The periodic table: classification of elements — IB Chemistry HL Revision Notes & Topical Questions

Section 1 of 62Layout of the periodic table

Layout of the periodic table

Order

• Order elements are arranged by increasing atomic number, 1 to 118.

Periods

• Periods rows; the period number equals the number of occupied electron shells.

Groups

• Groups columns (numbered 1–18);
  • elements in a group have the same number of valence electrons and similar chemistry.

Hydrogen

• Hydrogen placed alone because its properties match no single group.

Helium

• Helium placed in Group 18 because it is unreactive like the noble gases.

Periodicity

• Periodicity physical and chemical properties follow repeating patterns across periods and down groups.

Section 2 of 62Metals, non-metals and blocks

1Metals

• Metals left and centre; good conductors, malleable and ductile.

2Non-metals

• Non-metals right-hand side (p block); poor conductors, brittle solids, often forming acidic oxides.

3Metalloids

• Metalloids along the diagonal boundary, with properties of both, such as boron, silicon, arsenic and antimony.

4s block

• Groups 1 and 2 plus helium.

5p block

• Groups 13 to 18.

6d block

• d block transition metals, Groups 3 to 12.

7f block

• f block lanthanides and actinides.

8Know without the table

• Group 1 alkali metals, Group 17 halogens, Group 18 noble gases and Groups 3–12 transition elements.

Section 3 of 62Configuration and position

Configuration and position

Period

• Period the highest occupied shell number.

Group

• Group gives the number of valence electrons.

Block

• Block shown by the last subshell filled, and the position along the p block shows how many p electrons.

Shorthand configuration

• Shorthand configuration a noble gas core in brackets, with 4s written before or after 3d.

Ions

• Ions cations lose electrons from the outer shell; for example Mg²⁺ is [Ne].

Reading a configuration

• 1s² 2s² 2p⁶ 3s² 3p⁵ is period 3, 7 valence electrons, p block: chlorine.

Section 4 of 62Atomic and ionic radius

Atomic and ionic radius

Atomic radius

• Atomic radius the typical distance from the nucleus to the outermost electrons.

Across a period

• Across a period it decreases because nuclear charge rises, extra electrons enter the same shell and shielding is roughly constant.

Down a group

• Down a group it increases because of extra shells and more shielding.

Noble gas to alkali metal

• Noble gas to alkali metal a sharp increase, because of one extra shell.

Cations

• Cations smaller than the atom because electrons are lost and repulsion falls.
  • (often a whole shell)

Anions

• Anions larger than the atom because extra electrons increase repulsion.

Ionic radius down a group

• Ionic radius down a group increases with extra shells and shielding.

Section 5 of 62Ionisation energy and electron affinity

1

First ionisation energy

First ionisation energy the energy to remove one mole of electrons from one mole of gaseous atoms (kJ mol⁻¹).

2

Across a period

• Across a period it increases because of greater nuclear charge, constant shielding and smaller radius.

3

Down a group

• Down a group it decreases because the outer electron is further away and more shielded.

4

Electron affinity

Electron affinity the energy released when one mole of electrons is gained by one mole of gaseous atoms.

5

First electron affinity

• First electron affinity usually exothermic for non-metals, for example Cl(g) + e⁻ → Cl⁻(g), −349 kJ mol⁻¹.

6

Positive first EA

• Group 2, 12 and 15 elements, because the electron enters a stable or half-filled subshell.

7

Second electron affinity

• Second electron affinity endothermic, for example O⁻ → O²⁻, +753 kJ mol⁻¹, as the electron is added to a negative ion.

8

Trends in EA

• Trends in EA more negative across a period and less negative down a group.

Section 6 of 62Electronegativity

Electronegativity

Definition

• Definition the ability of an atom to attract the bonding pair of electrons in a covalent bond.

Across a period

• Across a period it increases because nuclear charge rises, shielding stays the same and radius falls.

Down a group

• Down a group it decreases because of an extra shell, more shielding and a larger radius.

Not electron affinity

• Not electron affinity electronegativity describes chemical character in covalent bonds,
  • while electron affinity is a measurable energy for forming negative ions.

Electropositivity

• Electropositivity the tendency to form positive ions.

Section 7 of 62Group 1 and Group 17

Group 1 and Group 17

Group 1

• Group 1 alkali metals with outer configuration ns¹;
  • soft, shiny when cut, low density, low melting points that fall down the group.

With water

• With water metal + water → metal hydroxide + hydrogen, for example 2Na + 2H₂O → 2NaOH + H₂;
  • reactivity increases down the group.

Reason

• Reason the outer electron is further away and more shielded, so it is lost more easily.

Group 17

• Group 17 diatomic poisonous non-metals with ns²np⁵;
  • colour gets darker and melting and boiling points rise down the group.

Reactivity

• Reactivity halogens become less reactive down the group because incoming electrons are attracted less.

Displacement

• Displacement a more reactive halogen displaces a less reactive one from its halide solution, with chlorine > bromine > iodine.

Colours

• Colours chlorine pale green gas, bromine orange-brown liquid, iodine grey-black solid giving purple vapour.

Section 8 of 62Period 3 oxides

Trend

• Trend oxides change from basic through amphoteric to acidic across a period.

Na₂O and MgO

• Na₂O and MgO giant ionic, basic, giving alkaline solutions (O²⁻ + H₂O → 2OH⁻).

Al₂O₃

• Al₂O₃ ionic with covalent character, amphoteric, reacting with both acids and bases.

SiO₂

• SiO₂ giant covalent and acidic.

P₄O₁₀, SO₂ and SO₃

• P₄O₁₀, SO₂ and SO₃ simple molecular, acidic, giving oxoacids such as H₃PO₄ and H₂SO₄.

Reason

• Reason the electronegativity difference with oxygen falls across the period, so bonding becomes more covalent.

Environment

• SO₂ and NO₂ cause acid rain; CO₂ dissolving forms carbonic acid and acidifies oceans.

Section 9 of 62Oxidation numbers

1Oxidation

Oxidation gain of oxygen, loss of hydrogen or loss of electrons.

2Reduction

Reduction loss of oxygen, gain of hydrogen or gain of electrons.

3Oxidation number

• Oxidation number the charge an atom would have if bonding were completely ionic, written with the sign first.

4Rules

• Rules uncombined element 0; Group 1 +1; Group 2 +2; F −1; H +1 ; O −2 (−1 in peroxides, +2 in F₂O).
  • (−1 in metal hydrides)

5Sums

• Sums a neutral compound sums to 0 and a polyatomic ion sums to its charge.

6More electronegative atom

• More electronegative atom takes the negative number.

7Fractional values

• Fractional values can arise as averages, such as +2.5 for S in S₄O₆²⁻.

8Stock notation

• Roman numerals show the oxidation state of a variable metal, for example iron(III) oxide.

Section 10 of 62Transition elements (HL)

Properties

• Properties variable oxidation states, high melting points, magnetism from unpaired d electrons, catalytic behaviour, coloured compounds and complex ions.

Configurations

• [Ar] 3dⁿ 4s², with Cr [Ar] 3d⁵ 4s¹ and Cu [Ar] 3d¹⁰ 4s¹; 4s is lost first in ions.

Colour

• Colour ligands split the d orbitals,
  • and absorbing light of energy equal to the gap promotes an electron;
    • the colour seen is the complement of that absorbed.

Catalysts

• Catalysts heterogeneous and homogeneous (haemoglobin).
  • (Fe in the Haber process, Pt/Rh in catalytic converters)

Complex ion

• Complex ion a central metal ion surrounded by ligands that donate lone pairs.

Section 11 of 62More from the specification

More from the specification

Electrochemical redox quantification

• Electrochemical redox quantification determining standard Gibbs free energy from cell potentials provides an analytical means to assess chemical equilibria.

Periodic law and valence configuration

• Elements in the same group possess identical valence electron configurations, leading to predictable similarities in chemical reactivity.

Oxidation states and overall charge

• The sum of oxidation numbers of all atoms in a neutral compound must equal zero, or equal the overall charge in a polyatomic ion.

Group 1 trends

• Alkali metals like sodium readily lose their single valence electron to achieve a noble gas configuration with a +1 oxidation state.

Periodicity across Period 3

• Gradual transition from metallic bonding through giant covalent (silicon) to simple molecular structures (phosphorus, sulfur, chlorine).
  • (sodium, magnesium)

Section 12 of 62Position from electron configuration

Position from electron configuration

Period

• Period the highest occupied principal energy level (n).

Group

• Group found from the valence electrons.

Block

• Block shown by the last subshell filled.

Germanium

• [Ar] 4s² 4p² is period 4, p block, Group 14 with 4 valence electrons.

Transition elements

• Groups 3 to 12.

Periodic trends

• Periodic trends regular patterns across a period (left to right) and down a group (top to bottom).

Section 13 of 62Atomic radius

Atomic radius

Definition

• Definition the typical distance from the nucleus to the outermost electrons.

Measurement

• Measurement taken as half the mean distance between two adjacent bonded atoms, giving comparative values.

Across a period

• Across a period it decreases.
• The number of protons rises, so nuclear charge increases.
• Electrons enter the same shell and shielding stays about constant.
• Attraction to the outer electrons increases, pulling them closer.

Down a group

• Down a group it increases.
• An extra principal shell is added each time.
• Outer electrons are further away and more shielded.
• Attraction weakens even though nuclear charge increases.

Period boundary

• Period boundary the radius jumps sharply between a noble gas and the next alkali metal, which has one extra shell.

Section 14 of 62Ionic radius

	Cations	Anions
Compared with atom	Smaller: electrons are lost, often a whole shell, so there is less repulsion and remaining electrons are pulled in more	Larger: extra electrons increase repulsion, which spreads the electron cloud
Across a period	Decreases as nuclear charge rises with constant shielding	Increases sharply where anions begin, then decreases as nuclear charge rises
Down a group	Increases: more shells, more shielding, greater distance	Increases: more shells, more shielding, greater distance

Section 15 of 62First ionisation energy

1Definition

Definition the energy to remove one mole of electrons from one mole of gaseous atoms to give one mole of gaseous 1+ ions.

2Conditions

• 298 K and 100 kPa, in kJ mol⁻¹.

3Example

• Ca(g) → Ca⁺(g) + e⁻, ΔH = +590 kJ mol⁻¹.

Section 16 of 62Ionisation energy trends

Ionisation energy trends

Across a period

• Across a period it generally increases.
• Nuclear charge increases.
• Electrons are added to the same shell and shielding stays about constant.
• The outer electrons are held more strongly, so more energy is needed.

Down a group

• Down a group it decreases.
• The outer electron is in a higher shell, further from the nucleus and more shielded.
• Attraction is weaker, so less energy is needed.

Between periods

• Between periods a rapid fall occurs from the last element of one period to the first of the next,
• because the greater distance and shielding outweigh the greater nuclear charge.

Section 17 of 62More from the specification

More from the specification

Cell electromotive force

• Cell electromotive force calculating cathode reduction potential minus anode reduction potential provides a straightforward means to predict reaction spontaneity.

Periodic trends in oxide acid-base character

• Period 3 oxides transition from basic ionic oxides (Na₂O, MgO) to amphoteric aluminum oxide (Al₂O₃),
• and finally to acidic covalent oxides.
  • (SiO₂, P₄O₁₀, SO₃)

Electronegativity trends across the periodic table

• Electronegativity increases across periods and decreases down groups, governed by effective nuclear charge and atomic shielding.

Oxidation states of transition metal elements

• Transition metals exhibit variable positive oxidation states due to the close proximity of 3d and 4s electronic energy levels.
  • (e.g. Fe²⁺, Fe³⁺, Mn7+)

Period 3 oxide acidity

• Period 3 oxide acidity basic metal oxides like sodium oxide dissolve to yield alkaline solutions, while non-metal oxides form acidic solutions.

Electronegativity trend

• Pauling electronegativity values increase across a period as effective nuclear charge attracts bonding electron pairs.

Transition metal oxidation states

• Transition metal oxidation states d-block metals exhibit variable oxidation states due to the close energy levels of 3d and 4s orbitals.

Oxidation

• Gain of oxygen, loss of electrons or an increase in oxidation number.

Oxide

• A compound of oxygen with one other element.

Electronegativity

• The ability of an atom to attract the bonding pair of electrons in a covalent bond.

Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

Property

• A feature that describes what a substance is like and what it does.

Bonding

• The joining of atoms.

Transition metal

• Metal in the central block, often forming coloured compounds.

Electron affinity

• Enthalpy change when electrons are added to gaseous atoms or ions.

Section 18 of 62Electron affinity

1Definition

Definition the energy released when one mole of electrons is gained by one mole of gaseous atoms to form one mole of gaseous ions.

2Conditions

• 298 K and 100 kPa, in kJ mol⁻¹; it is the opposite process to ionisation.

3First EA

• First EA usually exothermic for non-metals, for example Cl(g) + e⁻ → Cl⁻(g), −349 kJ mol⁻¹.

4Positive first EA

• Groups 2, 12 and 15,
• because the electron must enter a stable or half-filled sublevel where repulsion is significant, for example N(g) + e⁻ → N⁻(g), +7 kJ mol⁻¹.

5Second EA

• Second EA endothermic, for example O⁻(g) + e⁻ → O²⁻(g), +753 kJ mol⁻¹,
  • because the electron must overcome repulsion from an ion that is already negative.

Section 19 of 62Electron affinity trends

Electron affinity trends

Across a period

• Across a period more negative, because nuclear charge rises and radius falls so added electrons are held more strongly.
  • (more exothermic)

Down a group

• Down a group less negative, because of more shells, a larger radius and more shielding.

Irregularities

• Groups 2 and 15 deviate because of sublevel configuration and extra electron repulsion.

Section 20 of 62Electronegativity

Electronegativity

Definition

• Definition the ability of an atom to attract the bonding pair of electrons towards itself in a covalent bond.

Origin

• Origin the attraction of the positive nucleus for electrons in the outer shell.

Across a period

• Across a period it increases because nuclear charge rises, there is no new shell and radius falls.

Down a group

• Down a group it decreases because an extra filled shell adds shielding and distance.

Versus electron affinity

• Versus electron affinity electronegativity is chemical character in covalent bonds;
  • electron affinity is a measurable thermodynamic quantity for negative ion formation.

Electropositivity

• Electropositivity the tendency to form positive ions, useful for metals.

Section 21 of 62The alkali metals

Members

• Members lithium, sodium, potassium, rubidium, caesium and francium, in the first column of the s block.

Configuration

• Configuration all have an outer ns¹ electron, for example Li 2s¹ and K 4s¹.

Physical properties

• Physical properties soft and easy to cut, shiny when freshly cut, low density, low melting points and good conductors through delocalised electrons.

Melting point trend

• Melting point trend it decreases down the group as atomic radius increases and metallic bonding weakens.

In air

• In air they react with oxygen and water vapour, so they are stored under oil.

Section 22 of 62Reaction with water

General

• Group 1 metal + water → metal hydroxide + hydrogen, giving alkaline solutions of high pH.

Lithium

• Lithium floats, reacts slowly with hydrogen bubbles; 2Li + 2H₂O → 2LiOH + H₂.

Sodium

• Sodium more vigorous; melts into a silvery ball that moves on the surface; 2Na + 2H₂O → 2NaOH + H₂.

Potassium

• Potassium violent; melts, dashes about, and the hydrogen ignites with a lilac flame; 2K + 2H₂O → 2KOH + H₂.

Product

• Product a clear colourless solution of the hydroxide.

Section 23 of 62Reactivity trend and reaction with halogens

Reactivity trend and reaction with halogens

Reactivity down Group 1

• Reactivity down Group 1 it increases because atoms are larger, the outer electron is further from the nucleus and more shielded,
  • so it is lost more easily.

Metallic character

Metallic character the ease with which an atom loses electrons to form positive ions.

With halogens

• With halogens all alkali metals react vigorously to give metal halide salts, for example 2Na(s) + Cl₂(g) → 2NaCl(s).

Vigour

• Vigour the reaction becomes more vigorous down Group 1 for the same reason.

Most vigorous pair

• Most vigorous pair the lowest Group 1 metal with the highest halogen, caesium with fluorine, since halogen reactivity decreases down Group 17.

Section 24 of 62More from the specification

More from the specification

Standard hydrogen electrode

• Standard hydrogen electrode coupling platinum black electrodes under standard conditions provides a universal means to reference half-cell redox voltages.

Transition Metal Definition

• A transition metal is defined as a d-block element that forms at least one stable ionic species with an incomplete d subshell of orbitals.

Variable Oxidation States

• Transition elements display multiple stable oxidation states because 4s and 3d electrons have similar ionization energies and can both participate in bonding.

Transition Metal Oxides

• Oxides formed by transition elements range from basic ionic lattices at lower oxidation numbers to covalent acidic oxides at higher oxidation numbers.

Complex Ion Charge Balance

• The overall ionic charge of a transition metal complex represents the algebraic sum of the central metal oxidation state and the charges of all attached ligands.

Transition Metal Definition

• A transition metal is defined as a d-block element that forms at least one stable ionic species with an incomplete d subshell of orbitals.

Variable Oxidation States

• Transition elements display multiple stable oxidation states because 4s and 3d electrons have similar ionization energies and can both participate in bonding.

Transition Metal Oxides

• Oxides formed by transition elements range from basic ionic lattices at lower oxidation numbers to covalent acidic oxides at higher oxidation numbers.

Complex Ion Charge Balance

• The overall ionic charge of a transition metal complex represents the algebraic sum of the central metal oxidation state and the charges of all attached ligands.

Ionic radius

• Distance from the nucleus to the outermost electrons of an ion.

Oxidation

• Gain of oxygen, loss of electrons or an increase in oxidation number.

Element

• A substance made of only one type of atom.

Oxide

• A compound of oxygen with one other element.

Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

Transition metal

• Metal in the central block, often forming coloured compounds.

Transition elements

• The central block of metals between Group II and Group III.

Oxidation number

• The number of electrons an atom has lost or gained compared with the uncombined element.

Section 25 of 62Group 17: the halogens

Group 17: the halogens

Members

• Members fluorine, chlorine, bromine, iodine and astatine, in the p block.

Nature

• Nature poisonous non-metals forming diatomic molecules.

Electrons

• Electrons seven outer electrons (ns²np⁵); they form halide ions by gaining one electron.

Colour trend

• Colour trend colours get darker down the group.

Boiling points

• Boiling points melting and boiling points increase down the group because of larger masses and stronger van der Waals forces.

Section 26 of 62Appearance of the halogens

Halogen	State and colour at room temperature	Colour in water
Fluorine	Yellow gas	Not usually tested (too reactive)
Chlorine	Pale green gas	Green-blue
Bromine	Orange-brown liquid, evaporating to a brown gas	Orange
Iodine	Grey-black solid, subliming to a purple vapour	Dark brown

Section 27 of 62Reactivity of the halogens

Trend

• Trend reactivity decreases down the group.

Atomic radius

• Atomic radius increases down the group.

Shielding

• Shielding increases with the extra shells.

Electron affinity

• Electron affinity becomes less negative.

Effect

• Effect the attraction for an incoming electron is weaker, so the halogens are less effective at gaining electrons.

Meaning

• Meaning this reflects decreasing non-metallic character.

Section 28 of 62Displacement reactions

Rule

• Rule a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide.

Order

• Order chlorine > bromine > iodine.

Chlorine and bromide

• Chlorine and bromide pale green chlorine added to colourless KBr turns orange as bromine forms: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq).

Bromine and iodide

• Bromine and iodide orange bromine added to colourless NaI turns brown as iodine forms: Br₂(l) + 2NaI(aq) → 2NaBr(aq) + I₂(aq).

No reaction

• No reaction iodine cannot displace chlorine or bromine, and fluorine displaces chlorine.

Section 29 of 62Period 3 oxides: structure and bonding

Period 3 oxides: structure and bonding

Trend

• Trend oxides change from basic through amphoteric to acidic across the period.

Amphoteric

Amphoteric able to act as both an acid and a base; aluminium oxide reacts with HCl and NaOH.

Oxide	Structure	Character
Na2O, MgO	Giant ionic, high melting point	Basic
Al2O3	Giant ionic with some covalent character (small, highly charged Al3+), very high melting point	Amphoteric
SiO2	Giant covalent, very high melting point	Acidic
P4O10, SO2, SO3	Simple molecular, low melting point	Acidic

Section 30 of 62Electronegativity and oxide bonding

Oxygen

• Oxygen electronegativity 3.4.

Differences with oxygen

• Na (0.9) 2.5, Mg (1.2) 2.2, Al (1.5) 1.9, Si (1.8) 1.6, P (2.1) 1.3, S (2.5) 0.9.

Large difference

• Large difference electrons transfer from metal to oxygen, giving purely ionic oxides.

Across the period

• Across the period the difference falls, so bonding becomes more covalent.

Section 31 of 62Oxides with water

Basic oxides

• Basic oxides form hydroxides and alkaline solutions: Na₂O(s) + H₂O(l) → 2NaOH(aq), pH about 14.

Magnesium oxide

• MgO(s) + H₂O(l) → Mg(OH)₂(aq), weakly alkaline, pH about 10.

Oxide ion

• O²⁻ + H₂O → 2OH⁻.

Acidic oxides

• Acidic oxides covalent oxides of P and S react with water to give oxoacids, for example P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq), pH about 2.

Ionic oxides

• Ionic oxides give hydroxide ions and alkaline solutions.

Covalent oxides

• Covalent oxides give oxoacids and acidic solutions.

Overall

• Overall the pH of the solution falls from sodium oxide to sulfur trioxide.

Section 32 of 62More from the specification

1

Galvanic potential measurement

• Galvanic potential measurement recording high-impedance voltmeter readings across half-cells provides an accurate experimental means to establish standard reduction potentials.

2

Periodic trends in first ionisation energy

• First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form gaseous 1+ ions,
• increasing across a period as nuclear charge in the nucleus increases while shielding remains similar.

3

Transition elements and variable oxidation states

• Transition elements are d-block metals that form one or more stable ions with incomplete d-subshell orbitals,
• allowing them to exhibit variable oxidation states and form coloured complex ions.

4

Electron configuration of d-block elements

• The electron configuration of transition elements involves filling 3d orbitals after the 4s subshell, with anomalous stable configurations observed in chromium ([Ar] 3d5 4s1) and copper ([Ar] 3d10 4s1).

5

Nuclear attraction and orbital shielding

• Across a period of elements, successive valence electrons occupy the same principal quantum shell, experiencing greater effective nuclear pull from the nucleus,
• which decreases atomic radius.

6

Electron configuration

• The arrangement of electrons in the shells, subshells and orbitals of an atom.

7

Electronic configuration

• The number of electrons in each shell, written with commas.

8

Oxidation

• Gain of oxygen, loss of electrons or an increase in oxidation number.

9

Element

• A substance made of only one type of atom.

10

Nucleus

• Central part of an atom with protons and neutrons.

11

Ionisation

• The removal of an electron from an atom or molecule.

12

Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

13

Ionisation energy

• The energy required to remove one mole of electrons from one mole of gaseous atoms.

14

Transition metal

• Metal in the central block, often forming coloured compounds.

Section 33 of 62Predicting oxide properties

Predicting oxide properties

Metals

• Na and Mg form ionic, basic oxides.

Aluminium

• Aluminium forms an ionic oxide with covalent character that is amphoteric.

Non-metals

• Si, P and S form covalent, acidic oxides.

Down a group

• Down a group oxides become more ionic as electronegativity falls.

Across a period

• Across a period oxides become less ionic as electronegativity rises.

Use

• Use these trends deduce the bonding and acid–base behaviour of unfamiliar oxides.

Section 34 of 62Acid rain and ocean acidification

Acid rain

• Acid rain non-metal oxides dissolve in atmospheric water to form strong acids that lower the pH of rain and harm ecosystems.

Sulfur dioxide

• SO₂(g) + H₂O(l) → H₂SO₃(aq).

Nitrogen dioxide

• 2NO₂(aq) + H₂O(l) → HNO₃(aq) + HNO₂(aq).

Ocean acidification

• CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq), and the weak acid dissociates to H⁺ and HCO₃⁻.

Effect

• Effect the extra H⁺ lowers ocean pH and harms marine life.

Section 35 of 62Oxidation and reduction

	Oxidation	Reduction
Oxygen	Gain of oxygen: 2Mg + O2 → 2MgO	Loss of oxygen: 2CuO + C → 2Cu + CO2
Hydrogen	Loss of hydrogen: CH3OH → CH2O	Gain of hydrogen: C2H4 + H2 → C2H6
Electrons	Loss: Al → Al3+ + 3e−	Gain: F2 + 2e− → 2F−

Section 36 of 62Oxidation numbers

Oxidation number

Oxidation number the charge an atom would have if all bonding were completely ionic.

Uses

• Uses to tell whether oxidation or reduction occurred, to identify what was oxidised or reduced, and to balance redox equations.

Sign

• Sign written before the number, for example +3 and −2.

Interchangeable

• Interchangeable oxidation number and oxidation state mean the same.

Change

• Change an increase in oxidation number means oxidation and a decrease means reduction.

Section 37 of 62Rules for oxidation numbers

Rules for oxidation numbers

Uncombined element

• 0.

Monatomic ion

• Monatomic ion equal to its charge.

Group 1 and Group 2

• +1 and +2.

Fluorine

• −1.

Hydrogen

• +1, or −1 in metal hydrides.

Oxygen

• −2, except −1 in peroxides and +2 in F₂O.

Neutral compound

• Neutral compound the numbers sum to zero.

Polyatomic ion

• Polyatomic ion the numbers sum to the ion charge.

More electronegative atom

• More electronegative atom takes the negative value.

Section 38 of 62Fractional values and Stock notation

Fractional numbers

• Fractional numbers an oxidation number can be an average, such as +2.5 for sulfur in S₄O₆²⁻, where the sulfur atoms are in two environments.

Single atoms

• Single atoms always have whole-number oxidation states.

Stock notation

• Roman numerals show the oxidation number of a variable metal, for example iron(II) oxide FeO and iron(III) oxide Fe₂O₃.

Further names

• KMnO₄ is potassium manganate(VII), Na₂CrO₄ sodium chromate(VI) and Na₂Cr₂O₇ sodium dichromate(VI).

Non-metals

• Non-metals not usually named with Roman numerals, for example sulfur dioxide.

In exams

• In exams only the transition metal needs the Stock notation.

Section 39 of 62More from the specification

1Aqueous electrolyte solution

• Aqueous electrolyte solution dissolving soluble salts in water produces mobile ions, facilitating ionic current transport through the conducting solution.

2Electronic Configurations and Subshells

• Electrons in atoms occupy principal energy levels divided into s, p, d,
• and f subshells and orbitals, filling in order of increasing energy according to the Aufbau principle.
  • (1s < 2s < 2p < 3s < 3p < 4s < 3d)

3Hund's Rule and Pauli Exclusion Principle

• Orbitals in the same subshell are singly occupied with parallel electron spins before pairing occurs (Hund's rule), with each orbital accommodating a maximum of two electrons with opposite spins.

4Atomic Radius and Shielding Effects

• Atomic radius represents the distance from the nucleus to the outer valence electrons;
  • greater shielding by complete inner shells decreases effective nuclear attraction on outer electrons.

5Periodic Trends in Chemical Reactivity

• Chemical properties of elements are dictated by outer valence shell configurations;
  • metals with low ionisation energies lose electrons readily,
• while non-metals with high electron affinities gain electrons.

6Electron configuration

• The arrangement of electrons in the shells, subshells and orbitals of an atom.

7Atomic radius

• Half the distance between the nuclei of two bonded atoms of an element.

8Electron affinity

• Enthalpy change when electrons are added to gaseous atoms or ions.

9Electronic configuration

• The number of electrons in each shell, written with commas.

10Ionic radius

• Distance from the nucleus to the outermost electrons of an ion.

11Nucleus

• Central part of an atom with protons and neutrons.

12Ionisation

• The removal of an electron from an atom or molecule.

13Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

14Property

• A feature that describes what a substance is like and what it does.

15Shell

• An energy level around the nucleus that holds a limited number of electrons.

16Shielding

• The reduction in attraction to the nucleus caused by inner electron shells.

17Ionisation energy

• The energy required to remove one mole of electrons from one mole of gaseous atoms.

18Outer electrons

• The electrons in the outermost shell, which decide chemical behaviour.

Section 40 of 62First ionisation energy across Period 3

General trend

• IE₁ increases because nuclear charge rises, electrons enter the same shell and the radius decreases.

Sublevel evidence

• Sublevel evidence two dips support the existence of energy sublevels.

Mg to Al

• Mg to Al the Al electron is in the higher-energy 3p sublevel and shielded by the 3s electrons, so it is removed more easily.

P to S

• P to S the fourth 3p electron in S shares an orbital, and the repulsion makes it easier to remove.

Period 2

• Period 2 the same pattern gives Be (900 kJ mol⁻¹) to B (801) and N (1402) to O (1314).

Down a group

• IE decreases as shells and shielding increase.

Section 41 of 62Atomic radius and shielding

1Atomic radius

• Atomic radius decreases across a period and increases down a group.

2Noble gas to alkali metal

• Noble gas to alkali metal the radius jumps because the alkali metal has one extra principal shell and more shielding.

Section 42 of 62Explaining reactivity

Alkali metals

• Alkali metals reactivity increases down the group because the ns¹ electron is further away.

Halogens

• Halogens reactivity decreases because incoming electrons are attracted less.

Halogen displacement

• Halogen displacement fluorine displaces chlorine and bromine displaces iodine, but iodine cannot displace chlorine.

Three factors

• Three factors nuclear charge, distance of the outer electrons from the nucleus and shielding explain almost every trend.

Section 43 of 62Oxides and oxidation numbers recap

Oxides and oxidation numbers recap

Oxides

• Oxides change from basic to amphoteric to acidic across a period.

Sodium oxide

• Sodium oxide the large electronegativity difference with oxygen (3.4) gives an ionic oxide.

Phosphorus and sulfur oxides

• Phosphorus and sulfur oxides covalent, forming acids.

Oxidation numbers

• Oxidation numbers uncombined element 0, simple ion equals its charge, compounds sum to zero, the more electronegative element is negative.

Stock notation

• Roman numerals show variable oxidation states.

Section 44 of 62Transition elements: key properties

1Cause

• Cause partly filled d sublevels.

2Variable oxidation states

• Variable oxidation states more than one stable ion.

3High melting points

• High melting points many delocalised d and s electrons.

4Magnetism

• Magnetism from unpaired d electrons.

5Catalysis

• Catalysis transition metals and their compounds catalyse reactions.

6Coloured compounds

• Coloured compounds ligands split the d orbitals so electrons absorb part of visible light.

7Complex ions

• Complex ions a central metal ion bonded to ligands.

8Configurations

• Configurations written [Ar] 3dⁿ 4s², with exceptions Cr and Cu;
  • +3 is more stable in early metals and +2 in later ones.

Section 45 of 62Catalyst types and haem

Catalyst types and haem

Heterogeneous catalyst

• Heterogeneous catalyst in a different phase from the reactants, with reaction at active sites on its surface, for example iron in the Haber process.

Haem

• Haem an iron(II) ion at the centre of a porphyrin ring, bonded to four nitrogen atoms in a square planar arrangement.

Section 46 of 62Structuring a trend explanation

1

State the trend using the data given.

2

Name the factor that changes: nuclear charge, distance or shielding.

3

Link it to the force of attraction on the outer electron.

4

Conclude with the effect on the energy needed or the radius.

1Ionisation energy comparisons

• Ionisation energy comparisons compare the sublevel of the electron removed, the shielding it experiences and whether it is paired.

2Reactivity comparisons

• Reactivity comparisons compare the distance of the outer electron from the nucleus.
  • (or the incoming electron)

Section 47 of 62More from the specification

1

Electrolytic decomposition analysis

• Electrolytic decomposition analysis driving non-spontaneous redox reactions with external power supplies provides a reliable means to isolate pure elements.

2

First electron affinity definition

• First electron affinity is the enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of singly charged negative ions (X(g) + e- → X-(g)).

3

Ionic lattice enthalpies

• Lattice enthalpy measures the electrostatic attraction within an ionic crystal lattice, becoming more exothermic with smaller ionic radii and higher ionic charges.

4

Hydration enthalpy of aqueous ions

• Exothermic hydration enthalpy occurs when gaseous ions dissolve in water, stabilized by ion-dipole interactions with polar water molecules.

5

First electron affinity trend

• First electron affinity trend energy released when one mole of gaseous atoms gains an electron to form singly charged negative anions.

6

Reactivity of alkali metals with water

• Reactivity of alkali metals with water group 1 metals react vigorously with water, evolving hydrogen gas and forming alkaline metal hydroxide solutions.

7

Ionisation of outermost electron

• Ionisation of outermost electron first ionisation energy measures the minimum energy required to remove the outermost valence electron from a gaseous atom.

8

Bonding

• The joining of atoms.

9

Electron affinity

• Enthalpy change when electrons are added to gaseous atoms or ions.

Section 48 of 62Variable oxidation states

Variable states

• Variable states unlike most metals, transition elements form more than one positive ion.

Roman numerals

• Roman numerals iron forms Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)), but sodium forms only Na⁺.

Cause

• Cause the 3d and 4s subshells have similar energies, so several oxidation states are accessible.

Section 49 of 62Electron configurations of the first series

1Aufbau

• 4s fills before 3d.

2Abbreviated

• Ti [Ar] 3d² 4s², V 3d³ 4s², Mn 3d⁵ 4s², Fe 3d⁶ 4s², Co 3d⁷ 4s², Ni 3d⁸ 4s².

3Exceptions

• Cr is [Ar] 3d⁵ 4s¹ and Cu is [Ar] 3d¹⁰ 4s¹, because half-filled and filled d sublevels are more stable.

Section 50 of 62Magnetism

Source

• Source magnetism comes from unpaired d electrons; each spinning electron is a tiny magnetic dipole.

Paired electrons

• Paired electrons spin oppositely and cancel, so substances with only paired electrons are non-magnetic.

Iron, cobalt, nickel

• Iron, cobalt, nickel their unpaired d electrons align in an external field, giving strong magnetism.

Steel

• Steel magnetic because it contains iron.

Domains

• Domains regions of aligned dipoles that keep the magnetism after heating and cooling in a field.

Loss

• Loss heating or striking a permanent magnet disrupts the alignment.

Not assessed

• Not assessed different types of magnetism.
  • (diamagnetism, paramagnetism, ferromagnetism)

Section 51 of 62Transition metals as catalysts

1Why

• Why they have more than one stable oxidation state, so they gain and lose electrons easily and can be oxidised and reduced repeatedly.

2Heterogeneous

• Heterogeneous different phase from the reactants, with reaction at active sites on the surface.

3Haber process

• Fe catalyses N₂ + 3H₂ ⇌ 2NH₃.

4Hydrogenation

• Ni catalyses CH₂=CH₂ + H₂ → CH₃CH₃ and hardens vegetable oils.

5Decomposition

• MnO₂ catalyses 2H₂O₂ → 2H₂O + O₂.

6Homogeneous

• Homogeneous same phase as the reactants;
  • many enzyme-catalysed reactions use transition metals, such as the iron in haemoglobin.

Section 52 of 62Catalytic converters

Metals

• Metals platinum and rhodium.

Design

• Design finely divided on a ceramic honeycomb to maximise surface area and reduce the amount of precious metal.

Pollutants

• Pollutants carbon monoxide, nitrogen dioxide and unburnt hydrocarbons.

Reactions

• 2NO(g) + 2CO(g) → N₂(g) + 2CO₂(g) and CH₃CH₂CH₃(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(g).

Section 53 of 62Haemoglobin, colour and complex ions

Haem

• Haem an iron(II) ion at the centre of a porphyrin ring, bonded to four nitrogens in a square planar arrangement with coordination number four.

Oxygen transport

• Oxygen transport haemoglobin has four porphyrin rings, so it can carry four O₂ molecules.

Complex ion

Complex ion a central metal ion surrounded by ligands that donate lone pairs.

Chromium(III) complexes

• [Cr(NH₃)₆]³⁺, [Cr(OH)₆]³⁻ and [Cr(H₂O)₆]³⁺.

Colour

• [Cr(OH)₆]³⁻ is dark green and [Cr(NH₃)₆]³⁺ is purple, so different ligands give different colours with the same oxidation state.

Section 54 of 62More from the specification

1

Hydrolysis of non-metal oxides in solution

• Hydrolysis of non-metal oxides in solution dissolving acidic covalent oxides in aqueous solution provides a chemical means to produce acidic solutions.

2

First Ionisation Energy Across a Period

• First ionisation energy increases across a period because nuclear charge increases and atomic radius decreases while electron shielding remains approximately constant.

3

Ionisation Energy Trends Down a Group

• Down any group, the energy required to remove an outer electron decreases because additional electron shells increase atomic radius and electron shielding.

4

Nuclear Charge and Valence Electrons

• Greater positive charge in the nucleus draws outer shell electrons closer, increasing the Coulombic attraction felt by valence electrons.

5

Subshell Shielding and Discontinuities

• Small dips in ionisation energy across a period reveal extra stability in filled subshells and the onset of electron spin-pair repulsion in p orbitals.

6

First Ionisation Energy Across a Period

• First ionisation energy increases across a period because nuclear charge increases and atomic radius decreases while electron shielding remains approximately constant.

7

Ionisation Energy Trends Down a Group

• Down any group, the energy required to remove an outer electron decreases because additional electron shells increase atomic radius and electron shielding.

8

Nuclear Charge and Valence Electrons

• Greater positive charge in the nucleus draws outer shell electrons closer, increasing the Coulombic attraction felt by valence electrons.

9

Subshell Shielding and Discontinuities

• Small dips in ionisation energy across a period reveal extra stability in filled subshells and the onset of electron spin-pair repulsion in p orbitals.

10

Electron configuration

• The arrangement of electrons in the shells, subshells and orbitals of an atom.

11

Atomic radius

• Half the distance between the nuclei of two bonded atoms of an element.

12

Electron affinity

• Enthalpy change when electrons are added to gaseous atoms or ions.

13

Electronic configuration

• The number of electrons in each shell, written with commas.

14

Ionic radius

• Distance from the nucleus to the outermost electrons of an ion.

15

Electronegativity

• The ability of an atom to attract the bonding pair of electrons in a covalent bond.

16

Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

17

Property

• A feature that describes what a substance is like and what it does.

18

Trend

• General pattern down a group or across a period.

19

Bonding

• The joining of atoms.

20

Shielding

• The reduction in attraction to the nucleus caused by inner electron shells.

21

Ionisation energy

• The energy required to remove one mole of electrons from one mole of gaseous atoms.

22

Outer electrons

• The electrons in the outermost shell, which decide chemical behaviour.

23

Period

• A row of the Periodic Table, equal to the number of occupied shells.

24

Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

25

Nucleus

• Central part of an atom with protons and neutrons.

26

Ionisation

• The removal of an electron from an atom or molecule.

27

Oxide

• A compound of oxygen with one other element.

Section 55 of 62Electron configurations of ions

1Rule

• Rule transition metals lose 4s electrons before 3d.

2Reason

• Reason once filled, the 4s sublevel lies above 3d because of electron repulsion, so 4s is outermost.

3Consequence

• +2 is a common oxidation state, from losing two 4s electrons.

4Cu

• 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹.

5Order of writing

• 3d may be written before or after 4s, though 3d first is conventional.

Electron configurations of ions

Write the configuration of the neutral atom.

Remove electrons starting from 4s, then 3d.

Check that electrons = atomic number − charge.

Section 56 of 62Common oxidation states

Element	Common states
Ti	+3, +4
V	+3, +5
Cr	+3, +6
Mn	+2, +4, +7
Fe	+2, +3
Co	+2, +3
Ni	+2
Cu	+1, +2

Section 57 of 62Explaining variable states

1Ionisation energies

• Ionisation energies the first few for titanium and vanadium are small and close together, then successive values increase.

2Consequence

• Consequence the small energy difference lets the oxidation state change easily.

3Common states

• +2 and +3 occur in all transition elements;
  • +3 is more stable in early metals up to chromium and +2 in later ones.

4High oxidation states

• High oxidation states ions of +3 and above are polarising with some covalent character.

5Maximum state

• Maximum state equals the total number of 4s and 3d electrons, reached at Mn (+7), as in the strong oxidising agent MnO₄⁻.

Section 58 of 62Crystal field theory

Free ion

• Free ion the five d orbitals have the same energy, so they are degenerate.

Complex ion

• Complex ion ligand lone pairs repel the d electrons, splitting the d orbitals into two non-degenerate sets.

Model

• Model crystal field theory treats ligands as point charges.

Shapes

• Shapes the shapes of the d orbitals need not be learnt.

Section 59 of 62Why complexes are coloured

Promotion

• Promotion an electron in the lower set absorbs light with energy equal to the gap ΔE and moves to the higher set.

Colour seen

• Colour seen the complement of the colour absorbed, as complementary colours are opposite on the colour wheel.
  • (data booklet section 15)

Absorbed

• Absorbed the complex absorbs a part of the visible spectrum, and the rest is transmitted.

Nickel(II) sulfate

• Nickel(II) sulfate absorbs red light and appears green.
  • (about 647–700 nm)

Larger splitting

• Larger splitting needs more energy, so light of shorter wavelength and higher frequency is absorbed.

Relationships

• Relationships c = fλ with c = 3.00 × 10⁸ m s⁻¹, and E = hf with h = 6.63 × 10⁻³⁴ J s.

Section 60 of 62More from the specification

More from the specification

Faradaic quantitative electrolysis

• Faradaic quantitative electrolysis measuring electric current passed over elapsed duration provides a direct means to calculate deposited product moles.

Periodic Trends in Atomic Radius

• Atomic radius decreases across a period from left to right because increasing positive nuclear charge in the nucleus pulls valence electrons closer,
• while atomic radius increases down a group due to additional principal electron shells.

Electronegativity Across Periods and Down Groups

• Electronegativity increases across a period as greater nuclear charge exerts a stronger pull on shared electron pairs,
• but decreases down a group due to increased shielding and distance from the nucleus.

Acid-Base Character of Period 3 Oxides

• Metallic oxides on the left of Period 3 form alkaline solutions containing hydroxide ions,
  • (such as Na₂O and MgO)
• whereas non-metallic oxides on the right react with water to form acidic solutions.
  • (such as SO₂ and P₄O₁₀)

Amphoteric Oxides and Periodic Classification

• Aluminum oxide (Al₂O₃) exhibits amphoteric behavior by reacting with both acids and bases, marking the structural and chemical transition between basic metal oxides and acidic non-metal oxides across the periodic table.

Trend

• General pattern down a group or across a period.

Shielding

• The reduction in attraction to the nucleus caused by inner electron shells.

Atomic radius

• Half the distance between the nuclei of two bonded atoms of an element.

Electron affinity

• Enthalpy change when electrons are added to gaseous atoms or ions.

Ionic radius

• Distance from the nucleus to the outermost electrons of an ion.

Period

• A row of the Periodic Table, equal to the number of occupied shells.

Nucleus

• Central part of an atom with protons and neutrons.

Oxide

• A compound of oxygen with one other element.

Electronegativity

• The ability of an atom to attract the bonding pair of electrons in a covalent bond.

Property

• A feature that describes what a substance is like and what it does.

Section 61 of 62Key terms

Term	Meaning
Periodicity	Repeating trends in properties across periods and down groups.
Shielding	Repulsion of outer electrons by inner electrons.
Electron affinity	Energy released when a mole of gaseous atoms gains electrons.
Amphoteric	Able to react as both an acid and a base.
Oxidation number	The charge an atom would have if bonding were completely ionic.
Displacement reaction	A more reactive element replaces a less reactive one in a compound.
Complex ion	A central metal ion bonded to ligands.

Term	Meaning
Atomic radius	Typical distance from the nucleus to the outermost electrons.
Ionic radius	Distance from the nucleus to the outermost electrons of an ion.
Nuclear charge	Positive charge of the nucleus, equal to the number of protons.
First ionisation energy	Energy to remove one mole of electrons from one mole of gaseous atoms.
Shielding	Repulsion of outer electrons by inner electrons.

Term	Meaning
Electron affinity	Energy released when a mole of gaseous atoms gains a mole of electrons.
Electronegativity	Ability of an atom to attract a bonding pair of electrons.
Electropositivity	Tendency to form positive ions.
Alkali metals	Group 1 metals, which form alkaline solutions with water.
Metallic character	Ease of losing electrons to form positive ions.

Term	Meaning
Halogens	Group 17 non-metals that form diatomic molecules.
Displacement reaction	A more reactive element replaces a less reactive one from its compound.
Amphoteric oxide	An oxide that reacts with both acids and bases.
Basic oxide	An oxide that forms hydroxide ions in water.
Acidic oxide	An oxide that forms an oxoacid in water.
Sublimation	Change from solid directly to gas.

Term	Meaning
Oxidation number	The charge an atom would have if bonding were completely ionic.
Stock notation	Roman numerals giving the oxidation state of a metal in a name.
Acid rain	Rain made acidic by dissolved non-metal oxides.
Ocean acidification	Fall in ocean pH as CO2 dissolves to form carbonic acid.
Oxidation	Loss of electrons, loss of hydrogen or gain of oxygen.
Reduction	Gain of electrons, gain of hydrogen or loss of oxygen.

Term	Meaning
Sublevel evidence	Dips in ionisation energy show that sublevels exist.
Shielding	Repulsion of outer electrons by inner electrons.
Heterogeneous catalyst	A catalyst in a different phase from the reactants.
Haem	An iron(II)-containing porphyrin unit in haemoglobin.
Complex ion	A metal ion surrounded by ligands.

Term	Meaning
Magnetic domain	A region of aligned magnetic dipoles.
Heterogeneous catalyst	A catalyst in a different phase from the reactants.
Homogeneous catalyst	A catalyst in the same phase as the reactants.
Porphyrin	A large heterocyclic ring that binds iron(II) in haem.
Ligand	A species that donates a lone pair to a metal ion.
Complex ion	A central metal ion bonded to ligands.

Term	Meaning
Degenerate	Having the same energy.
Crystal field theory	A model of ligands as point charges that split the d orbitals.
Complementary colour	The colour opposite the absorbed colour on the colour wheel.
ΔE	The energy gap between the split d orbitals.
Polarising ion	A small highly charged ion that distorts electron clouds.

Section 62 of 62Exam technique

1

Explaining a trend only with nuclear charge — distance and shielding must be considered.

2

Confusing electronegativity with electron affinity.

3

Saying oxidation numbers are always whole numbers — averages can be fractional.

4

Saying Group 17 reactivity increases down the group — it decreases.

5

Giving the group of a halogen as 7 without the IUPAC number 17.

1

Saying atomic radius increases across a period because there are more electrons.

2

Forgetting that anions are larger than their atoms and cations smaller.

3

Explaining the fall in ionisation energy down a group by decreasing nuclear charge — nuclear charge increases but is outweighed by distance and shielding.

4

Ignoring the sharp jump in radius between a noble gas and the next alkali metal.

1

Confusing electronegativity with electron affinity.

2

Saying the first electron affinity is always exothermic — Groups 2, 12 and 15 can be endothermic.

3

Saying the second electron affinity is exothermic — it is endothermic due to repulsion.

4

Saying Group 1 metals become less reactive down the group.

1

Saying halogen reactivity increases down the group — it decreases.

2

Writing that a less reactive halogen displaces a more reactive one.

3

Calling aluminium oxide acidic or basic — it is amphoteric.

4

Saying SiO₂ is a simple molecule — it is giant covalent.

1

Giving an oxidation number without the sign.

2

Assuming oxygen is always −2 — it is −1 in peroxides and +2 in F₂O.

3

Saying acid rain is caused by metal oxides — non-metal oxides dissolve in rain water.

4

Treating a fractional oxidation number as a half electron — it is an average over different atoms.

1

Explaining the Mg to Al dip by shielding only — the electron is also in a higher-energy sublevel.

2

Explaining the P to S dip by distance — it is due to repulsion between paired electrons.

3

Saying the nuclear charge decreases down a group — it increases but is outweighed by distance and shielding.

4

Giving only the trend without the reason.

1

Saying all metals are magnetic — only those with unpaired d electrons, such as iron, cobalt and nickel.

2

Thinking a catalyst is used up — transition metals are regenerated by changing oxidation state.

3

Saying the iron in haem has coordination number six — it is bonded to four nitrogens in the porphyrin.

4

Assuming colour depends only on the metal ion — the ligand also matters.

1

Removing 3d electrons before 4s when forming transition metal ions.

2

Saying the colour seen is the colour absorbed — it is the complement.

3

Saying d orbitals split in energy in the free ion — splitting occurs when ligands bond.

4

Forgetting that a larger splitting means absorption of shorter wavelength light.