Mike Whittlesey Flashcards
Explain the ground state electronic configurations of Cr and Cu
In other first row metals, 4s fills before 3d, and then the d-orbitals fill sequentially
There is a perturbation at Cr (3d54s1) and Cu (3d104s1), where half-filled s-orbitals are generated
This configuration is most stable due to maximisation of the exchange energy
i.e. the number of parallel spins has been maximised
Exchange energy for 3d44s2
6K
also e-e repulsion between the electrons in 4s
Exchange energy for 3d54s1
10K
General trend in metallic radius for 1st row metals
Metallic radius decreases across the period due to increasing Zeff and the relatively poor shielding ability of d electrons (poor penetrating power - diagram)
Blip in the trend at Mn and Zn (metallic radius increases)
This is because the electrons are in a more spherical arrangement due to the extra electron being placed into an s-obrital (Mn and Zn are after Cr and Cu respectively)
General trend in heat of atomisation for 1st row metals
Similar to that for metallic radii
Increases across the period but not a constant trend (blips at Cr/Mn, Cu/Zn)
Heat/enthalpy of atomisation
The energy required to remove 1 atom of an element from the bulk of the metal
Measured in kJ/mol
An indicator of M-M bond strength
Why are s-electrons better at shielding than p/d/f?
S-electrons can get very close to the nucleus (high penetrating power) so can shield Zeff well
This is not the same for d/f electrons so they are poor shielders
Electronic configurations of 1st row metals
3d/4s orbitals are of similar energy so fill sequentially
Electronic configurations of 2nd row metals
4d/5s orbitals are similar in energy but not completely the same
They are most similar at the end of row but there is a big difference in their relative energies at the start of the period
Electronic configurations of 3rd row metals
There isn’t really a crossover in the energy of 5d/6s orbitals so the pattern in electronic configurations is different
Why is the pattern for the electronic configurations of 1st, 2nd and 3rd row metals different?
Because the relative energies of the s- and d-orbitals are different in each row
Trend in metallic radii from 1st to 2nd/3rd row metals
Metallic radii of 2nd and 3rd row metals are bigger than those in the 1st row
The electrons are in orbitals of higher n - i.e. the orbitals are larger and further away from the nucleus
Trend in metallic radius from 2nd to 3rd row
Not necessarily an increase in metallic radius/size from 2nd to 3rd row
This is due to the lanthanide contraction
The additional row of f-elements between the filling of the 4d and 5d orbitals are poor shielders due to orbital penetration
5s and 5p orbitals penetrate the 4f orbitals, meaning electrons in 4f orbitals are not shielded from the increasing nuclear charge, leading to a decrease in atomic radius across the row of lanthanides
Trend in DeltaHatom from 1st to 2nd to 3rd row metals
Enthalpy of atomisation increases from 1st to 2nd to 3rd row metals because the overlap of 5d-5d orbitals is more effective than 4d-4d and 3d-3d overlap
5d orbitals are larger (“greater spatial extent”) so the electron distribution is shifted further out and there is an increased probability of finding electrons further from the nucleus
i.e. bonds increase in strength down the group for d-block elements (in contrast to non-d-block elements, where bonds become weaker down the group)
Metal-metal bonding in d-block elements
The trend in enthalpies of atomisation for d-block metals suggests metal-metal bonding is more extensive for the 2nd and 3rd row
Might also anticipate that the highest incidence of metal-metal bonding occurs around the middle of the periods, where enthalpies of atomisation are highest
Heavy d-block elements like M-M bonds
Why is metal-metal bonding extensive in low oxidation state, polynuclear metal complexes of 2nd and 3rd row elements?
A low oxidation state means the electrons are readily available in diffuse d-orbitals
i.e. the d-orbitals are more diffuse when the metal is in a low oxidation state so metal-metal bonding is more common (“spatial diffusivity of d-orbitals”)
Metal-metal bonding can be illustrated by comparing the calculated and experimental values for DeltaHf in CrCl2, MoCl2 and WCl2
DeltaHf assumes the compound is ionic
DeltaHf(calc) and DeltaHf(exp) are very similar for CrCl2 so this compound can be treated as ionic
There is a discrepancy between the calculated and observed values for MoCl2 and WCl2 so cannot be regarded as ionic solids
(They actually form M-M clusters/aggregates)
Oxidation states
In general, higher oxidation states in the 2nd and 3rd row are more stable than those in the 1st row (e.g. WF6 is known but CrF6 is not, WO3 is not readily reduced where CrO3 is a strong oxidising agent)
Low oxidation state organometallic complexes are also increasingly stabilised upon descending a group
e.g. Ru3(CO)12 undergoes substitution of the CO ligands at room temp, whereas Os3(CO)12 only undergoes substitution upon heating, due to better M-L orbital overlap due to the spatial diffusivity of the d-orbitals
Why are higher oxidation states more common on the left hand side of the d-block?
Higher oxidation states decrease in stability as you go across the d-block because Zeff increases
(i.e. the nuclear charge is ‘higher’ so wants to retain its electrons)
Group oxidation state
Can be reached up to group 8 (e.g. OsO4)
Beyond that, the sum of the ionisation energies is too high to be compensated for by M-L bond formation
Typical coordination numbers for 2nd and 3rd row metals
The larger size of the 2nd and 3rd row elements favours CN > 6, and CN = 7, 8, 9 is even more common
High CNs are achieved with small/non-polarisable ligands e.g. OH-, F- or with chelating ligands with low steric requirements
CN = 7,8
There are several limiting geometries
And most complexes are fluxional on the NMR timescale, so their geometries must be established in the solid state by X-ray crystallography
Examples of 7-coordinate geometries
Pentagonal bipyramid
Capped trigonal prism
Capped octahedron
(draw)
Examples of 8-coordinate geometries
Square anti prism
Dodecahedron
Cube (generally unfavourable due to L-L interactions