Notes · ChemistryUK · A-Levels

Period 3 elements, transition metals and reactions of ions in solution (A-level)

This A-level-only inorganic topic covers the reactions of the Period 3 elements and the acid-base character of their oxides, then the distinctive chemistry of the transition metals - complex ions and their shapes and isomerism, coloured ions and their d-d transitions, variable oxidation states and catalysis - and finishes with the reactions of metal-aqua ions in aqueous solution.

5 sections~15 min reading time3 competenciesLevel Advanced 5

T·151515 / 18

Exam profile

AO1 · Recall the Period 3 reactions, the definition and properties of transition metals and the metal-aqua ion reactionsAO2 · Write electron configurations and complex-ion equations and carry out redox-titration and colorimetry calculationsAO3 · Interpret colour changes, redox-titration data and reaction observations to identify ions

Operators:statedescribeexplaindeducecalculatepredict

basic level

This entire topic is A-level only (A2); it is not assessed at AS.

higher level

The full A-Level treats the Period 3 oxides, the definition and properties of transition metals, complex-ion shape and isomerism, the origin of colour, variable oxidation states, catalysis and the reactions of metal-aqua ions, with quantitative work on redox titrations and colorimetry.

Depth

Text

Contents · 5 sections

Period 3 elements, transition metals and reactions of ions in solution (A-level)

Period 3 elements and their oxides#

AdvancedAQA 7405 3.2.4.1DfE GCE Chemistry - Period 3 oxides

pH of the Period 3 oxides in water

Fig. 1The oxides change from basic (Na2O) to acidic (SO3) across the period; Al2O3 is amphoteric and SiO2 insoluble.

Sodium and magnesium react with water: sodium vigorously with cold water to give a strongly alkaline solution of sodium hydroxide and hydrogen, and magnesium slowly with cold water (but readily with steam to give magnesium oxide). Across Period 3 the elements also react with oxygen to form oxides in which each element commonly shows its group oxidation state -

\text{Na}_2\text{O}

\text{MgO}

\text{Al}_2\text{O}_3

\text{SiO}_2

\text{P}_4\text{O}_{10}

and

\text{SO}_2/\text{SO}_3

The oxides change in structure and bonding across the period, from ionic (giant lattices) in

\text{Na}_2\text{O}

\text{MgO}

and

\text{Al}_2\text{O}_3

, through the giant covalent

\text{SiO}_2

, to the simple molecular

\text{P}_4\text{O}_{10}

and the sulfur oxides. This change in structure controls both their melting points and their behaviour with water and acids.

The acid-base character of the oxides changes from basic to acidic across the period. The metal oxides

\text{Na}_2\text{O}

and

\text{MgO}

are basic: they react with water to give alkaline hydroxide solutions (sodium oxide giving a strongly alkaline solution around pH

13

14

, magnesium oxide a weakly alkaline one around pH

9

because it is only sparingly soluble). Aluminium oxide is amphoteric - insoluble in water but reacting with both acids and bases. The non-metal oxides are acidic:

\text{P}_4\text{O}_{10}

and the sulfur oxides dissolve to give acidic solutions (phosphoric and sulfuric/sulfurous acids), with low pH values.

This basic-to-acidic trend is a periodic property that reflects the metal-to-non-metal change across the period, and silicon dioxide sits between them - insoluble in water but reacting with hot concentrated alkali, so it is classed as an acidic oxide. When asked about an oxide, state its structure, then predict its reaction with water and its approximate pH, and give the equation - for example

\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH}

(basic) versus

\text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4

(acidic).

\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH}

A basic oxide

Sodium oxide dissolves to give a strongly alkaline solution (about pH 13-14).

\text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4

An acidic oxide

Sulfur trioxide dissolves to give a strongly acidic solution (about pH 1).

Worked example

Comparing two Period 3 oxides

Compare the reactions with water, and the resulting pH, of sodium oxide and sulfur trioxide.

Sodium oxide
A basic ionic oxide: Na2O + H2O -> 2NaOH, giving a strongly alkaline solution of about pH 13-14.
Sulfur trioxide
An acidic molecular oxide: SO3 + H2O -> H2SO4, giving a strongly acidic solution of about pH 1.
Explain the difference
The metal oxide is basic and the non-metal oxide acidic - the basic-to-acidic trend across Period 3.

Result: Sodium oxide gives a strongly alkaline solution and sulfur trioxide a strongly acidic one, illustrating the trend across the period.

Typical mistakes

Forgetting that Al2O3 is amphoteric (reacts with both acids and bases) and that SiO2 is insoluble but acidic.
Predicting MgO gives a strongly alkaline solution - it is only sparingly soluble, so only weakly alkaline (about pH 9).

Active revision

Predict the approximate pH and write an equation for the reaction with water (where one occurs) of (a) sodium oxide and (b) phosphorus(V) oxide.

Active recall

Recall the key points — then reveal.

Sources: GCE AS and A level subject content for the sciences (Department for Education) · AQA A-level Chemistry 7405 specification (AQA)

The transition metals: definition and properties#

AdvancedAQA 7405 3.2.5.1DfE GCE Chemistry - transition metals

Why scandium and zinc are not transition metals

Fig. 2Only elements forming an ion with a partially filled d sub-shell are transition metals, which excludes Sc and Zn.

A transition metal is a d-block element that forms at least one stable ion with a partially filled d sub-shell. This precise definition matters, because it excludes two d-block elements: scandium (whose only ion,

\text{Sc}^{3+}

, has an empty

3d

sub-shell) and zinc (whose only ion,

\text{Zn}^{2+}

, has a full

3d^{10}

sub-shell). The transition metals of Period 4 therefore run from titanium to copper.

Their electron configurations fill the

3d

sub-shell after the

4s

; chromium and copper are the two exceptions, being

[\text{Ar}]\,3d^5\,4s^1

and

[\text{Ar}]\,3d^{10}\,4s^1

because a half-filled or full

3d

sub-shell is extra stable. When these metals form ions, the

4s

electrons are lost before the

3d

electrons, so, for example,

\text{Fe}

(

[\text{Ar}]\,3d^6\,4s^2

) forms

\text{Fe}^{2+}

(

[\text{Ar}]\,3d^6

) and

\text{Fe}^{3+}

(

[\text{Ar}]\,3d^5

The partially filled d sub-shell gives the transition metals four characteristic properties: they form complex ions, they form coloured compounds, they show variable oxidation states, and they (or their compounds) act as catalysts. These four properties recur throughout transition-metal chemistry and all stem from the availability of the d electrons and the empty d orbitals of similar energy.

The transition metals also have the general properties of metals - high melting points, high density, and good electrical conductivity - because the involvement of both

4s

and

3d

electrons in metallic bonding makes that bonding strong. But it is the four special properties above, arising from the partially filled d sub-shell, that make transition-metal chemistry distinctive and that this topic develops in turn.

Worked example

Configurations of a transition-metal ion

Write the electron configurations of a manganese atom (Z = 25) and of the Mn2+ ion.

Manganese atom
Filling to 25 electrons: [Ar] 3d5 4s2.
Form Mn2+
Remove the two 4s electrons first: [Ar] 3d5.
Confirm it is a transition metal
Mn2+ has a partially filled 3d5 sub-shell, so manganese is a transition metal.

Result: Mn is [Ar]3d5 4s2 and Mn2+ is [Ar]3d5, a partially filled d sub-shell confirming transition-metal character.

Typical mistakes

Removing 3d electrons before 4s when forming a transition-metal ion.
Calling scandium or zinc a transition metal despite their empty or full d sub-shell in their ions.

Active revision

Explain, with reference to electron configurations, why zinc is not classified as a transition metal.

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Complex ions, shape and isomerism#

AdvancedAQA 7405 3.2.5.2DfE GCE Chemistry - complex ions

An octahedral complex ion

Fig. 3Six monodentate ligands give an octahedral complex with 90-degree bond angles, as in [Cu(H2O)6]2+.

A complex ion consists of a central metal ion surrounded by ligands bonded to it by coordinate (dative covalent) bonds - each ligand donates a lone pair into an empty orbital of the metal. A ligand may be monodentate (donating one lone pair, such as

\text{H}_2\text{O}

\text{NH}_3

\text{Cl}^-

) or bidentate/multidentate (donating two or more, such as 1,2-diaminoethane 'en' or the hexadentate EDTA). The coordination number is the number of coordinate bonds to the central metal ion.

The shape of a complex depends on the coordination number and the size of the ligands. Six-coordinate complexes, the most common (with small ligands such as water and ammonia), are octahedral with

90^{\circ}

bond angles, for example

\text{[Cu(H}_2\text{O)}_6\text{]}^{2+}

. Four-coordinate complexes are usually tetrahedral with

109.5^{\circ}

angles (with larger ligands such as chloride, for example

\text{[CuCl}_4\text{]}^{2-}

), but some, notably the platinum complex cisplatin, are square planar.

Complexes show two kinds of stereoisomerism. Cis-trans (E-Z) isomerism occurs in square planar complexes such as

\text{[Pt(NH}_3)_2\text{Cl}_2\text{]}

and in octahedral complexes with two different ligands, where the two identical ligands can be adjacent (cis) or opposite (trans). Optical isomerism occurs in octahedral complexes containing bidentate ligands, such as

\text{[Ni(en)}_3\text{]}^{2+}

, which exist as non-superimposable mirror images.

This stereochemistry has real importance: cisplatin, the cis isomer of

\text{[Pt(NH}_3)_2\text{Cl}_2\text{]}

, is an anti-cancer drug that binds to DNA and prevents cancer cells dividing, whereas the trans isomer is far less effective - a striking example of how the arrangement in space, not just the formula, determines chemical behaviour. When drawing a complex, show the shape correctly and, where relevant, both stereoisomers.

Worked example

Predicting the shape of a complex

Predict and explain the shapes of [Cu(H2O)6]2+ and [CuCl4]2-.

Coordination numbers
[Cu(H2O)6]2+ has six small water ligands; [CuCl4]2- has four larger chloride ligands.
Assign shapes
Six ligands give an octahedral shape (90 degrees); four larger ligands give a tetrahedral shape (109.5 degrees).
Explain the difference
The larger chloride ligands cannot fit six around the copper, so only four bond, giving the tetrahedral complex.

Result: [Cu(H2O)6]2+ is octahedral (90 degrees) and [CuCl4]2- is tetrahedral (109.5 degrees), the coordination number set by the ligand size.

Typical mistakes

Assuming every four-coordinate complex is tetrahedral - some, such as cisplatin, are square planar.
Failing to show both cis and trans arrangements, or confusing which is the active anti-cancer isomer (the cis).

Active revision

Draw the two stereoisomers of the square planar complex $\text{[Pt(NH}_3)_2\text{Cl}_2\text{]}$ and state which is the anti-cancer drug cisplatin.

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Colour, variable oxidation states and catalysis#

AdvancedAQA 7405 3.2.5.3AQA 7405 3.2.5.4AQA 7405 3.2.5.5DfE GCE Chemistry - colour, oxidation states and catalysis

d-orbital splitting and a d-d transition

Fig. 4An electron absorbs a photon of energy equal to the gap and jumps between the split d levels; the complement of the absorbed colour is seen.

Transition-metal complexes are coloured because their d electrons can absorb visible light. In a complex the five d orbitals are split into two energy levels of slightly different energy; an electron can absorb a photon of visible light of exactly the energy gap

\Delta E

and jump from the lower to the higher level (a d-d transition). The light that is absorbed is removed from white light, and the colour seen is the complement of the absorbed colour. Because

\Delta E

depends on the metal's oxidation state, the ligands and the coordination number, changing any of these changes the colour.

The size of the energy gap, and hence the colour and the frequency of light absorbed, is related to the absorbed frequency by

\Delta E = h\nu

(with the frequency and wavelength linked by

c = \nu\lambda

). Colorimetry uses this: the concentration of a coloured complex can be found by measuring how much light of the absorbed colour it absorbs, using a calibration graph of absorbance against concentration.

Transition metals show variable oxidation states because the

4s

and

3d

electrons are close in energy, so different numbers of electrons can be lost. Iron forms

+2

and

+3

, manganese ranges from

+2

+7

, and this makes transition-metal ions useful oxidising and reducing agents. Their redox reactions are used in redox titrations - most importantly the titration of iron(II) with potassium manganate(VII), which is self-indicating (the purple

\text{MnO}_4^-

turns colourless until the first excess drop gives a permanent pink).

Transition metals and their compounds are effective catalysts, in two modes. A heterogeneous catalyst is in a different phase from the reactants and works by adsorbing them onto active sites on its surface (iron in the Haber process, vanadium(V) oxide in the Contact process, and platinum/palladium in catalytic converters). A homogeneous catalyst is in the same phase and works via an intermediate of a different oxidation state; autocatalysis, where a product of the reaction catalyses it, is seen in the manganate(VII)-ethanedioate reaction, catalysed by the

\text{Mn}^{2+}

it produces. Both modes rely on the ability of the metal to change oxidation state or to bind reactants using its d orbitals.

\Delta E = h\nu = \dfrac{hc}{\lambda}

Energy of a d-d transition

The energy gap between the split d orbitals equals the energy of the photon absorbed; it fixes the colour seen.

\text{MnO}_4^{-} + 8\text{H}^+ + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} + 5\text{Fe}^{3+}

Manganate(VII)-iron(II) titration

The 1:5 ratio underlies the redox-titration calculation; the reaction is self-indicating.

Redox titration of iron(II) with manganate(VII)

Fig. 5The titration is self-indicating: the purple manganate(VII) decolourises until the first excess drop turns the solution permanently pink.

Worked example

Manganate(VII) redox titration

25.0 cm^3 of an acidified iron(II) solution is titrated with 0.0200 mol dm^-3 KMnO4 and requires 24.0 cm^3. Calculate the concentration of Fe2+.

Moles of MnO4-
$n = cV = 0.0200 \times 24.0/1000 = 4.80\times10^{-4}\ \text{mol}$ .
Moles of Fe2+
The ratio is 1 MnO4- : 5 Fe2+, so $n(\text{Fe}^{2+}) = 5 \times 4.80\times10^{-4} = 2.40\times10^{-3}\ \text{mol}$ .
Concentration
$c = n/V = \dfrac{2.40\times10^{-3}}{25.0/1000} = 0.0960\ \text{mol dm}^{-3}$ .
$c = \dfrac{2.40\times10^{-3}}{0.0250} = 0.0960\ \text{mol dm}^{-3}$

Result: The iron(II) concentration is 0.0960 mol dm^-3.

Typical mistakes

Saying the colour seen is the colour absorbed - it is the complementary colour.
Using the wrong ratio (it is 1 MnO4- to 5 Fe2+) in a redox-titration calculation.

Active revision

In a redox titration, $25.0\ \text{cm}^3$ of an iron(II) solution reacts exactly with $24.0\ \text{cm}^3$ of $0.0200\ \text{mol dm}^{-3}$ potassium manganate(VII). Calculate the concentration of the iron(II) solution.

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Reactions of metal-aqua ions in aqueous solution#

AdvancedAQA 7405 3.2.6.1DfE GCE Chemistry - reactions of ions in solution

Reactions of the hexaaquacopper(II) ion

Fig. 6The aqua ion gives a blue hydroxide precipitate with base, and undergoes ligand substitution with excess ammonia and with concentrated HCl.

In aqueous solution transition-metal ions exist as metal-aqua ions,

\text{[M(H}_2\text{O)}_6\text{]}^{n+}

, an octahedral complex of six water ligands. These ions are acidic because the positive charge on the metal draws electron density away from the coordinated water, weakening the O-H bonds so that a water ligand can release a proton (hydrolysis). The higher the charge density on the metal, the more acidic the ion, so

3+

ions such as

\text{[Fe(H}_2\text{O)}_6\text{]}^{3+}

are noticeably more acidic than

2+

ions such as

\text{[Fe(H}_2\text{O)}_6\text{]}^{2+}

Adding a base removes protons from the aqua ion in successive steps, forming a neutral metal hydroxide precipitate. With a small amount of sodium hydroxide or ammonia, both

2+

and

3+

aqua ions give coloured hydroxide precipitates - for example

\text{[Cu(H}_2\text{O)}_6\text{]}^{2+}

gives a pale blue precipitate of

\text{Cu(OH)}_2

, and iron(III) gives a red-brown precipitate. The colours of these precipitates are used to identify the metal ion present.

With an excess of the base, further reactions distinguish the ions. Excess sodium hydroxide redissolves amphoteric hydroxides (such as aluminium hydroxide) but not others. Excess ammonia acts as a ligand and can substitute for the water: copper(II) forms the deep-blue

\text{[Cu(NH}_3)_4\text{(H}_2\text{O)}_2\text{]}^{2+}

by ligand substitution. Adding sodium carbonate distinguishes

2+

from

3+

ions:

2+

ions give a metal carbonate precipitate, whereas the more acidic

3+

ions are acidic enough to react with the carbonate to release carbon dioxide and give a metal hydroxide.

Ligand substitution reactions - in which one ligand replaces another around the metal - underlie several of these observations and are accompanied by colour changes. Concentrated hydrochloric acid replaces the six water ligands of copper(II) with four chloride ligands, changing the octahedral blue

\text{[Cu(H}_2\text{O)}_6\text{]}^{2+}

to the tetrahedral yellow

\text{[CuCl}_4\text{]}^{2-}

. When you answer these questions, state the observation (colour, precipitate, effervescence), name the type of reaction, and give the equation.

\text{[Cu(H}_2\text{O)}_6\text{]}^{2+} + 4\text{NH}_3 \rightarrow \text{[Cu(NH}_3)_4\text{(H}_2\text{O)}_2\text{]}^{2+} + 4\text{H}_2\text{O}

Ligand substitution with ammonia

Four ammonia ligands substitute for four water ligands, giving the deep-blue ammine complex.

Worked example

Distinguishing 2+ and 3+ ions with carbonate

Explain, in terms of acidity, why iron(III) aqua ions react with sodium carbonate to give carbon dioxide, whereas iron(II) aqua ions give only a precipitate.

Acidity of the ions
The 3+ iron(III) aqua ion has a higher charge density, so it is far more acidic than the 2+ iron(II) aqua ion.
Reaction of iron(III)
It is acidic enough to react with carbonate ions, releasing CO2 and forming iron(III) hydroxide (a red-brown precipitate).
Reaction of iron(II)
The weakly acidic iron(II) aqua ion simply forms an iron(II) carbonate precipitate, with no effervescence.

Result: The greater acidity of the 3+ ion lets it release CO2 from carbonate, whereas the 2+ ion only precipitates the carbonate.

Typical mistakes

Saying a 2+ aqua ion reacts with carbonate to give CO2 - only the more acidic 3+ ions do; 2+ ions give the carbonate precipitate.
Confusing ligand substitution (colour change from a new complex) with simple precipitation.

Active revision

Describe what is observed, and write equations, when sodium carbonate solution is added separately to solutions of $\text{[Fe(H}_2\text{O)}_6\text{]}^{2+}$ and $\text{[Fe(H}_2\text{O)}_6\text{]}^{3+}$ .

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

References & sources

Sources

Department for Education

GCE AS and A level subject content for the sciences

AQA

AQA A-level Chemistry 7405 specification

Period 3 elements, transition metals and reactions of ions in solution (A-level)

5 sections~15 min reading time3 competenciesLevel Advanced 5

Period 3 elements, transition metals and reactions of ions in solution (A-level)

Key points

Comparing two Period 3 oxides

Key points

Configurations of a transition-metal ion

Key points

Predicting the shape of a complex

Key points

Manganate(VII) redox titration

Key points

Distinguishing 2+ and 3+ ions with carbonate

Sources

Period 3 elements, transition metals and reactions of ions in solution (A-level)

Key points

Comparing two Period 3 oxides

Key points

Configurations of a transition-metal ion

Key points

Predicting the shape of a complex

Key points

Manganate(VII) redox titration

Key points

Distinguishing 2+ and 3+ ions with carbonate

Sources