Chemistry · Chemistry of the elements · Year 1

d- and f-block elements: colour, redox

Why does a single element, vanadium, show four colours, and why are Zr and Hf almost twins?

The situation

The Almalyk Mining and Metallurgical Complex produces copper cathodes, zinc, molybdenum, silver and gold; the Navoi complex brings Muruntau gold into cyanide solution, and the Chirchik plant makes tungsten and molybdenum products. All of these are d-elements. In the laboratory, if zinc is dropped into a yellow solution of ammonium vanadate, the solution turns sky blue, green and finally violet; if tin is used, it stops at green. Why does one element appear in four different colours, how can we predict in advance where the reduction will stop, and why are zirconium and hafnium almost twins?

Topic

d- and f-block elements: oxidation states and coloured chemistry. Almalyk Mining and Metallurgical Combine (copper, zinc, molybdenum, gold, silver), Navoi Mining and Metallurgical Combine gold and the Chirchik refractory metals plant (tungsten, molybdenum)…

Key ideas

Oxidation states and colour
Frost diagram
Lanthanide contraction

Formulas

E = E^{\circ} + \frac{0.0592}{n} \lg \frac{[Ox] {[H^{+}]}^{m}}{[Red]}

The Nernst equation (25 °C): the potential depends on the ratio of concentrations and on pH.

E_{1 \to 3}^{\circ} = \frac{n_{1} E_{1}^{\circ} + n_{2} E_{2}^{\circ}}{n_{1} + n_{2}}

In a Latimer diagram, for non-adjacent pairs the E° values are added weighted by the number of electrons ( $Δ G^{\circ} = - n F E^{\circ}$ is what adds, not E°).

N E^{\circ} (X^{N} / X^{0}) = - Δ G^{\circ} / F

Frost diagram: the volt-equivalent for each state; the slope between two points is the E° of that pair.

2 VO_{2}^{+} + Zn + 4 H^{+} \overset{}{\to} 2 VO^{2 +} + Zn^{2 +} + 2 H_{2} O

The vanadium ladder: V(V) → V(IV) → V(III) → V(II), one electron at each step.

μ = \sqrt{n (n + 2)} μ_{B}

Spin-only magnetic moment (3d ions): $n$ is the number of unpaired electrons. For 4f ions the orbital contribution is large and the formula does not work.

Theory

d-elements: rows and oxidation states

d-elements are those in which the inner $(n - 1) d$ shell is being filled: 3d (Sc–Zn), 4d (Y–Cd), 5d (La, Hf–Hg). The valence electrons are in $(n - 1) d$ and $n s$ , so they form many oxidation states: up to the middle of a row the highest equals the group number (Sc +3, Ti +4, V +5, Cr +6, Mn +7), then it decreases; in 4d and 5d, Ru and Os reach +8 (RuO₄, OsO₄). Group 3 (Sc, Y, La) gives only +3 and its chemistry is close to that of Al and the alkaline-earth metals: the oxides are basic. How it differs from p-elements: in p-groups, going down, the lower states become more stable (inert pair: Pb(II), Tl(I)), whereas in d-groups it is the opposite — the higher states become more stable: CrO₃ is a strong oxidising agent, while MoO₃ and WO₃ are stable. The metallic radius first decreases along a row (Sc 162 → Ni 124 pm), then increases again near d¹⁰ (Zn 134 pm); the melting point is highest in groups 5–6 — the d-electrons take the greatest part in bonding (W 3422 °C) — and low for d¹⁰ metals (Hg −38.83 °C). The element card and trend graphs are in the “Rows” mode.

Acid–base properties of oxides and hydroxides

The rule for a single element: a low oxidation state is basic, an intermediate one amphoteric, a high one acidic: CrO → Cr₂O₃ → CrO₃; MnO → MnO₂ → Mn₂O₇. Group 4 (Ti, Zr, Hf): TiO and Ti₂O₃ are basic, Ti³⁺ is violet and a strong reducing agent; TiO₂ is amphoteric (titanates). E(OH)₄ are in fact hydrated oxides EO₂·nH₂O — along Ti → Zr → Hf basicity increases. Group 5: VO and V₂O₃ are basic, VO₂ is amphoteric, V₂O₅ is amphoteric (acidic character predominates): in alkali it gives vanadates (VO₄³⁻ and polyvanadates), in strong acid yellow VO₂⁺; Nb₂O₅ and Ta₂O₅ are inert and weakly acidic. Group 6: CrO₃ gives chromic acids in water; as acidity increases, polychromates form: CrO₄²⁻ → Cr₂O₇²⁻ → Cr₃O₁₀²⁻ → Cr₄O₁₃²⁻; MoO₃ and WO₃ are acidic oxides (molybdates, tungstates; reduced to the metal by hydrogen: $WO_{3} + 3 H_{2} \overset{}{\to} W + 3 H_{2} O$ ). Group 12: Zn(OH)₂ is amphoteric (in excess alkali [Zn(OH)₄]²⁻), Cd(OH)₂ is basic — it does not dissolve in excess alkali but dissolves in ammonia. The coloured compounds of chromium and manganese are covered in detail in the “Chromium and manganese: coloured ions” experiment.

Redox: Nernst, Latimer and Frost diagrams

The Nernst equation links potential to concentration and pH: for $VO_{2}^{+} + 2 H^{+} + e^{-} \overset{}{\to} VO^{2 +} + H_{2} O$ , $E = 1.00 + 0.0592 \lg ([V O_{2}^{+}] {[H^{+}]}^{2} / [V O^{2 +}])$ — when pH rises by 1, $E$ falls by 0.118 V. When Ox and Red are equal, $E = E^{\circ}$ : during reduction the potential of a Pt electrode in the solution “pauses” around these values. A Latimer diagram is the series of oxidation states with the E° of each step; the E° of a non-adjacent pair is the average weighted by the number of electrons: MnO₄⁻/Mn²⁺ = (0.56 + 2·2.26 + 0.95 + 1.54)/5 = 1.51 V. In a Frost diagram ( $N E^{\circ}$ against N) the lowest point is the most stable state (Mn²⁺, V³⁺, Cr³⁺), and the slope between two points is E°. A species above the line joining its neighbours disproportionates: Mn³⁺ ( $2 Mn^{3 +} + 2 H_{2} O \overset{}{\to} Mn^{2 +} + MnO_{2} + 4 H^{+}$ ), MnO₄²⁻ in acid, and Cu⁺ ( $2 Cu^{+} \overset{}{\to} Cu + Cu^{2 +}$ ). A reducing agent passes only the steps whose slope is greater than its own E°: Zn (−0.76 V) reduces vanadate to V²⁺, whereas Sn (−0.14 V) reduces it only to V³⁺. The products of permanganate in the three media (Mn²⁺, MnO₂, MnO₄²⁻) also depend on these potentials and on pH: $E (M n O_{4}^{-} / M n^{2 +}) = 1.51 - 0.0947 pH$ (MnO₄⁻ and Mn²⁺ at 1 mol/L).

The iron family, platinum metals, copper and gold

Manganese has +2, +3, +4, +7: Mn²⁺ is the most stable, Mn³⁺ and MnO₄²⁻ disproportionate in acid, MnO₄⁻ is a strong oxidising agent. Iron: Fe(0) in carbonyls (Fe(CO)₅), +2 and +3 are the main states, +6 — ferrates: in strong alkali $2 Fe (OH)_{3} + 3 ClO^{-} + 4 OH^{-} \overset{}{\to} 2 FeO_{4}^{2 -} + 3 Cl^{-} + 5 H_{2} O$ ; FeO₄²⁻ ( $E^{\circ}$ ≈ 2.20 V) is a stronger oxidising agent than permanganate and is used in water treatment. Complexes of iron: [Fe(CN)₆]⁴⁻ and [Fe(CN)₆]³⁻, Prussian blue $4 Fe^{3 +} + 3 [Fe (CN)_{6}]^{4 -} \overset{}{\to} Fe_{4} [Fe (CN)_{6}]_{3}$ (in detail — “Iron and its compounds”). Cobalt (+2, +3): Co³⁺ is a very strong oxidising agent in water (1.92 V), but is stable in ammine complexes — complex formation changes the potential (“Complex compounds and ligands”). Rhodium and iridium are mainly +3, +4. Nickel is +2 (Ni(OH)₂ is green, in ammonia blue-violet [Ni(NH₃)₆]²⁺); Pd and Pt are catalysts: hydrogenation (Pd), oxidation of ammonia to NO (Pt–Rh gauze), car catalytic converters (Pt, Pd, Rh). Copper: Cu(I) disproportionates in water, Cu(II) is stable. Silver: Ag₂O dissolves in ammonia as [Ag(NH₃)₂]⁺. Gold dissolves only in aqua regia: $Au + HNO_{3} + 4 HCl \overset{}{\to} H [AuCl_{4}] + NO ↑ + 2 H_{2} O$ — Cl⁻ binds gold as [AuCl₄]⁻ and lowers its potential; in Navoi, though, it is done with cyanide and atmospheric oxygen: $4 Au + 8 NaCN + O_{2} + 2 H_{2} O \overset{}{\to} 4 Na [Au (CN)_{2}] + 4 NaOH$ (Elsner). In detail — “Copper, silver, gold and zinc”.

f-elements: lanthanoids and actinoids

In the lanthanoids (La–Lu) the 4f shell is filled: Ce³⁺ 4f¹ … Lu³⁺ 4f¹⁴. The 4f electrons lie deep inside and hardly take part in chemical bonding — so the main state of all of them is +3 and their properties are very similar. Internal periodicity: an empty (f⁰), half-filled (f⁷) and full (f¹⁴) shell is stable, which is why Ce⁴⁺ (f⁰), Tb⁴⁺ (f⁷), Eu²⁺ (f⁷) and Yb²⁺ (f¹⁴) occur; the number of unpaired electrons is greatest in Gd³⁺ (7). Lanthanoid contraction: f-electrons shield the nuclear charge poorly, and the radius of Ln³⁺ decreases from La³⁺ 103.2 pm to Lu³⁺ 86.1 pm (CN 6). As a result the 5d metals are almost the same size as their 4d partners above: Zr⁴⁺ 72 / Hf⁴⁺ 71 pm, Nb⁵⁺ and Ta⁵⁺ 64 pm, Y³⁺ 90.0 ≈ Ho³⁺ 90.1 pm. Ln₂O₃ are basic oxides; as the radius shrinks, the basicity of Ln(OH)₃ decreases from La to Lu. The colours of Ln³⁺ ions (Pr³⁺ green, Nd³⁺ violet-pink, Er³⁺ pink) come from f–f transitions — their lines are narrow and weak; the magnetic moment does not obey the spin-only formula. In the actinoids (Ac–Lr) the 5f shell is filled; in the first elements the 5f electrons take more part in bonding and the states are varied: U +3 … +6 (uranyl UO₂²⁺), Np and Pu up to +7. Uranium is mined in the Navoi region.

Laboratory (works 21–24), safety and the model

The Test tubes mode repeats lab works 21–24: to salts of Cr³⁺, Mn²⁺, Fe²⁺, Fe³⁺, Co²⁺, Ni²⁺, Cu²⁺, Ag⁺, Zn²⁺, Cd²⁺ add a little and an excess of NaOH, an excess of NH₃, then H₂O₂ in an alkaline medium, and watch the colours of the precipitates, amphoterism, the formation of ammine complexes and oxidation; the results go into the log on the board. Safety: wear safety goggles and gloves; work with ammonia in a fume cupboard and do not sniff it. Cadmium compounds and Cr(VI) (chromates) are poisonous and carcinogenic — do not pour the waste down the sink. AgNO₃ stains the skin black. Aqua regia and ferrates are strong oxidising agents, to be used only in the fume cupboard. The model: the results in the test tubes follow a stoichiometry table (0.2 mmol of salt, one drop = 0.1 mmol of reagent); “excess” means 1.5 times the stoichiometric amount; oxidation in air is speeded up 10 times; to obtain the Co²⁺ ammine complex in practice NH₄Cl is also added. Vanadium reduction is a series of first-order steps (the rate constants are a model, time is speeded up 20 times), at pH 0 and 25 °C; the Pt electrode potential comes from the Nernst equation for the dominant pair. Colours are approximate, from the Beer–Lambert law. Data: metallic radii (CN 12) and melting points — CRC/Wikipedia tables, ionic radii — Shannon (1976), E° — Housecroft and CRC tables (Cr₂O₇²⁻/Cr³⁺ — 1.33 V; some newer tables give 1.36 V).

Did you know?

Hafnium was “hiding” in zirconium ores in a few per cent: the two are so alike that it was discovered only in 1923, in Copenhagen, by D. Coster and G. de Hevesy using X-ray spectra.
Vanadium was named in 1830 by the Swedish chemist N. Sefström, for the variety of colours of its compounds, after Vanadis, the Scandinavian goddess of beauty; before that, in 1801, A. M. del Río had found it in a Mexican ore, but his discovery was not accepted.

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

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
Patterns in s- and p-block elements — Year 1
Complex compounds and ligands — Year 1
Chemical plant: reactors and workshops — Year 4