Is it a redox reaction if the oxidation state of the element has changed only in one of the products? I’m not sure on which elements to compare to ensure that this is a redox reaction:

$$\ce{2 H2SO4 (aq) + 2 NaBr (s) -&gt; Br2 (l) + SO2 (g) + Na2SO4 (aq) + 2 H2O}$$

I’m trying to indicate whether this involve oxidation/reduction. I’m looking at the oxidation state of $\ce{S}$. 

What do I compare $\ce{H2SO4}$ with: $\ce{SO2}$ or $\ce{Na2SO4}$? 

I tried both. Oxidation state of $\ce{S}$ in $\ce{SO2}$ is $+4$ and in $\ce{Na2SO4}$ is $+6$. One of them tells me it’s reduced, and the other tells me that no redox reaction is happening. 

Which do I compare with and why? I’m not sure.
One sulphuric acid is used to form sodium sulphate and the other one is used for oxidizing. 
So you have a reduction:

$\ce{4H+ + SO4^2- +2e- -&gt; SO2 + 2H2O}$

and an oxidation:

$\ce{2Br- -&gt; Br2 +2e^-}$

The sodium sulphate is formed with:

$\ce{H2SO4 + 2NaBr -&gt; Na2SO4 + 2HBr}$

This last one is not a redox reaction.