Everything from the 9 lessons of this topic on a few pages: the key idea of each lesson, the diagram to remember, the points to learn and the traps examiners set. Read it, cover it, test yourself.
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself. It increases with nuclear charge and decreases as the atomic radius and shielding increase, so it rises across a period and falls down a group; fluorine (4.0 on the Pauling scale) is the most electronegative element. When two bonded atoms have equal electronegativities the bond is pure covalent. A small difference gives a polar covalent bond with δ+ and δ− ends; a large difference (roughly 2 or more) means the electron transfer is essentially complete and the bonding is ionic.
Electronegativity is about electrons in a bond. Do not confuse it with electron affinity or ionisation energy, which describe isolated gaseous atoms. And a molecule can contain polar bonds yet be non-polar overall if the dipoles cancel by symmetry, as in CO2; that idea returns in the intermolecular forces lesson.
Ionic bonding is the electrostatic attraction between oppositely charged ions. The ions form when electrons transfer, usually from a metal to a non-metal, leaving each with a stable (often noble gas) configuration. The ions arrange themselves in a giant lattice in which every ion is surrounded by ions of opposite charge, so the attraction acts in all directions. The bonding is stronger when the ions carry larger charges and when they are smaller, because the charges get closer together; that is why MgO melts at about 2850 °C and NaCl at 801 °C.
Ionic solids conduct nothing because the ions cannot move, and it is ions, not electrons, that carry the current in a melt or solution. Writing “free electrons” for molten NaCl loses the mark. Also avoid the word “molecule” for an ionic compound: NaCl is a formula unit in a giant lattice.
Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons. The outer electrons of each atom are released into a “sea” that belongs to the whole giant lattice. The bonding gets stronger as more electrons per atom are delocalised, as the ionic charge rises and as the ions get smaller, so melting points rise from sodium to magnesium to aluminium. The mobile electrons carry current and heat, and because the bonding is non-directional the layers of ions can slide without the lattice breaking.
Say positive ions in a metal lattice, not atoms, and describe the electrons as delocalised. Answers that describe “metal atoms sharing electrons” or “ions attracting each other” do not score: the bond is ion to electron sea.
A covalent bond is the electrostatic attraction between two nuclei and a shared pair of electrons. Atoms can share one, two or three pairs to form single, double and triple bonds. In a dative covalent (coordinate) bond both electrons of the shared pair come from the same atom, which must have a lone pair, and the other atom must have an empty orbital to receive it; it is drawn as an arrow from donor to acceptor. Once formed, a dative bond is identical to any other covalent bond. Period 3 elements can use d orbitals to hold more than eight outer electrons (PCl5, SF6), while boron in BF3 has only six.
When you draw a dative bond, the arrow points from the atom that donates the lone pair to the atom that accepts it. In a dot-and-cross diagram, show both electrons of the dative pair with the donor's symbol. And remember the finished bond is not weaker or different: all four bonds in NH4+ have the same length and strength.
A σ (sigma) bond forms by direct, end-on overlap of orbitals, putting the electron density on the line between the nuclei. A π (pi) bond forms by sideways overlap of two p orbitals, putting the density above and below that line. A single bond is one σ bond; a double bond is one σ + one π; a triple bond is one σ + two π. Carbon mixes its 2s and 2p orbitals into hybrid orbitals: sp3 (four, tetrahedral, 109.5°), sp2 (three, trigonal planar, 120°, one p left for a π bond) and sp (two, linear, 180°, two p left). The π bond is weaker than the σ bond because sideways overlap is smaller.
Count σ and π bonds in a whole molecule by counting every bond as one σ, then adding one π for each double bond and two for each triple bond. Ethene, H2C=CH2, has 5 σ and 1 π, not 4 and 2. Forgetting the C–H bonds is the most common slip.
In VSEPR theory the electron pairs around a central atom repel one another and arrange themselves as far apart as possible, which fixes the shape. Count the bonding pairs (a double or triple bond counts as one region) and the lone pairs. Lone pairs are held closer to the nucleus and repel more than bonding pairs: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Each lone pair closes the remaining bond angles by roughly 2.5°, so CH4 is 109.5°, NH3 107° and H2O 104.5°. The shape is named from the positions of the atoms, not the lone pairs.
When explaining a bond angle, always name both the number of bonding pairs and the number of lone pairs, and state that lone pairs repel more than bonding pairs. “Water is bent because of its lone pairs” is not enough for the marks. Also, name the shape from the atoms: NH3 is pyramidal, not tetrahedral, even though its four pairs point towards the corners of a tetrahedron.
Intermolecular forces act between molecules and are much weaker than covalent bonds. Instantaneous dipole-induced dipole (id-id) forces act between all molecules: a momentary uneven electron distribution in one molecule induces a dipole in its neighbour. They grow with the number of electrons and with the contact area between molecules. Permanent dipole-permanent dipole (pd-pd) forces act as well between polar molecules, whose δ+ end of one attracts the δ− end of the next. A molecule with polar bonds is only polar if the bond dipoles do not cancel by symmetry. The strongest intermolecular force, hydrogen bonding, has its own lesson.
Never say that boiling a simple molecular substance “breaks covalent bonds”. It overcomes intermolecular forces. And do not assume the more polar molecule always boils higher: id-id forces grow so quickly with electron number that they often dominate, as HBr against HCl shows.
A hydrogen bond forms between a hydrogen atom bonded to a very electronegative atom (N, O or F) in one molecule and a lone pair on N, O or F in a neighbouring molecule. The H carries a large δ+ charge and is tiny, so the lone pair can approach it closely. Hydrogen bonds are the strongest intermolecular forces, which explains the unusually high boiling points of H2O, HF and NH3 and the open, less dense structure of ice. Inside molecules, bond energy is the energy needed to break one mole of a covalent bond in the gas phase; shorter bonds are generally stronger, so C≡C > C=C > C–C.
A hydrogen bond is an intermolecular force, not a covalent bond; the covalent O–H bond inside water is about 20 times stronger. When drawing hydrogen bonding you must show the lone pair and the partial charges, and the dashed line must go from the lone pair to the hydrogen, not to the other atom.
Dot-and-cross diagrams show outer-shell electrons only, using dots for one atom's electrons and crosses for the other's. For an ionic compound, draw each ion separately in square brackets with its charge, showing the electrons transferred. For a covalent molecule, place each shared pair in the overlap between atoms (one dot and one cross), two pairs for a double bond and three for a triple; draw lone pairs on the atom they belong to. In a dative bond both electrons of the shared pair carry the donor's symbol. Check the finish: every atom (except H, B and expanded-octet atoms) has eight outer electrons, and the total matches the number you started with, adjusted for any charge.
Count the total before and after. NH4+ has 5 + 4 − 1 = 8 outer electrons, which is four pairs and no lone pairs. The most common mistakes are leaving the lone pairs off the outer atoms, forgetting the square brackets and charge on an ion, and drawing a dot and a cross in a dative bond where both electrons should come from the donor.