Transcript Volumetric Analysis Oxidation

Volumetric Analysis
Oxidation- Reduction
Chapter 15
Potassium Permanganate
KMnO4
• Potassium Permanganate is an oxidising agent.
• It is a purple coloured solid.
• It is not a primary standard as it decomposes in the
presence of sunlight.
• It acts as an oxidising agent by gaining 5 electrons in
acidic solution.
MnO4- + 8H+ + 5e-
Mn2+ + 4H2O
Ammonium Iron (II) Sulphate
Primary standard for KMnO4
Primary standard for redox titration’s: Ammonium Iron (II) Sulphate.
 Stable-does not react with the gases in the air and has good solubility.
 High relative molecular mass of Ammonium Iron (II) Sulphate(392)
ensures a high degree of accuracy when weighing.
 When dissolving in water, sulphuric acid must be added to prevent it
from reacting with the water (HYDROLYSIS) and oxygen present.
Fe2+ oxidised to Fe3+.
Solutions should be made up with Distilled water as deionised water may
contain chlorine
Oxidising agent-KMnO4
Potassium Permanganate (KMnO4) is a good Oxidising Agent. It itself
is reduced(gains 5 electrons).
 Because KMnO4 is not very soluble in water a dilute solution (0.02
mol L-1) is used
 The MnO4- ion is purple and Mn2+ is colourless, so no indicator is
needed. The endpoint is a permanent pale pink colour.
 The meniscus is read from the top.
 The H+ ions are provided by H2SO4.
o HCl would be oxidised
o HNO3 is an oxidising agent.
 If sufficient H2SO4 is not added the MnO4- is reduced to MnO2 which
is a brown solid. More H2SO4 should be added.
Experiments
1. (Expt 4.5) To prepare a
standard solution of
Ammonium Iron(II)
Sulphate. And to use this
solution to standardise a
solution of KMnO4 by
titration. (Page 193 of
Book)
Questions and Answers
Solutions to student question
1. Why is ammonium iron(II) sulfate suitable as a primary standard?
Because it is stable and available in a highly pure form.
2. Why is sulfuric acid added to the iron(II) solution prior to titration? Could hydrochloric
acid or nitric acid be used instead of sulfuric acid? Explain.
Acidic conditions are necessary, because in neutral or alkaline conditions
Mn+7 is reduced only as far as Mn+4.
Hydrochloric acid is not suitable as it would react with the KMnO4, and
chlorine gas would be evolved. Nitric acid is not suitable as it is itself a very
powerful oxidising agent - the NO3- ion is readily reduced.
3. In preparing for the titration, explain (a) why the pipette and burette were rinsed with
deionised water followed by a little of the solutions they were to contain, (b) why the
conical flask was rinsed with deionised water only.
(a) Deionised water washes out any residual solutions in the burette and pipette
respectively. The second step was taken to remove any residual water, and so avoid
dilution of the solutions when they are added to the burette and pipette
respectively.
(b) Deionised water washes out any residual solution in the conical flask. If it were then
washed out with the solution it was to contain, traces of it would remain, and there
would not be a precisely known amount of the solution in the flask.
4. During the titration the sides of the conical flask were washed down with deionised
water from a wash bottle. Explain why this procedure is necessary and why it can be
carried out without affecting the result of the titration.
The washing process was carried out to ensure that all of the manganate(VII)
solution added from the burette reacted with the iron(II) solution. It did not affect
the result of the titration because only deionised water was added – no extra
reactants were introduced into the flask.
5. One of the products of this reaction acts as a catalyst for the reaction. Which product is
this? How could you demonstrate what substance is acting as the catalyst?
The reaction is catalysed by Mn2+ ions. This can be shown by taking a clean conical
flask, pipetting the Fe2+ solution into it, acidifying it and then before starting to
titrate adding some MnSO4 solution (a convenient source of Mn2+). Now the first
droplet of MnO4- added decolourises immediately as there is Mn2+ in place to act as
catalyst.
6. Why was sulfuric acid added in making up the ammonium iron(II) sulfate solution?
Iron(II) is very susceptible to air oxidation, forming iron(III), under neutral or
alkaline conditions but this oxidation is inhibited in the presence of acids. The
ammonium iron(II) sulfate solution is made up in dilute acid solution to make
it stable towards air oxidation.
Iron in Iron Tablet
2. (Expt 4.6) Using the
standardised KMnO4 to
determine the amount of
Iron in an Iron tablet.
(Page 197 of Book).
(Q 1.Leaving Cert 2003)
Questions and Answers
Solutions to student questions
1. In this experiment why is dilute sulfuric acid used rather than deionised water to dissolve
the iron tablets?
If deionised water were used, the Fe2+ in the tablets would be almost immediately
oxidised to Fe3+. The sulfuric acid prevents this occurring.
2. Why are burette readings taken from the top of the meniscus?
Because the very dark colour of the manganate(VII) solution makes the meniscus
difficult to see.
3. How is the end-point of the titration detected?
When the first permanent pale pink colour forms in the solution in the conical flask.
4. Why is a rough titration carried out?
To determine the approximate end-point. This can then be used to get accurate
results in the subsequent titrations.
5. Why is more than one titration carried out subsequently?
To reduce experimental error, by getting the mean of the accurate titres
.
6. Prior to the titration, what steps are taken to minimise error?
All glassware is washed with deionised water. The burette and pipette respectively are
rinsed with the solution they are to contain. The tap of the burette is opened briefly to
fill the part of the burette below the tap.
7. If a brown precipitate appears during the titration, what does this indicate,
and how can it be remedied?
Mn(IV) is formed, because of incomplete reduction of the Mn(VII). This should
only happen if there is insufficient sulfuric acid in the conical flask. The remedy
is to add more dilute sulfuric acid to the flask, or, preferably, to repeat the
experiment with sufficient acid present in the flask.
The reaction of Sodium Thiosulfate (Na2S2O3)
Sodium Thiosulphate (Na2S2O3) is an important reducing agent used in
chemistry.
It reacts with iodine molecules and reduces them to Iodide ions (I-) and
tetraionate ion.
I2 + 2S2O32-  S4O62- + 2IThe colour change is from the brown/red colour of iodine, I2, to a straw
yellow colour to a colourless iodide ion, I- .
Starch is used as an indicator to detect the end point. It goes from blue/ black
to colourless. The starch is added near the end point (when the solution is a
straw yellow colour) because the solution changes colour so abruptly. If
the starch is added at an early stage, the iodine present may become
strongly adsorbed on to the starch and make the titration less accurate.
Standardising Sodium Thiosulfate
Sodium thiosulphate is not a primary standard. It is standardised by
titration it against acidified potassium permanganate with an excess
of acidified potassium iodide. When potassium permanganate
reacts with potassium iodide, the potassium iodide is oxidised to
free iodine (I2).
2MnO4- + 10I- + 16H+ 2Mn2+ + 5I2 + 8H2O
I2 + 2S2O32- S4O62- + 2IThe free iodine can then be titrated against the unknown thiosulphate solution
using starch as an indicator.
The overall reactions are:
2 mol KMnO4 = 5mol I2 = 10 mol Na2S2O3
The iodine cannot be used as a primary standard because it does not dissolve
readily in water and iodine sublimes.
To determine the percentage (w/v) of
sodium hypochlorite in bleach
Many commercial bleaches are simply solutions of hypochlorite salts such as
sodium hypochlorite (NaOCl) or calcium hypochlorite (Ca(OCl)2).
Hypochlorite ion reacts with excess iodide ion in the presence of acid to
generate an iodine solution:
ClO- + 2I- + 2H+
Cl- + I2 + H2O
The liberated iodine solution can be titrated against sodium thiosulfate
solution using a freshly prepared starch solution as indicator. The titration
reaction may be represented by the equation:
I2 + 2S2O322I- + S4O62Starch indicator is added during the titration when the colour of the solution
in the conical flask fades to a pale yellow colour. The solution becomes blue
black, and the titration is continued until it goes colourless.
Experiments
1.(Expt.4.7) To prepare a
solution of sodium
thiosulfate and
standardise it by
titration
against a solution of
iodine.
(See page 202 book)
Questions and Answers
Solutions to student questions
1. Why is hydrated sodium thiosulfate not suitable as a primary standard?
It loses water of crystallisation readily, and it is not stable.
2. Why are iodine solutions made up using potassium iodide solution?
Iodine is sparingly soluble in water. However it reacts with iodide forming I3- ions,
which are very soluble. In this way the iodine is kept in solution.
3. Why does starch solution have to be freshly prepared?
It deteriorates quickly on standing.
4. Which of the three pieces of titration apparatus, the pipette, the burette or the conical
flask, should not be rinsed with the solution it is to contain? Give a reason for your
answer.
The conical flask should not be rinsed with the solution it is to contain. If it were
washed out with the solution it was to contain, then traces of the solution would
remain, and there would not be a precisely known amount of the solution in the
flask.
5. Why is starch indicator added close to the end-point?
To give a sharp end-point, while avoiding the formation of excess starch-iodine
complex, which would be difficult to decompose.
6. What happens at the end-point?
The colour changes from blue-black to colourless.
2. (Expt.4.8) To determine
the percentage (w/v) of
sodium hypochlorite in
household bleach.
(See Page 204 Book)
Questions and Answers
Solutions to student questions
1. Give one reason why, in making up the solution of diluted bleach, a volumetric flask is
preferable to a graduated cylinder.
A volumetric flask is quite an accurate measuring vessel, whereas a graduated
cylinder is not.
2. A burette, a pipette and a conical flask were used in the titration. State the correct washing
procedures for each of these items before starting the titration.
Rinse the burette, pipette and conical flask respectively with deionised water. Rinse
the burette with sodium thiosulfate solution, and rinse the pipette with diluted
bleach solution.
3. Why could you not use hydrochloric acid when acidifying the bleach?
Hydrochloric acid is not suitable, as it will react with hypochlorite to liberate
chlorine gas.