Bonding and structure — OCR A-Level Chemistry

Ionic bonding is the electrostatic attraction between positive and negative ions, formed when electrons are transferred from a metal to a non-metal so that each ion gains a stable noble-gas configuration. Sodium loses one electron to chlorine: Na gives Na⁺ and Cl gains an electron to give Cl⁻, and the oppositely charged ions attract in a giant ionic lattice. The strength of the attraction depends on the charges and the distance between the ion centres, so a 2⁺ ion beside a 2⁻ ion attracts more strongly than a 1⁺ beside a 1⁻ at the same separation. In multiple-choice questions you must recognise the definition, predict the ions formed from a formula or electron configuration, and distinguish ionic attraction from covalent sharing.

This statement covers two linked ideas. First, ionic bonding is the electrostatic attraction between positive and negative ions, produced when a metal atom transfers electrons to a non-metal atom. Second, you must be able to construct and interpret 'dot-and-cross' diagrams that show this transfer. In such a diagram the outer-shell electrons of one atom are drawn as dots and those of the other as crosses, so the electron gained by the non-metal appears as a cross inside its outer shell and the ion formed is enclosed in square brackets with its charge shown. For example, in sodium chloride the sodium atom loses its single outer electron, drawn as a cross, which appears in chlorine's outer shell, giving Na⁺ and Cl⁻. Multiple-choice items test whether you can read the diagram, count electrons correctly and identify the ions and their charges.

A giant ionic lattice forms when oppositely charged ions pack in a regular three-dimensional array, as in sodium chloride where Na⁺ and Cl⁻ alternate. Each ion is surrounded by ions of opposite charge, and the electrostatic attraction acts strongly in all directions, not just towards one neighbour. This extended attraction gives high melting and boiling points because many strong attractions must be overcome. The lattice is brittle: when a layer is displaced, ions of like charge are brought alongside each other, and the resulting repulsion splits the crystal. Solid lattices do not conduct because the ions are held in fixed positions, but the melt or aqueous solution conducts as ions become mobile.

Ionic compounds have high melting and boiling points because strong electrostatic attractions throughout the giant lattice must be overcome, and this takes a large amount of energy. Many ionic solids dissolve in polar solvents such as water, where ion–dipole interactions between water molecules and ions help to separate the lattice. Ionic solids do not conduct electricity because the ions are held in fixed positions and cannot move. When molten or dissolved in water, the ions become free to move and carry charge, so the liquid and aqueous states conduct. Conductivity depends on ions being mobile, not on electrons.

Covalent bonding is the electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms. Each atom contributes one electron to the shared pair, giving both atoms a full outer shell in most simple molecules such as H₂, Cl₂, H₂O and CH₄. A coordinate (dative covalent) bond forms when one atom donates both electrons of the shared pair, as in the H₃N→BF₃ adduct or the H₃O⁺ ion, so a shared pair need not come from one electron per atom. Learners should draw dot-and-cross diagrams, count shared pairs to give single, double or triple bonds, and explain why molecules such as CO₂ contain two double bonds while N₂ contains a triple bond.

A covalent bond is defined as the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms. The two nuclei are both positively charged, and the shared pair sits between them, so each nucleus is attracted to the same pair. This attraction holds the atoms together and makes the bond strong. The definition has three parts to check: the attraction is electrostatic, it is between the shared pair and the nuclei (plural), and it is strong. In H₂, one pair is shared between two hydrogen nuclei; in Cl₂, one pair is shared between two chlorine nuclei. A correct answer names the shared pair, both nuclei and the electrostatic attraction.

A dot-and-cross diagram shows the outer-shell electrons of each bonded atom, using dots for one atom's electrons and crosses for the other's, so shared pairs sit in the overlap. To construct one, count valence electrons for each atom, pair up the electrons needed to fill outer shells, then place any remaining electrons as lone pairs on the correct atom. For an ion, add one electron for each negative charge or remove one for each positive charge before drawing, and enclose the whole diagram in square brackets with the charge shown outside. The diagram then describes which electrons are shared and which are lone pairs, and shows the ion's overall charge.

A single covalent bond forms when two atoms share one pair of electrons, one electron from each atom, so that both atoms gain a filled outer shell. In a dot-and-cross diagram the shared pair is drawn once, between the nuclei, with one dot and one cross to show each atom's contribution. Each shared pair counts as one bond, so a single bond is one shared pair. Lone pairs are outer-shell pairs not involved in bonding and are drawn on the atom that holds them. The number of single bonds an atom forms depends on how many electrons it needs to complete its outer shell, which is why hydrogen forms one bond and carbon forms four.

Multiple covalent bonding occurs when two atoms share more than one pair of electrons, forming double bonds (two shared pairs) or triple bonds (three shared pairs). Each shared pair contributes one electron from each bonded atom, so a double bond contains four shared electrons and a triple bond six. Common examples are O₂ (O=O), N₂ (N≡N), CO₂ (O=C=O) and ethene, C₂H₄. Multiple bonds are shorter and stronger than the corresponding single bond between the same atoms, which affects reactivity and shape. In an MCQ, decide how many electron pairs are shared, count the electrons each atom contributes, and check that each atom reaches a stable electron arrangement.

A dative covalent (coordinate) bond is a covalent bond in which both shared electrons come from the same atom, the donor. The acceptor atom supplies an empty orbital or a vacant position and shares the electron pair. Once formed, a dative bond is indistinguishable from an ordinary covalent bond. Classic examples are the ammonium ion, NH₄⁺, where the nitrogen lone pair is donated to an H⁺ ion, and the formation of H₃O⁺ from H₂O and H⁺. In an MCQ, identify the donor (the atom with a lone pair) and the acceptor (the atom or ion with a vacant orbital), then confirm the shared pair originates from one atom only.

Average bond enthalpy is the enthalpy change when one mole of a specified gaseous covalent bond is broken, averaged over the same bond in a range of similar compounds. It is measured in kJ mol⁻¹ and is always positive because bond breaking is endothermic. Because the value is an average, it measures covalent bond strength only approximately: the same bond has slightly different enthalpies in different molecules. In methane, the successive bond dissociation enthalpies differ because the remaining fragment changes after each bond breaks, so the average is used instead. A larger average bond enthalpy indicates a stronger bond, so C≡C is stronger than C=C, which is stronger than C–C.

The shape of a simple molecule or ion is predicted from the number of electron pairs around the central atom using electron-pair repulsion theory. Count the bonding pairs and lone pairs around the central atom, then arrange them so that repulsion is minimised. Two bonding pairs give a linear shape with a 180° bond angle; three give trigonal planar with 120°; four give tetrahedral with 109.5°. Lone pairs repel more strongly than bonding pairs, so they compress bond angles: one lone pair with three bonding pairs gives a trigonal pyramidal shape with an angle near 107°, and two lone pairs with two bonding pairs give a bent shape with an angle near 104.5°. Learners should apply this reasoning to molecules such as CH₄, NH₃ and H₂O and to ions such as NH₄⁺.

Electron pair repulsion treats every pair of electrons around a central atom as a charge cloud that pushes other pairs away, so the pairs spread out as far as possible. Count all pairs, bonding and lone, to fix the total arrangement: two pairs give a linear base arrangement, three trigonal planar, four tetrahedral, five trigonal bipyramidal and six octahedral. Lone pairs repel more strongly than bonded pairs, so each lone pair compresses the remaining bond angles. For example, methane has four bonded pairs and an H–C–H angle of 109.5°, ammonia has three bonded pairs plus one lone pair and an H–N–H angle of 107°, and water has two bonded pairs plus two lone pairs and an H–O–H angle of 104.5°. Multiple bonds count as one region of electron density but repel slightly more than a single bond.

Electron pair repulsion explains each named shape by the number of electron regions around the central atom and how many are lone pairs. Two regions give linear, as in beryllium chloride BeCl₂ with a 180° angle. Three regions give trigonal planar, as in boron trifluoride BF₃ with 120° angles. Four regions give tetrahedral when all four are bonded, as in methane with 109.5°. When one of four regions is a lone pair the shape is pyramidal, as in ammonia with 107°, and when two are lone pairs the shape is non-linear, as in water with 104.5°. Six regions give octahedral, as in sulfur hexafluoride SF₆ with 90° angles. The shape name describes the atom positions, not the lone pairs.

Electronegativity measures how strongly an atom attracts the bonding electrons in a covalent bond. In a bond between two atoms of different electronegativity, the shared pair is pulled towards the more electronegative atom, which gains a partial negative charge (δ⁻) while the less electronegative atom gains a partial positive charge (δ⁺). This separation of charge makes the bond polar. A bond between identical atoms, such as in Cl₂ or H₂, is non-polar because the pull is equal. The greater the difference in electronegativity, the more polar the bond. For example, in HCl chlorine is more electronegative than hydrogen, so chlorine carries δ⁻ and hydrogen carries δ⁺. Bond polarity is a property of a bond; whether a whole molecule is polar also depends on its shape and on whether the individual bond dipoles cancel.

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. Pauling electronegativity values put this on a numerical scale, so you can compare atoms and judge how polar a bond will be. A larger value means a stronger attraction for the bonding electrons. For a bond between two different atoms, subtract the smaller value from the larger to find the difference: a bigger difference means a more polar bond, with the more electronegative atom taking δ⁻. For example, hydrogen is about 2.1 and chlorine about 3.0, giving a difference of about 0.9, so the H–Cl bond is polar with chlorine δ⁻. A bond between identical atoms has a difference of zero and is non-polar. The values are comparative, not measured in units.

This row is a lead-in stem, not a self-contained criterion: it announces that the following sub-statements (i) and (ii) are to be explained, so read it as a signpost rather than a separate target. Treat the explanation of a polar bond and permanent dipole, and of the origin of permanent dipole–dipole forces between molecules, as the assessed content. Build your notes around cause and effect: differing electronegativity draws bonding electron density towards the more electronegative atom, creating partial charges δ⁺ and δ⁻; an asymmetric arrangement leaves a permanent dipole; neighbouring molecules then attract between opposite partial charges. Use a worked example such as HCl and contrast it with a symmetrical molecule to test your reasoning.

A polar bond arises when two covalently bonded atoms have different electronegativities, so the shared electron pair is pulled towards the more electronegative atom. That atom gains a partial negative charge δ⁻ and the other a partial positive charge δ⁺. The bond dipole is a vector; where bond dipoles do not cancel, the molecule has a permanent dipole. In HCl the single bond is polar and the molecule is polar. In CO₂ each C=O bond is polar, but the linear symmetry makes the two dipoles cancel, so the molecule has no permanent dipole. Decide polarity by combining electronegativity difference with shape, not by counting polar bonds alone.

A bond is polar when the two bonded atoms differ in electronegativity, so the shared electron pair sits closer to the more electronegative atom and the bond carries a permanent dipole. A molecule is polar only if these bond dipoles do not cancel. Molecular shape decides this: the dipoles are vectors, so they add tip-to-tail. In CO₂ the two C=O dipoles point in opposite directions along a linear molecule and cancel, giving no overall dipole. In H₂O the two O–H dipoles both point towards oxygen, but the bent shape (about 104.5°) means they add to a resultant pointing from the hydrogen side towards oxygen. A molecule with polar bonds can therefore be non-polar overall, and a molecule with an overall dipole is described as polar. Compare shapes such as linear, trigonal planar, tetrahedral and bent when judging cancellation.

This row is a study-guidance heading within bonding and structure, so treat it as a prompt to read the specification and your textbook section on intermolecular forces rather than as a single assessed statement. Build a comparison of the three named types: induced dipole–dipole (London) forces, permanent dipole–dipole forces and hydrogen bonding. For each, note what causes it, its typical relative strength and the evidence it explains, such as boiling points and solubility. Link hydrogen bonding to an H atom bonded to N, O or F and to a lone pair on a nearby N, O or F. Read worked examples that compare molecules of similar electron number, and check each claim against the specification wording.

Intermolecular forces act between molecules rather than within them, and they are much weaker than covalent bonds. A permanent dipole–dipole interaction arises between molecules that already have a permanent dipole, because one region is δ⁺ and another is δ⁻; neighbouring molecules align so that opposite partial charges attract. An induced dipole–dipole interaction (London dispersion force) arises in all molecules: at any instant the electron cloud may be unevenly distributed, creating a temporary dipole that induces a dipole in a neighbour, so the molecules attract. Induced dipoles are the only intermolecular force in non-polar substances such as CH₄, and they strengthen as the number of electrons and the contact area increase.

Hydrogen bonding is a strong intermolecular attraction that occurs when a hydrogen atom covalently bonded to a highly electronegative N, O or F atom is attracted to a lone pair on an N, O or F atom of a neighbouring molecule. The H atom must be part of an –NH, –OH or HF group, so the bond is highly polar and the H carries a large δ⁺ charge. The small size of H allows close approach, making hydrogen bonds stronger than ordinary permanent dipole–dipole forces. Water shows extensive hydrogen bonding, which explains its relatively high boiling point. Hydrogen bonds are still much weaker than covalent bonds.

Water's hydrogen bonds arise because each O–H bond is strongly polar: the oxygen atom carries a partial negative charge and each hydrogen a partial positive charge, so a lone pair on one oxygen attracts a hydrogen on a neighbouring molecule. Each H₂O molecule can donate two hydrogen bonds through its hydrogens and accept two through its two lone pairs, giving a dynamic network rather than isolated pairs. This explains the anomalies: the energy needed to separate molecules is unusually large, so melting and boiling points are high for a molecule of relative molecular mass 18; the solid is less dense than the liquid because the tetrahedral hydrogen-bonded arrangement holds molecules further apart in an open lattice; and the high enthalpy of vaporisation follows from the same strong intermolecular attraction.

In ice, each water molecule forms up to four hydrogen bonds, two through its hydrogen atoms and two through the lone pairs on oxygen. These directional bonds hold the molecules in a rigid, open hexagonal lattice with large spaces between them, so a given mass of ice occupies a larger volume than the same mass of liquid water. Density is mass divided by volume, so the lower density of ice follows directly from that larger volume. When ice melts, some hydrogen bonds break and the lattice collapses, allowing molecules to pack more closely, so liquid water is denser than ice at 0 °C. This is why ice floats on water.

This clause is part of the explanation of water's anomalous properties caused by hydrogen bonding. Water has a small relative molecular mass of 18, yet its melting point is 273 K and its boiling point is 373 K at standard pressure, far higher than expected from London forces alone. The reason is that each water molecule forms hydrogen bonds: the polar O–H bonds leave hydrogen partially positive and oxygen partially negative, so lone pairs attract neighbouring hydrogen atoms. Melting and boiling require energy to overcome these intermolecular attractions, not to break the covalent O–H bonds. A correct answer identifies hydrogen bonding as the cause and links it to the energy needed to separate molecules.

A simple molecular lattice is a regular arrangement of discrete molecules. Within each molecule, atoms are held together by strong covalent bonds; between neighbouring molecules there are weak intermolecular forces. In solid iodine, I₂ molecules pack in a lattice held together by weak intermolecular forces, while each I–I bond is covalent. In ice, each water molecule has two O–H covalent bonds, and hydrogen bonds between molecules hold the lattice together; these intermolecular forces are much weaker than the covalent bonds. A strong answer distinguishes clearly between the covalent bonds inside molecules and the intermolecular forces between them, and explains that melting or boiling overcomes the weaker intermolecular forces first, leaving molecules intact.

Simple molecular covalent substances, such as iodine, ice and carbon dioxide, consist of discrete molecules held together in a lattice by weak intermolecular forces, while the atoms within each molecule are joined by strong covalent bonds. Melting and boiling require energy to overcome the weak intermolecular forces between molecules, not to break the covalent bonds, so these substances have low melting and boiling points that rise with stronger or more extensive intermolecular forces. Most simple molecular substances do not conduct electricity because they contain no free charged particles that can move. Solubility depends on the balance between the energy needed to separate the particles and the energy released when new solute–solvent intermolecular forces form, so a substance dissolves well when its intermolecular forces are similar in type and strength to those of the solvent.