The Periodic Table and Periodicity

Elements are arranged in order of increasing proton number. As you go along, properties repeat at regular intervals. This repeating pattern is periodicity.

Chapter 3 showed that the group gives the number of valence electrons and the period gives the number of shells. Two more facts complete the picture:

• The table divides into blocks named after the subshell being filled (Chapter 2): Groups 1 and 2 are the s-block, Groups 13 to 18 the p-block and Groups 3 to 12 the d-block.
• Elements in a group have the same outer electron configuration, so they have similar chemical properties.

Across a period the elements change from metals to non-metals, and their oxides change from basic through amphoteric to acidic (Chapter 5). Down a group, metallic character increases. Figure 8.1 sums up the trends that the next section explains.

Every trend in this section comes from three things you met in Chapter 2: the nuclear charge, the number of shells, and shielding by inner electrons.

• Across a period, the nuclear charge increases, but the electrons go into the same outer shell, so shielding stays about the same. The outer electrons are pulled in more strongly.
• Down a group, each element has one more filled shell. The outer electrons are further from the nucleus and more shielded, so they are held less strongly, even though the nuclear charge is larger.

8.2.1Atomic and ionic radius

The atomic radius is half the distance between the nuclei of two neighbouring atoms of the element. It decreases across a period and increases down a group (Figure 8.2).

• A cation is smaller than its atom: it has lost its outer shell, and the same nuclear charge pulls on fewer electrons.
• An anion is larger than its atom: the extra electrons repel one another and spread out.
• Na⁺ to Si⁴⁺ are isoelectronic (Chapter 2), so the radius falls as the nuclear charge rises. The same is true of P³⁻ to Cl⁻. P³⁻ is much larger than Si⁴⁺ because it has one more shell.

8.2.2First ionisation energy

The first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms, to form 1 mole of singly charged gaseous cations (Chapter 2). For sodium, Na(g) → Na⁺(g) + e⁻.

• Across a period it generally increases. The nuclear charge rises, but the electrons enter the same shell, so the shielding stays about the same. The outer electron is held more strongly, so more energy is needed to remove it.
• Down a group it decreases. The outer electron is in a shell further from the nucleus and more shielded by the inner shells. It is held less strongly, even though the nuclear charge is larger.

Period 3 has two dips in this rise (Chapter 2 drew the graph):

8.2.3Electronegativity

In Chapter 3 we met electronegativity briefly, as a rough picture: metals have few valence electrons and attract a shared pair weakly, while non-metals are close to a full shell and attract it strongly. Here we look at what electronegativity really is, how it varies and why.

The shared pair is pulled by the nucleus of each bonded atom, through that atom’s inner electrons. So electronegativity depends on the same things as ionisation energy: the nuclear charge, the distance of the shared pair from the nucleus, and the shielding.

• Across a period it increases. The nuclear charge rises with about the same shielding, and the atom gets smaller, so the shared pair is closer to the nucleus. Both pull the shared pair more strongly.
• Down a group it decreases. Each element has one more filled shell, so the shared pair is further from the nucleus and more shielded. This outweighs the larger nuclear charge.

Fluorine, top right, is the most electronegative element.

The melting point and conductivity of each Period 3 element follow from its structure (Chapter 3 and Figure 8.3).

• Na, Mg, Al: the number of delocalised electrons per atom rises (1, 2, 3) and the cation gets smaller and more highly charged, so metallic bonding gets stronger (Chapter 3). Melting point and conductivity rise.
• Si: a giant covalent structure. Very many strong covalent bonds must break, so it has the highest melting point. It is a semiconductor: it conducts a little, more at higher temperature.
• P, S, Cl: simple molecules held by weak instantaneous dipole–induced dipole forces, so low melting points. S₈ > P₄ > Cl₂, the order of the number of electrons in each molecule. They do not conduct: no mobile electrons or ions.

In its highest oxide or chloride, each element's oxidation state equals its number of valence electrons: +1 in Na₂O up to +6 in SO₃, and +1 in NaCl up to +5 in PCl₅. It is always positive, because oxygen and chlorine are more electronegative than every Period 3 element.

Across the period the electronegativity difference between the element and oxygen or chlorine shrinks, so the bonding changes from ionic to covalent. Structure decides behaviour: ionic oxides are basic, covalent oxides are acidic (Figure 8.4).

8.4.1The oxides

8.4.2The chlorides

Hydrolysis is the breaking of bonds by water. A small, highly charged cation such as Al³⁺ pulls electron density from the water molecules around it so strongly that they release H⁺:

Na⁺ is too large and weakly charged to do this, so NaCl gives a neutral solution. The covalent chlorides react completely with water, giving white fumes of HCl:

Lithium, sodium, potassium, rubidium and caesium each have one valence electron, which they lose to form M⁺ ions.

• They are soft (they can be cut with a knife), have low densities (lithium, sodium and potassium float on water) and low melting points.
• Their freshly cut surface is shiny but tarnishes quickly, so they are stored in oil.
• Melting points decrease down the group: Li 180 °C, Na 98 °C, K 63 °C, Rb 39 °C, Cs 29 °C. The metallic bond weakens as the atoms get larger.

Reactivity increases down the group. The atom gets larger, so its single valence electron is further from the nucleus and more shielded, and is lost more easily. Each metal is a reducing agent (Chapter 6), and the reducing power increases down the group.

Uses of Group 2 compounds: MgO lines furnaces (high melting point, low reactivity), CaCO₃ (limestone) is a building material, CaO (quicklime) treats acidic soil, and BaSO₄ is the barium meal in X-ray diagnosis.

Magnesium to barium have two valence electrons and form only M²⁺ ions: removing a third electron would mean breaking into an inner shell. Like Group 1, they become more reactive down the group, as the two outer electrons are lost more easily: their reducing power increases.

8.6.1Thermal stability of the carbonates

Group 2 carbonates break down on heating: MCO₃(s) → MO(s) + CO₂(g). The temperature needed rises down the group (Figure 8.5), so the carbonates become more thermally stable.

Why: down the group the M²⁺ ion gets larger, so its charge density (charge ÷ size) falls. A small, highly charged ion such as Mg²⁺ pulls on the electron cloud of the large carbonate ion and distorts it (polarises it), weakening the C–O bonds inside it. The carbonate then breaks up more easily. Ba²⁺ distorts it least, so BaCO₃ is the most stable.

Uses of the halogens: fluoride goes in toothpaste, chlorine disinfects water and ClO⁻ is the active part of bleach, silver bromide is in photographic film, and tincture of iodine is an antiseptic.

The halogens are non-metals made of diatomic molecules. Each atom has seven valence electrons and gains one to form a halide ion, X⁻.

Down the group the colours get darker and the melting and boiling points increase. The molecules have more electrons, so the instantaneous dipole–induced dipole forces between them are stronger. The halogens become less volatile.

8.7.1Reactivity and displacement

A halogen reacts by gaining an electron, so it is an oxidising agent. Reactivity decreases down the group: the larger the atom, the further the incoming electron is from the nucleus and the more it is shielded, so it is attracted less strongly.

8.7.2Hydrogen halides Core · A Level

The hydrogen halides become less stable to heat down the group. HF does not decompose even on strong heating, but HI gives purple fumes of iodine with a hot rod: 2HI(g) → H₂(g) + I₂(g). The larger halogen atom overlaps less well with hydrogen, so the bond gets weaker:

Helium, neon, argon, krypton, xenon and radon are colourless, monatomic gases with low melting and boiling points. Helium has two valence electrons and the others eight: their outer shells are full, so they rarely lose, gain or share electrons. They are very unreactive and are used where an inert atmosphere is needed: argon in steel-making and light bulbs, helium in balloons, neon in coloured lamps.

Metals can be placed in order of reactivity from how they react with water, steam and dilute acid (Figure 8.6). A more reactive metal loses its electrons more readily to form positive ions. Carbon and hydrogen are included, because they can reduce some metal oxides.

• Metals above hydrogen react with dilute hydrochloric acid to give hydrogen. Lead seems not to, because insoluble lead(II) chloride coats it (Chapter 5).
• Aluminium seems unreactive because a thin layer of aluminium oxide protects it.

Rusting is the corrosion of iron: its slow oxidation to hydrated iron(III) oxide.

Rusting needs both oxygen and water (Figure 8.7). Salt speeds it up, and so do acidic gases such as sulfur dioxide, so iron near the sea or in industrial areas rusts faster. Rust is flaky and falls off, exposing fresh iron, so rusting carries on until the metal is gone.