Classification of Elements and Periodicity in Properties

In this chapter: Dobereiner's triads, Newlands' octaves, Mendeleev's periodic law, the modern periodic law, IUPAC names for elements with Z > 100, electronic configuration and the s, p, d and f blocks, metals, non-metals and metalloids, trends in atomic and ionic radii, ionisation enthalpy, electron gain enthalpy, electronegativity and valence, anomalous properties of second period elements, and periodic trends in chemical reactivity.

How the periodic table developed

• Dobereiner's triads (1829): groups of three similar elements in which the middle element has an atomic weight close to the average of the other two. Li (7), Na (23), K (39): (7 + 39)/2 = 23. Other triads: Ca, Sr, Ba and Cl, Br, I. Only a few elements fitted.
• Newlands' law of octaves (1865): when elements are arranged by increasing atomic weight, every eighth element has properties similar to the first, like the eighth note of a musical scale. It worked only up to calcium.
• Mendeleev's periodic law (1869): the properties of elements are a periodic function of their atomic weights. Mendeleev gave more weight to similarity in properties than to strict order of atomic weight, so he placed iodine after tellurium. He left gaps for elements not yet discovered and predicted their properties: eka-aluminium and eka-silicon turned out to be gallium and germanium.

In 1913 Moseley showed that the frequency of X-rays emitted by an element varies with its atomic number (a plot of √ν against Z is a straight line), and not smoothly with atomic mass. This led to the modern periodic law: the physical and chemical properties of elements are periodic functions of their atomic numbers.

The long form of the periodic table

The table has 18 vertical columns called groups and 7 horizontal rows called periods. The period number is the highest principal quantum number (n) of the valence shell. The number of elements in each period is twice the number of orbitals available in the energy level being filled.

Naming elements with Z above 100

Until IUPAC approves a permanent name, a new element gets a temporary name built from roots for its digits, followed by "ium".

Worked example: Give the IUPAC temporary name and symbol for the element with Z = 120.
Solution: 1, 2, 0 gives un + bi + nil + ium = unbinilium, symbol Ubn (first letters of the roots). In the same way, Z = 115 was ununpentium (Uup) before it was named moscovium (Mc).

Electronic configuration and the four blocks

Metals make up more than 78% of known elements and lie on the left. Non-metals sit at the top right. Elements along the zig-zag border such as silicon, germanium, arsenic, antimony and tellurium show properties of both and are called metalloids or semi-metals. Metallic character increases down a group and decreases across a period.

Periodic trends in physical properties

Atomic radius

Covalent radius is half the distance between the nuclei of two atoms joined by a single bond (for chlorine, half of 198 pm = 99 pm). Metallic radius is half the distance between neighbouring atoms in a metal crystal. Across period 2 the radius falls from Li (152 pm) to F (64 pm); down group 1 it rises from Li (152 pm) to Cs (262 pm). Noble gas radii are van der Waals radii and are not compared with covalent radii of the other elements.

Ionic radius

A cation is smaller than its parent atom because it has fewer electrons for the same nuclear charge (Na 186 pm, Na⁺ 95 pm). An anion is larger (F 64 pm, F⁻ 136 pm). Isoelectronic species have the same number of electrons; among them, the greater the nuclear charge, the smaller the ion. For the 10-electron series, O²⁻ > F⁻ > Na⁺ > Mg²⁺.

Ionisation enthalpy

Ionisation enthalpy is the energy needed to remove an electron from an isolated gaseous atom in its ground state: X(g) → X⁺(g) + e⁻. It is always positive, and the second ionisation enthalpy is always higher than the first, because the electron is removed from a positive ion. It increases across a period and decreases down a group. Two exceptions in period 2 need explanations:

• B is lower than Be: boron loses a 2p electron, which is more shielded from the nucleus by the inner 2s electrons than the 2s electron that beryllium loses. s electrons penetrate closer to the nucleus than p electrons.
• O is lower than N: nitrogen has three singly occupied 2p orbitals (a half-filled, stable arrangement), while in oxygen two electrons share one 2p orbital, and their repulsion makes one of them easier to remove.

Worked example: The first ionisation enthalpy of sodium is 496 kJ mol⁻¹. How much energy is needed to ionise 2.3 g of gaseous sodium atoms, and what is the energy per atom?
Solution: 2.3 g Na = 2.3/23 = 0.10 mol. Energy = 0.10 × 496 = 49.6 kJ. Per atom = 496 × 10³ J ÷ 6.022 × 10²³ = 8.24 × 10⁻¹⁹ J.

Electron gain enthalpy

Electron gain enthalpy, Δ_egH, is the enthalpy change when an electron is added to an isolated gaseous atom: X(g) + e⁻ → X⁻(g). Halogens have large negative values because one electron completes their octet. Noble gases have large positive values because the electron must enter a new shell. It generally becomes more negative across a period and less negative down a group, with one important exception: chlorine has the most negative value (−349 kJ mol⁻¹ against −328 for fluorine). The incoming electron in fluorine enters the small, crowded 2p subshell and is repelled by the electrons already there. For the same reason, sulphur (−200) is more negative than oxygen (−141).

Electronegativity

Electronegativity is the tendency of an atom in a chemical compound to attract a shared pair of electrons towards itself. It is not a measurable quantity like ionisation enthalpy; it is given on scales such as the Pauling scale, where fluorine, the most electronegative element, is 4.0. It increases across a period and decreases down a group, and higher electronegativity goes with more non-metallic character.

Periodic trends in chemical properties

Valence of representative elements is usually the number of valence electrons, or eight minus that number. This is why the hydrides of period 2 go LiH, B₂H₆, CH₄, NH₃, H₂O, HF, and the oxides of period 3 go Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₃, Cl₂O₇.

Anomalous behaviour of second period elements: Li, Be, B, C, N, O and F differ from the rest of their groups because of their small size, high charge/radius ratio, high electronegativity and because they have only four valence orbitals (2s and 2p), so their maximum covalency is 4. Boron forms only [BF₄]⁻, while aluminium forms [AlF₆]³⁻. They also form pπ-pπ multiple bonds easily (C=C, C≡C, N≡N, O=O). Some of them resemble the element diagonally below and to the right: Li and Mg, Be and Al, B and Si (the diagonal relationship).

Nature of oxides: elements on the left form basic oxides and those on the right form acidic oxides.

Chemical reactivity is highest at the two ends of a period: the alkali metals on the left lose an electron easily, and the halogens on the right gain one easily.

Common mistakes: (1) Saying fluorine has the most negative electron gain enthalpy; it is chlorine. (2) Forgetting the Be/B and N/O exceptions in ionisation enthalpy. (3) Treating electronegativity as a measurable energy with units; it has no unit. (4) Ranking isoelectronic ions by number of electrons instead of by nuclear charge. (5) Using Mendeleev's "atomic weight" wording for the modern periodic law, which is based on atomic number.

JEE and NEET focus

• Order of atomic and ionic radii, especially isoelectronic series and cation/anion comparisons.
• Order of ionisation enthalpies within a period and group, with the Be/B and N/O exceptions, and successive ionisation enthalpies.
• Electron gain enthalpy order among halogens and group 16, and why noble gases are positive.
• IUPAC temporary names and symbols for elements with Z > 100.
• Identifying block, period and group from an electronic configuration.
• Acidic, basic, amphoteric and neutral oxides, and the diagonal relationship pairs.

Practice questions