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Core 4b > Metal-ligand bonding and inorganic mechanisms > Flashcards

Metal-ligand bonding and inorganic mechanisms Flashcards

1
Q

Explain the trend in the radii and electropositivity of the elements in the periodic table. Also discuss the oxidation states and bonding.

A

The covalent/ ionic radius increases from right to left and down a group. The electropositive character increases from right to left and down a group. These trends are a result of the effective nuclear charge (Zeff), which is a result of shielding and penetration (s>p>d>f). The earlier metals exhibit the greatest variety in oxidation state. The 2nd and 3rd row metals can more commonly exhibit higher oxidation states. Bonding becomes stronger down a group.

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2
Q

What is the lanthanide contraction?

A

An electron in an f-orbital is very poorly shielded, which results in a steady decrease in the radii of the lanthanides. The 2nd and 3rd row transition metals have very similar radii.

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3
Q

Explain the bonding in 1st row transition metal complexes.

A

The 3d orbitals in the 1st row metals are not as diffuse as the 4d and 5d orbitals in the 2nd and 3rd row metals. This leads to a larger ionic component in the bonding of 1st row metal complexes.

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4
Q

Explain the bonding in lanthanides.

A

The 4f orbitals of the lanthanides are essentially core orbitals and cannot participate significantly in covalent bonding. The bonding in lanthanides can be considered almost completely ionic.

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5
Q

What is n-hapticity?

A

The number of contiguous atoms of a ligand attached to a metal.

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6
Q

What is k-denticity?

A

The number of non-contiguous atoms coordinating from a ligand (often a chelating ligand).

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7
Q

What is u?

A

The number of metal atoms bridged by a ligand.

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8
Q

How do you calculate the metal oxidation state?

A

Oxidation state = the charge on the complex - sum of the charges of the ligands

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9
Q

What is the electroneutrality principle?

A

The electronic structure of substances is done so that each atom has essentially zero resultant charge. No atoms will have an actual charge greater than +/- 1. Therefore, the formal charge is not the actual charge distribution.

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10
Q

How do you calculate the metal d-electron count?

A

d-electron count = group number - oxidation state

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11
Q

How do you calculate the TVEC at the metal?

A

TVEC = d-electron count + electrons donated by ligands + number of metal-metal bonds

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12
Q

What are the 4 enthalpic effects of TM complex formation?

A
  1. The greater the number of ligands, the more stable the complex. The formation of strong bonds stabilises the complex. For TMs, common coordination numbers are 4 and 6.
  2. The number of ligands are limitted by ligand-ligand repulsions (usually up to 6 for TMs).
  3. Large negative charges are limitted by electron-electron repulsion, and large positive charges are limitted by ionisation energy.
  4. Electrostatic attraction has the biggest contribution to the thermodynamic stability, then destabilisation from core electrons, then destabilisation from valence electrons, and then CFSE.
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13
Q

What are the 2 entropic effects of TM complex formation?

A
  1. The number and size of any chelating ligands. 5 and 6 membered rings are most stable due to less ring strain.
  2. The requirement for ordered solvent cages lowers the entropy (solvation).
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14
Q

Why do most metals obey the 18 electron rule?

A

The valence shell of a TM has the structure nd + (n+1)s + (n+1)p, where n = 3-5. There are a total of 9 orbitals in the valence shell (5 d-orbitals, 3 p-orbitals and 1 s-orbital). A maximum of 2 electrons can occupy each orbital, so the shell can have a maximum of 18 electrons in its valence shell. For many complexes, this electronic configuration is the most thermodynamically stable.

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15
Q

What are the 4 exceptions to the 18 electron rule?

A
  1. 1st row coordination complexes where the bonding is predominantly ionic
  2. Square planar d8 complexes (16 electrons)
  3. Early metal complexes with π-donor ligands
  4. Paramagnetic complexes
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16
Q

Describe σ-donor ligands.

A

The bond between the ligand and the metal is a σ-bond. Examples include: H-, R- (any alkyl or aryl), H2O and NH3.

17
Q

How do you characterise a metal hydride?

A

For IR, v(M-H) is approximately 1750cm-1. For NMR, the hydride resonance is at high field, so less that 0ppm. Neutron diffraction is needed to locate the hydrogen nuclei.

18
Q

What is the MO diagram for an ML6 complex, where L is a σ-donor?

A
19
Q

How are the d-orbitals ordered in different complexes?

A
20
Q

Describe σ-donor, π-acceptor ligands.

A

The complex has a σ-donor bonding component plus an additional π interaction between the empty ligand orbitals and the occupied metal orbitals. Examples include: CO, CN, H2, alkenes, N2, O2 and PR3.

21
Q

What is the MO diagram for CO?

A
22
Q

Show the σ-donor and π-acceptor interactions for CO.

A
23
Q

What is the experimental evidence for the CO bonding model?

A
24
Q

Explain the trend in v(CO) as the oxidation state of the metal decreases.

A

As the oxidation state decreases, the electron density around the metal increases, so more electron density is available for more π-acceptor bonding to the CO ligands.

25
Q

Explain the trend in v(CO) as the number of coordinating CO ligands decreases.

A

As the number of coordinating CO ligands decreases, there are fewer π-accepting ligands to relieve the negative charge build up.

26
Q

What are the π-acceptor properties of the phosphine ligands?

A

PF3 > PCl3 > PMe3

27
Q

Explain the trend in v(CO) as the number of metals increases.

A

The more metals bonded to the CO ligand, the more electron desnity there is available for π-acceptor bonding.

28
Q

What is the MO diagram for N2?

A
29
Q

What is the MO diagram for O2?

A
30
Q

Show the n1 and n2 bonding for O2.

A
31
Q

Explain the trend in v(O-O) between the different forms of O2.

A

As the electron density in the π* orbitals increases (O22- > O2- > O2) the O-O bond distance increases and the vibrational frequency decreases.

32
Q

Why is n1-O2 bent when CO is linear?

A

O2 has to accommodate an extra pair of electrons in the 1πg (π*) orbital. These occupy 1πgx (to form the σ-bond through one lobe of the 1πgx orbital) leaving 1πgy to form a π-acceptor interaction.

33
Q

How many electrons does NO donate?

A
34
Q

What is the strategy for determining bent or linear NO, electron count and oxidation state?

A
35
Q

Show the σ-donor and π-acceptor interactions for H2.

A

If sufficient electron density is transferred from the metal to the σ* orbital of H2 the H-H σ-bond will break and give two M-H (metal-hydride) σ-bonds (oxidative addition).

36
Q

What is the difference in characterisation between dihydrogen M(H2) and dihydride M(H)2 complexes?

A
37
Q

What is the MO diagram for an alkene?

A
38
Q

Show the σ-donor and π-acceptor interactions for an alkene.

A
39
Q
A

Core 4b (4 decks)

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