Periodicity and Bonding

• Generally, the bonds become weaker and longer down a group • This is due to less of the valence orbital volume overlaps and thus sharing of e- density is less effective

Larger valence orbital volume but only ever 2e- per orbital, so more diffuse orbitals with lower charge density going down a group

• 2p-π orbitals (used in multiple bonding) where there is good overlap - due to small covalent radius of E • 3p-π orbitals used with larger covalent radius of E - poor p-π overlap and single bonds favoured - too big essentially to form multiple bonds

• Si is a larger atom, hence has larger valance orbitals BUT there is also a bigger difference in electronegativity • Silicon is not very electronegative compared to carbon • So there is a stronger electrostatic interaction between silicon and oxygen • This means that σ-bonds are favoured for silicon compounds over π

pπ-pπ overlap

• Electron donation from a filled p-orbtial on the oxygen, to an vacant d-orbtial on the phosphrous • The phosphorus already has 4 bonds but can expand it octet using the d-orbitals • Resulting in pπ-dπ bonding

• Generally believed that planarity of N(SiH₃)₃ comes from a pπ-dπ interaction • Due to overlap of N p-orbtials with d/σ’ orbitals on Si + enhanced steric factors = trigonal planar • (NMe₃ cannot do the same due to no d-orbitals)

This is the number of valence electrons used by an atom to form bonds to other things

Halogens have 7 valence e- Hence will form 1 bond to give a full valance ocetet Heavier elements can expand their ocetet however

Will form two bonds to give a full valence octet Heavier elements can expand their ocetet

3 bond gives the valance octet (3 bond pairs and 1 lone pair - lewis bases) Heaier elements again can expand their octet

4 bonds gives full valance octet Multiple bonding common for C (can be π or σ, so many geometries possible) Pi-bonding less common down the group

• only 3 valance electrons, so can only have a valance of 3 bonded pairs giving only 6 electrons • So the triel compounds are electron deficient (Lewis acids)

The second period is an exception due to greater electrostatic repulsion Which is pushing the atoms apart, hence making the end weaker For smaller atoms, N to F, with more charge-dense orbitals

H-H is an exception despiting having high electronegativity Smaller orbitals, good overlap, and effective sharing of electron density There is no electrostatic repulsion (only 2e- present) = very strong bond

Going from left to right, the atom gets smaller and more electronegative (from 1-14) Leading to more effective orbital overlap and sharing of electron density, hence stronger bonds

• Small orbtials and effective overlap meaning good sharing of electron density • The bond is also polar covalent so there is additional electrostatic attraction • There is also pi, donation from lone pairs of F to empty p-orbitals on B is possible

• Al is a larger atom, hence the valence orbitals will be more diffuse • Therefore will have poorer sharing of electron density

• In-F has an O.S. of (iii) • While Tl-F has an O.S of (i) • Due to the inert pair effect • BUT due to the big difference in electronegativity between Tl and F, the bonding will be more ionic which weakens the strength of the bond (hence not really a fair comparision)

• C-F bond is weaker than the B-F bond • There is a smaller difference in electronegativity going left to right (lower ΔΧ) • Hence reduced electrostatic interaction • (no pi-donation can occur for these period 2 elements as all C orbitals are filled)

• Si-F stronger than C-F • Larger difference in electronegativity for Si-F (higher ΔΧ) meaning increased electrostatic interaction • Silicon has the ability of expanding it’s ocetet due to having d-orbitals, so can undertake pπ-dπ bonding (from F LPs)

• P-F stronger than N-F • N is highly electronegative, meaning the difference in electronegativity is lower (smaller ΔΧ) and decreased electrostatic interaction • N and F have LPs and are smaller so there will also be increased electrostatic repulsion • AND phosphorus has d-orbitals available for pπ-dπ bonding

• P-F is stronger than S-F • There is a lower difference is electronegativity for S-F (lower ΔΧ), meaning decreased electrostatic interactions • S and F both have multiple LPs and are small = increased electrostatic repulsion greater than for P

• Enthalpy of formation is negative because energy is released through forming bonds • Generally there is a decreases as we go down the group = less exothermic as the bonds get weaker • This is due to larger valance orbitals and hence weaker bonds • Apart from C-Cl as carbon is weakest due to smaller difference in electronegativity - less electrostatic interactions