NAME: Dimitri Siewnarine
CLASS: 6N1
LAB #: 8
DATE: 11/01/2024
TITLE: REDOX TITRATION
AIM: To determine the value of n in FeSO4.nH2O by titrating against KMnO4(aq).
APPARATUS/ MATERIALS:
3 conical flasks, burette, 2 beakers, pipette, funnel, measuring cylinder, retort stand and clamp,
pipette filler, 0.02M KMnO4(aq), hydrated iron (II) sulphate solution of concentration 28 g dm-3,
1M H2SO4(aq), distilled water
METHOD:
1. The burette, measuring cylinder and the pipette were rinsed. The burette was conditioned
with KMnO4 and the pipette was conditioned with hydrated iron (II) sulphate.
2. The burette was then clamped onto the retort stand and filled over the “0” mark with the
KMnO4, using the funnel. Afterwards the burette was zeroed.
3. Using the pipette filler, 25cm3 of hydrated iron (II) sulphate was filled into the pipette
and transferred to a conical flask.
4. Using the measuring cylinder, 25cm3 of H2SO4 was measured and transferred into the
conical flask with the hydrated iron (II) sulphate.
5. The conical flask was placed under the burette, and the burette’s tap was gradually
opened, allowing only a small amount of KMnO4 to drip into the flask at a time, while the
conical flask was simultaneously swirled.
6. Step 5 was executed until a permanent colour change from colourless to pink was
observed.
7. The burette reading was recorded and the volume of KMnO4 used was calculated.
8. The procedure from steps 2 to 7 was repeated until concordant values of the volume of
KMnO4 used were achieved.
RESULTS:
TABLE SHOWING THE RESULTS OF TITRTAING 28 g dm-3 HYDRATED IRON (II)
SULPHATE AND A 1M SOLUTION OF SULPHURIC ACID AGAINST 0.02M POTASSIUM
MANGANATE (VII)
BURETTE READINGS/ cm3
ROUGH
1
2
3
Final Volume
Initial Volume
Volume of KMnO4 Used
25.60
0
25.60
25.30
0
25.30
25.10
0
25.10
25.10
0
25.10
CALCULATIONS:
Concordant Results: 25.10 cm3, 25.10 cm3
Average Titre Volume:
(25.10+25.10)
2
= 25.10 cm3
Equation: MnO4-(aq) + 5Fe2+(aq) + 8H+(aq)  Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
Concentration of KMnO4: 0.02 mol dm-3
Volume of KMnO4 used: 25.10 cm3
Volume of KMnO4 used (dm3): 25.10 ÷ 1000 = 0.0251 dm3
Number of moles of KMnO4 used: Concentration x Volume
= 0.02 x 0.0251 = 0.000502 moles
Therefore the reaction contained 0.000502 moles of MnO4- ions.
Since the mole ration between MnO4- and Fe2+ is a 1:5 ratio, 1 mole of MnO4 - will react with 5
moles of Fe2+.
CALCULATIONS:
Moles of Fe2+ = 0.000502 x 5 = 0.00251 moles
Therefore there are 0.00251 moles of FeSO4.
Concentration of FeSO4.nH2O: 28 g dm-3
Volume of FeSO4.nH2O used: 25 cm3
Volume of FeSO4.nH2O used (dm3): 25 ÷ 1000 = 0.025 dm3
Mass of FeSO4.nH2O = Concentration x Volume
= 28 x 0.025 = 0.7g
FeSO4 molar mass: 152 g mol-1
Mass of 0.00251 moles of FeSO4: Molar Mass x Number of Moles
= 152 x 0.00251 = 0.38152g
Mass of FeSO4.nH2O = Mass of FeSO4 + Mass of nH2O
0.7g
=
0.38152g
+ Mass of nH2O
Mass of nH2O = 0.7 – 0.38152 = 0.31848
H2O molar mass: 18 g mol-1
Number of moles of H2O = Mass ÷ Molar Mass
= 0.31848 ÷ 18 = 0.01769 moles
0.00251 moles of FeSO4 is hydrated by 0.01769 moles of H2O.
Therefore 1 mole of FeSO4 is hydrated by 0.01769 ÷ 0.00251 = 7.049 (≈7) moles of H2O
(FeSO4.7H2O).
DISCUSSION:
A redox titration is an experimental procedure used to determine the concentration of a given
substance by causing a redox reaction between the titrant and the analyte. It includes both
reduction and oxidation. Reduction can come about from either the addition of hydrogen, the
removal of oxygen, the acceptance of electrons or a reduction in oxidation state. Oxidation
comes about from either the removal of hydrogen, the addition of oxygen, the donation of
electrons or an increase in oxidation state.
Oxidation Half Equation: MnO4 -1 + 5 e- + 8H+ = 4H2O + Mn2+
Reduction Half Equation: Fe2+ - e- = Fe3+
In this experiment the Fe+ ion was the reducing agent ad it was oxidized by releasing one
electron, producing an iron (III) ion and its oxidation number increased from 2+ to 3+. The
MnO4- ion was the oxidizing agent as it was reduced by accepting 5 electrons to get an oxidation
number of 2+ from previously being +7. This reaction was also self-indicating and required no
indicator to determine the end point as a permanent colour change from colourless to pink was
observed.
PRECAUTIONS:

It was ensured that all apparatus was properly washed and conditioned.

It was ensured that a while tile was used to observe the colour change for accuracy.

It was ensured that the burette readings were taken from the top of the meniscus.

A dropper was used when measuring the H2SO4 with the measuring cylinder for
accuracy.
SOURCES OF ERROR:

The solution in the burette (KMnO4) could have dripped onto the walls of the conical
flask instead of directly into hydrated iron (II) sulphate, causing a higher and inaccurate
titre value.

The colour change of the reaction could have been misinterpreted and too little or too
much KMnO4 could have been used.
REFLECTION:
During this experiment, I learnt the proper procedure to carry out a redox titration and improved
on my precision. Redox titrations are used in metallurgical processes for extracting metals from
ores and combustion of fuel and in the food industry to determine the concentration of different
substances in food.
CONCLUSION: The value of n in FeSO4.nH2O was found to be 7 by titrating it against
KMnO4(aq), therefore its formula will be FeSO4.7H2O.