Chemical Bonding and Structure

The Periodic Table (Chapter 2) lists the elements in order of proton number. Its rows are periods and its columns are groups. An element's position gives its electron arrangement without working it out (Figure 3.1):

• The period number is the number of occupied electron shells.
• For Groups 1 and 2, the group number is the number of valence electrons. For Groups 13 to 18, take away 10: Group 13 has 3 valence electrons and Group 17 has 7. Group 18 has a full outer shell of 8, and helium has 2.

Chlorine is in Period 3 and Group 17, so it has three shells with 7 electrons in the outer one: 2.8.7. Magnesium is in Period 3 and Group 2: 2.8.2.

A zigzag line, the staircase, runs down from boron. Elements to its left are metals and elements to its right are non-metals. Hydrogen, at the top left, is a non-metal. Elements along the staircase, such as silicon and germanium, are metalloids, with some properties of each.

Which way an atom reaches its octet depends on how many valence electrons it starts with.

• Metals have few valence electrons, usually 1 to 3. Losing those few is far easier than gaining five or more, so metal atoms give electrons away and form positive ions.
• Non-metals have 4 to 7 valence electrons, already close to 8. Gaining a few more is far easier than losing many, so non-metal atoms pull electrons towards themselves.

How strongly an atom pulls electrons towards itself in a bond is its electronegativity. A useful first picture: non-metals are more electronegative because they are already close to an octet, while metals are less electronegative because their few valence electrons are easily lost. Fluorine is the most electronegative element of all. Chapter 8 explains the real reasons for these trends.

The noble gases (Group 18) hardly react, because their valence shells are already full, and they exist as single atoms: they are monoatomic. Every other atom bonds to reach the same octet (Section 3.1): metals by losing electrons, non-metals by gaining or sharing them.

These bonds are strong. Molecules are also attracted to one another by much weaker intermolecular forces (3.9). Which bonds and forces a substance has decides its structure, and its structure decides its properties (3.10).

A metal atom loses its valence electrons to become a positive ion, a cation. A non-metal atom gains electrons to become a negative ion, an anion. Each ion then has a noble gas configuration. Sodium (2.8.1) loses one electron to give Na⁺ (2.8): Na → Na⁺ + e⁻. Chlorine (2.8.7) gains one to give Cl⁻ (2.8.8): Cl + e⁻ → Cl⁻.

3.3.1Charges and formulae

An atom forms the ion that reaches an octet by moving the fewest electrons (Section 3.1). A Group 1 atom has one valence electron. Losing it leaves the full shell underneath, while gaining seven would be far harder, so it forms a 1+ ion. In the same way a Group 2 atom loses two electrons (2+) and aluminium in Group 13 loses three (3+). On the other side, a Group 17 atom gains one electron (1−), Group 16 gains two (2−) and Group 15 gains three (3−). A Group 14 atom would have to lose or gain four, so carbon and silicon share electrons instead (Section 3.5). The charge on a simple ion therefore follows from its group.

A polyatomic ion is a group of covalently bonded atoms that carries an overall charge. Its charge does not follow from a group, so it must be learnt.

A simple anion takes the ending -ide: chloride, oxide, nitride. Transition metals such as iron and copper can form more than one ion, so the charge is shown in brackets in the name: iron(II) is Fe²⁺ and iron(III) is Fe³⁺.

An ionic compound has no overall charge, so its formula combines the ions so that the charges add up to zero. Brackets go round a polyatomic ion when more than one is needed.

3.3.2Dot-and-cross diagrams and the ionic lattice

A dot-and-cross diagram shows the valence electrons only, the electrons of one element as dots and of the other as crosses. Each ion goes in square brackets with its charge outside. The number of each ion is written in front: calcium fluoride is shown as one [Ca]²⁺ with 2[F]⁻.

Each ion attracts every oppositely charged ion around it, in all directions. In sodium chloride each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ by 6 Na⁺ (Figure 3.3 shows one layer). This number is the coordination number.

3.3.3Properties of ionic compounds

• High melting and boiling points: a large amount of energy is needed to overcome the strong attraction between the ions.
• Conduct electricity only when molten or dissolved: the ions must be mobile to carry charge. In the solid they are held in fixed positions (Figure 3.4).
• Hard but brittle: a force can push one layer along by one ion, so that ions of the same charge meet and repel, and the lattice shatters (Figure 3.5).
• Often soluble in water, but insoluble in solvents such as hexane and oil.

How strong the attraction is, and so how high the melting point, depends on the charges and sizes of the ions (Section 3.8).

3.3.4Ionic bonds with covalent character Core · A Level

A small, highly charged cation pulls the electron cloud of a nearby anion towards itself. It polarises the anion. Electron density is drawn between the two nuclei, so the electrons are partly shared and the bond gains covalent character (Figure 3.6).

• The polarising power of a cation rises with its charge density, its charge divided by its radius. Al³⁺ polarises far more than Na⁺.
• A larger anion is more easily polarised: I⁻ more than Cl⁻, Cl⁻ more than F⁻.
• So aluminium chloride behaves as a covalent compound, while sodium chloride and aluminium fluoride are ionic.

Each metal atom gives its valence electrons to the sea. The electrons are not tied to any one ion, and they move throughout the lattice (Figure 3.7).

• Conduct electricity and heat in the solid and liquid: the delocalised electrons are mobile and carry charge and energy through the metal.
• Malleable and ductile: layers of ions can slide over one another, and the sea of electrons still holds them together, so the metal bends instead of breaking.
• Usually high melting points: a lot of energy is needed to overcome the strong metallic bonds. How strong they are varies from metal to metal (Section 3.8).

3.4.1Alloys

An alloy is a mixture of a metal with one or more other elements, such as brass (copper and zinc) and steel (iron and carbon). The atoms of the added element are a different size, so they disrupt the regular layers. The layers can no longer slide easily, which makes alloys harder and stronger but less malleable than the pure metal (Figure 3.8). Like other mixtures, alloys usually melt over a range of temperatures.

Two non-metal atoms each provide one electron to a shared pair, so both reach an octet (hydrogen reaches 2). The number of bonds an atom forms, its valency, is the number of electrons it still needs, 8 minus its valence electrons: C forms 4, N 3, O 2 and Cl 1. Hydrogen forms 1. Two shared pairs make a double bond and three make a triple bond. Pairs of valence electrons not used in bonding are lone pairs. In a structural formula each line is one shared pair: H–Cl, O=O, N≡N.

3.5.1Drawing dot-and-cross diagrams

1. Choose the central atom: the one that forms the most bonds, or the least electronegative (3.7). Hydrogen is never central.
2. For an ion, add one electron per negative charge to the more electronegative atoms, or remove one per positive charge. Put the ion in square brackets with its charge outside.
3. Give every atom an octet using single, double or triple bonds. A Period 2 atom can hold at most 8 valence electrons.

Figure 3.9 Dot-and-cross diagrams. Crosses are the central atom's electrons, dots are the other atoms' electrons.

3.5.2Exceptions to the octet rule Core · A Level

• Fewer than 8 around the central atom: BeCl₂ (4), BF₃ and AlCl₃ (6). These are electron-deficient.
• More than 8: PCl₅ (10), SF₆ (12). Atoms of Period 3 and below have empty 3d orbitals close in energy to their valence orbitals, so they can expand the octet. A Period 2 atom cannot, so NCl₅ does not exist.
• An odd number of electrons: NO and NO₂. A particle with an unpaired electron is a radical, and is very reactive.

3.5.3σ and π bonds Core · A Level

A covalent bond forms when two half-filled orbitals, one from each atom, overlap and their electrons pair up (Figure 3.10).

Figure 3.10 Orbital overlap in H₂, HCl and Cl₂ (σ), and the sideways overlap of a π bond.

• A single bond is one σ bond. A double bond is one σ and one π. A triple bond is one σ and two π.
• Sideways overlap is less effective, so a π bond is weaker than a σ bond. A π bond forms only after a σ bond has brought the atoms close.

3.5.4Dative bonds Core · A Level

In the ammonium ion, N donates its lone pair to H⁺. H₃O⁺ forms the same way from H₂O and H⁺, and in H₃N→BF₃ the lone pair goes into the empty orbital of the electron-deficient B. In Al₂Cl₆, a Cl atom of each AlCl₃ donates a lone pair to the Al of the other, so each Al reaches an octet. A dative bond is drawn as an arrow from donor to acceptor. Once formed, it is identical to any other covalent bond (Figure 3.11).

Valence shell electron pair repulsion (VSEPR) theory predicts shapes from the electron pairs around the central atom.

1. Electron pairs repel, so they arrange themselves as far apart as possible.
2. A double or triple bond counts as one region, like a single bond.
3. Lone pairs repel more strongly than bond pairs, because a lone pair is held closer to the nucleus: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
4. The shape is named from the positions of the atoms only, though the lone pairs still take up space.

Figure 3.12 Shapes and bond angles (bp: bond pairs, lp: lone pairs). A wedge points towards you and a hatched bond away from you. Lone pairs are shaded.

Each lone pair squeezes the bond angles: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°. With 5 pairs the lone pairs go in the middle (equatorial) positions, and with 6 pairs they sit opposite each other.

To count the pairs on an ion, take the central atom's valence electrons, add one for each singly bonded atom, take away a positive charge or add a negative charge, and halve the total. PCl₄⁺ has (5 + 4 − 1)/2 = 4 pairs, all bond pairs: tetrahedral. ICl₄⁻ has (7 + 4 + 1)/2 = 6 pairs, 4 bond pairs and 2 lone pairs: square planar.

An oxygen atom usually joins the central atom by a double bond, as in CO₂ and SO₂. In an oxoanion, the charge says how many O⁻ have a single bond instead. ClO₃⁻ has one O⁻ and two =O, which leaves Cl one lone pair: 3 bond pairs and 1 lone pair, trigonal pyramidal.

• It increases across a period and decreases down a group, so the most electronegative elements are at the top right, apart from the noble gases. Chapter 8 explains why.
• Fluorine is the most electronegative element, then oxygen, then nitrogen and chlorine.

3.7.1Polar bonds and polar molecules

When two different atoms share electrons, the more electronegative atom pulls the pair closer. It gains a partial negative charge, δ−, and the other atom a partial positive charge, δ+. The bond has a dipole and is polar (Figure 3.13). Bonds between identical atoms, such as H–H and Cl–Cl, are non-polar.

A molecule with polar bonds may still be non-polar overall: if its shape is symmetrical, the bond dipoles cancel. To decide, find the polar bonds, work out the shape, then see whether the dipoles cancel (Figure 3.14).

Figure 3.14 CO₂ and BF₃ are non-polar because their dipoles cancel. H₂O and NH₃ are polar.

Sections 3.3 to 3.5 describe the properties shared by all ionic, all metallic and all covalent substances. This section explains why one bond of a type is stronger than another. Every bond is an electrostatic attraction, so the same two ideas run through all three: the attraction is stronger when the charges are larger and when they are closer together.

3.8.1Ionic bonds

Not every ionic attraction is equally strong. It is stronger when the ions carry larger charges and when they are smaller, so their centres are closer. Down a group, ions have more shells and get larger. Across a period, positive ions with the same number of electrons get smaller as the nuclear charge rises: Na⁺ > Mg²⁺ > Al³⁺.

3.8.2Metallic bonds

The metallic bond is stronger when each atom gives more electrons to the sea and when the metal ions are smaller, with a higher charge. Sodium, magnesium and aluminium give 1, 2 and 3 electrons each, so their melting points rise: Na (98 °C) < Mg (650 °C) < Al (660 °C). Down Group 1 the ions get larger, so the bond gets weaker and the melting point falls from lithium to caesium.

3.8.3Covalent bonds Core · A Level

Breaking a bond always takes in energy, so bond energies are positive: Cl₂(g) → 2Cl(g) needs +244 kJ mol⁻¹. A stronger bond has a larger bond energy and is usually shorter.

In a simple molecular substance, strong covalent bonds hold the atoms together inside each molecule, but only weak intermolecular forces hold the molecules to one another. Melting and boiling overcome the intermolecular forces only, not the covalent bonds, so simple molecular substances have low melting and boiling points. Many are volatile: they evaporate easily. They do not conduct electricity, because their molecules carry no overall charge. Hydrogen chloride is different in water: it forms H⁺ and Cl⁻ ions, which are mobile, so the solution conducts.

3.9.1Instantaneous dipole–induced dipole forces Core · A Level

The electrons in a molecule are always moving. At any instant the electron cloud may sit more on one side, giving a temporary dipole. This instantaneous dipole distorts the electron cloud of a neighbouring molecule, inducing a dipole in it, and the two attract (Figure 3.15). These id–id forces act between all molecules, and they are the only intermolecular forces between non-polar molecules.

• More electrons give a larger, more easily distorted electron cloud and stronger id–id forces: F₂ boils at −188 °C, Cl₂ at −34 °C and Br₂ at 59 °C.
• More surface contact gives stronger id–id forces. Long chain-shaped molecules touch along their length, compact ball-shaped ones only at a point: pentane (a chain of five carbon atoms) boils at 36 °C, its compact isomer 2,2-dimethylpropane at 10 °C.

3.9.2Permanent dipole–permanent dipole forces Core · A Level

Polar molecules have permanent dipoles. The δ+ end of one molecule attracts the δ− end of the next. These pd–pd forces act in addition to id–id forces. Between molecules with a similar number of electrons, the polar one boils higher: propene (non-polar, M_r 42) boils at −48 °C but ethanal (polar, M_r 44) at 20 °C. When the numbers of electrons differ greatly, id–id forces can win: CH₃Br (M_r 94.9, 4 °C) boils higher than the more polar CH₃Cl (M_r 50.5, −24 °C).

3.9.3Hydrogen bonding Core · A Level

F, O and N are so electronegative that the H bonded to them carries a large δ+, almost a bare proton, and their lone pairs are small and concentrated. A hydrogen bond is the strongest kind of intermolecular force, though still far weaker than a covalent bond. To draw one, show the lone pair, the δ+ and δ− charges and the bond itself (Figure 3.16).

• The number of hydrogen bonds a molecule can form is the smaller of its number of H atoms on F, O or N and its number of lone pairs on F, O or N. Water forms 2 per molecule, HF and NH₃ only 1.
• So water (100 °C) boils higher than HF (20 °C), which boils higher than NH₃ (−33 °C), because the H–F bond is more polar than N–H.
• H₂O, HF and NH₃ boil far higher than the other hydrides in their groups, which have no hydrogen bonding (Figure 3.17).
• In ice, each water molecule forms hydrogen bonds to 4 others in an open structure, so ice is less dense than water and floats.
• Ethanol, ammonia and HF mix with water because they form hydrogen bonds with water molecules.

3.9.4Solubility

"Like dissolves like." A substance dissolves when the attractions it forms with the solvent make up for the attractions broken. Polar and hydrogen-bonded substances, and many ionic compounds, dissolve in water. Non-polar substances, such as oils and iodine, dissolve in non-polar solvents such as hexane, and not in water. That is why hexane and water are immiscible.

In a giant covalent (giant molecular) structure, every atom is joined to its neighbours by covalent bonds that extend throughout the solid. Melting means breaking many strong covalent bonds, so these substances have very high melting points and are insoluble in all solvents. Examples are diamond, graphite and silicon dioxide (Figure 3.18). Diamond and graphite are allotropes of carbon: different forms of the same element, with different arrangements of atoms.

Silicon dioxide, SiO₂, has a structure like diamond, with each Si bonded to 4 O and each O to 2 Si. It melts at 1710 °C, while carbon dioxide, a simple molecular substance, is a gas at room temperature.

3.10.1Structure and properties