Inorganic Chemistry Study

Study Notes for Inorganic Chemistry:

Part 1: All the d block elements can be found in nature except for Tc which is highly-radioactive. Of the Lanthanoids only Pm is not naturally occurring. However of the Actinoids Only Thallium (Th) and Uranium ARE naturally occurring.

The noble metals are named such as they don’t react with H+ under normal conditions.

Lanthanoids must be separated via ion exchange chemistry since they are too similar & hard to react to separate via physical methods.

In general the size of atoms goes increases as you go down the periodic table. The size of an atom decreases as you move from left to right. (1st year chem.) This is because of increasing attracting from the nucleus pulling the electrons closer. This causes the orbitals to contract. Ionization energy increases as a result (though there are many spots the don’t follow this trend such as places with full shells or full half-shells). Electronegativity also increases, with F being the most electronegative atom.
p orbitals tend to be comparable in size to s orbitals, though are a touch smaller. This allows for large amounts of sp mixing. d orbitals are much smaller then s orbitals, but are large enough to participate in bonding. f orbitals are much smaller then even d orbitals, and are generally too small to participate in bonding or to mix with other orbitals.

The Lanthanoids decrease in size from La to Lu, this is known as Lanthanoid Contraction.

Note that the AUFBU principle doesn’t work very well for the d & f block elements. In the d-block elements will get a half filled shell one element early by moving a s-electron into the d-block. In the f-block atoms will hold onto a half filled shell one element longer then normal by placing the new electron into a d-orbital rather then an f-orbital.

In complexes all the electrons are in the d-orbitals.

<Skipping over a ton of stuff he will not ask about)

Early transition 3d metals prefer oxidation state 3, late 3d transition metals prefer 2. 4d & 5d transition metals prefer higher oxidation states. f-block elements mostly will take +3, but no trends.

Due to steric effects the highest oxidation states are achieved with oxygen.

Part 2: Coordination Chemistry

A complex is a chemical entity consisting entirely of metal atoms or a metal center surrounded by ligands.

A ligand is an ion or molecule that binds to a metal.

A coordination complex, also known as a Werner complex is a complex that does not contain carbon-metal bonds

An orgometallic complex is one that does contain metal-carbon bonds

The Coordination sphere is the arragment of ligands around the metal center

The coordination number is the number of ligands that are directly connected to the metal center.

In VB theory the ligand uses its lone pairs to form bonds to the metal center. The metal center provides empty orbitals.

A chelating ligand is one that can make several bonds to one metal center. (From the greek word for claw)

Monodentate: 1 attachment, Bidentate =2, etc.

Larger metals allow higher coordination numbers, as do smaller ligands.

CN2: Always linear (in transition metals), almost exclusively observed in d10 complexs

CN3 is very rare. d10 with strong steric effects will form trigonal planer (most common), d0 will form trig pyramidal w/ strict sterics.

CN4 is very common, usually tetrahedral. D8 metals will sometimes for square planer.

CN5 is also common. Trigonal bypirimidal is the most common, but is very close to square pyramidal, so much so that you will sometimes see both at once.

CN6 is also very common. Octahedral is the most common form of it (and the one we deal with the most in this class), trigonal pyramidal is also possible but quite rare.

CN7 is relatively rare. Can be observed w/ very small ligands or early 3d & 4d transition metals (due to high acidity). Much more common with Lanthanoids however & Actinoids however. Forms 3 strange shapes.

CN8 makes a bunch of odd shapes.

CN9 is relatively rare, try d0, Sc, Y, & f-block elements.

CN10-12 are all f-block elements and we don’t really care much about them.

Summary: CN4:Tetrahedral, CN5 trigbyipyramidal or square pyramidal, CN6: Octahedral.

UNLESS: d8: often square planer, d10 often linear, d0: Strange stuff.


Ionization Isomers are when an anionic ligand from outside the coordination sphere swaps with an anion from the coordination sphere.

Hydration isomers happen when water swaps with a ligand outside the coordination sphere.

Coordination isomers happen when you have a complex ligand with multiple metal centers, and ligands swap between them.

Linkage isomers occure when a ligand has more then one point by which it can attach to the metal center. The isomers occure when you swap between these points. Ex: CN, NO2, SCN.

Diasteriomers: Same as organic. Cis/Trans, Mer/Fac.

Enaintomers: Same as organic but Δ and Λ instead of S/R.

Kf is the formation constant. Formed stepwise, Kf = [complex]/[metal][ligand].

β  is formed by multiplying each Kf together.

Chelate effect: Less entropy loss from chains binding to metal centers then free ligands. Therefore binding chains is favoured.

Macrocycle effect: Same as chelate but comparing cyles & chains.

Part 3:

Crystal Field Theory:

Ligands are modeled as point charges.

Purely electrostatic forces: Therefore ligand:Metal interactions are repulsive.

Therefore orbitals that point at ligands are higher in energy then ones that don’t.

Δo values are the difference between the upper and lower d orbitals in an octahedral feild. No clear trend across the table, higher oxidation states increase it, as does moving down the periodic table.

Crystal Field Stabilization Energy: (0.4* eg electrons – 0.5* tg electrons)Δo.

If CFSE is lower then the pairing energy then ‘high spin’ state generated and electrons moved to the tg orbitals.

If you have the right set up in the tg orbitals you can get Jahn-Teller distortion. Moves xy orbital up, xz, yz down (this doesn’t matter), moves x2 orbital down, x2-y2 up: This can stabilize the complex.

D orbitals of tetrahedral complex’s are opposite of octahedral. (3 up, 2 down). Δt is 4/9Δo due to inefficient splitting.

Part 4: Molecular Orbital Theory

Molecular Orbitals are formed from atomic orbitals through a linear combination of atomic orbitals (LCAO)

The number of molecular orbitals formed is the same as the original number of atomic orbitals.

Mixing s & p orbitals is known as hybridization, increases bonding & antibonding character. Can only mix if they have the same symmetry.

Electronegativity lowers the relative energy of the bond.

Bonds become more ionic as the difference between the orbitals it is made of increases.

This is just a general example, CO disproves.

In octahedral feilds:

  • there is no spd mixing.
  • t2g is nonbonding (unless π donating/accepting is involved)
  • 6 lowest orbitals are mostly on the ligands


  • t2 orbitals interact with ligand orbitals of the same symmetry.
  • e orbitals are nonbonding as there are no ligand orbitals with the same symmetry.
  • d orbital order is reversed.

The amount of splitting at the frountier orbital depends on the different between the electronegativities.

π-bonding in octahedral complexs: Δo is increased by π-donation, decreased by π-acceptation.

MO theory can explain the spectrochemical series, with good π donors having weak splitting, σ donors having strong splitting and π acceptors having very strong splitting.

Metals also contribute to splitting: higher oxidation state splits more. Splitting increases down a group. However no trends within a given oxidation state among transition metals.

d0 prefers cis, d2 prefers trans.

Octahedral complexs prefer 18 electrons, tetrahedral prefer 16 electrons.

Part 6. Metal-Metal Bonding and Transition Metal Clusters

Atoms are considered bonded when they are closer then the atomic radius or the metallic radius.

A polynuclear or polymetallic complex contains more then one metal center. 3 types exist: Chains, Clusters (Closed structures with metal-metal bonds), and Cage Compounds (Closed structures with bridging ligands between the metal centers)

To form these metal-metal complexes the metal atoms must have partially empty subshells (so they can accept electrons), have a low oxidation state (To prevent repulsion between cations), and posses the ability to form strong metal-metal bonds (No DUH).

In metal-metal bonds the σ bond is the strongest, followed by the π bonds, then of course the δ bonds.

Metal-metal double bonds exist, though are surprisingly weak: many non-metal bonds are stronger then them. This is due to small overlap, especially with the δ-orbitals, and heavy inter-ligand repulsion.

A Quintuple bond has been formed. That is just wrong. Moving along….

Now a bunch of examples of 1-4 bonded complexes. Not sure what to learn from that. Wasn’t feeling so hot that day, didn’t pay much attention. *sigh* At the start you just pull off the δ electrons, but after they are gone you have to add antibonding electrons.

d4 metals can make double bonds. That is Re(III), Cr(II), Mo(II), W(II), Ru(IV), Nb(I), Despite the fact it is d4 no known quadruple bonds of Mn(III) exist.

Multiple metal-metal bonds tend to be reaction centers.

Alright, moving on to more useful (& thus testable things): Structures of clusters with metal-metal bonds:

We name ligands that bridge atoms μX, as in μ1. The subscript is the number of metal atoms the ligand bridges.

A closo cluster is completely closed, convex, and single shelled. The atoms in it will form a polyhedron. If the polyheadron only had triangular faces it is known as a deltahedron.

There are 4 general bonding types in closo clusters:

  1. Electron precise clusters have exactly 1 electron pair per polyhedron edge.

For a transition metal the 18 electron rule should be followed. μ1, halogens, hydrogen atoms, and groups such as SiR3 all supply 1 electron. NH3, PR3 and CO each supply 2 electrons. Bridging and such will change all of this.

To get the number of metal-metal bonds simply divide unpaired the electrons by 2. As a formula:
M-M bonds = (18*metal atoms –the total electrons)/2
However this only works for metal-metal single bonds.

  1. Clusters with one 2 electron, 3 center (2e3c) bond per triangular face
    If a deltrahedra with no more then 4 edges/faces meeting at any   one vertex does not have enough electrons it forms 2e3c bonds at        the faces.
    The number of 2e3c bonds is: (18*metal atoms – total electrons)/4
  2. Clusters that follow the Wade Rules
    If you have even less electrons then above then you use Wade’s               rules.
    Wade’s Rules: You need 2n+2 skeleton electrons to form a stable   closo compound. n = # of metal atoms. Does not work for     tetrahedral compounds, they only need 8 electrons instead of 10.
  3. And of course: Special Unique Snowflakes, I mean none of the above.

Polyoxometallates: Oxo complexes have no metal-metal bonds, but use an oxo (O2-) ligand as a bridge. They are a result of acid/base chemistry. In 3d cations the coordination number must be 4 and only 1 oxygen can be shared between any 2 atoms. In 4d & 5d metals the larger sizes open up coordination number 6 and allow multiple oxygen per 2 metal atoms, allowing for highly complex structures.

Part 7: Orometallic chem. of the transition metals

Hapicity is the number of atoms in a single ligand bonded to one metal center. Denoted by superscript ηx. If no superscript then the maximum # of atoms is being used. Note this may be different then the number of bonds.

CO bonds w/ carbon, not oxygen. 1 bond w/o back donation, 2 with. σ donor, π acceptor.

Phosphines are not officially orgometallics but we will treat them as such. σ donor, π acceptor. Electrons transferred from metal in σ* bond.

Hydride directly bonds to the metal. Can serve as bridges.

H2 can bond to metals. The stronger the π backdonation the weaker the H-H bond until you have 2 hydrogen lignands instead of one H2 ligand.

η1 hydrocarbons are all σ no π.

η2 alkenes & alkynes will change to alkanes if there is enough backdonation.

η4 alkenes that are nonconjuated act as 2 independent double bonds.

Benzene on the other hand donates via 1 σ donation and 2 π donations. Dontes electrons back via δ orbitals.

Common orgometallic reactions:

  1. Ligand substitution. Thermally or photochemically activated. Preceeds in steps that each have 16 electrons, problems with solvent competition.
  2. Oxidative addition: Turn XY into MX + MY. Are common for 16 electron species. Ligands usually add cis. Reverse = reductive elimination
  3. 1-1 insertion. Atom one away inserts between ligand & metal.
  4. 1,4 insertion: η2 ligands become η1 ligands w/o oxidation state change on the metal.
  5. Reverse of that, β-hydride elimination.
  6. α-Hydrogen abstraction. Removes hydrogen from C, adds a bond to metal.

Electron counting:

Neutral-Ligand method: all ligands are nutral, you count how many e- then donate to get this way.

Electron count: Metal valance electrons + Donated by ligands + charge of complex

Ligands donate 1 or 2 electrons.

Is fast an easy, but overestimates character of bonds. However it overestimates covalent character, underestimates charge of metal, makes it hard to find oxidation state of the metal.

Donor Pair Method: Ligands donate electrons in pairs. They are either neutral or charged.            Oxidation # of the metal = total charge on complex – charge on ligands.

Number of electrons form metal is group # – oxidation state.

Total electrons is sum of electrons on metal and donated by ligands.

Oxidation numbers easy to get, however overesitmats charge, assumed reactivity might be wrong.

So sorry I don’t have any more gaming posts yet: I’m busy with exams. I might get some more of  The Wyzard’s stuff edited on the weekend. I’ll update this post as I do more study, so I don’t have a stream of ichem study posts. Anyway, off to get food now. Until next time Stay Geeky


Post updated as I head to bed: I’d better get a lot more done tomorrow! Ah well, I think I can be ready.

Updated with another section of notes.

Update: I’m now done this exam, but will leave these notes up for interests sake. Until I get embarrassed by how rushed the notes near the end are, or the fact I never got through all of them.  Anyway, I’m going to see if I can get some RPG posts up latter tonight before Quantum comes and eats me alive.

Published in: on December 8, 2009 at 10:07 pm  Comments (2)  
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2 CommentsLeave a comment

  1. f orbitals are much smaller than even s orbitals, but did you mean than even d orbitals?

    • Yes, that should indeed be smaller then even d orbitals. Looks to be about half the size of a d orbital on average.

      I’ve fixed this in my offline copy and will be fixed here when I upload the next batch of notes. Thanks Love.

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