Lecture Fourteen - Inorganic chemistry (periodic table and the elements) II

Question 1

Draw and explain the five equivilent 3d atomic orbitals.

Answer 1

The d orbitals are five fold degenerate.

Like the p orbitals, they are directional.

However, unlike p orbitals, they have two nodal planes and the sign change does not occur through the centre.

Sign change refers to the convention of labeling one side of the orbital + and the other -, or shading one bulb and not the other (means the same thing).

Question 2

What are the rules for neutral transition metal atoms?

Answer 2

General ground state electron configuration is [noble gas] ns^2 (n-1)d^x, n=4-7, x=1-10.

The ns orbital generally fills first.

Exceptions are Cr: [Ar] 4s13d5 and Cu: [Ar] 4s13d10

Due to change in relative energies of 4s and 3d orbitals and unusual stability of half filled or filled subshelles.

Aufbau principle = According to the principle, electrons fill orbitals starting at the lowest available (possible) energy levels before filling higher levels.

Pauli principle = In the case of electrons, it can be stated as follows: it is impossible for two electrons of a poly-electron atom to have the same values of the fourquantum numbers (n, ℓ, mℓand ms)

Hund’s rule = If two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs.

Question 3

What are the rules for transition metal ions?

Answer 3

In forming ions: S electrons are lost before d electrons.

E.g. Ti2+ electronic configuration is [Ar] 3d2 NOT [Ar} 4s2

Different metal ions with same electronic configuration have similar properties. E.g. Pale colour on Mn2+ and Fe3+ (3d5 ions).

Question 4

What are some of the trends seen in the periodic table?

Answer 4

Atomic size decreases across periods then stays constant after the metaloids.

Filled inner d-orbitals shield outer electrons (4s) from increased charge (c.f. main group elements where outer orbitals shield increased nuclear charge poorly).

Electrnegitivity (A) and ionization energy (B):

Both increase across whole periods but change is relatively small across transition series - small change in size.

Similar to values for large main group metals (e.g. Ga and Ge).

Question 5

Explain oxidation states in transition metals.

Answer 5

Multiple oxidation states are a feature:

Can accept and donate electrons - able to participate in redox reactions.

Since ns and (n-1)d electrons have similar energyies they can all be involved in bonding.

+2 and +3 oxidation states are most common:

ns2 electrons lost easily -> 2+ ion. Form ionic, basic compounds (like main group metals).

Up to Mn, highest oxidation state equals group number:

Remove valance electrons to form covalent, acidic compounds (like main group non-metals).

After Mn, high oxidation states are less common.

For 2+ ions, ability to act as a reducing agent decreases across the period.

All metals, except Cu, reduce H+ from aqueous acid to H2.