Bonding and structure

1. Overview

The properties of a solid depend on what its particles are and how strongly they are held together. This section sorts crystalline solids into four lattice types, giant ionic, simple molecular, giant molecular and giant metallic, using nine named examples. It then uses them to explain and predict melting and boiling points, electrical conductivity and solubility, and to deduce the structure of an unknown substance from data.

Key Definitions

• Lattice: a regular, repeating three-dimensional arrangement of particles (ions, atoms or molecules) that extends throughout a crystalline solid.
• Giant ionic lattice: a lattice of oppositely charged ions held together by strong electrostatic attraction acting in all directions; it contains no separate molecules, so its formula gives only the ratio of the ions.
• Simple molecular lattice: a lattice of separate small molecules held in place by weak intermolecular forces, while the atoms inside each molecule are joined by strong covalent bonds.
• Giant molecular lattice (giant covalent): a lattice of atoms joined to their neighbours by covalent bonds in a continuous network that runs through the whole crystal (in graphite, through each whole layer), so there are no small separate molecules.
• Giant metallic lattice: a regular arrangement of positive metal ions surrounded by delocalised electrons, held together by metallic bonding.
• Metallic bonding: the electrostatic attraction between positive metal ions and delocalised electrons.
• Delocalised electrons: electrons that are not held by one atom or one bond but are spread over many atoms or ions and can move among them: through the whole lattice in a metal, along each layer in graphite, or over one molecule in C₆₀.
• Intermolecular forces (van der Waals' forces): the attractions between separate molecules, namely instantaneous dipole-induced dipole forces and permanent dipole-permanent dipole forces, including hydrogen bonds; in general they are much weaker than ionic, covalent or metallic bonds.

Content

Four kinds of crystalline solid

In a crystalline solid, the particles sit in a regular pattern, the lattice, that repeats in all three directions.

Lattice type	Particles	What holds the lattice together	Examples you must know
Giant ionic	positive and negative ions	ionic bonding: electrostatic attraction between oppositely charged ions	sodium chloride, magnesium oxide
Simple molecular	small, separate molecules	weak intermolecular forces between molecules (strong covalent bonds inside each one)	iodine, buckminsterfullerene (C₆₀), ice
Giant molecular	atoms	strong covalent bonds throughout (in graphite, within each layer)	diamond, graphite, silicon(IV) oxide
Giant metallic	positive metal ions and delocalised electrons	metallic bonding	copper

Giant means the strong bonding continues through the whole crystal (in graphite, through each layer), with no small separate molecules, so the formula is only a ratio: SiO₂ means one silicon atom for every two oxygen atoms. Simple means separate molecules with a fixed formula, such as I₂ or C₆₀.

Giant ionic: sodium chloride and magnesium oxide

In sodium chloride, Na⁺ and Cl⁻ ions alternate along three directions at right angles, building up a cubic lattice. Each Na⁺ ion has six Cl⁻ ions as nearest neighbours, and each Cl⁻ ion has six Na⁺ ions. An ion is not bonded to one partner: it attracts every oppositely charged ion around it.

Magnesium oxide has the same arrangement, with Mg²⁺ and O²⁻ ions. Their charges are larger and the ions smaller, so the attraction is much stronger (topic 3.2 explains why).

Ionic crystals are brittle: if one layer of ions slides by one ion's width, ions with the same charge meet, repel, and the crystal splits.

Simple molecular: iodine, buckminsterfullerene and ice

A simple molecular solid has strong covalent bonds inside each molecule and weak intermolecular forces between molecules. Melting or boiling overcomes only the intermolecular forces.

Iodine, I₂. Each molecule is two atoms joined by a covalent I-I bond. The molecules are packed in a regular arrangement, held by instantaneous dipole-induced dipole (id-id) forces. With 106 electrons per molecule, these forces keep iodine solid at room temperature, as shiny grey-black crystals. Warmed gently, it gives a purple vapour that is still made of I₂ molecules.

Buckminsterfullerene, C₆₀. Each molecule is a hollow cage of 60 carbon atoms shaped like a football, made of 20 hexagons and 12 pentagons. Each carbon atom is covalently bonded to three others, and some electrons are delocalised over the surface of the ball. The balls pack into a regular lattice held by id-id forces. With 360 electrons per molecule, these forces are stronger than in iodine, but C₆₀ still turns to vapour far below diamond or graphite, because no covalent bonds are broken.

Ice, H₂O(s). Each water molecule forms hydrogen bonds to four neighbours: two through its own hydrogen atoms and two through the lone pairs on its oxygen atom. The neighbours sit tetrahedrally around each oxygen, giving an open network with a lot of empty space, so ice is less dense than liquid water (topic 3.6).

Giant molecular: diamond, graphite and silicon(IV) oxide

Diamond. Each carbon atom is covalently bonded to four others, arranged tetrahedrally (bond angle 109.5°), in a three-dimensional network through the whole crystal. All four outer electrons of every atom are held in these bonds.

Graphite. The carbon atoms form flat layers of joined hexagons. Each atom is covalently bonded to three others in its layer, at 120°. The fourth outer electron of each atom is delocalised and can move along the layer. The layers are held together only by weak id-id forces.

Silicon(IV) oxide, SiO₂. Each silicon atom is covalently bonded to four oxygen atoms, arranged tetrahedrally, and each oxygen atom is bonded to two silicon atoms. Every oxygen is shared between two silicon atoms, so the ratio is one Si to two O. The three-dimensional network is rather like diamond with an oxygen atom between each pair of neighbouring silicon atoms.

Giant metallic: copper

Solid copper is a regular, closely packed lattice of positive copper ions. The outer electrons are delocalised and free to move through the whole lattice. Metallic bonding, the attraction between the positive ions and the delocalised electrons, acts in all directions.

Metals are malleable and ductile: layers of ions can slide over one another while the delocalised electrons keep holding them, so the metal changes shape instead of shattering.

Worked example 1: counting bonds in giant and simple structures

Question: ($A_{\text{r}}$: C 12.0, O 16.0, Si 28.1; $L = 6.022 \times 10^{23}$ mol⁻¹)

(a) Calculate the number of Si-O bonds in 5.00 g of silicon(IV) oxide.

(b) Calculate the amount, in mol, of C-C bonds in a 0.300 g diamond.

(c) How many covalent bonds are broken when 0.300 g of C₆₀ turns completely to vapour?

(a) Step 1: amount of SiO₂. $$M(\text{SiO}_2) = 28.1 + (2 \times 16.0) = 60.1\ \text{g mol}^{-1} \qquad n = \frac{5.00}{60.1} = 0.08319\ \text{mol}$$

Step 2: bonds per formula unit. Each Si atom forms four Si-O bonds, and each bond contains exactly one Si atom, so 1 mol of SiO₂ holds 4 mol of Si-O bonds. $$n(\text{Si-O}) = 4 \times 0.08319 = 0.3328\ \text{mol} \qquad N = 0.3328 \times 6.022 \times 10^{23} = 2.004 \times 10^{23}$$

Answer (a): $2.00 \times 10^{23}$ Si-O bonds (3 s.f.).

(b) Each C atom forms four bonds, but each bond joins two atoms, so there are $4 \div 2 = 2$ C-C bonds per C atom. $$n(\text{C}) = \frac{0.300}{12.0} = 0.0250\ \text{mol} \qquad n(\text{C-C}) = 2 \times 0.0250 = 0.0500\ \text{mol}$$

Answer (b): 0.0500 mol of C-C bonds (3 s.f.).

(c) Turning C₆₀ into vapour overcomes only the intermolecular forces between the balls. However many molecules the sample holds, each one keeps all 90 of its C-C bonds ($60 \times 3 \div 2$), so the mass given does not matter.

Answer (c): zero covalent bonds are broken.

How structure decides physical properties

Two questions explain almost every property here: what must be overcome, and what is free to move?

Property	Giant ionic	Simple molecular	Giant molecular	Giant metallic
Melting overcomes	many strong ionic attractions	weak intermolecular forces only	many strong covalent bonds	strong metallic bonding
Melting point	high	low	very high	usually high, but varies widely
Conducts as a solid?	no: ions fixed	no	no (graphite: yes, along its layers)	yes
Conducts as a liquid?	yes: ions move	no	no	yes
Dissolves in water?	many do	most do not, unless they form hydrogen bonds with water or react with it	no	no (a few react with it)
Dissolves in hexane (non-polar)?	no	many do	no	no

Melting and boiling points

Giant structures melt and boil at high temperatures, and simple molecular substances at low ones, for the reasons in the first row of the table above. Ice melts at 0 °C and iodine at 114 °C, but sodium chloride needs 801 °C, and diamond and graphite are still solid above 3500 °C. Within each type, stronger attractions give higher values:

• Simple molecular: more electrons per molecule give stronger id-id forces (C₆₀ against I₂). For molecules with similar numbers of electrons, hydrogen bonds are stronger than other permanent dipole-permanent dipole (pd-pd) forces, which are stronger than id-id forces. Water and methane, CH₄, both have 10 electrons, yet ice melts at 0 °C and solid methane at −182 °C.
• Ionic: larger charges and smaller ions attract more strongly. MgO melts far higher than NaCl. Caesium iodide, CsI, melts lower than NaCl: its ions carry the same charges as Na⁺ and Cl⁻ but are larger, so their centres are further apart.
• Metals: more delocalised electrons per ion and smaller, more highly charged ions give stronger metallic bonding (topics 3.3 and 9.1 apply this to sodium, magnesium and aluminium).

A metal boils far above its melting point. On melting, the ions can slide past one another but are still held by the delocalised electrons; boiling overcomes the metallic bonding completely and separates the atoms.

Electrical conductivity

To conduct, a substance needs charged particles that are free to move: mobile ions or delocalised electrons.

• Metals conduct as solids and liquids: the delocalised electrons move through the lattice and carry charge.
• Graphite conducts along its layers, where its delocalised electrons move, but very poorly at right angles to them. Diamond and SiO₂ do not conduct: every outer electron is held in a localised bond or lone pair.
• Ionic solids do not conduct, because their ions are fixed. Molten or in aqueous solution, the ions move and carry the charge.
• Simple molecular substances have no ions and no electrons that can move between molecules, so they do not conduct in any state (pure water conducts only very slightly). The delocalised electrons in C₆₀ stay on their own ball.

Solubility

A solid dissolves well when solvent molecules can surround its particles and the new attractions roughly make up for those broken in the solid and in the solvent.

• Ionic solids often dissolve in water: its polar molecules attract the ions, δ− oxygen atoms to positive ions and δ+ hydrogen atoms to negative ions, and pull them out of the lattice. Where the ions attract each other very strongly, the solid barely dissolves: magnesium oxide is almost insoluble. Non-polar hexane cannot attract ions strongly, so ionic solids do not dissolve in it.
• Simple molecular solids with non-polar molecules dissolve in non-polar solvents, forming id-id forces like those they replace: iodine dissolves readily in hexane and C₆₀ in methylbenzene. Both barely dissolve in water, because they cannot attract water molecules as strongly as water molecules attract one another by hydrogen bonding. Molecules that hydrogen bond with water, such as ethanol, do dissolve.
• Giant molecular solids do not dissolve in any common solvent: no solvent can replace so many strong covalent bonds.
• Metals do not dissolve in water or hexane. Sodium reacts with water, which makes new substances; that is not dissolving.

Deducing structure and bonding from information

Conductivity is usually the most decisive test, so work through the evidence in this order:

1. Conducts as a solid (and when molten): giant metallic. The main exception is graphite, a non-metal that conducts but is giant molecular.
2. Conducts when molten but not as a solid: giant ionic, with ions fixed in the solid and free to move in the liquid. A conducting solution alone is weaker evidence: it shows that ions are present, not that the solid was ionic. Hydrogen chloride is simple molecular, but its solution conducts because it reacts with water to form H⁺(aq) and Cl⁻(aq).
3. Conducts in no state: covalent. A low melting point means simple molecular. A very high melting point (usually above 1000 °C) and insolubility in all common solvents mean giant molecular.

Other clues can confirm the answer. A gas or liquid at room temperature is almost always simple molecular (mercury, a liquid metal, is the common exception). A metal combined with a non-metal is usually ionic, but check the electronegativity difference (topic 3.1).

Worked example 2: deducing the structures of four unknown solids

Question: Four solids, numbered 1 to 4, are tested. Deduce the type of structure and the bonding in each.

Solid	Melting point / °C	Conducts as a solid?	Conducts when molten?	In water	In hexane
1	681	no	yes	dissolves; the solution conducts	does not dissolve
2	420	yes	yes	does not dissolve	does not dissolve
3	80	no	no	does not dissolve	dissolves
4	3300	no	no	does not dissolve	does not dissolve

Step 1: conducts as a solid. Solid 2 conducts as a solid and when molten, so it has delocalised electrons that can move. It melts at only 420 °C, far too low for graphite, so it is a metal: conductivity outweighs the moderate melting point.

Step 2: conducts only when molten. Solid 1 has ions that are fixed in the solid and free to move in the liquid. Dissolving in water but not in hexane fits.

Step 3: conducts in no state. Solids 3 and 4 are covalent. Solid 3 melts at only 80 °C, so only weak intermolecular forces are overcome; dissolving in hexane fits non-polar molecules. Solid 4 melts at 3300 °C and dissolves in neither solvent, so melting it must break many strong covalent bonds.

Answer:

• Solid 1: giant ionic; ionic bonding.
• Solid 2: giant metallic; metallic bonding.
• Solid 3: simple molecular; covalent bonds within the molecules, weak intermolecular forces between them.
• Solid 4: giant molecular; covalent bonding throughout.

Key Equations

Covalent bonds in a giant lattice of one element ($n$ = amount in mol; each bond joins two atoms, so diamond has $4 \div 2 = 2$ bonds per atom, and graphite $3 \div 2 = 1.5$ within its layers): $$n(\text{bonds}) = n(\text{atoms}) \times \frac{\text{bonds formed by each atom}}{2}$$

Si-O bonds in silicon(IV) oxide (each Si atom forms four Si-O bonds, and each bond contains one Si atom): $$n(\text{Si-O}) = 4 \times n(\text{SiO}_2)$$

Amount and number ($m$ = mass in g; $M$ = molar mass in g mol⁻¹; $N$ = number of bonds or particles; $L$ = Avogadro constant, $6.022 \times 10^{23}$ mol⁻¹): $$n = \frac{m}{M} \qquad N = n \times L$$

Common Mistakes to Avoid

• Mistake: saying covalent bonds break when a simple molecular solid melts or boils. Fix: only intermolecular forces are overcome. Iodine vapour is still I₂ molecules; in Worked example 1, every C₆₀ molecule keeps all 90 bonds.
• Mistake: using "intermolecular forces" to explain the melting point of a giant structure. Fix: a giant structure has no small molecules. Melting SiO₂ means breaking many strong covalent Si-O bonds; melting a metal means overcoming strong metallic bonding.
• Mistake: writing "copper conducts because it has delocalised electrons" and stopping there. Fix: "the delocalised electrons are free to move through the lattice and carry charge"; for graphite, add "along the layers".
• Mistake: saying molten sodium chloride conducts because electrons move, or that the solid conducts. Fix: ions carry the charge: in molten NaCl, Na⁺ ions move to the negative electrode and Cl⁻ ions to the positive; in the solid they are fixed.
• Mistake: calling the positive particles in a metal "atoms", "nuclei" or "protons", or drawing them scattered at random. Fix: draw regular rows of positive ions with delocalised electrons between them, and label both.
• Mistake: deciding the structure from one clue: "ionic" because a solution conducts, "metallic" because a solid conducts, or "giant" because a substance is pure carbon. Fix: check the other evidence: hydrogen chloride solution conducts, but HCl is simple molecular; graphite conducts but is giant molecular; C₆₀ is pure carbon but simple molecular.

Exam Tips

• Identify a lattice with one of the four exact terms: giant ionic, simple molecular, giant molecular (giant covalent) or giant metallic. For "structure and bonding", give both, and make them match: "giant molecular with ionic bonding" contradicts itself; for a simple molecular substance, name the covalent bonds inside the molecules and the intermolecular forces between them.
• Describe a lattice by giving the particles, their arrangement and what holds them, with numbers where you can: "each Si atom is covalently bonded to four O atoms, and each O atom to two Si atoms".
• Explain a property as a chain. Melting or boiling point: the structure, the bonds or forces overcome, how strong and how many they are, and so how much energy is needed; when comparing, say which attraction is stronger and why. Conductivity: which charged particles (ions or delocalised electrons), whether they are free to move, and in which state.
• Deduce or suggest a structure by linking each piece of evidence to its conclusion in the same sentence: "it conducts only when molten, so its ions are fixed in the solid: giant ionic". Repeating the data without the reasoning is not an explanation.