Molecular orbitals 3

An electron can absorb a photon, using the photon's energy to jump to a higher-energy molecular orbital.

In the case of a single-bonded diatomic as shown above, the excited state has one electron in a bonding orbital and one in an anti-bonding orbital, producing a net bond order of 0.

No net bonding means the atoms can drift apart. The energy of the absorbed photon would therefore equal the disassociation energy of the bond.

An example of the use of this is absorption of UV light by BR₂ to produce bromine radicals, used to produce the free radical reactions seen in A-level chemistry.

S and p character with hybridisation

A nucleophile will tend to donate electrons which are at the highest energy level. Generally:

Which makes sense. For example, we already expect ammonia to donate its lone pair rather than a hydride ion (H^-). 

There is a related pattern in the hybridisation of s and p orbitals. In ethyne, a carbon atom has two p orbitals used in pi bonding, and one s and one p orbital used in sigma bonding. The sigma bonds are therefore 50% s character.

For a carbon atom in ethene, one p orbital is used in a pi bond and one s and two p orbitals are used in sigma bonds, sigma bonds in ethene are therefore 33% s character.

In ethane, sigma bonds are 25% s character.

A s orbital is lower in energy than a p orbital, so molecular orbitals with high s character are lower in energy.

A visible result of this is the structure of CH₃⁺, which has three sp² orbitals and one empty pi orbital:

If someone asked "why isn't it sp³ hybridized with an empty sp³ orbital. Just remember that lower-energy orbitals only affect stability when electrons are contained in them. sp² hybridisation rather than sp³ in this molecule lowers the energy of the occupied orbitals and increases the energy of the unoccupied orbital.. 

In other words, it won't waste s character on an empty p orbital.

And obviously if you added a lone pair to make CH₃^- then it would sp³ hybridise and adopt a tetrahedral shape.

Bromine test, alkene bromination

The bromine test is able to detect unsaturated compounds:

1. Elemental bromine is dissolved in water, dichloromethane, or tetrachloromethane. The dissolved bromine will make the liquid brown.

2. The dissolved bromine is added to the alkene and shaken. If the alkene is a gas then it can be bubbled through the brown liquid.

3. Discolaration indicates unsaturated bonds. The ratio of bromine and alkenes can also indicate the amount of unsaturated bonds per molecule.

If the mixture contains phenol, you may also get a white precipitate of tribromophenol. Filter and dry this to get white needle-like crystals with an antiseptic smell. It is used as a fungicide and wood preserver.

Tribromophenol

Chlorine is faster at this reaction, iodine is slower. 

If bromine water is used then this complicates the reaction, since the cation created in the mechanism is much more likely to collide with a H₂O molecular than Br₂. For example, if the alkene used is ethene, then the most common product will be bromoethanol and hydrobromic acid, as shown below:

Step 1:

Step 2, without water:

Step 2, with water:

Factors affecting orbital overlap, boron halides

There are three main factors which affect orbital overlap, the first two have been described earlier.

1. Symmetry

2. Proximity of energy levels

3. Size of the orbitals

This third factor helps explain why boron trifluoride is a weaker lewis acid than boron trichloride:

From simple electronegativity arguments we would expect BF₃ to be a stronger acid, since electrons would be pulled away from the boron atom hence encouraging it to accept an electron pair. 

The reason BF₃ is weaker is that the halogen atoms contain electrons in p orbitals, which in boron trihalides have a small overlap with the empty p orbital of boron. This effect is weaker in heavier halogens because of the size difference:

The p-orbital overlap can also represented as resonance structures:

This overlap/resonance/extra stability cannot happen in an adduct, since in an adduct the orbital is already occupied by a lone pair.

Acids

Acids have two main definitions. The Bronsted-Lowly definition is the one used an A-level.

Bronsted-Lowly acid: A proton donor.

Lewis acid: A substance which accept a pair of electrons.

H+ is a Lewis acid, it has an empty 1s orbital to take in a lone pair:

Boron trichloride is also a Lewis acid. The boron atom in this is sp² hybridized, leaving an unoccupied p orbital. This p orbital can accept an electron pair.

A Lewis base is a lone-pair donor. When a Lewis base donates a lone pair to a Lewis atom it creates a complex called an adduct. An example of an adduct you would have seen is using AlCl₃ as a friedel-crafts catalyst in the electrophillic substitution of benzene:

Lewis acids in water are similar to Bronsted acids, because the adduct they form with water allows easy disassociation of H+:

Doping and corundum

Intrinsic defects occur in a pure crystal, without having to add other elements.

Extrinsic defects are from the presence of impurities, from atoms not contained in the normal stoichiometry of the pure crystal. These can be substitutional or interstitial. Even trace amounts of these impurities can radically alter macroscopic properties. When these impurities are purposely added, they are often called "dopants", and the process is called "doping". Non-crystalline compounds can also be doped.

Steel is an example of a doped crystal structure mentioned already. Another good example is aluminium oxide, also called alumina. This is a white power known for its strength and commonly used as an abrasive, such as in grinding tools, sandpaper, and toothpaste.

The most stable crystalline form of aluminium oxide is corundum. Pure corundum is transparent:

If corundum contains a fraction of Al3+ ions substituted for Cr3+, it will produce ruby:

Other colors of corundum are called sapphires. A blue sapphire is created by substituting a mixture of Fe²⁺ and Ti⁴⁺ ions into the place of Al³⁺:

Anti-site defect, coloured gold

An anti-site defect can also be called "atom interchange". It consists of an interchanged pair of atoms:

It is a common defect in metal alloys, which consist of neutral atoms, such as in CuAu above. It is very unfavourable in binary ionic compounds because it would cause similarly-charged ions to be neighbouring.

Gold, silver and copper blend together smoothly in a range of alloys varying in colour.

It is interesting that the color many people associate with gold is not the color of pure gold, since marketplace gold usually has copper and silver added, making it yellow or yellowish. 

Many ancient cultures had high copper content in their gold, so they associated it more with red than yellow.

Gold can also be made white by mixing it with about 10% nickel, palladium or manganese. This alloy is what "white gold" means in jewellers, it is unlikely to refer to a gold/silver alloy, though they might mix a small amount of silver in.

Generally a small touch of zinc is added to harden the alloy.

Pure gold is 24 carat. Mixing 14 parts of gold to 10 parts something else will create 14 carat gold, 18 to 6 creates 18K gold, and so on.