Why do real gases not behave exactly like ideal gases?

The molecules have (a) volume and (b) attractive forces  

Step-by-step explanation:

At ordinary conditions, the molecules are so far apart that the gases behave almost ideally.

However, if you use high pressure and/or low temperatures, you force the molecules to be close together.

There are two competing effects:

n+h→nh3

we must remember that nitrogen and hydrogen are both diatomic molecules in their standard gas form. this adjusts the equation to

n2+h2→nh3

now we need to adjust coefficients in order to balance the atoms on each side of the equation. currently we have 2 atoms of nitrogen and 2 atoms of hydrogen on the reactant side and 1 atom of nitrogen and 3 atoms of hydrogen on the product side.

we can balance the hydrogens by placing a coefficient of 2 in front of ammonia and a coefficient of 3 in front of the hydrogen.

n2+3h2→2nh3

this gives us 6 hydrogen on each side and coincidentally the nitrogens now equal 2 on each side.