Chemical Bonding

Chemical bonding: particles, electrons and properties

Chemical bonding explains why atoms join and why substances have different melting points, conductivities and structures. Atoms form bonds using outer-shell electrons. A useful starting model is that atoms tend toward a stable outer electron arrangement, but exam answers must name the actual process: electrons are transferred, shared, or delocalised.

Ions and ionic bonding

An ion is a charged particle formed when an atom gains or loses electrons. Metals generally lose electrons to form positive ions (cations); non-metals generally gain electrons to form negative ions (anions). The charge is not created by electrons disappearing: losing a negative electron leaves more protons than electrons, so the ion is positive.

Sodium has electron arrangement 2,8,1 and chlorine has 2,8,7. Sodium transfers one outer electron to chlorine:

[ ext{Na} ightarrow ext{Na}^+ + e^- ext{Cl}+e^- ightarrow ext{Cl}^-. ]

The oppositely charged ions attract strongly. This electrostatic attraction is ionic bonding. In dot-and-cross diagrams, show only outer electrons, distinguish the source of transferred electrons, and include brackets and charges around the resulting ions.

Ionic compounds form a giant, repeating lattice rather than separate molecules. This gives them high melting and boiling points because many strong attractions must be overcome. Solid ionic substances do not conduct electricity: ions are fixed in position. When molten or dissolved in water, ions can move and carry charge, so they conduct. Do not say that the solid “has no ions”; it has ions, but they cannot move.

Writing ionic formulae

The total charge in a compound is zero. Magnesium forms ( ext{Mg}^{2+}); chloride is ( ext{Cl}^-), so one magnesium ion needs two chloride ions: ( ext{MgCl}_2). Aluminium ( ext{Al}^{3+}) and oxide ( ext{O}^{2-}) combine in the ratio 2:3, giving ( ext{Al}_2 ext{O}_3). Write the simplest whole-number ratio; do not write charges in the final neutral formula.

Covalent bonding

A covalent bond is a shared pair of electrons between non-metal atoms. Each shared pair counts toward both atoms’ outer shells. In hydrogen, two hydrogen atoms share one pair, forming ( ext{H}_2). In water, oxygen shares one pair with each of two hydrogens, forming two single covalent bonds. Oxygen also has two lone pairs.

Covalent substances may be simple molecules or giant covalent structures. Simple molecules such as methane or carbon dioxide have strong covalent bonds within each molecule but weak intermolecular forces between molecules. Relatively little energy is needed to overcome the weak forces, so they often have low melting and boiling points. They generally do not conduct electricity because they have no mobile ions or delocalised electrons.

Diamond is different: each carbon atom forms four covalent bonds in a giant three-dimensional network. The many strong bonds make diamond hard and give it a high melting point. It does not conduct because all outer electrons are held in bonds. Graphite has each carbon bonded to three others in layers. One electron per atom is delocalised, so graphite conducts; weak forces between layers allow them to slide, making graphite soft and useful as a lubricant. Link every property to the structure, not merely to the word “covalent”.

Metallic bonding

Metals consist of a regular lattice of positive metal ions surrounded by delocalised electrons. Metallic bonding is the strong electrostatic attraction between these ions and the sea of delocalised electrons. The electrons move through the structure, explaining electrical and thermal conductivity. Layers of ions can slide while attraction remains, explaining malleability and ductility. Strong attractions usually give metals high melting points, although values vary between metals.

Comparing structures in an exam answer

For “Explain why sodium chloride has a high melting point”, write: it has a giant ionic lattice with strong electrostatic attractions between oppositely charged ions; much energy is required to overcome these attractions. For “Explain why it conducts when molten but not solid”, add: in the solid, ions are fixed; when molten, ions are free to move and carry charge.

For “Why is carbon dioxide a gas at room temperature?”, do not say “because covalent bonds are weak”. Its covalent bonds are strong. Say that there are weak forces between simple molecules, so little energy is needed to separate them.

Exam practice and errors

Draw dot-and-cross diagrams for magnesium oxide and water. State the bonding and conductivity of copper. Explain the difference between diamond and graphite using number of bonds, electron mobility and layers.

Common mistakes include calling a molecule an ionic lattice, mixing up intermolecular forces with covalent bonds, and claiming that all covalent substances have low melting points. Begin from the particles present, describe their arrangement, identify forces or mobile charged particles, then link this chain to the observed property.

Formula and diagram check

Before finalising an ionic formula, count total positive and negative charge separately. Before finalising a dot-and-cross diagram, count the outer electrons around each atom or ion and show brackets for ions. These two small checks prevent the most frequent representation errors without replacing an explanation of the bonding.