Atoms, the periodic table and chemical bonding

Lesson 1 treated atoms and molecules as units to be counted. Why do they combine in the ratios they do, and why do the substances they form behave so differently? Why is steel strong and ductile while glass is strong and brittle, why does water boil at 100 °C while methane, a similar-sized molecule, boils at −162 °C, and why does salt dissolve in water but wax does not? The answers lie in the structure of atoms and the way their outer electrons form chemical bonds. This lesson describes the structure of the atom, the arrangement of electrons that gives the periodic table its shape, the trends in atomic properties across it, the three main types of chemical bond (ionic, covalent and metallic), molecular shape and polarity, and the weaker forces between molecules that control boiling points, solubility and much of the behaviour of engineering materials.

The structure of the atom

An atom has a tiny, dense nucleus of protons (positive) and neutrons (neutral), surrounded by electrons (negative) that occupy most of its volume. The number of protons, the atomic number Z, defines the element; a neutral atom has Z electrons. Atoms of the same element with different numbers of neutrons are isotopes: carbon-12 and carbon-14, uranium-235 and uranium-238. The atomic masses in the periodic table are averages over the natural mixture of isotopes, which is why chlorine's is 35.45 rather than a whole number.

Electrons occupy energy levels (shells), labelled n = 1, 2, 3 and so on, each divided into subshells (s, p, d, f) holding up to 2, 6, 10 and 14 electrons. Electrons fill the lowest available levels first. The electrons in the outermost shell, the valence electrons, take part in bonding and decide an element's chemistry. Sodium (Z = 11) has the configuration 1s² 2s² 2p⁶ 3s¹: one valence electron that it readily loses. Chlorine (Z = 17), 1s² 2s² 2p⁶ 3s² 3p⁵, needs one more electron to complete its outer shell.

The periodic table

Arranging the elements by atomic number, with a new row (period) each time a new shell begins, places elements with the same number of valence electrons in the same column (group), and they behave alike. Group 1, the alkali metals, all have one valence electron and react vigorously; group 17, the halogens, all lack one; group 18, the noble gases, have full shells and are almost inert, which is why argon is used to shield welds. The transition metals in the middle, including iron, copper, nickel, chromium and titanium, fill d subshells and form the structural and functional metals of engineering.

Several properties change systematically across the table:

• Atomic radius decreases from left to right across a period (the growing nuclear charge pulls the electrons in) and increases down a group (more shells).
• Ionization energy, the energy to remove an electron, increases across a period and decreases down a group. Metals, on the left, lose electrons easily; non-metals, on the right, hold them tightly.
• Electronegativity, an atom's pull on the electrons in a bond, follows the same trend: fluorine is the most electronegative element (4.0 on the Pauling scale), oxygen 3.4, carbon 2.6, hydrogen 2.2, sodium 0.9.

These trends predict which elements form which kinds of bonds, how reactive metals are (and so how readily they corrode, Lesson 9), and why oxygen and fluorine compounds are so stable.

Ionic bonds

When a metal with low ionization energy meets a non-metal with high electronegativity, electrons transfer: sodium gives its electron to chlorine, forming Na⁺ and Cl⁻ ions with full outer shells. The oppositely charged ions attract strongly (Coulomb's law, PHYS-131) and pack into a crystal lattice. Ionic compounds are typically hard, brittle and high-melting (sodium chloride melts at 801 °C); they conduct electricity when molten or dissolved, because the ions can then move, but not as solids. The charges determine the formula: Ca²⁺ and Cl⁻ give CaCl₂; Al³⁺ and O²⁻ give Al₂O₃ (alumina, a hard ceramic). As a rule of thumb, bonds between atoms differing in electronegativity by more than about 1.7 to 2 are mainly ionic.

Common ions and naming compounds

Many ions that engineers meet are polyatomic: small covalently bonded groups that carry an overall charge and behave as a unit in ionic compounds. Recognizing them makes formulas readable and calculations in later lessons straightforward:

Ionic compounds are named cation first, then anion: sodium chloride, calcium carbonate, aluminium sulphate. Metals that form more than one ion show the charge in Roman numerals: iron(II) oxide, FeO, and iron(III) oxide, Fe₂O₃, the main component of rust and of hematite ore. Molecular compounds of two non-metals use Greek prefixes for the number of atoms: carbon monoxide (CO), carbon dioxide (CO₂), sulphur trioxide (SO₃), dinitrogen oxide (N₂O). The formula of any ionic compound follows from balancing charges, so that the compound is electrically neutral overall.

Covalent bonds

Between non-metals, neither atom can take electrons from the other, so they share pairs of electrons in covalent bonds. Hydrogen shares its single electron with another hydrogen to form H₂; carbon, with four valence electrons, forms four bonds, as in methane (CH₄). Atoms can share two or three pairs, forming double and triple bonds, which are shorter and stronger: the triple bond in N₂ (945 kJ/mol) makes nitrogen gas very unreactive, which is why making ammonia for fertilizer requires high pressure, high temperature and a catalyst (Lesson 8).

When the atoms differ in electronegativity, the shared electrons sit closer to the more electronegative one, making a polar covalent bond with partial charges, written δ+ and δ−. In water, oxygen pulls electrons from both hydrogens. Covalent solids can be molecular (soft, low-melting, made of discrete molecules, like wax or ice) or network solids, in which covalent bonds extend through the whole crystal, giving the extreme hardness and high melting points of diamond, silicon carbide and quartz.

Metallic bonds

In metals, the valence electrons are not tied to particular atoms but move freely through a lattice of positive ions, a "sea" of delocalized electrons. This explains the characteristic properties of metals. The mobile electrons carry current and heat, so metals are good electrical and thermal conductors (PHYS-131). The bonding is not directional, so planes of atoms can slide past one another without breaking the solid apart, which makes metals ductile and malleable: steel can be rolled, drawn and bent, and it deforms visibly before it breaks, a crucial safety property in structures. Alloying, mixing in other elements such as carbon in steel or tin in bronze, disrupts this sliding and makes metals stronger but less ductile (Lesson 10).

Molecular shape and polarity

The shape of a molecule follows from the repulsion of its electron pairs, which arrange themselves as far apart as possible (the VSEPR model). Four pairs around a central atom point to the corners of a tetrahedron, about 109.5° apart: methane is tetrahedral; ammonia, with one non-bonding pair, is pyramidal; water, with two, is bent, at 104.5°. Two pairs give a straight line, as in CO₂.

Shape decides whether polar bonds add up to a polar molecule. In CO₂ the two polar C=O bonds point in opposite directions and cancel, so the molecule is non-polar. In bent water they do not cancel, so water is strongly polar, with a negative oxygen end and positive hydrogen end. Polarity has large practical consequences: it is why water dissolves salts and polar substances but not oils, why microwave ovens heat water-containing food, and why CO₂, although non-polar, absorbs infrared radiation when it bends and stretches, making it a greenhouse gas (Lesson 11).

Intermolecular forces

Bonds within molecules are strong; the forces between molecules are much weaker, but they decide whether a substance is a gas, a liquid or a solid at room temperature, how much energy it takes to boil it, and what it dissolves in. In increasing strength:

• London dispersion forces act between all molecules, from momentary fluctuations in electron distribution. They grow with molecular size: methane (CH₄) boils at −162 °C, octane (C₈H₁₈, in gasoline) at 126 °C, and long-chain waxes are solids.
• Dipole-dipole forces attract polar molecules to one another.
• Hydrogen bonds, a strong special case, form when hydrogen bonded to N, O or F is attracted to a lone pair on another such atom. They give water its unusually high boiling point, large specific heat and latent heat (Lesson 5), its surface tension, and the fact that ice is less dense than liquid water, so it floats, expands as it freezes and cracks rocks, pipes and concrete.

The rule "like dissolves like" follows: polar and ionic substances dissolve in polar solvents such as water; non-polar substances such as oils, greases and many polymers dissolve in non-polar solvents. Detergents and surfactants have a polar head and a non-polar tail, which lets them carry grease into water, the basis of cleaning, of some oil spill treatments and of the air-entraining admixtures that protect concrete from frost damage.