Chemistry · Chemistry of the elements · Year 1

Patterns in s- and p-block elements

Why does caesium explode in water, lithium fizz gently, and fluorine act as the strongest oxidising agent?

The situation

At Kungrad soda is made from table salt, at Dehkanabad fertiliser from potassium salts, at Mubarek yellow sulfur from sour gas, and at Navoi nitric acid from the nitrogen in the air. All these elements stand on the left (s) and right (p) wings of the periodic table. Why do the elements of one group resemble each other, yet their properties change so sharply as we go down: caesium explodes in water, lithium fizzes slowly, fluorine even oxidises water, and iodine is a much weaker oxidising agent? Could a chemist who knows the pattern predict the properties of a compound nobody has seen yet?

Topic

s- and p-block elements: trends down the groups. The left and right wings of the periodic table on a map of Uzbekistan: Kungrad Soda Plant (Na), Dehkanabad Potash Plant (K), Kyzylkum Phosphorite Complex (P), Mubarek…

Key ideas

Trends down a group
Diagonal relationship
Inert pair effect

Formulas

I_{1} : Li 520 > Na 496 > K 419 > Rb 403 > Cs 376 kJ/mol

Down the group the ionisation energy falls — metallic character and reactivity increase.

2 M + 2 H_{2} O \overset{}{\to} 2 MOH + H_{2} ↑

Alkali metals (M = Li … Cs) with water: alkali and hydrogen.

Cl_{2} + 2 KBr \overset{}{\to} 2 KCl + Br_{2}

The halogen displacement series: F₂ > Cl₂ > Br₂ > I₂.

E_{i \to j}^{\circ} = \frac{\sum n_{k} E_{k}^{\circ}}{\sum n_{k}}

On a Latimer diagram, the potential between species that are not neighbours ( $n_{k}$ — the electrons in each step).

PbO_{2} + 4 HCl \overset{}{\to} PbCl_{2} + Cl_{2} ↑ + 2 H_{2} O

The inert-pair effect: Pb(IV) → Pb(II) — lead(IV) oxide oxidises chloride to chlorine.

Theory

Abundance, geochemistry and the elements as simple substances

In the Universe, hydrogen (≈ 74 %) and helium (≈ 24 %) predominate by mass: they were formed in the Big Bang, while heavier nuclei (up to iron) were synthesised in stars by thermonuclear reactions; elements with even $Z$ are more abundant than their neighbours (the Oddo–Harkins rule) — this is cosmochemistry. In the Earth's crust (mass %): O 46.1 · Si 28.2 · Al 8.2 · Fe 5.6 · Ca 4.2 · Na 2.4 · Mg 2.3 · K 2.1 — eight elements make up 99 %, mostly as silicates and aluminosilicates. Geochemistry groups the elements by their companions: lithophile (in silicates: Na, K, Mg, Al, Si), chalcophile (in sulfides: Cu, Zn, Pb, Ag), siderophile (with iron: Ni, Co, Au, Pt), atmophile (N₂, noble gases). In Uzbekistan: the Kyzylkum phosphorites, the Dehkanabad potassium salts (sylvinite), the Kungrad table salt, the Mubarek sour gas (H₂S → sulfur) — “Mineral resources”. Structure of the elements as simple substances: s-block metals — metallic lattice; N₂, O₂, the halogens and noble gases — molecular; B, C (diamond), Si — atomic (covalent) lattice; allotropy: O₂/O₃, white and red phosphorus, rhombic and monoclinic sulfur, diamond and graphite. Compounds of variable composition (berthollides) do not obey the law of multiple proportions: wüstite $F e_{1 - x} O$ (x ≈ 0.05–0.15), many oxides and sulfides; molecular substances of fixed composition are daltonides.

Internal and secondary periodicity, oxidation states

Across a period the nuclear charge increases: the radius decreases, and the ionisation energy and electronegativity increase; down a group the reverse is true. Within a period there are “jumps” (internal periodicity): the $I_{1}$ of N ( $2 p^{3}$ , half-filled) is larger than that of O, and the $I_{1}$ of Be is larger than that of B. Secondary periodicity: properties do not change smoothly down a group but in a “saw-tooth” pattern — after the 3d subshell is filled, Ga, Ge, As, Se and Br “feel” a poorly shielded nucleus: the covalent radius of Ga is almost equal to that of Al (122 and 121 pm), its $I_{1}$ is even slightly larger (579 and 578 kJ/mol), and compounds of As(V), Se(VI) and Br(VII) are less stable than those of P, S and Cl (HBrO₄ was first obtained only in 1968). Kainosymmetry: the 2p elements (B…F) fill a p-orbital for the first time, have no d-orbitals and are small — they form strong π-bonds (N≡N, C=O) and have a coordination number of at most 4 (BF₄⁻, CF₄, NO₃⁻), while in period 3 it can reach 6 (AlF₆³⁻, SiF₆²⁻, PCl₆⁻, SF₆) — that is why N₂ is a gas but phosphorus is P₄, and CO₂ is a gas but SiO₂ is quartz. Oxidation state: the highest equals the group number (in groups 13–18 the number − 10), the lowest equals the group number − 18; in heavy elements the stable state is lower by 2 — the inert pair (6p: the relativistically contracted $6 s^{2}$ ). The periodic table itself is in the experiment “Elements and the periodic table”.

s-block elements: groups 1 and 2

The alkali metals ( $n s^{1}$ ) are the most reactive metals: $I_{1}$ falls from Li 520 to Cs 376 kJ/mol and reactivity increases — K ignites in water, and Rb and Cs explode. Yet $E^{\circ} (Li^{+} / Li)$ = −3.04 V is even more negative than that of Cs: the hydration energy of the small Li⁺ ion is very large. The thermodynamic “strength” (E°) and the rate of reaction (kinetics: melting point, ionisation energy) are different things. The basicity of the hydroxides increases LiOH < NaOH < KOH < RbOH < CsOH. Lithium resembles Mg (diagonal relationship): Li₂CO₃, LiF and Li₃PO₄ are sparingly soluble, and in air it forms the nitride Li₃N. Group 2 ( $n s^{2}$ ): Be(OH)₂ is amphoteric (like Al(OH)₃ — a diagonal relationship), Mg(OH)₂ is a weak base, and Ca(OH)₂, Sr(OH)₂ and Ba(OH)₂ are alkalis; the solubility and basicity of the hydroxides increase down the group, while those of the sulfates decrease (MgSO₄ is very soluble, BaSO₄ practically insoluble). Water hardness — Ca²⁺ and Mg²⁺: temporary (hydrogencarbonates; CaCO₃ precipitates on boiling) and permanent (sulfates and chlorides; softened with soda or by ion exchange). Hydrogen (L10): prepared by $Zn + 2 HCl \overset{}{\to} ZnCl_{2} + H_{2} ↑$ ; H₂O₂ is both an oxidising and a reducing agent. For the school-level details, see “Alkali metals and soda”, “Water purification and softening” and “Hydrogen — the lightest gas”.

Halogens and chalcogens (groups 17 and 16)

Hydrogen halides: the H–X bond weakens down the group (565 → 299 kJ/mol), and the acid strength and reducing power increase: HF is a weak acid (p $K_{a}$ 3.2), boils at 19.5 °C because of hydrogen bonding, and attacks glass; HI is the strongest. The oxidising power of the halogens F₂ > Cl₂ > Br₂ > I₂ is the displacement series. Oxides: OF₂, Cl₂O, ClO₂, Cl₂O₇, I₂O₅. Oxoacids HClO → HClO₂ → HClO₃ → HClO₄: the strength and stability increase, and the practical oxidising power decreases; their salts are hypochlorites (bleach), chlorates (Berthollet's salt KClO₃) and perchlorates. The Latimer diagram (on the board) shows disproportionation: in alkali $Cl_{2} + 2 NaOH \overset{}{\to} NaCl + NaClO + H_{2} O$ . Interhalogen compounds: ClF, ClF₃, BrF₅, IF₇, ICl — by VSEPR T-shaped, square pyramidal and pentagonal bipyramidal. Group 16: in the H₂E hydrides the acidity increases H₂O < H₂S < H₂Se < H₂Te, the bond angle approaches 90° from 104.5°, and the stability decreases; the boiling point of H₂O is anomalous (hydrogen bonds). H₂SO₃ (reducing agent) → H₂SeO₃ (oxidises SO₂, releasing selenium) → H₂TeO₃; H₂SO₄ → H₂SeO₄ (the strongest oxidising agent, dissolves gold when heated) → H₆TeO₆ (a weak acid). School level — “Non-metals and halogens”, “Chlorine and hydrochloric acid”, “The oxygen group and sulfur”, “Sulfuric acid and the contact process”, “Oxygen and ozone”.

Groups 15 and 14

Nitrogen: its hydrides are NH₃, N₂H₄ (hydrazine) and HN₃ (hydrazoic acid); its oxides are N₂O (+1), NO (+2), N₂O₃ (+3), NO₂/N₂O₄ (+4) and N₂O₅ (+5); HNO₂ is weak and both oxidising and reducing. In nitric acid the oxidising agent is NO₃⁻ (N⁺⁵), so hydrogen is hardly released with metals: the more dilute the acid and the more reactive the metal, the deeper the nitrogen is reduced (NO₂ → NO → N₂O → N₂ → NH₄⁺); concentrated HNO₃ passivates Fe, Al and Cr. Phosphorus: white (P₄, toxic, ignites spontaneously) and red; its oxoacids are H₃PO₄ (tribasic), H₄P₂O₇, (HPO₃)ₙ, H₃PO₃ (dibasic) and H₃PO₂ (monobasic) — the basicity equals the number of P–OH groups. As, Sb, Bi: metallic character increases down the group, the oxides go from acidic to amphoteric to basic, the stability of EH₃ decreases, and for Bi the +3 state is stable. Group 14: CO (toxic), CO₂, carbonates, carbides; SiO₂ has a giant atomic lattice; silicates; silicic acid (weaker than carbonic acid); Ge, Sn and Pb have amphoteric oxides and hydroxides, Sn(II) is a reducing agent and Pb(IV) an oxidising agent. EH₄ hydrides: CH₄ is stable, SiH₄ ignites spontaneously in air, GeH₄ and SnH₄ are unstable, and PbH₄ has hardly been obtained. School level — “Nitrogen, ammonia and nitric acid”, “Phosphorus and mineral fertilisers”, “Carbon and its compounds”, “Silicon and silicates”.

Groups 13 and 18, safety and the model

Boron is a non-metal: B₂O₃ and H₃BO₃ are acidic; boric acid does not donate a proton but takes up OH⁻ (a Lewis acid, p $K_{a}$ 9.24); borates: orthoborate BO₃³⁻, metaborate (BO₂⁻)ₙ and polyborates (borax Na₂B₄O₇·10H₂O); boranes (B₂H₆) — electron-deficient three-centre bonds. Al → Tl: the basicity of E(OH)₃ increases — Al(OH)₃ and Ga(OH)₃ are amphoteric, In(OH)₃ is a base, and Tl(I) is stable (TlOH is a strong base). Noble gases: a full $n s^{2} n p^{6}$ shell, the highest $I_{1}$ and no electron affinity — that is why they are inert; the $I_{1}$ of Xe (1170 kJ/mol) is close to that of the O₂ molecule — in 1962 N. Bartlett prepared Xe⁺[PtF₆]⁻; XeF₂ (linear), XeF₄ (square planar), XeF₆, XeO₃. Safety: alkali metals only from under kerosene with tweezers, a small piece, behind a protective screen; compounds of Be, Tl, As, Sb, Pb and Hg are toxic — in a school laboratory they are used only under the teacher's supervision or not at all; use a fume cupboard for chlorine, bromine, H₂S, SO₂ and NO₂; wear goggles and gloves with concentrated acids; add acid to water, never the reverse; never point a test tube at yourself or anyone else. The model: the properties of the elements are reference values (NIST, CRC, Pauling); the observations in the test tubes are idealised, and the product shown is the main product; the nitric acid scheme is simplified; the Be(OH)₂ experiment is a model; the map is schematic, and the locations of the plants are approximate.

Did you know?

In 1962 Neil Bartlett prepared the first noble-gas compound — xenon hexafluoroplatinate: seeing that an oxygen molecule combines with PtF₆, he noticed that the ionisation energy of xenon is very close to that of O₂.
Almost half the mass of the Earth's crust is oxygen (≈ 46 %), and more than a quarter is silicon (≈ 28 %): together they form quartz and the silicates — the basis of rocks.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

More experiments in this section

Alkali metals and soda — Grade 9
Aluminium and its compounds — Grade 9
Copper, silver, gold and zinc — Grade 9
Chromium and manganese: coloured ions — Grade 9
Iron and its compounds — Grade 9
d- and f-block elements: colour, redox — Year 1
Complex compounds and ligands — Year 1
Chemical plant: reactors and workshops — Year 4