Chemical Bonding

In NaCl, sodium loses one electron and chlorine gains one, forming Na⁺ and Cl^-. In MgO, magnesium loses two and oxygen gains two, forming Mg²⁺ and O^2-. Both products have full outer shells and zero overall charge. The formula gives an ion ratio; neither solid consists of isolated two-ion molecules.

An ionic lattice is held by many attractions in all directions. Larger charge magnitudes and smaller interionic distances strengthen electrostatic attraction. MgO therefore has much stronger lattice attraction than NaCl: its ions have charges of magnitude two rather than one. Compare the whole structure rather than imagining a single Na-Cl covalent bond.

Dots and crosses track which atom supplied each electron; they do not identify different kinds of electron. A shared pair contributes to the outer-shell count of both atoms. Hydrogen needs a duet in its first shell; the C, N, O and Cl examples below reach octets. The diagram counts electrons and does not necessarily show molecular shape.

In a co-ordinate (dative covalent) bond, one atom supplies both electrons of the shared pair. The donor must have a lone pair, and the acceptor must have an available orbital. An arrow points from donor to acceptor when showing bond formation. It does not mean that the finished bond contains a different kind of electron.

AlCl₃ is electron deficient in its simple monomer description: Al has three shared pairs, only six electrons around it. Two units can associate to make Al₂Cl₆. A chlorine lone pair from each unit donates to the other aluminium. Two chlorine atoms bridge the Al centres, and each Al then has four shared pairs.

A sigma (σ) bond has electron density along the line joining the nuclei. It forms by end-on overlap: s-s in H₂, s-p in a simple H-Cl description, or p-p. A pi (π) bond forms by sideways overlap of parallel p orbitals, producing electron density above and below that line. The two pi regions together constitute one bond.

A carbon-carbon single bond can usually rotate without destroying its sigma overlap. Rotation about a double bond requires loss of the pi overlap and is restricted. Bond rotation is therefore not the same as rotating the entire molecule in space. Detailed d-orbital overlap is outside the overlap requirement here.

VSEPR places regions of electron density around a central atom as far apart as possible. A single, double or triple bond counts as one region for this purpose. A lone pair is another region. The electron-region arrangement includes lone pairs; the molecular shape names the positions of the atoms only.

A lone pair is held by only one nucleus and occupies more space near the central atom than a bonding pair. The usual repulsion order is lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair. The tetrahedral electron arrangement therefore gives progressively smaller bond angles from CH₄ to NH₃ to H₂O. Exact angles also depend on the atoms and bonds: do not assign water's angle to every bent molecule.

Why can sulfur have more than eight outer electrons when oxygen, in the same group, cannot? 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 regions. A Period 2 atom such as N, O or F has only 2s and 2p orbitals, so it is limited to four regions 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 regions, 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 regions 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.

Electronegativity is an atom's ability to attract the shared electrons of a covalent bond. Unequal attraction creates partial charges, δ+ and δ-. In H-Cl, chlorine attracts the pair more strongly, so H is δ+ and Cl is δ-. These are partial charges within a covalent bond, not a declaration that HCl is an ionic lattice.

At any instant an electron cloud may be uneven, creating an instantaneous dipole. It induces a dipole in a neighbour, and opposite partial charges attract. These instantaneous dipole-induced dipole attractions occur in all atoms and molecules, including non-polar Br₂ and the monatomic liquid noble gases. A larger, more easily distorted electron cloud is generally more polarizable and gives stronger attractions.

Permanent dipole-dipole attractions also occur between polar molecules, such as CHCl₃. Favourable orientations bring a δ+ region near a δ- region. CHCl₃ still has instantaneous-dipole attractions as well; do not choose exactly one force for every substance.

Hydrogen bonding is a particularly strong, directional interaction involving H covalently bonded to N, O or F and a suitable lone pair on N, O or F of another molecule (or another part of the same molecule). In water, O-H groups are donors and oxygen lone pairs are acceptors. In ammonia, N-H groups donate and the nitrogen lone pair accepts.

In ordinary ice, water molecules form an open hydrogen-bonded network. Each water molecule can donate two hydrogen bonds through its O-H groups and accept two through oxygen lone pairs. The roughly tetrahedral arrangement leaves more empty space than in liquid water.

On melting, enough of the network is disrupted for molecules to pack more closely. Ice is therefore less dense than liquid water and floats. Liquid water still contains hydrogen bonds: melting does not remove every attraction, and it does not break water into H and O atoms.

Ammonia also hydrogen-bonds, but its available lone-pair sites and the strength and extent of the network differ from water. Do not explain every boiling-point comparison by molar mass alone, or claim that any molecule containing hydrogen can hydrogen-bond.

Bond length is the equilibrium distance between two bonded nuclei. Bond energy is the energy required to break one mole of specified covalent bonds in gaseous species. A mean bond energy averages the same bond type across compounds; its use in a calculation is approximate, not an exact measurement for every molecule.

For bonds between the same elements, higher bond order generally means a shorter, stronger bond: C≡C is shorter and has greater total bond energy than C=C, which is shorter and stronger than C-C. But the pi part of a multiple bond can react while the sigma framework remains. A large total double-bond energy does not make an alkene generally less reactive than an alkane.

Crystalline means a regular, repeating arrangement; it does not mean that every crystalline material is ionic. The unit-cell construction is not required. Electrical conductivity needs mobile charged particles, so identify whether the carrier is an electron or ion and whether it can move in the stated physical state.

Solubility depends on the balance of attractions broken and formed, as well as disorder, not a universal "ionic means soluble" rule. Some ionic compounds dissolve very little. Molecular substances capable of strong interactions with water can dissolve; non-polar substances often dissolve better in non-polar solvents.