surface, as a
result the reaction of 1, with N%!&O3titrant is very slow. 'Iheref&e, very dose to the
: end point, when the d o u r is very light blue, add 1g potassium thiocyanate, KSCN.
~teaddedatthisstagcrerdswilthCulandfams~u~~~displacia%iodin
from the surface, making it available for the reaction.

However, if thiocyanate is added earlier during the titration, it will be slowly oxidiscd
to sulphate by iodine. At the end point, the blue d o u r of the solution d q q x a m
and the pmzipitote eppears white, or slightly grey, vheo allowed to settle. A ~ w
standing for a couple of minutes at the end point, the precipitate should bacame pure
white. Rcxord the burette readings in ohmation Tabk 11. Repeat the srme nadae
to get at least two conco~dnrrtreadings.

SAQ 3
During iodometric titrations, starch is added only towards the end of the titmion.
WY?
 Estimation of Copper
SAQ 4
Why is sodium hydrogen carbonate or sodium bicarbonate added in the
standardisation of sodium thiosulphate using potassium dichromate as titrand?

.....................................................................................................................................................
5.3.4 Observations
Mass of the weighmg bottle =
+
Mass of bottle potassium dichromate crystals =
Mass of the bottle (after transferring K2Cr207) =
Mass of potassium dichromate transferred I

Molar mass of potassium dichromate

Volume of K,Cr207 prepared (
Molarity of K2Cr20, solution
mX4
mol dm-3
Molar mass

........... mol dm-3

Observation Table I
Potassium diehumate solution vs. sodium tliosulpl.te solution

Volume of K,Cr,O, Burette reading Volume of Na2S20J

solution in cm' solution in c d
Initial Final (Final - Initial)
--
20
20
3 20
1
Observation Table U
Sodium thiosulphate solution vs. copper (11) solntioo prepared frola copper wire

Volume of Copper (11) Burette reading Volume of Na,S,O,

solution in cm'
Initial (Final - Initial)

5.3.5 Calculations
Estimation of the strength of sodium thiosulphate sMution
Molarity of K2Cr20, solution = MI = ........... mol dm-3
Volume of K2Cr,0, solution = V, = 20 ~ r n ~
Volume of Na2S203solution used = -=I, ...-cm3 ..I
(From Table I)
Molarity of Na2S203solution M 2 = ?
Using the molarity equation,
MI
Vl =PM,Vz
Molarity of Na,S,03 solution = M2 = -
6~ v,
- ........... mol dm-3

Molarity of Na2S203solution = M, = M2 ...........mol dm-3

Volume of Na2S203solution used = V, = ........... cm3
-
&timation of the strength of Copper (II) solution prepared from copper wire

Volume of copper (II) solution taken -- V, = 20 cm3

 Quantitative Analysis-11
- -
Molarity of copper (11) solution M4 ?
Using moLarity equation M4V4 M3V3,
- -
Molarity of copper (II) sdution M, -
M3 v3
v4

-
...........mol dm-3
-
Detemhtion of the amount of copper present in copper wire
Mass of copper present in 1dm3of the solution Molarity of the sdution X atomic
prepared from copper wire mass of copper

5.3.6 Result
-
The percentage of copper in the given copper wire ...O/O. You can compare the
above value with the actual one which you can get from your counsellor.

&dorimetry is based on the measurement of the intensity of colour to find the

amumtration of a given sdution. Intensity of the cdlour depends on the
concentration of the species which may be ions, molecules or a complex causing it.
The species to be determined may p<wsess an intrinsic capacity to impart colour to
the solution or it may give a distinct d o u r on being complexed with a suitable
reagent.

When light of an appro-te wavelength is passed through a coloured solution

amtahed in a cell, a M o n of the light is absorbed depending on the concentration
of the absorbing species and the thickness of the absorbkg medium, and the rest of
the light is transmitted.Though some light is reflected back from the solution, its
amount is neghgibiy small and is eliminated by using a control. For all practical
wPo= we may say,
4-Z+Z

---
where,
4 Intensity of incident light
I; Intensity of light absorbed
I; Intensity of transmitted light.
The relationship between the intensity of incident radiation and that of the
transmitted one is best given by Lambed's and Beer's laws which correlate I, with
dl
the thickness and cuucetltration of the medium, respectively. Let us understand these
m i i i w s first.
v
Lambed's Law
According to this law, when a light beam passes through a medium/solution, equal
fractions of the incident light are absorbed by layers of equal thickness or we may say
that the differential decrease in intensity with thickness of the absorbing medium is
proportional to the intensity of the incident light. Mathematically,
dl
-
- d l =kI
where,
-
k proportionality constant
I -- thickness
Rearranging, we get,
-- dl
1 =kdl
26

-
A
 Integrating and taking the condition that, when 1= 0. Estimation of Copper
I = L, we get,

log,, !o is called 'absorbance' while K(k/2.303)is referred to as the absorption

I

Beer's law
This law states that the intensity of a beam of light decreases exponentially as the
concentration of the medium decreases arithmetically. We may say that the
differential decrease in the intensity of light as a function of concentration is directly
proportional to the intensity of the incident light.
, -dl
k ' I
dc
Rearranging, we get,
-dl
I
- kdc
-

Integrating and putting the condition that when c

log, I/ lo -- -kc
- 0, I = lo,we get,

or log,, Io/I=-- kc '

a
2.303
log IJI= A, i.e., absorbance
K = absorption coefficient
AS you K-Q

5.4.1 Beer-Lambert Law

The two laws explained above are combined to give the commonly known
Beer-Lambert law which states that the fraction of light absorbed by a given
absorbing medium is directly proportional to the thickness of the medium and the
concentration of the absorbing species. Solving the mathematical expression similar
to the one in Lambert's law and Beer's law, we get,

where,
I = thickness of the medium
c -- concentration in mol dm-3
E = molar absorption coefficient
E the m o b absorption coefficient is the absorbance of a solution having unit
concentration, lM, placed in a cell of unit thickness, 1 cm. Absorbance is also called
f
epticrb density (OD).
According to Eq.5.16, the absorbance or OD of a solution in a container of fixed
path length is directly proportional to the concentration. A plot between absorbance
and concentration is expected to be linear and a solution showing such a behaviour is
said to obey Beer-Lambert law. Dilute solutions obey the law over a considerable
concentnition range, the upper limit varies from system to system. At higher
amcentnitions discrepancies are found which are attributed to the changes occurring
in the colbured solute, which may undergo association at higher concentration. It is,
therefore, advisable to prepare a calibration curve using a series of standards of
h o r n concentration.
auantitative Analysis-I1 There are a number of instruments in which a colorimetric determination can be
made. We will make use of a simple instrument called colorimeter. The details of the
instrument and the instructions for its use are discussed in the instruction manual.
Further, the use of the instrument would also be explained by your counsellor. The
basic principle on which the instrument is based is briefly given here. Before going
over to that try the following SAQ.
SAQ 5
Tick V in front of the right statements and put X in front of the wrong statements
given below :
i) Transmittance of a sample increases with a decrease in absorbancb
ii) Absorbance of a sample decreases with an increase in its concenti.at$on.
iii) Absorbance of a sample is independent of its length.
iv) An air bubble in the sample will not affect the value of absorbance.

5.4.2 Principle of Colorimeter

Generally one determines the intensity of a given colour by the use of one's eyes, i.e.
we have a visual estimate of the colour. We can compare two colours and within the
limits of human error we may differentiate between the deeper aqd the lighter colour.
But it is difficult to quantify colour VISUALLY. For this we need the help of a
measuring device. A photoelectric colorimeter is such a device. This, too, gives an
indirect estimate. It does not measure the colour, rather it measures the amount of
light which comes out after passing through the solution. Knowing the initial intensity
of light, we can work out the amount of light absorbed.

A colorimeter consists essentially of a light source, a cell/cuvette for holding the

solution, a photoelectric cell to capture the radiation transmitted by the solution and
a measuring device to detect the response of a photocell.

A schematic diagram is given in Fig. 5.1. There are three light emitting diodes
(LEDs) in the colorimeter which you are going to use. These emit light of different
colours. You would be using one of them depending on the colour of absorbing
medium. The light from the source is made to pass through a slit so that we get a thin
ray, which falls on the cell containing the solution. Some of the light is absorbed and
the rest is transmitted. The transmitted light falls on the photocell where a current is
generated, whose magnitude is proportional to the intensity of the light falling on it.
This current signal is suitably amplified and then measured by the help of an
ammeter. The deflection on the meter is proportion&io the light intensity. The
intensity of incident light is measured by putting only distilled water in the cuvette,
when no light is absorbed and the whole of it falls on the photocell. In case the
solution is made in a solvent other than water, the reference sample taken as the pure
solvent. The difference of the two readings gives the amount of light absorbed.

Light i
7
%
- Cuvette L Detector
F 7
source containing or
solution , photocell

Fig.S.1 : Schematic diagram of a yolorimeter

-
, -

5.5 EXPERIMENT 8: DETERMINATION OF

PERCENTAGE OF COPPER IN THE GWEN
SOLUTION USING COLORIMETRIC METHOD

l n t h e previous experiment, you estimated copper using an indicator titrimetric

metrod. In this experiment, you will use the instrument, colorimeter for the same.
Colorimetric determinations are possible only when the colours of the solutions are !:5iimation of Coppe
not too intense. Extremely dilute solutions can be used for such determinations when
volumetric methods do not work. These methods are also used widely because of
their high speed. You will know the advantages of colorimetry when you do this
experiment. You have already read the Beer-Lambert law on which colorimetric
determinations are based. Now you will read and learn the principle on which this
experiment is based, and about the instrument and its calibration, procedural details
and plotting a calibration curve in the following subsections.

5.5.1 Principle
The colorimetric determination of copper in a given solution is based on a simple
principle. As you know, the blue colour of copper salts is due to hydrated Cu2+ions.
The intensity of the colour can be used as a measure of concentration of Cu2+ions in
the solution. Here you will prepare a number of solutions containing known but
variable amounts of Cu2+ions and measure their absorbance in the colorimeter to
make a concentration-absorbance calibration curve. The concentration of the
unknown solution is determined by the help of this calibration curve.

The colorimeter on which you will perform your experiment is shown in Fig. 5.2.

Fig. 5.2 : Cdorimeter

..
Dewnptm d Colrtrds
Power Sllitch: 'Ihis is a SPST toggle switch, located on the back panel of the
instnunent, which turns the instrument OFF/ON. When ON, the LED on the front
panel will glow if the instrument is plugged to a 220 V supply.
Colov Sdsdor: 'Ibis is a 2 pole 3 way rotary selector switch used to switch on a
particular LED in the cell holder and also to bring into the amplifier circuit a
particular parallel resistor. It has three positions R, Y, and G signifying red, yellow
and green cdours.
Set Zao:This knob is used to set the meter reading to zero when the reference
dution is introduced in the ceU holder.
!3cdMQ: 'Ihis knob is used to adjust the meter sensitivity.
Six pi. pkg: 'Ibis is located at the right hand side of the back panel. The six pin
s g k e t from the cell hdder is inserted in this plug. (Wires fromthe LEDs and the
W R are connected to the socket).

Before using the colorimeter, you will have to calibrate it by the procedure given
below and plotting a graph to check the linearity.
 Quantitative A ~ ~ I Y S ~ S - I I 5.5.3 Calibration of Calorimeter for Colorimetric Measurement
Requirements
Apparatus Chemicals
Colorirneter with cuvettes CuSO, . 5H,O
Burette -1
Test tubes - 10
Test tube stand -1
Camtioa: DO not use a plastic '
Take a clear dry cuvette and fill it w ~ t hdistilled water or the reference solution.
cuvette for organic solvents like Note that the cuvette has two plane sides and two striated sides. Mark one of the
chloroform, acetone, etc.
plane sides with a pen and insert the cuvette in the cell holder with the marked
side facing the LEDs inside the holder.
4 Always iosert the cuvette the same way. Close the lid of the cell holder. *
, Use the Set Zero knob to adjust the meter reading to zero.
Remove the cuvette, pour off the reference solution, rinse and dry it.
Prepare 100 cm3 of 8% copper sulphate stock solution. Fill the cukette with the
stock solution. Insert the cuvette in the cell holder in the same orientation as in
Step 1. Close the lid of the cell holder.
Set the Selector on R. (A copper sulphate solution has an absorption maximum in
the red region. For an unknown solution, choose the LED which gives the highest
meter reading, i.e. the largest absorbance.) Use the Sensitivity knob to adjust the
meter reading near to the end of the scale (say 0.9).
Repeat Steps 1 and 2 for setting the zero with distilled water (or the reference
solution.)

Linearity Check
Take ten clean, dry test tubes and add 10.0 cm3, 9.0 cm3, 8.0 cm3, 7.0 cm3, 6.0 cm3,
5.0 cm3, 4.0 cm3, 3.0 cm3, 2.0 cm3 and 1.0 cm3 of the CuSO, 5 H 2 0 stock
solution in them respectively. Dilute each with distilled water to make 10.0 cm3 of
total volume.
Take the same cuvette as used for calibration. Measure the meter reading, which is
proportional to absorbance, for each of the solutions making sure that the cuvette
is rinsed properly before pouring the solution. Also make sure that the set zero
and sensitivity knobs are not disturbed throughout this set of measurements.
Plot the meter readings against the volume of stock solution taken in each of the
test tubes. A linear graph is expected as CuSO, solution is known to obey the
Beer-Lambert law in this concentration range. (A linear graph also shows that the
vlaue of parallel resistor for red LED is correct.)

1 5.5.3 Requirements
Apparatns
Colorimeter 1-
Volumetric flask (100 cm3) - 1
Test tubes - 3 5
Test tube stand - 1
Measuring cylinder - 1
Beaker - 2
Burette (XI cm3) - 1
Burette stand - 1

Solutions Provided
Cu2+ion sdutioa,prepared from copper wire using the same procedure as in the
iodometry experiment. However, here the mass of Cu wire taken is 1.7 g, and the
volume of solution prepared is 100 cm3.
Stock solution of copper nitrate (looh m/v), prepared by dissolving 10 g of
Cu(N03), . 3H,O in water and making the volume upto 100 cm3.

5.5.5 Procedure
Before starting the experiment you will have to prepare copper nitrate solution of
varying concentrations as you did for the linearity check of the instrument.

For this purpose take six test tubes and label them 1 to 6. Put Cu (NO,), . 3 H 2 0
stock solution and water in the marked test tubes with the help of a burette as given
in the following table :
S. No. Volume of C0(NO3), . 3Hz0 Volume of '10 Cu(NO,), 3H,O Estimation of Copper
stock sobtion distilkd water
1 0 10 0
2 2 8 2
3 4 6 4
4 6 4 6
5 8 2 8
6 10 0 10

Thus, you will get six solutions where the concentration of Cu(NO,), 3H20 varies
from 0 - 10% as given in the table.
Before estimating Cu2+ions in an unknown solution, a calibration curve will have to
be plotted between the concentration and the meter response in the instrument. For
this, clean the cuvette thoroughly and fill it with solution no. 1, after rinsing it with
the same solution. Place the cuvette in the cuvette holder in the instrument and
record the response in the meter in observation Table I. Then remove the solution
and rinse the cuvette with solution no.2, fill it with the solution and once again note
and record the meter response in the table. Repeat the same procedure with the rest
of the solutions too. Plot the calibration curve in the graph sheet.
Wash the cuvette again and fill it with the solution whose concentration has to be
determined. Place tbe cuvette in the cuvette holder and note the meter response.
Using the calibration curve, measure the concentration corresponding to this reading.
5.5.6 Obsenations
Obeenation Table I
Mda rrrp.rc M a hnction of coaeenhrtion of copper nihrte

I ..No. I .
Strength of CU(NO,)~ 3 H 2 0 in
Oh m/v
I Meter Response
I

--- - -

5.5.7 Calibration Curve

A sample reading, Table 11, and a calibration curve, reproduced from the readings is
shown in the graph given in Fig. 5.3 for your guidatice. See it carefully, it will help
you to plot the one with the readings you have noted.

Table U
Meter r c s p u c ur a f m c t b d cowentptloe of copper snlphnte ,

-. % o f CuSc- 5H,O Solul~on

Fir T 3 : Samale c a l i h t i o n curve drawn from the readings given in Table II

Quantitative Analysis-I1 Now plot the observations recorded in Table I in a graph.

5.5.8 Calculations
From the graph, the of copper nitrate solution is x0/0 = ...........1'0
63.5 =,z g of CU/IOO cm3
187.5
, mass of copper wire taken = 1.7 g/100 crn3-

5.5.9 Result
The percentage of copper in the given copper solution = ...O/O. You can compare
the above value with the correct value which you can get from your counsellor.

In this unit you have used two methods for the determination of percentage of
copper in a given solution. As you know, one of these methods is titrimetric indicator
.method, the other colorimetric which is an instrumental method. You can very well
compare the two methods after having used them. The comparison can be made in
terms of:
a) accuracy : Which method is more accurate? Generally the instrumental methods
are more accurate because of the very obvious errors which we can
make in titrimetric methods, e.g., errors of distinguishing a colour
change and thus the end point, etc.
b) facility : Instrumental method is more facile.
c) time : This you can judge yourself and we are sure-that you will find that the
instrumental-method has taken lesser time.
d) cost : For this particular experiment, KI is so expensive that one will like to
avoid its use. Instead of this, you are using a low cost instrument so in
terms of cost, again, instrumental method is supposed to be better.

You can discuss these experiments in the light of above points with your counsellor.

1) i) I, is almost insoluble in water.

ii) I, is volatile in nature and is lost from an open container in a short period.
It requires standardisation every few days.

3) The iodine-starch complex is only slightly dissociated and a diffuse end point will
result if large amount of iodine were absorbed on starch.

4) Sodium bicarbonate produces CO, in a solution containing KI, K2Cr20, and

acid and displaces the air present in it. Air present in the solution, otherwise,
would oxidise iodidz to iodine and cause an error in the titration.

5) i) . d ii) X iii) X iv) d