14.3

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Alexandra Carpenter 1G
Posts: 32
Joined: Sat Jul 22, 2017 3:00 am

14.3

Postby Alexandra Carpenter 1G » Mon Feb 05, 2018 1:39 pm

In class we learned that if something was oxidized, that means it lost an electron, causing the oxidation number to increase. In 14.3 part a, the equation reads as: Cl2 (g) + S2O32- (aq) → Cl- (aq) + SO42-
If the chlorine ion is going from an oxidation state of 0 to -1, why is it being oxidized?

RyanS2J
Posts: 32
Joined: Thu Jul 27, 2017 3:00 am

Re: 14.3

Postby RyanS2J » Mon Feb 05, 2018 1:59 pm

Since sulfur S is going from a state of +2 in the reactants to +6 in the products, it looks to me like sulfur is the one losing the electrons and being oxidized.

RyanS2J
Posts: 32
Joined: Thu Jul 27, 2017 3:00 am

Re: 14.3

Postby RyanS2J » Mon Feb 05, 2018 3:35 pm

Also, to clarify, the question is asking for the oxidizing and reducing agents. The oxidizing agent oxidizes another element or molecule by taking the electron(s) from this other element or molecule (which then loses the electron(s)), and by taking the electron(s), the oxidizing agent itself is reduced. By similar logic, the reducing agent reduces another element or molecule and is oxidized. Therefore, since chloride gains an electron and is reduced, it is the oxidizing agent, and brings about the oxidation of sulfur by removing its electrons.

Ashley Davis 1I
Posts: 57
Joined: Fri Sep 29, 2017 7:04 am

Re: 14.3

Postby Ashley Davis 1I » Wed Feb 07, 2018 6:30 pm

Yeah, basically it's reversed when it comes to the agents: The oxidizing agent is the species that is being reduced. The reducing agent is the species that is being oxidized.


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