# Write an equation for the disproportionation of chlorate(I) ions and show the oxidati

Watch
Announcements
Thread starter 7 years ago
#1
Write an equation for the disproportionation of
chlorate(i) ions and show the
oxidation states of chlorine in the products.
Thanks
0
7 years ago
#2
(Original post by Rhubarb96)
Write an equation for the disproportionation of
chlorate(i) ions and show the
oxidation states of chlorine in the products.
Thanks
See question 1bii in

http://www.ocr.org.uk/Images/60448-q...-resources.pdf
1
7 years ago
#3
Forgive me, but did the question not seem to imply that ClO- was a reactant?
0
7 years ago
#4
(Original post by PythianLegume)
Forgive me, but did the question not seem to imply that ClO- was a reactant?
Sorry
0
7 years ago
#5
(Original post by krisshP)
Sorry
That's fine. I'm not sure how to answer it, because I don't know what the products would be. At a guess, I'd say Cl2 and ClO3, but it would be a guess.
0
7 years ago
#6
(Original post by PythianLegume)
That's fine. I'm not sure how to answer it, because I don't know what the products would be. At a guess, I'd say Cl2 and ClO3, but it would be a guess.
The chlorate(I) disproportionates (simultaneously oxidises and reduces) to chloride ions and chlorate(V) ions ...
1
7 years ago
#7
Hi Rhubarb

The easiest way to write any disproportionation is to write single electron reactions for the reduction and the oxidation and then add the two equations to each other.

Assuming it disproportionates to chloride ions and chlorate ions, you have the following half equations.

Reduction of ClO(-) into Cl(-)

0.5 ClO(-) + H(+) + e(-) --> 0.5 Cl(-) +0.5H2O

Oxidation of ClO(-) into ClO3(-)

0.25ClO(-) + 0.5 H2O --> 0.25ClO3(-) + H(+) + e(-)

Now add the two equations, and everything simplifies nicely with the hydrogen ions, electrons and waters cancelling out.

0.75 ClO(-) --> 0.5 Cl(-) + 0.25 ClO3(-)

This can also be written with integer coefficients of 3, 2 and 1 respectively.

EDIT: Just saw you asked for oxidation states too. ClO(-) has chlorine in oxidation state +1, Cl(-) has chlorine in oxidation state -1, ClO3(-) has chlorine in oxidation state +5. You can work this out as the oxidation state of oxygen in all non-peroxide compounds is (-2), and the sum of...

oxidation state of chlorine + ( (-2) * number of oxygen atoms) = overall charge on the ion.

This is why the reaction is a disproportionation, because 0.75 moles of chlorine at oxidation state 1 is simultaneously reduced to 0.5 moles of chloride ions at oxidation state -1 and oxidised to 0.25 moles of chlorate (V) ions at oxidation state +5.

Hope this helped!
Matt

Posted from TSR Mobile
0
7 years ago
#8
(Original post by Rhubarb96)
Write an equation for the disproportionation of
chlorate(i) ions and show the
oxidation states of chlorine in the products.
Thanks
3ClO- -> ClO3 - + 2Cl-
+1------> +5 + -1
0
7 years ago
#9
(Original post by Coral Reafs)
3ClO- -> ClO3 - + 2Cl-
+1------> +5 + -1
yeah nice one ...

.. you have successfully re-answered a question.

value much not
1
7 years ago
#10
(Original post by charco)
yeah nice one ...

.. you have successfully re-answered a question.

value much not
oops..didn't check entire page
0
4 years ago
#11
Brilliant, thanks! Eve of Kent
0
X

new posts
Back
to top
Latest
My Feed

### Oops, nobody has postedin the last few hours.

Why not re-start the conversation?

see more

### See more of what you like onThe Student Room

You can personalise what you see on TSR. Tell us a little about yourself to get started.

### Poll

Join the discussion

#### Are you travelling in the Uni student travel window (3-9 Dec) to go home for Christmas?

Yes (120)
28.1%
No - I have already returned home (57)
13.35%
No - I plan on travelling outside these dates (84)
19.67%
No - I'm staying at my term time address over Christmas (40)
9.37%
No - I live at home during term anyway (126)
29.51%