Chemical Bonding
and Structure

• Period number = number of occupied shells
• Group number gives the valence electrons: Groups 1 and 2 directly, Groups 13 to 18 take away 10
• Cl: Period 3, Group 17, so 2.8.7
• The staircase: metals on the left, non-metals on the right, metalloids along it

Atoms lose, gain or share valence electrons to reach a noble gas configuration. Every bond is an electrostatic attraction.

Molecules also attract one another by much weaker intermolecular forces (section 7).

3.2Dot-and-cross diagrams

• Mg loses its 2 valence electrons (crosses), O gains them.
• Show valence electrons only, as dots and crosses.
• Each ion in square brackets, charge outside.
• Number of each ion in front: [Ca]²⁺ 2[F]⁻.

3.3Properties

• High melting points: a lot of energy is needed to overcome the strong attraction between ions.
• Conduct only when molten or dissolved: the ions must be mobile. In the solid they are fixed.
• Hard but brittle: pushed along one ion, like charges meet and repel.
• Often soluble in water, insoluble in hexane and oil.

3.4Covalent character

• A small, highly charged cation polarises the anion's electron cloud: electrons are partly shared.
• Polarising power rises with charge density, charge ÷ radius.
• Larger anions are polarised more easily. AlCl₃ is covalent, NaCl and AlF₃ are ionic.

AnalogyEach atom puts its valence electrons into a shared pool. The pool is the glue that holds the positive ions together.

• Conduct electricity and heat as solid and liquid: the delocalised electrons are mobile.
• Malleable and ductile: layers slide, and the sea of electrons still holds them together.
• Usually high melting points.

5.1Drawing steps: CO₂

1. Central atom: forms the most bonds, or is the least electronegative. Never H.
2. For an ion, add or remove electrons for the charge. Square brackets, charge outside.
3. Give every atom an octet with single, double or triple bonds: C forms 4, N 3, O 2, Cl 1 and H 1.
4. A Period 2 atom holds at most 8 valence electrons.

5.4Dative bonds

• Drawn as an arrow from donor to acceptor. Once formed, it is identical to any covalent bond.
• H₃O⁺ forms as NH₄⁺ does. In H₃N→BF₃ the pair goes to the electron-deficient B.
• Al₂Cl₆: a Cl of each AlCl₃ donates a lone pair to the other Al, so each Al gets an octet.

Your turnDraw the dot-and-cross diagrams of (a) H₂O and (b) N₂. Show outer electrons only.

1. Electron pairs round the central atom are arranged as far apart as possible.
2. Repulsion: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair

• A lone pair is held closer to the nucleus, so it repels more.
• A double or triple bond counts as one region.

AnalogyRival campers on a tiny round planet pitch their tents as far apart as they can.

• Increases across a period and decreases down a group: the most electronegative elements are at the top right.
• Chapter 8 explains why.
• Different atoms: the more electronegative one gets δ−, the other δ+. The bond is polar.

1. Find the polar bonds
2. Work out the shape
3. Do the dipoles cancel? If not, the molecule is polar

AnalogyA tug of war: equal teams pulling opposite ways go nowhere (CO₂). Pull at an angle and the rope moves (H₂O).

8.1Ionic bonds

The attraction is stronger when the ions have larger charges and are smaller.

• Down a group, ions get larger (more shells).
• Na⁺ > Mg²⁺ > Al³⁺: same electrons, more protons.

1. Both are giant ionic lattices.
2. Mg²⁺ has a higher charge and a smaller radius than Na⁺.
3. The attraction to O²⁻ is stronger, so more energy is needed to overcome it.

8.2Metallic bonds

The metallic bond is stronger when

1. each atom gives more electrons to the sea
2. the ions have a higher charge and a smaller radius

Na (98 °C) < Mg (650 °C) < Al (660 °C): 1, 2 and 3 electrons per atom.

8.3Covalent bonds: bond energy and length

1. Larger atoms: more diffuse orbitals overlap less, so the bond is weaker
2. More shared pairs: stronger. C=C is less than twice C–C as π is weaker than σ
3. Polarity adds attraction: N–H 390 < O–H 460 < F–H 562

9.1Instantaneous dipole–induced dipole (id–id)

1. Moving electrons give an instant when one side has more electron density
2. This instantaneous dipole induces a dipole in a neighbour
3. The two attract. id–id acts between all molecules, and is the only force between non-polar ones

What makes id–id stronger?

• More electrons: a more easily distorted cloud. F₂ −188 °C, Cl₂ −34 °C, Br₂ 59 °C
• More surface contact: pentane (a chain) 36 °C, its compact isomer 2,2-dimethylpropane 10 °C

AnalogyTwo strips of tape stick along their length, two balls touch at one point.

9.3Hydrogen bonding

F, O and N are so electronegative that the H is almost a bare proton, and their lone pairs are small and concentrated.

Draw the lone pair, the δ+ and δ− charges, and label the bond.

AnalogyEach H on F, O or N is a hand, each lone pair a handle. One hydrogen bond is one hand holding one handle, so the number per molecule is the smaller of the two.

So H₂O (100 °C) > HF (20 °C) > NH₃ (−33 °C): H–F is more polar than N–H.

Going furtherDimers

Ethanoic acid dissolved in benzene has an apparent M_r of 120, not 60.

Its molecules pair up, held by hydrogen bonds.

Going furtherWithin one molecule

2-nitrophenol boils at 216 °C, its isomer 4-nitrophenol at 279 °C.

The groups of 2-nitrophenol are close enough to hydrogen bond within the molecule, so fewer sites are left for hydrogen bonds between molecules.

Ice and solubility

1. In ice each water molecule forms hydrogen bonds to 4 others in an open structure
2. The molecules sit further apart than in the liquid, so ice is less dense and floats

Like dissolves like. Polar, hydrogen-bonded and many ionic substances dissolve in water. Non-polar substances (oils, iodine) dissolve in hexane. Hexane and water do not mix.