Redox Titrations Flashcards
Describe the procedure To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
1- filling the burette with the potassium managanate vii solution.
fill KMnO4 in burette, fill above 0 cm mark first reading @ eyelevel.
UNTIL TOP OF MENISCUS ON 0CM MARK
Making the hydrated ammonium iron (II) sulfate up into a standard solution
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
Dilute sulfuric acid (H2SO4) is added to the deionised water when making up the solution of hydrated amonium sulfate.
describe the process of carring out the titration
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
BEFORE TITRATION BEGINS, EXCESS DILUTE SULFURIC ACID H2SO4 IS ADDED TO THE HYDRATED AMMONIUM IRON II SULFATE IN THE CONICAL FLASK
the purple drops will decolourise during the titration, at the end point a permanent pink colour is observed in the conical flask.
what is the suitable indicator for this reaction?
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
no indicator required, potassium manganate acts as its own indicator (self indicating)
describe the colour chage observed in this reaction
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
during the titration- the potassium mananate is added to the acidified hydrated ammonium iron II sulfate, its purple colour decolourises
at end point- a permanent pink colour is observed in the conical flask
explain the colour change that occours during and at end point
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
during titration- Mn 7+ ions cause the purple colour in KMnO4.
Fe2+ ions are oxidised to Fe3+ ions : the Mn7+ ions are reduced to Mn 2+ ions
as Mn2+ ions are formed the purple colour decolourises
at end point- There is no more Fe2+ ions to reduce the MN7+ ions The last drop of KmnO4 is added and leaves a permanent pink colour.
Explain why a standard solution of potassium manganate (VII) CANNOT be directly made up i.e. why must potassium manganate (VII) be standardised by titration?
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
Potassium manganate VII is not a primary standard,it cannot be obtained in a pure state
Explain why a standard solution can be directly made up from hydrated ammonium iron (II) sulfate
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
Hydrated ammonium iron ii sulfate is a primary standard, it can be obtained in a pure form that is stable in air and can be dissolved in water to make up a solution of accurately known conc.
Why is ammonium iron (II) sulfate used as a primary standard instead of iron (II) sulfate?
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
Iron ii sulfate is easily oxidised by oxgen in the air, not a suitable primary standard
Ammonium II sulfate can be obtained in a pure state that is stable in air and has a high mm and can be dissolved directly in water to make a solution of known conc
It is noted during the titration that the first few drops of KMnO4 are slow to decolourise, but
subsequent drops decolourise rapidly. Explain
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
Mn2+ ions act as a auto catalyst, as more are formed they speed up thye reaction and cause the purple colour to decolourise more rapidly.- auto catalysis
At what two occasions is dilute sulfuric acid required to be added and explain why it is required on each occasion.
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
1- sulfuric acid is added when making up and dissolving the hydrated ammonium iron II sullfate into solution- Prevent Fe2+ ions being oxidised to Fe3+ ions by oxygen in the air
2- excess sulfuric is added to the hydrated ammonium II sulfate solution in the conical flask before titrating against potassium manganate VII - only in an acidic enviorment are MN 7+ ions reduced to Mn2+ ions and NOT Mn 4+ ions
if a brown precipitate forms during the titrations, what conclusion can be drawn?
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
not enough dilute sulfuric acid has been added to the hydrated ammonium iron II sulfate before the titration_ Mn+7 ions have ONLY been reduced to Mn+4 ions and an insoluble brown percipitate has been formed
How is the potassium permanganate read in the burette during the titrations? And why?
To use a standard solution of hydrated ammonium iron (II)
sulfate to standardise a solution of potassium manganate (VII) by titration
The potassium permanganate is read from the top of the meniscus eye level .
The intense purple colour of potassium manganate (VII) makes it difficult to read from the bottom of the meniscus.
(Also for this reason, dilute solutions of potassium manganate (VII) are used)
Describe the procedure for dertiming the amount of iron in an iron tablet
1 fill burette with standard solution of potassium manganate VII. Read from top of meniscus at eye level
2. making the iron tables up into solution. The iron tablets need to be crushed with a pestle and mortar and dilute sulfuric acid is added to the deionied water when making up the iron tablet solution.
Describe how you carried out the titration
To dertimine the amount of iron in an iron tablet
before titration begins, excess dilute sulfuric acid is added to the iron tablet solution in the conical flask.
the purple KMnO4 decolourises and a permant pink colour forms.
what is the average amount of iron in an iron tablet
To dertimine the amount of iron in an iron tablet
25- 50%
what is the active compound in iron tables
To dertimine the amount of iron in an iron tablet
Iron (II) sulfate
why are iron tablets somtimes medically percibed?
To dertimine the amount of iron in an iron tablet
Iron is part of haemoglobin in our red blood cells which carry oxugen and provide energy. A person suffering from anameia may be perscribes iron tablets for tirednedd and fatigue
how was it possible to have a standard solution of KMnO4 to use in this titration despite the fact its not a promary standard
To dertimine the amount of iron in an iron tablet
KMnO4 perviously titred against a standard sln of Hydrated ammonium iron II sulfate
At what two occations is diltute sulfuric acid added and why is it required
To dertimine the amount of iron in an iron tablet
1- sulfuric acid is added when making up and dissolving the hydrated ammonium iron II sullfate into solution- Prevent Fe2+ ions being oxidised to Fe3+ ions by oxygen in the air
2- excess sulfuric is added to the hydrated ammonium II sulfate solution in the conical flask before titrating against potassium manganate VII - only in an acidic enviorment are MN 7+ ions reduced to Mn2+ ions and NOT Mn 4+ ions
Why Is it important to use the previously standardised KMnO4 immediately to dertimine the % Fe in Iron tablets?
To dertimine the amount of iron in an iron tablet
KMnO4 is unstable, it decoposes in the presence of light and heat
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
Procedure: making the dosium thiosulfate up into a solution if required.
2: filling te burette with the sodium thisulfate solution.
3. making a standard soultion of iodine using potassium permanganate - irodine is not a primary standard.
A set volume of perviously standardised solution of potassium maganate is reacted with EXCESS potassium iodide (KI) and EXCESS sulfuric acid (H2SO4) to form iodine.
The colour change at THIS point is purple to red-brown. (2:5) ratio.
Carry out titration, Colour change during titration. goes from red-brown colour to yellow to pale yellow. When pale yellow, add starch solution indicator. The solution in conical flask turns blue- black. At the end point all iodine is used up and solutio becomes colourless.
what is suitable indicator for this solution?
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
startch solution is added when close to the end point of a titration, when a pale yellow colour forms
why should the indicator be freshly prepared.
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
starch is biodegradable
describe the colour chage during this titration
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
red-brown to yellow to pale yellow. Now add starch - a blue black colour forms. Blue-black to colourless
explain the colour change during and at the end point of the titration.
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
During the titration the sodium thosufate is added to and reacts with the iodine, whos colour becomes less intense as its being used up in the reaction.
At the end point starch solution indicator is added and a blue black colour forms due to small amount of iodine left, as soon iodine has been conpeltely used up the blue-black colour decolourises.
Why is the stach only added when a pale-yellow colour forms in the conical flask.
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
1) waiting until pale yellow colour tells us end point is very near.the sodium thiosulfate can be added in slowly in small drops, resulting in accurate end point
2) Iodine absorbs onto starch, preventing it reactig with sodium thiosulfate, adding too early results in innacuratly large end pt
Explain why a standard solution of sodium thiosuldate cannot be directly made up, why must it be standardised by titration
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
sodium thiosulfate is not a primary standard- it is a hydrated compound and loses some of their water of crystallisation in dry air- cannot be obtained in a pure state, so a percise mass cannot be weighed out
why cant a standard solution of iodine be directly made up?
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
iodine is not a primary standard
Iodine sublimes (changes directly drom a solid to a vapour at room temperatire)
poorly soluble in water
How is a standard aqueous solution of iodine obtained? What colour change occours in conical flask as a redult of this reaction?
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
reacted with excess potassium iodide and excess sulfuric acid (it has molar ratio 2:5)
Colour change is purple to red brown.
why when making up the iodine, is excess sulfuric acid added to the conical flask?
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
Only in an acidic enviorment are Mn7+ ions fully reduced to Mn2+ ion, meaning that the I- ions will be fully oxidised to form iodine Produces the max amount of iodine
In making up iodine, why cant nitric acid or hydrocloric acid be used to provide an acidic enviorment
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
nitric acid is an oxidising agent- will interfere
Hydrocloric acid will be oxidised by potassium manganate forming Cl2
Why is excess potassium iodide added to the conical flask?
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
1- potassium iodide in excess ensure that the potassium manganate is the limiting reagent- ensuring all of the potassium mangante reacts- produces max amount of iodine
2- the iodine that is produced will react with potassium iodine in excess forming the triiodide ion I3- a soluble version iodine- keeps iodine in aqueous solution.
explain how iodine, a non-polar substance of very low water solubility , is brought into aqueous solution
Describe the procedure of using a standard solution of iodine to standardise a solution of sodium thiosulfate
reacting the acdified potassium manganate with EXCESS potassium iodine. The iodine that is produced will react with the potassium iodide in excess forming the triiodide ion I3-. keeps iodine in aqeous solution
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
1.making the sodium thiosulfate up into a standard solution (if required)
2.Filling the burette with the sodium thosulfate solution.
3. diluting the bleach
4.making a solution of iodine using the diltued bleach- a set volume of the diluted bleach containing sodium hypochloride (NaCLO) is pipetted to a conical flask and reacted excess potassium (KI) and excess sulfuric acid (H2SO4) The colour wil change from colourless to red-brown, the molar ratio means the moles of sodium hypocloride can be known.
The titration is carried out red-brown to yellow to pale yellow. Starch indicator is added, solution turns blue black, at ent point, titration, solution becomes colourless.
what is the typical w/v of sodium hypochlorite in bleach.
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
3-10%
how was it possible to have a standard solution of sodium thiosulfate to use in this titration despite the fact it is not a primary standard
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
the sodium thisolfate was previously standardised by titration it against a standard solution of iodine
why is the bleach diltuted before use in this titration?
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
the orginal bleach was too concentrated meaming the end point of the titration would take too long to occour. a large volume of sodium thiosulfate would be required.
how is an aqueous solution of iodine obtained from the bleach
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
the diluted bleach containing sodium hypochlorite is reacted with excess potasium iodide and excess sulfuric acid
when making up the iodine what colour change occours in the conical flask as a result of the reaction?
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
colourless to red-brown
why is excess sulphuric acid added to the conical flask when making the I2
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
Only in an acidic enviorment are CLO- ions in sodium hypochlorite fully reduced to Cl- ion. I- ions will be fully oxidised to form iodine (i2) produces the maxium amount of iodine
why is excess potassium iodide added to the conical flask
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
1.Potassium iodide in excess ensures that the sodiim hypochlorite is the limiting reagent - the ensures all of the sodium hypochlorite reacts- produces max iodine.
- The iodine that is produces will react with the extra potassium iodide in excess forming the triiodine ion I3- keeps iodine in aqueous solution
explain how iodine, a non-polar substance of very low water solubility, is brought until aqueous solution
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
reacting the acidified sodium hypochlorite with excess potassium iodide. The iodine that is produces will react with the potassium iodide in excess forming the triiiodide ion I3-
Explain why the use of the distilled water instead of deionised water throught this expeirment would be more likely to ensure a more acctuate result.
Describe the procedure to dertime the % w/v of sodium hyprochlorite in bleach
Deionised water has all ions removed but could still contain chlorine- chlorine is an deoxidising agent.
Distilled water is the most pure form of water
