Electron Configuration and the Periodic Table: Blocks, Properties and Trends

The periodic table is an electron-configuration chart

Mendeleev arranged the elements by their properties in 1869, long before anyone knew about electrons. The quantum mechanical model explains why his arrangement works: elements in the same group have the same valence configuration, and the shape of the table follows the sublevels.

Remember from Grade 11: in SCH3U you explained the trends in atomic radius, ionization energy, electron affinity and electronegativity using effective nuclear charge (Zeff) and the distance of valence electrons from the nucleus. You also saw small "exceptions" in ionization energy that were promised an explanation in Grade 12. Here it is.

The four blocks

• s block (groups 1 and 2, plus helium): the last electron enters an ns orbital. Valence configuration ns¹ or ns².
• p block (groups 13 to 18): the last electron enters an np orbital. Valence configuration ns² np¹ to ns² np⁶.
• d block (groups 3 to 12, the transition metals): the last electron enters an (n − 1)d orbital. Period 4 transition metals fill 3d.
• f block (the lanthanides and actinides, usually printed below the table): the last electron enters an (n − 2)f orbital.

The widths of the blocks match the sublevel capacities: 2, 6, 10 and 14. Period lengths (2, 8, 8, 18, 18, 32, 32) follow from which sublevels fill in each period. You can read a configuration straight off the table: walk across the periods to the element, writing each block you pass. For example, selenium (period 4, fourth p column) is [Ar] 4s² 3d¹⁰ 4p⁴.

Characteristic properties of each block

s block: soft, reactive metals with low ionization energies. They lose their one or two valence electrons easily to form 1+ (group 1) or 2+ (group 2) ions. Group 1 metals react vigorously with water. (Hydrogen, though placed above group 1 because it is 1s¹, is a non-metal.)

p block: the most varied block. It contains metals (aluminium, tin, lead), metalloids (silicon, germanium) along a staircase line, non-metals (carbon, nitrogen, oxygen, the halogens) and the noble gases, whose full ns² np⁶ configuration makes them very unreactive. Non-metals here tend to gain electrons (forming anions such as Cl⁻ and O²⁻) or share them.

d block (transition metals): hard, dense metals with high melting points (except mercury). Because 4s and 3d electrons are close in energy, many form ions with several different charges (iron: 2+ and 3+; copper: 1+ and 2+; manganese: from 2+ up to 7+). Many of their compounds are coloured (copper(II) sulfate solution is blue) and many are paramagnetic because of unpaired d electrons. Many are good catalysts (iron in ammonia manufacture, platinum and palladium in car catalytic converters).

f block: the lanthanides are very similar to one another chemically (their differences are buried in the inner 4f sublevel), which makes them hard to separate. Some are used in strong magnets, such as neodymium magnets in wind turbines, electric motors and earbuds. The actinides are all radioactive; uranium and plutonium are nuclear fuels.

Explaining the trends with sublevels

The general trends from Grade 11 still hold: across a period, Zeff increases, so atoms get smaller and ionization energy (IE) and electronegativity increase; down a group, valence electrons are in higher n levels, farther from the nucleus, so atoms get larger and IE decreases.

Now look at the period 2 ionization energies in the graph. There are two dips:

1. Be → B (899 → 801 kJ/mol). Boron's outermost electron is in a 2p orbital, which is higher in energy than 2s and partly shielded by the 2s electrons. It is easier to remove, even though boron's nuclear charge is higher.
2. N → O (1402 → 1314 kJ/mol). Nitrogen's three 2p electrons are each alone in an orbital. Oxygen's fourth 2p electron must pair with another in the same orbital. The extra electron-electron repulsion in that orbital makes it easier to remove.

The same two dips appear in period 3 (Mg → Al and P → S).

Successive ionization energies show the shells

Removing electrons one after another takes more energy each time, but there is a huge jump when you start removing core electrons from a full inner level. For magnesium ([Ne] 3s²):

• IE₁ = 738 kJ/mol, IE₂ = 1451 kJ/mol, then IE₃ = 7733 kJ/mol.

The jump after the second electron shows magnesium has 2 valence electrons, which is why it forms Mg²⁺ and never Mg³⁺. Successive ionization energies are some of the strongest experimental evidence that electrons are arranged in levels.