Charge

Oxidation States

= (total charge on complex - total charge on X-type ligands)

Oxidation states are based on treating a covalent compound as if it was an ionic compound, i.e. exaggerate ionic character. On that note, it is also based on as if all bonds to it are broken such that it has either electrons paired or not, i.e. closed shell (H⁺ 1s⁰ or H^- 1s², but not H•).

Oxidation state can be thought of as the “charge” on the metal when all ligands removed heterolytically in their closed form, with the electrons being transferred to the more electronegative partner.

Oxidation states can be written as

• Fe⁺²

  • Not Fe²⁺, as the number followed by the sign indicates the actual charge, not an oxidation state.

• Fe^III

• Fe(III)

Trends in metal oxidation states:

• Group oxidation state (oxidation state aligning with group number) achieved for left side, but not right. This is because the late transition metals have more compact atoms, leading to less ligands to oxidize the metal, which wouldn’t correspond to the group number.

• In general, +3 is more common with the early metals and +2 for the middle and late, because as size decreases, there will be less ligands to oxidize the metal (sterics-woven concept).

• Higher oxidation states are more stable as you descend the d-block because larger atoms can fit more ligands around them. This is unlike the p-block because p-block’s nuclear charge compacts the metal, leading to less ligands and lower oxidation states.

• First row transition metals are less electronegative than second and third row elements for later groups in the d-block. Because the later groups are more electronegative, they will not give up electrons readily. If they will rather keep those electrons, then they won’t be oxidized as easily. This leads to lower oxidation states because it doesn’t want to be oxidized.

• High oxidation states are stabilized strongly by electronegative elements. If I had a highly oxidized metal (a highly positive metal), and I can get stabilized by four ligands. Would I want each ligand to be more negatively charged or less negatively charged to stabilize me? Answer: I would want four ligands, all of them being more negatively charged than four of the other option with less negative charge.

  • Case in point: MnF₄ and MnO₄^-. Each fluorine offers -1, while each oxygen offers -2. This leads to the former having Mn(IV) (a lower oxidation state) and the latter having Mn(VII) (a higher oxidation state with more electronegative ligands for the same number of ligands).

• pi-acceptor ligands stabilize lower oxidation states.

• Oxidation states usually change from m to m-2, m to m+2 in reactions (i.e. in 2e^- steps).

For group n or n+10:

• Oxidation state (OS) is never greater than +n or lower than -n.

• OS is usually even for n even, odd for n odd, e.g. group 6 will usually have even number OS while group 7 will usually have odd number OS.

• For metals, OS is usually greater than or equal to 0.

• For electropositive metals, the OS is usually +n, e.g. Yttrium (EN = 1.22 with group 3) will have an oxidation state of +3.

Ending this, oxidation states do not indicate real charge at metal centre. It doesn’t tell you how much charge there is at a metal centre. We are only exaggerating ionic character - imagining if ionic character was prevalent. However, it does give us an indication whether a structure or composition is reasonable. Indications are what we discussed in the “trends” section!

Formal Charge

= (# of electrons in valence shell of free atom - # of e’s remaining on atom in molecule after breaking all bonds homolytically)

The actual charge on an atom is typically considered to be less than |±1|, a concept that is known as the Pauling electronegativity principle.

Valence

= (# of electrons in valence shell of free atom - # non-bonding electrons an atom in molecule)

Unfortunately, this breaks down when formal charge is included. The revised formula then becomes…

= (# of bonds) + (formal charge)

The number of electrons that an atom uses in bonding, it tells you about the number of non-bonding electrons available for additional bonding, which would result in valence < group valence).

dⁿ configuration of transition metals are critical to electronic structure. You can also think of this as “the number of electrons used in bonding and number of electrons not used in bonding” in a d system of some configuration is critical to electronic structure, magnetic properties, reactivity, etc.

The number of d-electrons tells us anything about the electrons of the metal centre in a complex.

Electron counting

For elements within the p-block, there is the octet rule. It isn’t really a rule. It is merely a rule of thumb to predict stability and reactivity. If you were to say that a molecule was electron deficient, then it would have fewer than 8 electrons.

For elements within the d-block, there is the 18-electron rule. Again, just a rule of thumb. If a complex had fewer than 18 electrons, then it would be electron deficient. Transition metal compounds with a full complement of 18 valence electrons are considered electron precise, same as for main-group compounds with a full octet, with some exceptions owed to other explanations in chemistry, e.g. electronics, sterics, MO theory, etc.

• For early TMs, 18 valence electron counts are often unattainable for steric reasons, as the required number of ligands would not fit.

• For later TMs, 16 valence electron counts are often quite stable, in particular square planar d⁸.

• For open-shell complexes, every valence orbital wants to be used for at least one electron. (Hund’s rule).

TIP: One of the questions that could pop up is to decide which reaction would occur. If any reaction scheme has an electron-sufficient count, then that would be the more favourable route because it would like to have that full shell.

Electron-Rich Organometallics

In a complex, there are fewer covalent bonds that “should” be present, as there is not enough valence orbitals available for all electrons present. These excess-electron compounds are relatively rare, especially for transition metals, and are often generated by reduction of lower electron count species.

Electrons around a metal can be in metal-ligand bonding orbitals or in metal-centred lone pairs. These lone pairs will be fairly high in energy, cf. MO theory’s HOMO. A metal atom with a lone pair is a sigma-donor and can be susceptible to electrophilic attack or itself be a powerful reductant/reducing agent.