For the first question, I believe what you'd want to do is find all the products between the isotope masses and their relative abundances (so for the first one 32 x 0.511, second one 33 x 0.325, etc.) and just add them all together to find the average atomic mass, which would be the defined atomic mass of the element. The answer would then be 33.19 (2dp)
What's actually happening here is you're using the abundance values (which represent how much of that element in nature is that particular isotope) to find the average atomic mass of the element in general, which covers all of its isotopes.
For example lets say we have a made up element of Queenslandium. This element has two isotopes, one with a atomic mass of 10 (we'll call Q1) and another with an atomic mass of 12 (we'll call Q2). Now, 75% (three quarters) of all Queenslandium on earth is Q1, but the remaining 25% (quarter) is Q2. To find the atomic mass of Queenslandium itself, we multiply the masses of each isotope with how much they represent in the real world, so that we end up with
(3/4) x mass Q1 + (1/4) x mass Q2
= [(3/4) x 10] + [(1/4) x 12]
So the atomic mass of Queenslandium is 10.5