Inorganic chemistry

for 2. think about what is required to go to a higher oxidation state, (successive ionisation energies), and then how ionisation energies vary across the period

d block elements have similar ionisation energy?

Think about periodic trends. The d block elements do not have the same IE. Moving across the group, lets take the 1st row transition metals, from Ti --> Zn , you are increasing the effective nuclear charge (Zeff). You are filling up the 3d orbital, d1 --> d10 electrons, as you go across the series.

So as an example, Ni d8 has greater Zeff than Ti d2, meaning its valence electrons experiences greater attraction to the nucleus and therefore is bound more tightly by the nucleus and harder to be ionised, hence a higher IE.

Increase in Zeff across the series = increase in IE

oh i get it, because of higher IE, electrons are harder to remove, so the oxidation state do not increase across the period?

Are you a chemistry undergraduate?

If so I imagine you need more detail, ie showing the pattern of oxidation states that are stable across the series, it is not just a straight decrease across the period, it increases first, peaking at Mn, where you get some Mn7+ species (these are highly oxidising), then decreases again, so you need to look into that - probably explained on your lectures

What have you tried so far?

is it because the trend decreases as the tendency for d electrons to participate in bonding decreases?

is it because the metals are coated with a layer of metal oxide which prevent the reaction?

Why do you think d-block metals don't produce H₂ when reacted with acids?

they are lower than hydrogen in the ECS?