Chemical Bonding

Electron transfer explains ion formation; attraction between the resulting ions is the ionic bond. Sodium transfers one outer electron to chlorine, making Na⁺ and Cl^-. Magnesium transfers two to oxygen, making Mg²⁺ and O^2-. Each compound is neutral overall and extends as a lattice, not a collection of NaCl or MgO molecules.

In a covalent bond, electron density between nuclei attracts both nuclei. A shared pair can come one electron from each atom or both from one atom. Dots and crosses record origin; electrons themselves are identical. In a metal, electrons are delocalised through the structure, so attractions are not restricted to one fixed atom pair.

Count the total outer electrons first. A single bond uses one shared pair, a double bond two pairs, and a triple bond three. Hydrogen has a duet; the other simple molecules below achieve octets. An electron shared in a bond counts towards both bonded atoms, but is counted only once in the total-electron audit.

A co-ordinate or dative covalent bond contains a shared pair supplied entirely by one atom. NH₃ has a lone pair on N; H⁺ has an empty orbital and no electron to contribute. Donation forms NH₄⁺. After formation all four N-H bonds are equivalent; the origin of the pair does not create a permanently weaker bond.

In gaseous AlCl₃ under suitable conditions, an Al atom has only six electrons in its three covalent bonds. Two units join to form Al₂Cl₆: a Cl from each unit donates a lone pair to the electron-deficient Al in the other unit. Two bridging Cl atoms result; each Al is surrounded by four shared pairs.

A sigma bond forms by head-on overlap along the line joining the nuclei. It can use s-s overlap (H₂), s-p overlap (HCl) or p-p overlap. A pi bond forms by sideways overlap of parallel p orbitals, with electron density above and below the internuclear axis. A single bond is sigma; a double bond is one sigma plus one pi; a triple bond is one sigma plus two pi.

VSEPR places electron domains as far apart as possible to reduce repulsion. Around a central atom, a single, double or triple bond counts as one bonding domain. Lone pairs also repel, but are omitted from the name of the molecular shape. Lone-pair/bond-pair repulsion is greater than bond-pair/bond-pair repulsion, compressing some angles.

The table follows the same repulsion argument throughout: two bonding domains point in opposite directions (linear), three spread in one plane (trigonal planar), and four point towards the corners of a tetrahedron. Six equivalent domains point along three perpendicular axes (octahedral). NH₃ and H₂O both start from four electron domains. Replace one tetrahedral bonding direction by a lone pair to obtain a pyramidal atom arrangement; replace two to obtain a bent arrangement. Lone pairs occupy more angular space near the central atom, so they compress the remaining bond angles.

Sulfur in SF₆ has twelve outer electrons, but oxygen in the same group never goes past eight. An element in Period 3 or below has empty d orbitals in its outer shell (3d for Period 3). They are close enough in energy to be used, so the central atom can hold five or six electron domains. A Period 2 atom such as N, O or F has only 2s and 2p orbitals, so it is limited to four domains and eight electrons. That is why PCl₅ and SF₆ exist, but NCl₅ and OF₆ do not. This is called expanding the octet.

Where the lone pairs go decides the shape. With five domains, the three equatorial positions are 120 degrees apart, while the two axial positions are at 90 degrees to three others. A lone pair takes an equatorial position, where it has fewer close neighbours at 90 degrees. So one lone pair gives a see-saw, two give a T-shape and three give a straight line. With six domains every position is the same, so one lone pair gives a square pyramid. A second lone pair goes opposite the first, as far from it as possible, which leaves a square plane.

Ethene contains five sigma bonds in total (four C-H and one C-C) plus one C-C pi bond. Sideways overlap requires the two p orbitals to remain aligned, so the double bond cannot rotate freely without disrupting that overlap. This explains later cis-trans isomerism.

Electronegativity describes an atom's attraction for a bonding electron pair. In H-Cl, Cl attracts the pair more strongly: H is partially positive and Cl partially negative. These are partial charges within a covalent bond, not full H⁺ and Cl^- ions. The molecular dipole is the combined effect of all bond dipoles in three dimensions.

Instantaneous dipole-induced dipole attractions occur in all particles: a momentary uneven electron distribution induces a dipole in a neighbouring particle, and the nearby opposite partial charges attract electrostatically. They are the principal intermolecular attraction in Br₂ and the attraction between noble-gas atoms. Larger, more easily polarised electron clouds usually give stronger attractions. CHCl₃ also has permanent dipole-permanent dipole attraction between oppositely charged ends of neighbouring molecules.

Hydrogen bonding is an electrostatic attraction between the partially positive H covalently bonded to N, O or F and a lone pair on N, O or F in a neighbouring particle. Water provides O-H donors and oxygen lone pairs; ammonia provides N-H donors and a nitrogen lone pair. Merely containing hydrogen is insufficient: methane cannot hydrogen-bond to itself.

The same possible attractions exist in a gas and its liquid, but gas particles are usually far apart. Cooling reduces kinetic energy and compression brings particles closer, allowing attractions to hold particles in a liquid. A larger attraction generally makes liquefaction easier. Pressure alone does not guarantee liquefaction at every temperature; this is a qualitative explanation, not an ideal-gas calculation.

Hydrogen bonding makes water's boiling point relatively high for such a small molecule because substantial energy is needed to separate molecules. In ordinary ice, each water molecule hydrogen-bonds to four neighbours in an open, approximately tetrahedral network: its two O-H groups donate two hydrogen bonds and its two oxygen lone pairs accept two. On melting, some of that open arrangement collapses, so liquid water is denser than ice and ice floats. The covalent O-H bonds remain intact.

Bond length is the equilibrium distance between the nuclei of two bonded atoms. Bond energy is the energy required to break one mole of a specified covalent bond in gaseous species; tabulated average bond energies average over molecular environments. Bond breaking is endothermic. For similar bonds, greater length usually means poorer overlap and a weaker bond.

Compare C-Cl, C-Br and C-I: the halogen atom becomes larger down the group, so the bond becomes longer and generally weaker. Breaking C-I is easier than breaking C-Cl, despite C-Cl being more polar. Polarity instead identifies electron-poor and electron-rich sites, which influences how a reagent approaches and interacts with a molecule.

For melting, identify the attraction that must be overcome. For conductivity, identify mobile charged particles. For solubility, consider whether new solute-solvent attractions can compensate for attractions disrupted in both substances; "ionic" is not a guarantee of water solubility. For malleability, explain whether a displaced layer remains held together or brings like charges into unfavourable proximity.