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Structure 3.1 · The periodic table

1Periods, groups and blocks 3.1.1 SL + HL Repeated

In 1869 Dmitri Mendeleev arranged the 63 elements then known in order of increasing atomic mass and placed elements with similar chemical properties in the same vertical column. He left gaps where the pattern demanded an element that had not yet been found, and predicted the properties of those missing elements from their neighbours. When gallium (1875) and germanium (1886) were discovered, their densities, formulas of their oxides and chlorides, and other properties matched his predictions closely. A classification that predicts successfully is more than a filing system: it reflects an underlying cause. That cause, discovered some fifty years later, is the arrangement of electrons.

The modern periodic table arranges the elements in order of increasing atomic number, Z, the number of protons in the nucleus. Ordering by atomic number removes the few anomalies that ordering by mass produced (argon, A_r = 39.95, correctly precedes potassium, A_r = 39.10, because argon has 18 protons and potassium 19).

The shape of the table is a direct consequence of sublevel capacities. An s sublevel holds two electrons, so the s-block is two columns wide; a p sublevel holds six, a d sublevel ten and an f sublevel fourteen, so the p-, d- and f-blocks are six, ten and fourteen columns wide. Period 1 contains only two elements because level 1 has only a 1s sublevel. Periods 2 and 3 contain eight each (s and p). Period 4 contains eighteen, because the 3d sublevel fills between 4s and 4p.

Metals, metalloids and non-metals

About three-quarters of the elements are metals. They occupy the left-hand side and the centre of the table: the s-block (except hydrogen and helium), the whole d- and f-blocks, and the lower left of the p-block (aluminium, gallium, indium, tin, thallium, lead and bismuth). Non-metals occupy the upper right of the p-block, together with hydrogen. Between them a diagonal band of metalloids — boron, silicon, germanium, arsenic, antimony and tellurium — shows intermediate properties: they look metallic and conduct electricity weakly (they are semiconductors), but chemically they behave more like non-metals.

Four families are named in the syllabus and should be known by name and position: the alkali metals (group 1), the halogens (group 17), the noble gases (group 18) and the transition elements (the d-block, with the qualifications discussed in section 8).

Exam focus · what the published papers show

2Position and electron configuration, in both directions 3.1.2 SL + HL Repeated

The syllabus statement: The period number shows the outer energy level that is occupied by electrons. Elements in a group have a common number of valence electrons. The skill: deduce the electron configuration of an atom up to Z = 36 from the element's position in the periodic table and vice versa.

Electrons occupy the lowest available energy sublevels first (the Aufbau principle). Each orbital holds at most two electrons of opposite spin (the Pauli exclusion principle), and electrons occupy the orbitals of a sublevel singly before pairing (Hund’s rule). The order in which sublevels fill up to Z = 36 is

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

The only surprise is that 4s fills before 3d. In an isolated atom of potassium or calcium the 4s sublevel is lower in energy than 3d, so the nineteenth and twentieth electrons enter 4s; the ten 3d electrons follow from scandium to zinc. Reading the periodic table from left to right along each period reproduces exactly this order, which is why the table can be used instead of memorising it.

The full electron configuration lists every occupied sublevel, for example sulfur, 1s²2s²2p⁶3s²3p⁴. The condensed form replaces the inner electrons by the symbol of the preceding noble gas in square brackets: [Ne]3s²3p⁴. Both are accepted unless a question asks for one of them by name. For atoms beyond argon the 3d sublevel may be written before or after 4s — [Ar]3d⁶4s² and [Ar]4s²3d⁶ describe the same iron atom.

Two atoms up to Z = 36 do not follow the simple filling order. Chromium is [Ar]3d⁵4s¹ and copper is [Ar]3d¹⁰4s¹. In both, one 4s electron occupies the 3d sublevel instead, giving a half-filled or completely filled 3d sublevel. These two exceptions must be learned.

Exam focus · what the published papers show

3Periodicity: the five trends 3.1.3 SL + HL Strong

The syllabus statement: Periodicity refers to trends in properties of elements across a period and down a group. One line asks a great deal: explain the periodicity of atomic radius, ionic radius, ionization energy, electron affinity and electronegativity. Five properties, and one explanation underneath all of them.

A periodic property repeats its pattern each time a new period begins. All five properties in this outcome are periodic because each is controlled by how strongly the nucleus attracts the outer electrons, and that attraction follows the structure of the table. Three factors decide it.

The combined effect of nuclear charge and shielding is sometimes described as the effective nuclear charge: the net positive charge experienced by an outer electron. Across period 3, from sodium to chlorine, the number of protons rises from 11 to 17 while the number of inner (shielding) electrons stays at ten, so the effective nuclear charge acting on the outer electrons increases steadily. Down group 1 the number of protons rises sharply, but each new element adds a complete inner shell, so the extra protons are almost exactly cancelled by extra shielding; what does change is the distance of the outer electron from the nucleus.

Atomic radius

An atom has no sharp edge, so its size is defined operationally. The atomic (covalent) radius is half the distance between the nuclei of two identical atoms joined by a single covalent bond. For metals the radius is half the distance between neighbouring nuclei in the metallic lattice.

Across a period atomic radius decreases. Electrons are added to the same outer energy level while the nuclear charge increases and shielding stays approximately constant, so the outer electrons are pulled closer to the nucleus. Down a group atomic radius increases, because each successive element has its outer electrons in a higher energy level, further from the nucleus; the increase in nuclear charge is offset by the extra shielding of the added inner shell.

Noble gases are usually omitted from atomic-radius comparisons. They form almost no covalent bonds, so their radii are measured differently (from non-bonded contact distances) and are not directly comparable with covalent radii.

Ionic radius

When an atom loses electrons to form a cation, the remaining electrons are attracted by an unchanged number of protons, and a whole outer energy level is often emptied. Cations are therefore smaller than their parent atoms: Na is [Ne]3s¹, but Na⁺ is [Ne], with its outer electrons in level 2 instead of level 3. When an atom gains electrons to form an anion, the extra electrons increase repulsion between electrons while the nuclear charge is unchanged, so anions are larger than their parent atoms.

Ionic radius follows the group trend in the same way as atomic radius — it increases down a group. Across a period the pattern breaks into two series. From Na⁺ to Si⁴⁺ (all with the neon configuration) the radius decreases as nuclear charge increases; from P³⁻ to Cl⁻ (all with the argon configuration) it also decreases for the same reason; but the anions are much larger than the cations because they have an extra occupied energy level. Ions with the same number of electrons are called isoelectronic, and within an isoelectronic series radius depends only on the number of protons.

First ionization energy

Ionization is always endothermic: energy must be supplied to overcome the attraction between the electron and the nucleus. The state symbol (g) is essential — the definition refers to isolated atoms, not to the solid element.

Across a period first ionization energy increases in general, because the nuclear charge increases while electrons are removed from the same energy level with similar shielding. Down a group it decreases, because the electron removed is further from the nucleus and more shielded, so less energy is needed despite the larger nuclear charge. The general rise across a period is interrupted by two small dips in each of periods 2 and 3; their explanation is additional higher level content (section 7), but you should expect them on any graph of first ionization energy.

Electron affinity

For most elements the first electron affinity is exothermic (negative), because the incoming electron is attracted by the nucleus. It becomes more exothermic across a period in general (the nuclear charge attracting the new electron increases) and less exothermic down a group (the electron enters a level further from the nucleus). There are well-known exceptions: fluorine’s value is less exothermic than chlorine’s because the added electron enters the small, crowded 2p sublevel where electron–electron repulsion is large, and the noble gases and group 2 elements have endothermic (positive) values because the electron would have to enter a new, higher-energy sublevel or level.

Second electron affinities, X⁻(g) + e⁻ → X²⁻(g), are always endothermic: the electron is being added to an ion that is already negative, so energy is needed to overcome the repulsion.

Electronegativity

Electronegativity increases across a period (greater nuclear charge, similar shielding, smaller radius) and decreases down a group (bonding pair further from the nucleus and more shielded). Noble gases are usually given no value because they form very few bonds. The difference in electronegativity between two bonded atoms determines bond polarity, which links this outcome to Structure 2.2.

Exam focus · what the published papers show

4Down a group: the alkali metals and the halogens 3.1.4 SL + HL Repeated

The syllabus statement: Trends in properties of elements down a group include the increasing metallic character of group 1 elements and decreasing non-metallic character of group 17 elements. The skill is to describe and explain the reactions of group 1 metals with water, and of group 17 elements with halide ions — describe and explain, which is two jobs.

A group is the clearest place to watch periodicity at work, because every member has the same outer electron arrangement and only the size of the atom changes. Going down a group, metallic character increases: atoms lose electrons more easily because their outer electrons are further from the nucleus and more shielded. In group 1 this makes the metals more reactive; in group 17, whose atoms react by gaining an electron, the same change makes the elements less reactive. The two groups therefore show opposite reactivity trends for the same underlying reason.

Group 1: the alkali metals

Lithium, sodium, potassium, rubidium and caesium each have a single electron in an s sublevel outside a noble-gas core (ns¹). They are soft, low-density metals with low melting points, and they are stored under oil because they react rapidly with air and water. Their melting points decrease down the group because the metallic bonding weakens as the cation radius increases (Structure 2.3).

Every alkali metal reacts with cold water to form a hydroxide and hydrogen gas. The general equation, and the ionic equation that shows what actually happens to the metal, are:

2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)     2M(s) + 2H₂O(l) → 2M⁺(aq) + 2OH⁻(aq) + H₂(g)

The reaction is a redox reaction: each metal atom is oxidized to M⁺ (oxidation state 0 → +1) and hydrogen in water is reduced to H₂ (+1 → 0). The hydroxide ions make the resulting solution alkaline; adding universal indicator shows a colour change to blue or purple.

Group 17: the halogens

The halogens exist as diatomic molecules, X₂, with an ns²np⁵ outer arrangement: one electron short of a noble-gas configuration. Their physical states change down the group because the London (dispersion) forces between the molecules increase with the number of electrons, not because of any change in reactivity.

Halogens react by gaining electrons: they are oxidizing agents. A more reactive halogen oxidizes the halide ions of a less reactive halogen, displacing it from solution. Chlorine displaces bromine and iodine; bromine displaces iodine; iodine displaces neither.

Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)    colourless solution turns yellow-orange

Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)    colourless solution turns brown

Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)    orange solution turns brown

In each equation the halogen is reduced (0 → −1) and the halide ion is oxidized (−1 → 0). The spectator ions (for example K⁺ from potassium bromide) are omitted from the ionic equation. Where the colours are difficult to distinguish, shaking the mixture with a non-polar solvent such as hexane helps: iodine dissolves in the organic layer to give a purple (violet) colour and bromine an orange colour.

Exam focus · what the published papers show

5The metal/non-metal continuum and the oxides 3.1.5 SL + HL Repeated

The syllabus statement: Metallic and non-metallic properties show a continuum. This includes the trend from basic metal oxides through amphoteric to acidic non-metal oxides. The skill: deduce equations for the reactions with water of the oxides of group 1 and group 2 metals, carbon and sulfur, and the scope note adds acid rain caused by gaseous non-metal oxides, and ocean acidification caused by increasing CO2 levels.

There is no sharp boundary between metals and non-metals. Across period 3, sodium, magnesium and aluminium are metals; silicon is a metalloid; phosphorus, sulfur, chlorine and argon are non-metals. The change is gradual because electronegativity rises steadily across the period, so an element’s tendency to lose electrons (metallic behaviour) turns gradually into a tendency to share or gain them (non-metallic behaviour). The oxides show this continuum most clearly, because both their bonding and their acid–base behaviour follow it.

A basic oxide reacts with water to form a hydroxide (an alkali if soluble), or reacts with an acid to form a salt and water. An acidic oxide reacts with water to form an acid, or reacts with a base to form a salt and water. An amphoteric oxide reacts with both acids and bases. Metal oxides are basic because the oxide ion, O²⁻, is a strong base: it accepts a proton from water, O²⁻ + H₂O → 2OH⁻. Non-metal oxides are acidic because the non-metal atom, bonded covalently to oxygen, reacts with water to form an oxoacid that releases H⁺ ions.

Aluminium oxide shows its amphoteric character by reacting with both hydrochloric acid and sodium hydroxide solution:

Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)     Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq)

Silicon dioxide does not react with water because its giant covalent structure is very stable, but it is classed as acidic because it reacts with hot concentrated alkali to form silicates.

Acid deposition and ocean acidification

Unpolluted rain is naturally slightly acidic, with a pH of about 5.6, because carbon dioxide dissolves in it to form carbonic acid. Acid deposition refers to rain, snow, fog or dry particles with a pH below this, caused by gaseous non-metal oxides released mainly by human activity:

• Sulfur oxides come chiefly from burning fossil fuels (especially coal) that contain sulfur: S + O₂ → SO₂. In the atmosphere SO₂ is oxidized to SO₃, and both react with water to form H₂SO₃ and H₂SO₄.
• Nitrogen oxides form when nitrogen and oxygen from the air combine at the high temperatures inside internal combustion engines and power stations. Nitrogen dioxide reacts with water to form nitric acid (and nitrous acid): 2NO₂(g) + H₂O(l) → HNO₃(aq) + HNO₂(aq).

Acid deposition damages buildings and statues made of limestone and marble (CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂), corrodes metals, lowers the pH of lakes, harming aquatic life, and leaches nutrients and toxic Al³⁺ ions from soils, damaging forests. It is reduced by removing sulfur from fuels before combustion, by scrubbing flue gases with calcium oxide or carbonate after combustion, and by catalytic converters that reduce NO_x emissions from vehicles.

Ocean acidification is a separate consequence of increasing atmospheric CO₂. As more carbon dioxide dissolves in sea water, the equilibria

CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)

shift to the right, increasing [H⁺] and lowering the pH. The additional H⁺ ions react with carbonate ions, H⁺ + CO₃²⁻ → HCO₃⁻, reducing the carbonate available to corals, shellfish and plankton that build their shells and skeletons from calcium carbonate.

Exam focus · what the published papers show

6Oxidation states 3.1.6 SL + HL Strong

The syllabus statement: The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge that atom would have if the compound were composed of ions. This outcome appeared in every session in the collection behind these notes — and one report names it by number, listing Structure 3.1.6 deducing oxidation states among the areas students were well prepared for.

An oxidation state is a bookkeeping device. For every atom in a species we pretend that each bond is fully ionic, with the bonding electrons assigned entirely to the more electronegative atom, and ask what charge the atom would then carry. In sodium chloride the oxidation states (+1 and −1) equal the real ionic charges. In water, oxygen is assigned both bonding pairs and receives an oxidation state of −2, each hydrogen +1, even though the O–H bonds are polar covalent and no ions exist. The value of the device is that it tracks electron transfer consistently across ionic and covalent substances, which is exactly what is needed to analyse redox reactions (Reactivity 3.2).

Many elements, particularly non-metals and transition elements, show several oxidation states. Names show the oxidation state as a Roman numeral in brackets after the element: iron(II) chloride, FeCl₂; iron(III) chloride, FeCl₃; manganese(IV) oxide, MnO₂. The same convention names oxyanions — sulfate(VI), SO₄²⁻; sulfate(IV), SO₃²⁻; nitrate(V), NO₃⁻; nitrate(III), NO₂⁻ — although the traditional names sulfate, sulfite, nitrate and nitrite remain acceptable.

Exam focus · what the published papers show

Structure 3.1 · Additional higher level

7Discontinuities in first ionization energy 3.1.7 HL only Strong

The syllabus statement: Discontinuities occur in the trend of increasing first ionization energy across a period. The skill is unusual and worth reading twice: explain how these discontinuities provide evidence for the existence of energy sublevels — the breaks in the trend are being used as evidence, not merely described.

The simple shell model of the atom (Structure 1.3) predicts that first ionization energy should rise smoothly across a period, because the nuclear charge increases while the electrons are removed from the same shell. The measured values rise overall, but in periods 2 and 3 they fall back twice. Those two falls cannot be explained by shells alone. They are evidence that each shell is divided into sublevels of different energy, and that the orbitals within a sublevel can hold electrons singly or in pairs.

The first discontinuity — group 2 to group 13 (Be → B, Mg → Al). Beryllium’s outermost electron is in 2s; boron’s is in 2p. The 2p sublevel is higher in energy than 2s, and the 2p electron is partly shielded by the filled 2s sublevel, so less energy is needed to remove it. The effect outweighs boron’s extra proton, and the first ionization energy falls.

The second discontinuity — group 15 to group 16 (N → O, P → S). Nitrogen, 2p³, has one electron in each of its three 2p orbitals. Oxygen, 2p⁴, must place its fourth 2p electron in an orbital that is already occupied. The two electrons in that orbital repel each other, raising the energy of the paired electron, so it is removed more easily than any of nitrogen’s unpaired electrons.

Successive ionization energies: evidence for shells

Successive ionization energies — removing a first, second, third electron and so on from the same atom — provide the complementary evidence for main energy levels. Each successive value is larger, because every electron is removed from an increasingly positive ion. A large jump appears whenever the next electron must come from a new, inner shell, much closer to the nucleus and far less shielded. For magnesium the first two values are 738 and 1451 kJ mol⁻¹, but the third is 7733 kJ mol⁻¹: two electrons are in the outer shell, so magnesium is in group 2. Reading the position of the first large jump is a reliable way to find the group of an unknown element.

Exam focus · what the published papers show

8Transition elements and their properties 3.1.8 HL only Repeated

The syllabus statement: Transition elements have incomplete d-sublevels that give them characteristic properties. You are asked to recognize properties, including: variable oxidation state, high melting points, magnetic properties, catalytic properties, formation of coloured compounds and formation of complex ions with ligands — six properties, all traceable to one structural fact.

All ten elements from scandium to zinc are in the d-block, but the definition is stricter than block membership. Zinc is not a transition element: the atom is [Ar]3d¹⁰4s² and its only ion, Zn²⁺, is [Ar]3d¹⁰ — the d sublevel is full in both. Scandium is debated: the atom, [Ar]3d¹4s², has an incomplete d sublevel, so it meets the definition, but its only common ion, Sc³⁺, has an empty d sublevel and its compounds show few of the characteristic properties (they are white, and scandium has a single oxidation state). The syllabus poses this as a nature-of-science question: the classification depends on whether the definition refers to the atom or to the ions.

The incomplete d sublevel gives transition elements a characteristic set of properties.

A complex ion consists of a central metal ion surrounded by ligands — molecules or anions with at least one lone pair of electrons, such as H₂O, NH₃, Cl⁻, OH⁻ and CN⁻. Each ligand donates a lone pair into an empty orbital of the metal ion, forming a coordination (dative) bond. In acid–base terms the ligand is a Lewis base and the metal ion a Lewis acid (Reactivity 3.4). The number of coordination bonds is the coordination number; six (octahedral) and four (tetrahedral or square planar) are the most common. The charge of the complex is the sum of the metal ion’s charge and the ligands’ charges: [Fe(H₂O)₆]³⁺ carries 3+, while [CoCl₄]²⁻ carries +2 + 4(−1) = 2−.

Exam focus · what the published papers show

9Variable oxidation states, and the configurations of the ions 3.1.9 HL only Repeated

The syllabus statement: The formation of variable oxidation states in transition elements can be explained by the fact that their successive ionization energies are close in value. The skill: deduce the electron configurations of ions of the first-row transition elements.

The successive ionization energies of calcium show a huge jump between the second and third values, because the third electron would have to come from the inner 3p sublevel. Calcium is therefore found only as Ca²⁺. In a transition element the 4s and 3d sublevels are close in energy, so removing the third, fourth or fifth electron does not require breaking into a new inner shell: successive ionization energies rise gradually, with no sharp jump to set a limit. The energy required to form a higher oxidation state can be recovered from lattice enthalpy, hydration enthalpy or covalent bond formation, so several oxidation states are stable in compounds.

Almost all first-row transition elements show +2, formed by losing the two 4s electrons. The maximum oxidation state rises from scandium (+3) to manganese (+7), equal to the total number of 4s and 3d electrons; beyond manganese the higher states become harder to reach as the nuclear charge increases.

When a transition element forms a positive ion, the 4s electrons are removed first, even though 4s was filled before 3d. Once the 3d orbitals are occupied they fall below 4s in energy, so the 4s electrons are the outermost and highest in energy. Only after the 4s sublevel is empty are 3d electrons removed.

Exam focus · what the published papers show

10Why transition element complexes are coloured 3.1.10 HL only Repeated

The syllabus statement: Transition element complexes are coloured due to the absorption of light when an electron is promoted between the orbitals in the split d-sublevels. The colour absorbed is complementary to the colour observed. The skill: apply the colour wheel to deduce the wavelengths and frequencies of light absorbed and/or observed. The scope note is generous — students are not expected to know the different splitting patterns and their relation to the coordination number.

In an isolated transition-metal ion the five 3d orbitals have the same energy. When ligands approach, their lone pairs repel the d electrons. Some d orbitals point more directly towards the ligands than others, so they are raised in energy more: the d sublevel splits into two groups separated by a small energy gap, ΔE. This gap corresponds to the energy of photons of visible light.

When white light passes through a solution of the complex, an electron in a lower-energy d orbital absorbs a photon whose energy exactly matches ΔE and is promoted to a higher-energy d orbital. Light of that frequency is removed from the beam; the light that is transmitted (or reflected) is the complementary colour, found opposite the absorbed colour on the colour wheel. The relationships needed, both given in the data booklet, are

ΔE = hf     and     c = λf

where h is the Planck constant (6.63 × 10⁻³⁴ J s), f the frequency in s⁻¹ (Hz), c the speed of light (3.00 × 10⁸ m s⁻¹) and λ the wavelength in m. A larger splitting absorbs higher-frequency, shorter-wavelength light.

The size of ΔE, and therefore the colour, depends on three factors:

• The identity of the metal — different nuclear charge and d-electron count: [Co(H₂O)₆]²⁺ is pink, [Cu(H₂O)₆]²⁺ blue; [NiCl₄]²⁻ and [CoCl₄]²⁻ differ in colour although ligand and charge are the same.
• The oxidation state of the metal — a higher charge attracts the ligands more strongly and increases the splitting: [Fe(H₂O)₆]²⁺ is green, [Fe(H₂O)₆]³⁺ orange.
• The identity of the ligand — ligands split the d orbitals by different amounts: adding concentrated ammonia to blue [Cu(H₂O)₆]²⁺ gives the deep blue [Cu(NH₃)₄(H₂O)₂]²⁺; [Cr(H₂O)₆]³⁺ is violet but [Cr(NH₃)₆]³⁺ yellow; adding concentrated hydrochloric acid to pink [Co(H₂O)₆]²⁺ gives blue [CoCl₄]²⁻.

The colour of a complex is also used quantitatively. In colorimetry (or spectrophotometry) light of the wavelength most strongly absorbed is passed through the solution; the absorbance is proportional to the concentration of the coloured species, so a calibration curve of absorbance against known concentrations lets an unknown concentration be read off (Tool 1, Inquiry 2).

Exam focus · what the published papers show

Past-paper practice · Practice set R · Additional higher level: ionization energies and transition elements HL only

Attempt these before opening the solutions below. References give the session, level, paper and question number of the original examination.

R1IB · May 2023 · HL Paper 1 · TZ1 · Q5 · [1]

Which statement best explains the first ionization energy of sulfur being lower than that of phosphorus?

A. Sulfur has more protons than phosphorus.B. Phosphorus does not have paired electrons in the outer p sub-level.C. Sulfur has an unpaired electron in the outer p sub-level.D. Phosphorus is more reactive than sulfur.

R2IB · May 2022 · HL Paper 1 · TZ2 · Q5 · [1]

What is the correct order for increasing first ionization energy?

A. Na < Mg < AlB. Na < Al < MgC. Al < Mg < NaD. Al < Na < Mg

R3IB · May 2019 · HL Paper 1 · TZ1 · Q8 · [1]

Which electrons are removed from iron (Z = 26) to form iron(II)?

A. two 3d electronsB. two 4s electronsC. one 4s electron and one 3d electronD. two 4p electrons

R4IB · May 2023 · HL Paper 1 · TZ1 · Q8 · [1]

Which element is not a transition metal?

A. CrB. MnC. NiD. Zn

R5IB · November 2023 · HL Paper 1 · TZ1 · Q7 · [1]

Which group of elements have the most similar atomic radii?

A. Li, Be, B, CB. Fe, Co, Ni, CuC. K, Ca, Br, KrD. Ne, Ar, Kr, Xe

R6IB · November 2022 · HL Paper 1 · Q6 · [1]

Which best explains why complexes of d-block elements are coloured?

A. Light is absorbed when electrons are promoted between d orbitals.B. Light is emitted when electrons are promoted between d orbitals.C. Light is absorbed when electrons return to lower energy d orbitals.D. Light is emitted when electrons return to lower energy d orbitals.

R7IB · May 2021 · HL Paper 1 · TZ2 · Q8 · [1]

Which factor does not affect the colour of a complex ion?

A. temperature of the solutionB. identity of the ligandC. identity of the metalD. oxidation number of the metal

R8IB · May 2022 · HL Paper 1 · TZ2 · Q8 · [1]

	Wavelength of light absorbed by [Cr(OH2)6]3+	d-level splitting caused by H2O compared to NH3 ligands
A.	λ = 5.8 × 10−7 m	H2O > NH3
B.	λ = 5.8 × 10−7 m	H2O < NH3
C.	λ = 4.1 × 10−7 m	H2O > NH3
D.	λ = 4.1 × 10−7 m	H2O < NH3

R9IB · November 2023 · HL Paper 2 · TZ1 · Q5(g) · [4]

Explain, in terms of nuclear charge, electron subshells and the shielding provided by filled electron shells, why the first ionization energy increases from Li to Be, but decreases from Be to B. [4]

R10IB · November 2023 · HL Paper 2 · TZ1 · Q5(e) · [4]

Iron(III) chloride also exists as a dimer in the vapour phase, but iron, unlike beryllium, is a transition element.

(i) Outline, in terms of its electronic structure, what identifies a transition element. [1]

(ii) The first four ionization energies of beryllium and iron are shown.

One common property of transition elements is that they have variable oxidation states. Discuss, referring to the graph, why iron, but not beryllium, displays this characteristic. [3]

R11IB · May 2023 · HL Paper 2 · TZ2 · Q8(c) · [5]

Cobalt also forms chlorides with the formula CoCl₂.

(i) State the full electron configuration of the cobalt(II) ion, Co²⁺. [1]
(ii) Hydrated cobalt(II) ions, Co(H₂O)₆²⁺, are pink. Describe the interaction between the cobalt ion and a water molecule in terms of the type of bond and how this bond is formed. [2]
CoCl₄²⁻ ions are blue.
(iii) Explain why the different ligands cause different coloured complexes. [2]

R12IB · May 2019 · HL Paper 2 · TZ1 · Q6(a)–(b) · [3]

This question is about iron.

(a) Deduce the full electron configuration of Fe²⁺. [1]
(b) Explain why, when ligands bond to the iron ion causing the d-orbitals to split, the complex is coloured. [2]

R13IB · May 2021 · HL Paper 2 · TZ1 · Q3(f)–(g) · [3]

(f) Outline why, unlike typical transition metals, zinc compounds are not coloured. [1]
(g) Transition metals like iron can form complex ions. Discuss the bonding between transition metals and their ligands in terms of acid–base theory. [2]

Solutions and mark-scheme guidance · Set R

R1 B

Phosphorus has 3p³, one electron in each 3p orbital. In sulfur, 3p⁴, one orbital holds a pair; repulsion between the paired electrons makes one of them easier to remove. A is true but predicts the opposite result.

R2 B

First ionization energies: Na 496, Al 578, Mg 738 kJ mol⁻¹. Aluminium’s electron is removed from 3p, higher in energy than magnesium’s 3s electron, so Al < Mg despite its extra proton.

R3 B

Fe is [Ar]3d⁶4s²; Fe²⁺ is [Ar]3d⁶. The 4s electrons are removed first.

R4 D

Zinc forms only Zn²⁺, [Ar]3d¹⁰: a full d sublevel in the atom and in its only ion, so it is not a transition element.

R5 B

Across the first transition series the added electrons enter the inner 3d sublevel, which shields the outer 4s electrons almost as fast as the nuclear charge rises, so the radii change very little.

R6 A

Colour is due to absorption as an electron is promoted between split d orbitals; the transmitted light is the complementary colour. Options involving emission describe a different phenomenon.

R7 A

The colour depends on the energy gap between the split d orbitals, which is set by the metal, its oxidation state and the ligands. Temperature is not one of these factors in the syllabus model.

R8 B

[Cr(OH₂)₆]³⁺ appears violet, so it absorbs the complementary colour, yellow, λ = 5.8 × 10⁻⁷ m. [Cr(NH₃)₆]³⁺ appears yellow, so it absorbs violet, shorter wavelength and higher energy: NH₃ causes the larger splitting, H₂O < NH₃.

R9 [4]

Nuclear charge / number of protons increases «for both» ✔; Li and Be «outer electrons have» the same subshell / shielding ✔; the electron in B is lost from a p-subshell whereas that in Be is lost from an s-subshell ✔; the «outer electron in» B / p-subshell experiences greater shielding / has higher energy ✔. The scheme does not accept explanations invoking the distance of electrons from the nucleus.

R10 [4]

(i) Has a partially filled d sub-shell «in a common oxidation state» ✔.
(ii) IE values of Fe increase gradually AND IE values of Be show a sudden rise (after the second) ✔; «Be:» first and second ionization energies close together, therefore it does not form a +1 oxidation state / singly charged ion (Be always loses two electrons, only +2) ✔; «Fe:» further IEs are close to the second, so the number of electrons Fe loses / its oxidation state can vary ✔. The answer must be built from the graph.

R11 [5]

(i) 1s²2s²2p⁶3s²3p⁶3d⁷ ✔
(ii) Type of bond: coordinate / dative / covalent ✔; how it forms: the oxygen of the water molecule (ligand) donates an electron pair to the cobalt(II) ion ✔.
(iii) The magnitude of ΔE / energy gap between the split d-orbitals differs according to the ligand ✔; ΔE determines the wavelength of light absorbed / colour of the complex ✔. A reference to the different energy gap or splitting is needed for the first mark.

R12 [3]

(a) 1s²2s²2p⁶3s²3p⁶3d⁶ ✔
(b) «Frequency/wavelength of visible» light is absorbed by electrons moving between d levels/orbitals ✔; the colour is due to the remaining frequencies OR the complementary colour is transmitted ✔.

R13 [3]

(f) «Zn²⁺» has a full d-shell OR does not form ions with an incomplete d-shell ✔. The scheme does not accept “Zn is not a transition metal” (that restates the fact without explaining it) or answers about zinc atoms instead of ions.
(g) Ligands donate pairs of electrons to metal ions OR form coordinate covalent/dative bonds ✔; ligands are Lewis bases AND metal ions are Lewis acids ✔.

Review · Structure 3.1

11Misconceptions, the examiner’s view, and the question types

Six question types cover most of this topic, and each has a fixed opening move.

12Quick check

13Summary and knowledge organiser

Essential knowledge

• Elements are arranged by atomic number in 7 periods and 18 groups; the period number is the outer energy level, the group gives the valence electrons, and the block is the sublevel being filled.
• Metals lie left and centre, non-metals upper right, metalloids (B, Si, Ge, As, Sb, Te) along the diagonal between them.
• Across a period: atomic radius decreases; ionization energy, electron affinity (in general) and electronegativity increase. Down a group the reverse. Cause: nuclear charge, shielding and energy level.
• Cations are smaller and anions larger than their atoms; in an isoelectronic series radius decreases as proton number increases.
• Group 1 reactivity with water increases down the group; group 17 oxidizing power decreases down the group.
• Oxides change from basic (metals) through amphoteric (Al₂O₃) to acidic (non-metals); SO_x and NO_x cause acid deposition; CO₂ causes ocean acidification.
• Oxidation states sum to the charge on the species; elements are 0; H is −1 in metal hydrides and O is −1 in peroxides.
• HL Dips in first IE (Be→B, Mg→Al; N→O, P→S) are evidence for sublevels; successive IE jumps are evidence for shells.
• HL Transition elements have an incomplete d sublevel (atom or common ion): variable oxidation states, high melting points, magnetism, catalysis, coloured compounds, complex ions. Zn is not one.
• HL 4s electrons are lost first; colour arises from d–d absorption, and the complementary colour is observed.

Essential definitions and equations

Knowledge organiser