CHEM 4680 / Topic 2: Valency

What is the formula for oxidation state?

OS = total charge on complex - total charge on X-type ligands

What are oxidation states?

Oxidation states are hypothetical charges assigned to a metal based on treating a covalent compound as if it was an ionic compound (i.e. exaggerate ionic character and electrons being transferred to the more electronegative partner) and removing the ligands heterolytically in their closed-shell form (pairs of electrons or no pairs of electrons).

What are the three ways you can write oxidation states?

• Fe⁺²

• Fe^III

• Fe(III)

Explain the trend between metal oxidation states and group 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.

Explain the trend between metal oxidation states and early, middle, and late metals.

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).

Explain the trend between metal oxidation states and the size of d-block elements.

Higher oxidation states are more stable as you descend the d-block because larger atoms can fit more ligands around them.

Explain the trend between metal oxidation states and the electronegativities of d-block elements.

For later groups in the d-block, first row transition metals are less electronegative than second and third row elements. Because the second and third rows 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.

Explain the trend between metal oxidation states and stabilization of electronegative elements.

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.

What’s the relationship between metal oxidation states and pi-acceptors?

pi-acceptor ligands stabilize lower oxidation states.

In relation to oxidation states, how many electrons are usually taken away or taken into for elements?

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

What are the four trends of group oxidation states?

• 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.

What is the formula for formal charge?

Formal charge = # of electrons in valence shell of free atom - # of electrons on atom in molecule after breaking all bonds homolytically

• What is the formula for valence when there is no formal charge?

• What is the formula for valence when there is formal charge?

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

• valence = # of bonds + formal charge

What is valency?

The number of electrons (or “hooks”) that an atom uses in bonding.

• What is the octet rule?

• What does it mean for a molecule to be electron-deficient in this case?

• 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.

• What is the 18-electron rule?

• What does it mean for a molecule to be electron-deficient in this case?

• What does it mean for a molecule to be electron-precise in this case?

• 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.

Are 18 valence electron-counts present for early metals? Why or why not?

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

Are 18 valence electron-counts present for later metals? Why or why not?

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

With electron-rich organometallics, there can be fewer covalent bonds than what you would expect. Explain why.

• In a complex, there are fewer covalent bonds that “should” be present, as there is not enough valence orbitals available for all electrons present for bonding.

• The electrons in excess 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.

• These excess-electron compounds are relatively rare, especially for transition metals, and are often generated by reduction of lower electron count species.

NOTE: 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.

When do you use formal charge and when do you use oxidation state?

• Formal charge is used in resonance structures and Lewis structures to determine the most stable arrangement of atoms by tracking electron distribution. Formal charges are dependent on resonance structures.

• Oxidation state is used in redox reactions and electron transfer to track how many electrons an atom has gained or lost. Oxidation states are independent of resonance structures.

What is the equation for bond order?

(# of electrons in BMO - # of electrons in ABMO) / 2

Explain how Cp₂WH₂ can be protonated and become stable when becoming Cp₂WH₃⁺.

This is because there are a total of two leftover d electrons from Cp₂WH₂ to attack the proton to become a product with a stable 18 electron count.

How can metal-metal bonding be possible?

If there are available d orbitals that can be used to form a bond by accepting one electron from another metal partner and donating one electron to said partner, then it is possible.

What are the two types of bridges?

• Bridges with classical 2c-2e interactions.

• Bridges with non-classical 3c-2e interactions.

• What allows 3c2e bonds occur?

• What does not allow 3c2e to occur?

• When can we say that there is a net “metal-metal” bonding interaction in a bridge?

• What allows 3c2e bonds to occur is when a covalent bond between an atom and a ligand (that has no lone pairs!) donates their two electrons to the vacant orbital of another atom.

• What does not allow 3c2e bonds to occur is when the ligand (involved in the bond wanting to be used) has a lone pair. If lone pairs are available, we use them.

• We can only say that there is a net atom-atom bond when there is 3c2e, because the electrons used by one metal is also being used by the other metal in the bridge. We cannot say that there is when 2c2e, because the electrons used by one of the metals is not the same as the electrons being used by the other metal in the bridge.

NOTE: If we have lone pairs, we can use them. If we do not have lone pairs, we use bonding pairs.

• What does bimetallic species mean?

• What does trimetallic species mean?

• What does homobimetallic species mean?

• What does heterobimetallic species mean?

• Bimetallic species means a species with two metal centres.

• Trimetallic species means a species with three metal centres.

• Homobimetallic species means a species with two of the same metal centres.

• Heterobimetallic species means a species with two different metal centres.

There are instances wherein the 18-electron rule is disobeyed. The instances are related to the following bullet points. Dictate what specific limitation it would be:

• d⁰ complexes.

• d⁸ complexes.

• d¹⁰ complexes.

• Complexes with bulky ligands.

• Metal centres which interact with ligands electrostatically.

• Metal clusters with more than 5 metals.

• d⁰ complexes: 18-electron rule is disobeyed because once you have reached d⁰, then chances are you cannot make any more bonds to increase your valence electron count because you have no electrons to share to make a bond. On that note, you can also say that further ligand interactions cannot occur because of electron absence.

• d⁸ complexes: 18-electron rule is disobeyed because some of these complexes prefer a square planar complex with 16 electrons. This results in the the metals not wanting to add as many ligands to reach higher electron counts.

• d¹⁰ complexes: 18-electron rule is disobeyed because they aren’t accessible to accept or be Lewis acidic enough due to fully occupied orbitals, then the metals can’t add as many ligands to reach higher electron counts.

• Complexes with bulky ligands: Because the steric bulk of ligands blocks the binding of other ligands, then the metals can’t add as many ligands to reach higher electron counts.

• Metal centres with electrostatic interactions: Because these electrostatic interactions are ionic rather than covalent, then there won’t be any sharing. Just give and take → ionic. This doesn’t add valence electrons to the complex to increase TVe^-, rather it takes valence electrons from the complex.