Chemistry of Chromium - Chemistry LibreTexts
Chemistry of Chromium - Chemistry LibreTexts
Chemistry of Chromium - Chemistry LibreTexts
+
+ H O
3
The complex ion is acting as an acid by donating a hydrogen ion to water molecules in
the solution. The water is, of course, acting as a base by accepting the hydrogen ion.
Because of the confusing presence of water from two different sources (the ligands
and the solution), it is easier to simplify this:
3 + 2 + +
Cr(H O) −
↽⇀
− Cr(H O) (OH) + H (aq)
2 6 2 5
However, if you write it like this, remember that the hydrogen ion isn't just falling off
the complex ion. It is being pulled off by a water molecule in the solution. Whenever
you write "H+(aq)" what you really mean is a hydroxonium ion, H3O+.
One of the water molecules is replaced by a sulfate ion. Notice the change in the
charge on the ion. Two of the positive charges are canceled by the presence of the two
negative charges on the sulfate ion.
Replacement of the water by chloride ions
In the presence of chloride ions (for example with chromium(III) chloride), the most
commonly observed color is green. This happens when two of the water molecules are
replaced by chloride ions to give the tetraaquadichlorochromium(III) ion -
[Cr(H2O)4Cl2]+. Once again, notice that replacing water molecules by chloride ions
changes the charge on the ion.
+ 3H O
2
But the process doesn't stop there. More hydrogen ions are removed to give ions like
[Cr(H O) (OH) ]
2 2 4
−
and [Cr(OH)
6
3 −
] . For example:
−
[Cr(H O) (OH) ](s) + 3 OH
2 3 3
3 −
⟶ [Cr(OH) ] (aq) + 3 H O
6 2
The precipitate redissolves because these ions are soluble in water. In the test-tube,
the color changes are:
+
⟶ [Cr(H O) (OH) ](s) + 3 NH (aq)
2 3 3 4
That precipitate dissolves to some extent if you add an excess of ammonia (especially
if it is concentrated). The ammonia replaces water as a ligand to give
hexaamminechromium(III) ions (this is an example of a ligand exchange reaction).
3 +
[Cr(H O) ] (aq) + 6 NH (aq)
2 6 3
3 +
⟶ [Cr(NH ) ] (aq) + 6 H O(l)
3 6 2
Apart from the carbon dioxide, there is nothing new in this reaction:
This is then oxidised by warming it with hydrogen peroxide solution. You eventually
get a bright yellow solution containing chromate(VI) ions.
−
+ 2 OH + 8 H O(l)
2
yellow chromate(VI) ion, CrO . Changing between them is easy; i f dilute sulfuric
2 −
4
acid is added to the yellow solution it turns orange. If you add sodium hydroxide
solution to the orange solution it turns yellow.
If you add extra hydrogen ions to this, the equilibrium shifts to the right, which is
consistent with Le Chatelier's Principle.
If you add hydroxide ions, these react with the hydrogen ions. The equilibrium tips to
the left to replace them.
The reason for the inverted commas around the chromium(III) ion is that this is a
simplification. The exact nature of the complex ion will depend on which acid you use
in the reduction process. This has already been discussed towards the top of the page.
H O + 3 CH CHO
2 3
If the oxidizing agent is in excess, and you do not allow the product to escape -e.g., by
heating the mixture under reflux (heating the flask with a condenser placed vertically
in the neck) - you get ethanoic acid.
2 − + 3 +
2 Cr O + 16 H + 3 CH CH OH → 4 Cr
2 7 3 2
+ 11 H O + 3 CH COOH
2 3
CH CH OH + [O] → CH CHO + H O
3 2 3 2
CH CH OH + 2 [O] → CH COOH + H O
3 2 3 2
The oxygen written in square brackets just means "oxygen from an oxidizing agent".
H O + 3 CH CHO
2 3
This ionic equation obviously does not contain the spectator ions, potassium and
sulfate. Feeding those back in gives the full equation:
K Cr O + 4 HSO + 3 CH CH OH
2 2 7 4 3 2
→ Cr (SO ) + K SO + 7 H O + 3 CH CHO
2 4 3 2 4 2 3
You will see that the chromium(III) sulfate and potassium sulfate are produced in
exactly the right proportions to make the double salt.
In the Lab
In practice
There are advantages and disadvantages in using potassium dichromate(VI).
Advantages
Potassium dichromate(VI) can be used as a primary standard. That means that it can
be made up to give a stable solution of accurately known concentration. That isn't true
of potassium manganate(VII).
Potassium dichromate(VI) can be used in the presence of chloride ions (as long as the
chloride ions aren't present in very high concentration).
Potassium manganate(VII) oxidises chloride ions to chlorine; potassium
dichromate(VI) isn't quite a strong enough oxidising agent to do this. That means that
you don't get unwanted side reactions with the potassium dichromate(VI) soution.
Disadvantage
The main disadvantage lies in the color change. Potassium manganate(VII) titrations
are self-indicating. As you run the potassium manganate(VII) solution into the
reaction, the solution becomes colorless. As soon as you add as much as one drop too
much, the solution becomes pink - and you know you have reached the end point.
Unfortunately potassium dichromate(VI) solution turns green as you run it into the
reaction, and there is no way you could possibly detect the color change when you
have one drop of excess orange solution in a strongly colored green solution.
With potassium dichromate(VI) solution you have to use a separate indicator, known
as a redox indicator. These change color in the presence of an oxidising agent.
There are several such indicators - such as diphenylamine sulfonate. This gives a
violet-blue color in the presence of excess potassium dichromate(VI) solution.
However, the color is made difficult by the strong green also present. The end point of
a potassium dichromate(VI) titration isn't as easy to see as the end point of a
potassium manganate(VII) one.
The calculation
The half-equation for the dichromate(VI) ion is:
2 − + − 3 +
Cr O + 14 H + 6e ⟶ 2 Cr + 7H O
2 7 2
3 +
+ 6 Fe + 7H O
2
You can see that the reacting proportions are 1 mole of dichromate(VI) ions to 6 moles
of iron(II) ions. Once you have established that, the titration calculation is going to be
just like any other one.
You can't rely on this as a test for chromate(VI) ions, however. It might be that you
have a solution containing an acid-base indicator which happens to have the same
color change!
This page titled Chemistry of Chromium is shared under a CC BY-NC 4.0 license and was
authored, remixed, and/or curated by Jim Clark.