Why classify elements?
With over a hundred elements, studying each one separately would be impossible. Grouping elements with similar properties makes chemistry systematic and lets us predict properties.
Early attempts
Döbereiner's triads (1829): groups of three similar elements in which the middle one's atomic mass is about the average of the other two, e.g. Li, Na, K.
Newlands' law of octaves (1865): arranged by atomic mass, every eighth element had properties like the first, like notes in music. It worked only up to calcium.
Mendeleev's periodic law (1869): the properties of elements are a periodic function of their atomic masses.
Mendeleev's great insight was to put similarity of properties ahead of strict mass order. He left gaps for undiscovered elements and predicted their properties. When gallium (his eka-aluminium) and germanium (eka-silicon) were found, their properties matched his predictions closely.
The modern periodic law
Henry Moseley showed that the atomic number, not the atomic mass, is the fundamental property. The modern periodic law states:
The physical and chemical properties of elements are periodic functions of their atomic numbers.
The long form of the periodic table has 18 groups (vertical columns) and 7 periods (horizontal rows). The period number equals the principal quantum number of the outermost shell. Elements in a group have the same outer electronic configuration, which is why they behave alike.
Naming new elements
Newly discovered elements get temporary IUPAC names from their atomic numbers, using roots: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn, followed by "ium". Element 120 would be unbinilium (Ubn).
The four blocks
Block | Last electron enters | Groups | Features |
|---|---|---|---|
s-block | s orbital (ns¹⁻²) | 1 and 2 | reactive metals with low ionisation enthalpies; alkali and alkaline earth metals |
p-block | p orbital (ns² np¹⁻⁶) | 13 to 18 | metals, metalloids and non-metals; group 18 noble gases have full shells |
d-block | d orbital ((n−1)d¹⁻¹⁰ ns⁰⁻²) | 3 to 12 | transition elements: metals, often coloured ions, variable oxidation states |
f-block | f orbital | placed below the table | lanthanoids (4f) and actinoids (5f), the inner transition elements |
Periodic trends
Atomic radius
Across a period, it decreases. Electrons are added to the same shell while the nuclear charge increases, pulling them closer.
Down a group, it increases, as new shells are added.
Ionic radius: a cation is smaller than its parent atom (fewer electrons, same nuclear charge); an anion is larger. In an isoelectronic series (same number of electrons), the greater the nuclear charge, the smaller the ion: O²⁻ > F⁻ > Na⁺ > Mg²⁺.
Ionisation enthalpy
The energy needed to remove the most loosely held electron from an isolated gaseous atom: X(g) → X⁺(g) + e⁻.
It generally increases across a period and decreases down a group.
Successive ionisation enthalpies increase: IE₂ > IE₁.
Irregularities: Be > B, because boron's electron is removed from a 2p orbital, which is shielded by the 2s electrons and higher in energy. N > O, because nitrogen's half-filled 2p³ is extra stable, while oxygen loses a paired electron that is repelled by its partner.
Electron gain enthalpy
The enthalpy change when an electron is added to a gaseous atom: X(g) + e⁻ → X⁻(g). It is usually negative (energy released).
It becomes more negative across a period and less negative down a group.
Halogens have the most negative values; noble gases have large positive values.
Chlorine has a more negative value than fluorine, because fluorine's small 2p subshell is crowded and repels the incoming electron.
Electronegativity
The tendency of an atom in a molecule to attract the shared pair of electrons towards itself. It is not a measured energy but a scale; on the Pauling scale fluorine (4.0) is the most electronegative.
It increases across a period and decreases down a group, so metallic character decreases across a period and increases down a group.
Valence and the diagonal relationship
The valence of s- and p-block elements generally equals the number of outer electrons, or 8 minus that number. Oxides become less basic and more acidic across a period: Na₂O is strongly basic, Al₂O₃ amphoteric and Cl₂O₇ strongly acidic.
The first element of each group in the s and p blocks differs from the rest (small size, high electronegativity, no d orbitals). It often resembles the element diagonally below and to its right, the diagonal relationship: Li with Mg, Be with Al, B with Si.
Summary
The modern periodic law is based on atomic number; 18 groups and 7 periods.
s, p, d and f blocks are set by the subshell receiving the last electron.
Across a period: radius ↓, ionisation enthalpy ↑, electronegativity ↑, electron gain enthalpy more negative. Down a group, the reverse.
Be > B and N > O in ionisation enthalpy; Cl has the most negative electron gain enthalpy.