# How do you balance this redox reaction?

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#As → H_2AsO_4^-1+ AsH#

What I got is:

#4H_2O + 6As → H_2AsO_4^-1+5 AsH + H^+#

What I got is:

##### 1 Answer

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#### Explanation

#### Explanation:

#### Answer:

Yes, it is accurate assuming this reaction happens in an acidic solution

#### Explanation:

Let's first identify the oxidation states of all the elements in the equation as it is...

Ok so Arsenic is apparently being reduced and oxidized

We will now write half reactions:

When we cancel out the 5 electrons on each side, we are left with:

So now if we rewrite the given equation with those coefficients:

Not only are the charges unbalanced as in the reactant has no charge and the products have an overall -1 charge, the equation lacks oxygen and hydrogen on the reactant side, so let's add H_2O on the reactant side to barely meet the oxygen requirement in the arsenate.

Your equation seems correct.

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