Across a period (left to right), atomic radius decreases because nuclear charge increases while electrons are added to the same shell, so the increased effective nuclear charge pulls the electron cloud closer. Down a group, atomic radius increases because a new shell is added at each step and the number of inner shielding electrons increases, outweighing the rise in nuclear charge.

Nitrogen has the configuration 1s2 2s2 2p3 with a half-filled 2p subshell, which is extra stable due to symmetry and exchange energy, so removing an electron requires more energy. Oxygen (1s2 2s2 2p4) has one paired electron in 2p; removing it relieves electron-electron repulsion and gives a stable half-filled configuration, so less energy is needed. Hence IE1 of nitrogen > oxygen.

Ionization enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. Across a period it increases because nuclear charge increases and atomic size decreases, so the electron is held more tightly. Down a group it decreases because atomic size increases and shielding by inner electrons increases, so the outer electron is more easily removed.

M(g) -> M+(g) + e-

Although fluorine is smaller, its 2p subshell is very compact, so the incoming electron experiences strong inter-electronic repulsion from the already crowded 2p electrons. In chlorine, the larger 3p subshell accommodates the incoming electron with less repulsion, so more energy is released. Hence chlorine has a more negative (more exothermic) electron gain enthalpy than fluorine.

Electronegativity is the tendency of an atom in a molecule to attract the shared pair of electrons in a covalent bond towards itself. It increases across a period (as size decreases and nuclear charge increases) and decreases down a group (as size increases). Fluorine is the most electronegative element. The Pauling scale is commonly used.

All four species have 10 electrons (isoelectronic). For isoelectronic species, the greater the nuclear charge (number of protons), the smaller the radius. Nuclear charges: O2- (8), F- (9), Na+ (11), Mg2+ (12). Increasing order of radius: Mg2+ < Na+ < F- < O2-.

Elements are classified by the subshell into which the last (differentiating) electron enters: s-block (last electron in s), p-block (in p), d-block (in d) and f-block (in f). d-block elements are called transition elements because they lie between the s-block (metals) and p-block (non-metals) and show a gradual transition in properties; their atoms or common ions have partially filled d orbitals.

An ionic (electrovalent) bond is formed by the complete transfer of one or more electrons from one atom to another, producing oppositely charged ions held by electrostatic attraction (e.g. NaCl). A covalent bond is formed by the mutual sharing of electron pairs between atoms (e.g. H2 or Cl2). Ionic bonds form between metals and non-metals; covalent bonds form between non-metals.

BF3: central B has 3 bond pairs and no lone pair -> trigonal planar, bond angle 120 deg. NH3: central N has 3 bond pairs and 1 lone pair -> pyramidal (trigonal pyramidal), bond angle about 107 deg. H2O: central O has 2 bond pairs and 2 lone pairs -> bent/angular, bond angle about 104.5 deg. Lone pairs repel more than bond pairs, reducing the angle from the ideal 109.5 deg in NH3 and H2O.

Both molecules are based on sp3 hybridisation with an ideal angle of 109.5 deg. NH3 has one lone pair, while H2O has two lone pairs. Lone pair-lone pair repulsion is stronger than lone pair-bond pair repulsion, and water has an extra lone pair, so its bond pairs are pushed closer together. Hence the bond angle of H2O (104.5 deg) is smaller than that of NH3 (107 deg).

In ethyne each carbon is sp hybridised. The two sp hybrid orbitals on each carbon form sigma bonds: one C-C sigma bond and one C-H sigma bond, giving a linear molecule with a bond angle of 180 deg. The remaining two unhybridised p orbitals on each carbon overlap sideways to form two pi bonds, so there is a triple bond (1 sigma + 2 pi) between the carbon atoms.

H-C(triple bond)C-H

O2 has 16 electrons. Configuration: sigma(1s)2 sigma*(1s)2 sigma(2s)2 sigma*(2s)2 sigma(2pz)2 pi(2px)2 pi(2py)2 pi*(2px)1 pi*(2py)1. Bonding electrons Nb = 10, antibonding Na = 6. Bond order = (Nb - Na)/2 = (10 - 6)/2 = 2. The two unpaired electrons in the pi* antibonding orbitals make O2 paramagnetic.

Bond order = (Nb - Na)/2

N2 has 14 electrons. Number of bonding electrons Nb = 10 and antibonding electrons Na = 4. Bond order = (Nb - Na)/2 = (10 - 4)/2 = 3. This corresponds to a nitrogen-nitrogen triple bond, which explains the very high stability and bond dissociation energy of N2.

Bond order = (Nb - Na)/2

A coordinate bond is a covalent bond in which the shared pair of electrons is contributed by only one of the two bonded atoms (the donor). In the ammonium ion (NH4+), nitrogen in NH3 has a lone pair which it donates to a proton (H+) that has no electrons, forming the fourth N-H bond as a coordinate bond. Once formed, all four N-H bonds are identical.

NH3 + H+ -> NH4+

A hydrogen bond is a weak electrostatic attraction between a hydrogen atom covalently bonded to a highly electronegative atom (F, O or N) and the lone pair of another electronegative atom. Conditions: hydrogen must be attached to a small, highly electronegative atom. Types: intermolecular hydrogen bonding (between different molecules, e.g. in water and HF) and intramolecular hydrogen bonding (within the same molecule, e.g. o-nitrophenol).

Oxygen is much more electronegative and smaller than sulphur, so water molecules form strong intermolecular hydrogen bonds, whereas H2S molecules are held only by weak van der Waals (dipole) forces. Extra energy is needed to break the hydrogen bonds in water, so water has a much higher boiling point than H2S even though H2S has a higher molar mass.

Fajans' rules: covalent character in an ionic compound increases when (1) the cation is small, (2) the anion is large, and (3) the cation has a high charge (all three increase the polarising power/polarisability). Also, cations with a pseudo-noble gas (18-electron) configuration cause greater polarisation than those with a noble gas configuration. Greater polarisation of the anion by the cation increases covalent character.

Dipole moment (mu) is the product of the magnitude of charge and the distance between the centres of positive and negative charge; it is a vector quantity (unit: debye). CO2 is linear (O=C=O) and its two C=O bond dipoles are equal and opposite, so they cancel and the net dipole moment is zero. H2O is bent/angular, so its two O-H bond dipoles do not cancel and add up to give a net dipole moment (1.85 D). Hence CO2 is non-polar but H2O is polar.

mu = q x d

A sigma bond is formed by the head-on (axial) overlap of orbitals along the internuclear axis, while a pi bond is formed by the sidewise (lateral) overlap of parallel p orbitals. A sigma bond is stronger because axial overlap is more effective and greater, giving a larger region of electron density between the nuclei; pi bonds have smaller lateral overlap and are weaker and more reactive.

Boyle's law: at constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure, i.e. V is proportional to 1/P, so PV = constant. Charles's law: at constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute (Kelvin) temperature, i.e. V is proportional to T, so V/T = constant.

PV = constant (Boyle); V/T = constant (Charles)