IB HL chemistry: 2.10-20; all bonding HL topics

Explain résonance structures

When a molecule has multiple valid VSEPR structures, those are resonance structures

• the actual structure is intermediate between the pair.

• It’s called a resonance hybrid

• Bond length/order and bond density is an average of the two structures

• There must be at least a PI bond present, as that the delocalised electron which goes from one bond to the other.

• Occurs when there are adjacent atoms of equivalent electro negativity

Evidence for resonance structure of benzene

• all bond angles and lengths are the same

• Enthalpy of hydrogenation is lower than expected ( only 210 instead of 360). More stable bc of pi bonds

• Saturation tests; it didn’t react as alkene would react; didn’t turn bromide water colourless as. Doesn’t undergo addition reaction with bromide p, but substitution

Expanded octets

• only smaller atoms, period 3 or above

• The energy level of 3p and 3d is very similar

Higher VSEPR

Fuck VSEPR - use drawn flash cards or kognity

Formal charge formula

Number of valence electrons - number of lone electrons - nr of electrons in bonding pairs

• The closer to 0 the more stable

• Can happen with coordination bond

• For an ion net formal charge has to be same as overall chance and for a molecule it has to be 0

Sigma bond

Head on overlap of orbitals; ie s-s, p-p, s-p.

Single bonds are Always sigma

Stronger than pi, as there is only one electron dense region on the inter nuclear axis

Pi bond

• sideways overlap of p orbitals

• Different Axial plane than sigma bonds as the electrons repel each other

• These are all the extra bonds after a single bond

• They are weaker than sigma bonds and are first to break, as they are located further away from the positive charge of the nucleus

• If there are multiple Pi orbitals close to each other, the electrons are delocalised and can move through there

Excitation of electron

Moving from ground state to excited state by energy absorption

Hybridisation

• mix of atomic orbitals to produce hybrid orbitals of intermediate energy; ie between S and between P

• total energy hasn’t changed -, but has been equally redistributed

carbon SP3 hybridisation

one electron from S orbital and 3 from P blend together

only single bonds as no P orbitals left

carbon SP2 hybridisation

one electron from S orbital and 2 from P blend together

double bonds can be formed as a P is left

carbon SP hybridisation

one electron from S orbital and one from P blend together

triple bonds can be formed as there still are 2 P orbitals left