Question

Can oxidation happen without reduction in a chemical reaction?

Answer 1

Nope. Oxidation occurs because an agent that causes the oxidation---the oxidizing agent---must itself get reduced.

HALF-REACTIONS

You may be confused as to what it means to have a half-reaction; in those, one atom or compound gets reduced or oxidized, but nothing else gets oxidized or reduced.

In a half-reaction, there is a theoretical transfer of electrons out from or into a compound or atom, changing its oxidation state. For example:

`"Fe"(s) -> "Fe"^(2+)(aq) + 2e^(-)`

Here, two electrons are removed from iron solid, hence the production of two lone `e^(-)`'s and an iron(II) cation of a more positive (less negative) charge. We've written out the reaction for the oxidation of iron.

The charges add up to neutrality on either side; `0 = 2 + (-2)`.

The thing is, where do the electrons go? They don't just escape the iron solid and go out into the world; the probability of an electron existing outside of orbitals is `0%`. It has to be donated to another compound or atom.

"RELEASED" ELECTRONS MUST GO TO AN ELECTRON ACCEPTOR

We say that iron must be "active" enough to donate this electron; so a less "active" metal that can accept this electron is silver, according to the activity series, as `"Ag"^(+)`.

So, we can write another theoretical half-reaction:

`"Ag"^(+)(aq) + e^(-) -> "Ag"(s)`

But one equivalent of silver(I) only wants one electron, so we need two equivalents of silver(I) to accept iron's two electrons.

THE FULL REDOX (REDUCTION-OXIDATION) REACTION

Now we can sum up the two reactions to get the full redox reaction:

`"Fe"(s) -> "Fe"^(2+)(aq) + cancel(2e^(-))`
`2"Ag"^(+)(aq) + cancel(2e^(-)) -> 2"Ag"(s)`
`"----------------------------------------"`
`color(blue)("Fe"(s) + 2"Ag"^(+)(aq) -> "Fe"^(2+)(aq) + 2"Ag"(s))`

Because the two electrons were transferred from iron to silver, no electrons are "out in limbo" (as they were in the theoretical half-reactions), and we can cancel out the donated electrons on the products and reactants sides.

In this full redox reaction:

• Because silver decreased in positive charge (or, increased in negative charge), silver got reduced. `"Ag"^(+1) -> "Ag"^(0)`
• Because iron increased in positive charge (or, decreased in negative charge), iron got oxidized. `"Fe"^(0) -> "Fe"^(+2)`

As for the roles:

• Iron is the reducing agent. It donated the electrons to silver, reducing it.
• Silver is the oxidizing agent. Silver(I) cation accepted the electrons from iron, oxidizing it.

---

As an aside, because iron is more active than silver, you can't reverse this reaction to get:

`"Fe"^(2+)(aq) + 2"Ag"(s) color(red)(cancel(color(black)(->))) "Fe"(s) + 2"Ag"^(+)(aq)`

and it doesn't occur.