L1 - origins of variable oxidation states

objectives

Understandings:

• Transition elements have variable oxidation states
• Zn is not considered to be a transition element as it does not form ions with incomplete d-orbitals.
• Transition elements show an oxidation state of +2 when the s-electrons are removed.

Applications and skills:

• Explanation of the ability of transition metals to form variable oxidation states from successive ionization energies.

Guidance:

• Common oxidation numbers of the transition metal ions are listed in the data booklet in sections 9 and 14.

Definitions

• The definition of a transition element is

'D-block elements forming one or more stable ions with partially filled (incomplete) d-sub shells'

• The first row runs from scandium to zinc filling the 3d orbitals.
• Properties arise from an incomplete d sub-shell in atoms or ions

Metallic Properties

• all the transition elements are metals
• They have strong metallic bonds due to small ionic size and close packing
• higher melting, boiling points and densities than s-block metals

Electronic Configuration of the d-block elements

In numerical terms one would expect the 3d orbitals to be filled next.

However, because the principal energy levels get closer together as you go further from the nucleus coupled with the splitting into sub energy levels, the 4s orbital is of a LOWER ENERGY than the 3d orbitals so gets filled first.

• As expected, the next electron in pairs up to complete a filled 4s orbital.
• This explanation, using sub levels fits in with the position of potassium and calcium in the Periodic Table. All elements with an -s¹ electronic configuration are in Group I and all with an -s² configuration are in Group II.

• With the lower energy 4s orbital filled, the next electrons can now fill p the 3d orbitals. There are five d orbitals. They are filled according to Hund’s Rule.

BUT WATCH OUT FOR TWO SPECIAL CASES.

• The 3d orbitals are filled according to Hund’s rule so the next electron doesn’t pair up but goes into an empty orbital in the same sub level.

Special Cases

In SL Topic 3 we looked at s,p,d,f electron configuration from the wave model of atomic theory. We also looked at Cu and Cr as two elements that do not follow the aufbau principle directly due to their electron transitions from 4s into 3d.

• One would expect the configuration of chromium atoms to end in 4s² 3d⁴.
• To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d to give six unpaired electrons with lower repulsion.

• One would expect the configuration of copper atoms to end in 4s² 3d⁹.
• To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d.

• You should remember that electrons normally fill orbitals separately to prevent repulsion and normally do so successively. The idea that electrons fill orbitals separately is Hund's rule and the idea that electrons normally fill lower energy orbitals first is called the aufbau principle.
• Chromium does not follow the aufbau principle because it can promote it's 4s2 electron into the 3d shell. The 4s shell is close the the 3d shell in terms of energy.
• In chemistry there are two further ideas that help to explain why this happens. These combine to explain that 'half-filled shells and full shells are more stable' than empty or incomplete shells. This is down to the Pauli exclusion Principle (where electrons can only spin in the same direction to reduce repulsion) and the Fermi correlation.
• To apply this to Cr and Cu then -the half filled 4s1 and 3d5 in Chromium are more stable than the full 4s2 and incomplete 3d4 shell. For Cu the half filled 4s1 and the full 3d10 is more stable than the full 4s2 and the incomplete 3d9.
• The take home message here is that electrons can move between 4s and 3d relatively easily and that the energy associated with them is very similar.

When electrons are removed they come from the 4s orbitals first

• When transition metals form ions the 4s shell changes to become higher in energy than the 3d shell.
• This means that the 4s electrons are always lost first from these elements when they ionise.
• This results in ALL Transition metals having a stable +2 ion.
• When this happens we would have 4s0 - this would normally be omitted but can be included in exam question answers.

• In all the above examples the electron are first lost from 4s and then 3d.

Variable Oxidation States

• As the transition metals are metals they prefer to Lose electrons to form a stable species.
• This process is called Oxidation.
• An Oxidation state is simply thought of as 'the number of electrons a species needs to gain in order for it to return to it's ground state.'
• The ground state of an element is when the element has no ionic charge.
• EG
• Zn⁺² requires 2 electrons to return to it's ground state

How is the Oxidation State determined?

• The Oxidation state is determined by 'the maximum number of electrons an element can lose without requiring so much energy to remove the electron that the energy cannot be recovered by bonding.'
• This idea about energy and bonding can be best displayed by Ionisation energies in.
• In Topic 3 we discussed how to identify an element's group from it's successive ionisation energies.
• We could identify these as there would be a large 'jump' in ionisation energy when we drop a period.

• In the above graphs we can see that the Na is in Group 1 as the large jump in Ionisation Energy (IE) is between the 1st and 2nd IE.
• Mg is in Group 2 as the large jump is between the 2nd and 3rd IE.
• Both these jumps are very high in Energy meaning the energy cannot be recovered by bonding and so the formation of a stable Na^{+2 or} Mg⁺³ ion is very unlikely.

• In the d block there a much larger number of valence electrons.
• As successive ionisation energies take more and more energy to remove the electrons (due to the increase in positive charge and decrease in ionic radius) it becomes unfeasible to remove all valence electrons like in the above Na or Mg examples.
• We have already explained why it is easy to remove 4s electrons however it is also difficult to predict how many electrons these elements will lose.
• This is because the successive IE's only rise gradually. There is no marked jump between energy levels.

• In Figure 3.30 you can see the overlay of Magnesium and Manganese successive IE's
• As you can see - the manganese has a gradual incline of IE's whereas the jump from 2 to 3 for Mg is obvious.
• As the difference of these IE's is so small this energy can sometimes be recovered by bonding.
• Thus - as these ions all have similar energies transition metals can exist at different (Variable) Oxidation States.

A general rule taken from the above oxidation states of transition metals table.

From Ti --> Mn these metals can lose their valence electrons it is important to state that it is very rare for elements to lose electrons beyond 3d5.

This is because the Attraction for the Nucleus and the effective nuclear charge is so strong that it takes too much energy to remove the electrons completely.

Which elements do not follow these rules?

ZINC is NOT classed as a Transition Metal

• This is because whilst other transition metals form ions with incomplete d shells Zn only forms stable ions when losing s shell electrons - it is colourless.

SCANDIUM is also not classified as a Transition metal

• This is because it again - does not form stable ions with an incomplete d-shell. Scandium only forms +3 ions where both the s and d shells are empty. Again this results in compounds that are colourless.