DM ghl: D-block electronic configurations; transition metal catalysis Flashcards
Indicate the orbital blocks on the periodic table.
Draw a diagram showing the energies of electron sub-shells from n = 1 to n = 4.
State the elements for which this diagram is correct.
Correct up to Ni in period 4, after which 4s sub-shell has higher energy than 3d.
Note: energy gaps should decrease with increasing principal quantum numbers, as below.
What defines d-block elements?
Elements for which the last-placed electron is in a d-orbital.
Not outermost shell; 4s already filled since lower energy.
What is the electronic configuration of argon?
1s22s22p63s23p6
Give the electronic configuration of all of the elements in period 4 of the d-block, in terms of the electronic configuration of argon.
Give the electronic configuration of chromium, and explain why it is is not [Ar]3d44s2, as might be expected.
[Ar]3d54s1.
This is an energetically favourable arrangement since it avoids repulsion of two electrons in 4s orbital.
Give the electronic configuration of copper, and explain why it is is not [Ar]3d94s2, as might be expected.
[Ar]3d104s1
This is an energetically favourable arrangement, since 2 electrons in the fifth 3d orbital experience less repulsion than 2 electrons in the 4s orbital.
Draw a diagram showing the arrangement of electrons in the ground state of elements in the first row of the d-block.
What is a transition metal?
A d-block element which forms one or more stable ions with incompletely filled d-orbitals.
From where are electrons first lost when transition metals form simple ions?
The 4s sub-shell (and then the 3d sub-shell).
When empty, 3d has more energy than 4s, but this reverses when electrons are populated (quantum physics…)
Which elements are transition metals?
Ti to Cu (not Sc or Zn).
All trans metals are d-block elements but not all d-block elements are trans metals.
Use the electronic configurations of the common ions of zinc and scandium to demonstrate why they are not transition metals.
- Zn [Ar] 3d104s2 →* Zn2+ [Ar] 3d10
- Sc [Ar] 3d14s2 →* Sc3+ [Ar]
Neither form stable ions with incompletely filled d-orbitals, so they are not transition metals.
Why do transition metals exist in variable oxidation states?
There are several stable arrangements of their 3d and 4s electrons.
- What are the common oxidation states of iron?
- Write their electronic configurations.
Fe [Ar] 3d64s2
Oxidation state = +2: Fe(II) or Fe2+ [Ar] 3d6
Oxidation state = +3: Fe(III) or Fe3+ [Ar] 3d5
With reference to electronic configuration, suggest and explain why Fe(II) is spontaneously oxidised to Fe(III).
- Fe2+ has 2 electrons in its 5th d-orbital, causing repulsion
- Fe3+ has only one electron in each d-orbital, so experiences less repulsion
- Fe3+ has the more energetically favourable configuration so is stabler