Chemical bonding

📚The chemistry

3.1 Electronegativity

Electronegativity is the power of an atom to attract electrons to itself — specifically, the pair of electrons in a covalent bond. It depends on the same three things as ionisation energy: nuclear charge (more protons, more pull), atomic radius (a bonding pair closer to the nucleus is pulled harder) and shielding by inner shells and sub-shells.

• Across a period it increases: nuclear charge rises, radius falls, and shielding hardly changes. Fluorine is the most electronegative element of all (4.0 on the Pauling scale).
• Down a group it decreases: each element has an extra shell, so the bonding pair is further away and more shielded.

The difference in electronegativity between two bonded atoms predicts the bond type. No difference gives a non-polar covalent bond (Cl–Cl); a moderate difference gives a polar covalent bond (H–Cl, 3.0 − 2.1 = 0.9); a large difference — as a working rule, more than about 1.8 — gives an ionic bond (NaCl, 3.0 − 0.9 = 2.1). Pauling values are given in the data section and in the question when you need them.

3.2–3.3 Ionic and metallic bonding

Ionic bonding is the electrostatic attraction between oppositely charged ions. Electrons transfer from metal to non-metal, and in the solid each ion is surrounded by ions of the opposite charge in a giant lattice. In sodium chloride, Na (2,8,1) gives one electron to Cl (2,8,7), forming Na⁺ and Cl⁻, each with a noble-gas arrangement. Magnesium oxide transfers two electrons, forming Mg²⁺ and O²⁻; the doubled charges make the attraction far stronger, which is why MgO melts at about 2850 °C and NaCl at 801 °C. In calcium fluoride each Ca gives one electron to each of two F atoms: Ca²⁺ and 2F⁻.

Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons. The outer electrons are not attached to any one atom but move through the whole lattice. The more electrons each atom contributes, and the smaller and more highly charged the ions, the stronger the bonding.

3.4 Covalent and coordinate bonding

A covalent bond is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. Both nuclei attract the same pair, and that is what holds them together. The syllabus names H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆ and C₂H₄; you should be able to draw every one. O₂ and C₂H₄ contain double bonds, N₂ a triple bond, CO₂ two double bonds.

Expanded octets. Elements in period 3 and beyond have empty 3d orbitals close in energy to 3s and 3p, so they can use more than four electron pairs around the central atom. Sulfur in SO₂ has two S=O double bonds and a lone pair (ten electrons around S). Phosphorus in PCl₅ has five bonding pairs; sulfur in SF₆ has six. Period 2 elements cannot do this: nitrogen never forms NCl₅, because it has no 2d sub-shell.

A coordinate (dative covalent) bond is a covalent bond in which both shared electrons come from the same atom. The donor needs a lone pair; the acceptor needs an empty orbital. Two examples are named:

• NH₃ + HCl → NH₄Cl. The lone pair on nitrogen is donated to an H⁺, forming NH₄⁺. Once formed, the four N–H bonds are identical — you cannot tell which was the dative one.
• Al₂Cl₆. In AlCl₃ the aluminium has only six outer electrons. At lower temperatures two molecules join: a lone pair on a chlorine of each molecule is donated to the aluminium of the other, forming two Cl→Al dative bonds that bridge the two Al atoms.

Draw a dative bond as an arrow from donor to acceptor.

σ and π bonds, and hybridisation

A σ (sigma) bond forms by direct, end-on overlap of orbitals along the line between the two nuclei. Every single bond is a σ bond. A π (pi) bond forms by the sideways overlap of adjacent p orbitals, giving electron density above and below the σ bond. A double bond is one σ plus one π; a triple bond is one σ plus two π.

The π bond is the weaker of the two, because sideways overlap is less effective. The data section shows it: C=C is 610 kJ mol⁻¹, less than twice C–C (2 × 350 = 700). That weakness is why alkenes react by breaking their π bond.

Hybridisation explains the geometry. The s and p orbitals of the central atom mix to form a set of equivalent hybrid orbitals:

• sp³ — one s + three p → four hybrids pointing to the corners of a tetrahedron, 109.5° apart. Carbon in CH₄ and C₂H₆: four σ bonds.
• sp² — one s + two p → three hybrids in a plane, 120° apart, with one unhybridised p orbital at right angles. Carbon in C₂H₄: three σ bonds, and the leftover p orbitals on the two carbons overlap sideways to form the π bond.
• sp — one s + one p → two hybrids at 180°, with two unhybridised p orbitals. Carbon in HCN: two σ bonds (to H and N) and two π bonds to N. N₂ is the same: one σ and two π.

Bond energy, bond length and reactivity

Bond energy is the energy required to break one mole of a particular covalent bond in the gaseous state. Bond length is the internuclear distance between two covalently bonded atoms. In general, the shorter the bond, the stronger it is, and the stronger the bonds in a molecule, the less reactive it is.

• N₂ has a triple bond of 944 kJ mol⁻¹. That is why nitrogen is so unreactive, and why the Haber process needs an iron catalyst and high temperature.
• The hydrogen halides get longer and weaker down the group: H–Cl 431, H–Br 366, H–I 299 kJ mol⁻¹. HI is the easiest to decompose into its elements; a hot wire will do it.

3.5 Shapes of molecules

VSEPR (valence shell electron pair repulsion) theory: electron pairs around a central atom repel each other and arrange themselves as far apart as possible. The method is always the same:

1. Count the outer electrons on the central atom; add one for each negative charge, subtract one for each positive charge.
2. Work out how many are used in bonds and how many are left as lone pairs.
3. Count the regions of electron density. A double or triple bond counts as one region, like a single bond.
4. Arrange the regions as far apart as possible, then name the shape from the positions of the atoms only.
5. Adjust the angle: lone pair–lone pair repulsion > lone pair–bond pair > bond pair–bond pair, so each lone pair squeezes the bond angle by about 2.5°.

The seven named shapes:

You are expected to predict the shapes of analogous species: NH₄⁺ is tetrahedral like CH₄; PH₃ is pyramidal like NH₃; H₂S is non-linear like H₂O; SiCl₄ is tetrahedral.

3.6 Intermolecular forces

Cambridge uses van der Waals’ forces as the generic name for all intermolecular forces — the forces between molecules other than those from bond formation. There are two types in this syllabus:

• Instantaneous dipole–induced dipole (id-id), also called London dispersion forces. Electrons are always moving, so at any instant the electron cloud may be uneven, creating a temporary dipole. That induces a dipole in a neighbouring molecule, and the two attract. They act between all molecules. They get stronger with more electrons (bigger molecules) and with more surface contact: pentane (a straight chain) boils at 36 °C, its branched isomer 2,2-dimethylpropane at 10 °C.
• Permanent dipole–permanent dipole (pd-pd), between polar molecules. The δ+ end of one molecule attracts the δ− end of the next.

Hydrogen bonding is a special, strong case of pd-pd attraction. It happens when hydrogen is bonded to a highly electronegative atom — in this syllabus, N–H and O–H groups. The H carries a large δ+ and is attracted to a lone pair on an N or O atom of another molecule. Draw it with the lone pair shown and the O–H···O atoms in a straight line.

Ionic, covalent and metallic bonds are, in general, much stronger than any intermolecular force. Boiling a molecular liquid breaks intermolecular forces, not covalent bonds: water boils to H₂O molecules, not to hydrogen and oxygen.

Polarity and dipoles

A bond between atoms of different electronegativity is polar: the more electronegative atom is δ−. Whether the molecule has a dipole depends on its shape. If polar bonds are arranged symmetrically, their dipoles cancel: CO₂ (linear) and CCl₄ (tetrahedral) have polar bonds but no overall dipole. H₂O (non-linear), NH₃ (pyramidal) and CHCl₃ (one H among three Cl) are polar molecules.

Why water is strange

Each water molecule has two H atoms and two lone pairs, so it can form up to four hydrogen bonds. That explains the three anomalies the syllabus lists:

• High melting and boiling points. H₂O boils at 100 °C, while H₂S, a bigger molecule with more electrons, boils at about −60 °C. Extra energy is needed to break the hydrogen bonds.
• High surface tension. Molecules at the surface are pulled inwards by hydrogen bonds to the molecules below.
• Ice is less dense than liquid water. In ice, every molecule forms four hydrogen bonds in a tetrahedral arrangement, holding the molecules in an open lattice with a lot of empty space. On melting, some hydrogen bonds break and the molecules pack more closely. So ice floats.

3.7 Dot-and-cross diagrams

Show the electrons of one atom as dots and the other as crosses, outer shells only unless asked. For ionic compounds draw separate ions in square brackets with the charge outside, e.g. [Na]⁺ and [Cl]⁻ with eight electrons, seven dots and one cross. For covalent molecules put the shared pairs in the overlap of the circles and show every lone pair. For a dative bond the shared pair is two of the same symbol.

You may be asked for species with an expanded octet (PCl₅, SF₆, SO₂) or an odd number of electrons, such as NO or NO₂, where one atom ends up with an unpaired electron — a free radical, as in topic 1.