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

Section 1 of 45Position and electron configuration

Position and electron configuration

Arrangement

• Arrangement elements are ordered by increasing atomic number (1 to 118), in periods (rows) and groups (columns).

Period

• Period the period number equals the number of occupied shells.
  • (the highest principal quantum number)

Group

• Group for s- and p-block elements the group number (1, 2, 13–18) indicates the valence electrons, adding 10 for groups 13–18 in IUPAC numbering.

Block

• Block the last sublevel filled gives the block: s, p (groups 13–18), d (groups 3–12) and f (lanthanoids and actinoids).
  • (groups 1–2 plus He)

Types

• Types metals are on the left and centre, non-metals on the right and metalloids lie on the diagonal boundary.
  • (B, Si, Ge, As, Sb, Te)

Special cases

• Special cases hydrogen has its own position and helium sits in group 18 because it is unreactive.

Naming groups

• Naming groups group 1 alkali metals, group 17 halogens, group 18 noble gases and groups 3–12 transition elements.

Reading a configuration

• [Ne] 3s² 3p³ is period 3, group 15, p-block, so it is phosphorus.

Section 2 of 45Periodic trends and their explanation

Property	Across a period	Down a group
Atomic radius	Decreases	Increases
Ionic radius	Cations decrease; anions are larger and decrease	Increases
First ionisation energy	Generally increases	Decreases
Electronegativity	Increases	Decreases
Electron affinity	Generally more exothermic	Generally less exothermic
Metallic character	Decreases	Increases

1

Identify the change in nuclear charge.

2

Identify the change in number of shells and distance.

3

Consider shielding by inner electrons.

4

Decide the net attraction on the valence electrons.

5

Relate it to radius, ionisation energy or reactivity.

Periodic trends and their explanation

Across a period

• Across a period nuclear charge rises while shielding stays roughly constant, so valence electrons are pulled closer.

Down a group

• Down a group extra shells increase distance and shielding, which outweigh the extra nuclear charge.

Ions

• Ions cations are smaller than their atoms because a shell is lost, and anions are larger because of extra electron repulsion.

Section 3 of 45Electron affinity and electronegativity

1Electron affinity

Electron affinity the energy change when one mole of gaseous atoms gains one mole of electrons to form gaseous 1− ions, in kJ mol⁻¹.

2First value

• First value usually exothermic, e.g. F(g) + e⁻ → F⁻(g), but endothermic for group 2 atoms and N.

3Second value

• Second value endothermic, e.g. O⁻(g) + e⁻ → O²⁻(g), because the electron is repelled by the negative ion.

4Electronegativity

Electronegativity the ability of an atom to attract the shared pair of electrons in a covalent bond; it has no units and is not directly measured.

5Comparison

• Comparison electron affinity is a measurable enthalpy change for forming anions, and both increase across a period and decrease down a group.

Section 4 of 45Group 1 metals with water and halogens

• Equation
  • 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g), giving an alkaline solution.

Metal	Observations with water
Lithium	Floats and fizzes steadily, forming colourless LiOH solution
Sodium	Reacts faster, melts into a ball that moves on the surface
Potassium	Reacts violently, ignites the hydrogen with a lilac flame

Group 1 metals with water and halogens

Reactivity trend

• Reactivity trend increases down the group because the single outer electron is further from the nucleus, more shielded and more easily lost.
  • (lower ionisation energy)

Physical properties

• Physical properties soft, shiny when cut, good conductors with low melting points that fall down the group and low densities.

Storage

• Storage kept under oil because they react with air and moisture.

With halogens

• With halogens form ionic halides, e.g. 2Na(s) + Cl₂(g) → 2NaCl(s), most vigorously for metals low in group 1 and halogens high in group 17,
  • such as Cs with F₂.

Section 5 of 45Group 17 and halide ions

Group 17 and halide ions

Structure

• Structure halogens are diatomic with seven outer electrons and form halide ions by gaining one electron.

States and colours

• F₂ yellow gas, Cl₂ pale green gas, Br₂ red-brown liquid and I₂ grey-black solid, darkening down the group.

Solutions in water

• Solutions in water chlorine green-blue, bromine orange and iodine dark brown.

Melting and boiling points

• Melting and boiling points rise down the group because larger molecules have more electrons and stronger London forces.

Reactivity

• Reactivity decreases down the group because the incoming electron is further away, more shielded and less strongly attracted.

Displacement

• Displacement a more reactive halogen displaces a less reactive halide, e.g. Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq), giving an orange solution.

Reactivity order

• Cl₂ > Br₂ > I₂, so Br₂ displaces iodide to give a brown solution.

Halide tests

• Halide tests acidified AgNO₃(aq) gives white AgCl, cream AgBr and yellow AgI.

Section 6 of 45Oxides across period 3

Oxide	Bonding and structure	Nature	With water
Na2O	Giant ionic	Basic	NaOH, strongly alkaline
MgO	Giant ionic	Basic	Mg(OH)2, weakly alkaline
Al2O3	Ionic with covalent character	Amphoteric	No reaction; reacts with acids and bases
SiO2	Giant covalent	Acidic	No reaction
P4O10, SO3	Simple molecular	Acidic	H3PO4, H2SO4

Oxides across period 3

Trend

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

Cause

• Cause the electronegativity difference with oxygen falls across the period, so bonding changes from ionic to covalent.

Basic oxide equations

• Na₂O + H₂O → 2NaOH and MgO + H₂O → Mg(OH)₂, and MgO + 2HCl → MgCl₂ + H₂O.

Acidic oxide equations

• P₄O₁₀ + 6H₂O → 4H₃PO₄ and SO₃ + H₂O → H₂SO₄.

Down a group

• Down a group oxides become more ionic.

Section 7 of 45Acid deposition and ocean acidification

Acid rain

• Acid rain non-metal oxides dissolve in atmospheric water, SO₂ + H₂O → H₂SO₃ and 2NO₂ + H₂O → HNO₃ + HNO₂, lowering rainwater pH.

Ocean acidification

• CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq), a weak acid that partly dissociates: H₂CO₃ ⇌ H⁺ + HCO₃⁻.

Effect

• Effect the extra H⁺ lowers seawater pH and harms organisms that build calcium carbonate shells.

Section 8 of 45Oxidation and reduction

Definition	Oxidation	Reduction
Oxygen	Gain of oxygen	Loss of oxygen
Hydrogen	Loss of hydrogen	Gain of hydrogen
Electrons	Loss of electrons	Gain of electrons
Oxidation number	Increase	Decrease

Memory aid

• OIL RIGOxidation Is Loss, Reduction Is Gain of electrons.

Redox reaction

• Redox reaction any reaction in which an element changes its oxidation number.

Section 9 of 45Assigning oxidation states

1

Give uncombined elements 0 and monatomic ions their charge.

2

Give group 1 +1, group 2 +2 and fluorine −1.

3

Give hydrogen +1, except −1 in metal hydrides.

4

Give oxygen −2, except −1 in peroxides and +2 in OF₂.

5

Make the sum equal zero for a compound or the charge for an ion.

6

Solve for the unknown and write Roman numerals in names, e.g. manganate(VII).

Electronegativity

• Electronegativity the more electronegative element takes the negative oxidation state,
  • so OF₂ has F = −1 and O = +2,
  • whereas Cl₂O has O = −2 and Cl = +1.

Fractional states

• Fractional states an average over several atoms can be fractional, e.g. +8/3 for Fe in Fe₃O₄, but a single atom always has a whole-number state.

Examples

• Mn is +7 in KMnO₄, Cr is +6 in Cr₂O₇²⁻ and S is +6 in SO₄²⁻.

Writing

• Writing the sign comes before the number, as +2.

Section 10 of 45More from the specification

1Effective nuclear charge (Zeff)

• The net positive nuclear attraction experienced by an outermost electron after accounting for inner shell electron shielding.

2Atomic radius periodicity

• Atomic radius decreases across a period as added protons increase Z_eff, pulling the outermost electron shell closer to the nucleus.

3First ionisation energy trends

• General increase across a period reflecting greater electrostatic pull on the outermost electron.

4Oxidation states and overall charge

• In polyatomic ions, the algebraic sum of individual oxidation states of all atoms equals the overall charge of the ion.

Section 11 of 45Factors behind periodic trends

Nuclear charge

• Nuclear charge the number of protons, which rises across a period.

Number of shells

• Number of shells the outer electrons are further away as shells are added.

Shielding

• Shielding inner electrons repel the outer electrons and reduce the attraction.

Effective nuclear charge

• Effective nuclear charge the net attraction felt by the outer electrons.

Periodic trend

• Periodic trend a predictable change in a property across a period or down a group.

Section 12 of 45Atomic radius

Atomic radius

Definition

• Definition the distance from the nucleus to the outermost shell, often taken as half the mean distance between adjacent bonded atoms.

Across a period (decreases)

• Across a period (decreases) the nuclear charge increases, electrons enter the same shell, shielding is roughly constant and the outer electrons are pulled closer.

Down a group (increases)

• Down a group (increases) each element has an extra shell, so distance and shielding increase and the attraction weakens.

Group 18 to group 1

• Group 18 to group 1 the radius rises sharply because the alkali metal has one extra principal shell.

Section 13 of 45Ionic radius

Ionic radius

Cations

• Cations are smaller than their parent atoms because electrons are lost and the remaining electrons are pulled in by greater effective nuclear attraction.
  • (often a whole shell)

Anions

• Anions are larger than their parent atoms because extra electrons increase repulsion, which outweighs the increased nuclear charge.

Across a period

• Across a period cation radius decreases, then jumps up at the first anion, then decreases again.

Down a group

• Down a group ionic radius increases because more shells are added.

Section 14 of 45Ionisation energy

1First ionisation energy

First ionisation energy the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

2Conditions and units

• Conditions and units measured at 298 K and 100 kPa, in kJ mol⁻¹, e.g. Ca(g) → Ca⁺(g) + e⁻, ΔH = +590 kJ mol⁻¹.

3Across a period (generally increases)

• Across a period (generally increases) nuclear charge rises with constant shielding, so more energy is needed.

4Between periods

• Between periods the value falls sharply from the last element of one period to the first of the next,
• because greater distance and increased shielding outweigh the higher nuclear charge.

5Small dips

• Small dips occur at boron and at oxygen (repulsion between paired 2p electrons).
  • (2p is higher in energy than 2s)

6Down a group (decreases)

• Down a group (decreases) the outer electron is further away and more shielded, so it is easier to remove.

Section 15 of 45Electron affinity

1First electron affinity

First electron affinity the enthalpy change when one mole of gaseous atoms gains one mole of electrons to form gaseous 1− ions.

2Sign

• Sign usually exothermic, e.g. Cl(g) + e⁻ → Cl⁻(g), ΔH = −349 kJ mol⁻¹.

3Endothermic exceptions

• Endothermic exceptions groups 2, 12 and 15, e.g. N(g) + e⁻ → N⁻(g), ΔH = +7 kJ mol⁻¹, because of repulsion in a stable or half-filled subshell.

4Second electron affinity

• Second electron affinity endothermic, e.g. O⁻(g) + e⁻ → O²⁻(g), ΔH = +753 kJ mol⁻¹, because the electron is repelled by the negative ion.

5Across a period

• Across a period becomes more exothermic as nuclear charge rises and radius falls.

6Down a group

• Down a group becomes less exothermic as the added electron is further away and more shielded.

Section 16 of 45Electronegativity

Electronegativity

Definition

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

Across a period (increases)

• Across a period (increases) nuclear charge rises, shielding is constant and radius falls, so bonding electrons are held more strongly.

Down a group (decreases)

• Down a group (decreases) extra shells increase shielding and distance, so the attraction weakens.

Distinction

• Distinction electronegativity concerns covalent bonds and has no units, whereas electron affinity is a measurable enthalpy change for forming anions.

Electropositivity

• Electropositivity describes the tendency to form positive ions and is useful for metals.

Section 17 of 45More from the specification

More from the specification

Acid-Base Character of Period 3 Oxides

• Metallic oxides like sodium oxide (Na₂O) and magnesium oxide form alkaline solutions in water,
• whereas non-metal oxides like sulfur and chlorine oxides form acidic solutions.

Electronic Configuration of Alkali Metals

• Sodium and potassium share an outer s¹ electron configuration, readily undergoing oxidation by losing this valence electron to form +1 cations.

Redox Reactions with Halogens

• Potassium reacts vigorously with chlorine to form ionic potassium chloride,
• where oxidation of potassium is coupled with the reduction of chlorine.

Overall Oxidation Number Rules

• The overall sum of oxidation numbers in a neutral compound must equal zero, while in a polyatomic ion it equals the overall ionic charge.

Oxidation

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

Oxide

• A compound of oxygen with one other element.

Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

Oxidation number

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

Electron configuration

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

Section 18 of 45Layout of the periodic table

Layout of the periodic table

Order

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

Periods

• Periods rows, with the period number equal to the number of occupied shells.

Groups

• Groups columns, whose members have the same number of outer electrons and so similar reactivity.

Special cases

• Special cases hydrogen is placed on its own and helium is in group 18 because it is unreactive like the noble gases.

Periodicity

• Periodicity properties change predictably across a period and repeat down groups.

Data booklet

• Data booklet the periodic table in Section 7 does not label groups, blocks or element types.

Section 19 of 45Blocks and electron configuration

Blocks and electron configuration

Blocks

• Blocks s, p (groups 13–18), d (groups 3–12) and f (lanthanoids and actinoids), named after the subshell holding the last electron.
  • (groups 1–2 plus He)

Period from configuration

• Period from configuration the highest principal quantum number gives the period.

Group from configuration

• Group from configuration for main-group elements the outer s plus p electrons give the group number.
  • (add 10 in groups 13–18 numbering)

Position in p-block

• Position in p-block the superscript of the last p term shows how far along the block the element is.

Germanium

• Germanium ending 4s² 3d¹⁰ 4p² is period 4, group 14 and the second element in the p-block.

Sulfur

• Sulfur period 3, group 16 has [Ne] 3s² 3p⁴ and forms S²⁻ by gaining two electrons.

Use of blocks

• Use of blocks they help predict reactivity, oxidation states and bonding.

Section 20 of 45Metals, non-metals and metalloids

Class	Position	Properties
Metals	Left and centre (s-, d- and part of p-block)	Good conductors, malleable, ductile, shiny; form positive ions
Non-metals	Upper right (p-block)	Poor conductors, brittle solids or gases; form negative ions or covalent bonds; often form acidic oxides
Metalloids	Diagonal boundary (B, Si, Ge, As, Sb, Te)	Intermediate properties; semiconductors such as silicon

• Ions by group
  • group 1 forms 1+ ions and group 17 forms 1− ions.

Section 21 of 45Predicting properties from position

Predicting properties from position

Formula and bonding

• Formula and bonding position predicts ion charge, formula and bonding type.

Across a period

• Across a period elements change from metallic to non-metallic and oxides from basic to acidic.

Down group 1

• Down group 1 reactivity increases because the outer electron is further from the nucleus, more shielded and lost more easily.

Down group 17

• Down group 17 reactivity decreases because an incoming electron is attracted less strongly.

Skill

• Skill state the group, period and block from a configuration, or give a configuration from a position.

Section 22 of 45Group 1 and group 17 physical trends

Group 1 and group 17 physical trends

Group 1 melting points

• Group 1 melting points fall down the group because metallic bonding weakens as the ions get larger.

Group 17 melting and boiling points

• Group 17 melting and boiling points rise down the group because London forces increase with electron number.

States

• States chlorine is a gas, bromine a liquid and iodine a solid.

Displacement

• Displacement a more reactive halogen displaces a less reactive halide, e.g. chlorine displaces bromide ions, which tests the reactivity order.

Section 23 of 45Period 3 oxides

Na₂O

• Na₂O ionic and basic.

Al₂O₃

• Al₂O₃ amphoteric.

SO₃

• SO₃ covalent and acidic.

Trend

• Trend oxides change from ionic and basic to covalent and acidic across the period.

Section 24 of 45More from the specification

More from the specification

Group 1 Alkali Metal Trends

• Down Group 1 from lithium to sodium and potassium, atomic radius increases as additional electron shells are added,
• increasing electron shielding and decreasing first ionisation energy.

Electronegativity Across the Periodic Table

• Electronegativity is the relative ability of an atom in a molecule to attract bonding electron pairs;
  • it decreases down groups and increases across periods toward fluorine.

Reactions of Alkali Metals with Water

• Group 1 metals react vigorously with water to produce hydrogen gas and an alkaline metal hydroxide solution (2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)), with reactivity increasing down the group.

Energy Changes and Metallic Properties

• Because valence electrons in heavier alkali metals like potassium are held less tightly by the nucleus, less energy is required to lose the outer electron,
• resulting in more vigorous chemical reactions.

Alkali metals

• The Group I metals, which form alkaline solutions with water.

Electronegativity

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

Ionisation

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

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.

Ionisation energy

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

Section 25 of 45Empirical and molecular formulae

Empirical formula

Empirical formula the simplest whole-number ratio of atoms in a compound, e.g. HO for hydrogen peroxide.

Molecular formula

Molecular formula the actual number of atoms of each element in a molecule, e.g. C₄H₁₀ for butane and C₄H₈ for butene.

Relationship

• Relationship the empirical formula of butane is C₂H₅ and that of glucose, C₆H₁₂O₆, is CH₂O.

Confusion

• Confusion do not mix up the two; many organic compounds have different empirical and molecular formulae.

Section 26 of 45Structural and condensed formulae

Structural and condensed formulae

Structural (displayed) formula

• Structural (displayed) formula shows all the atoms and all the bonds in the molecule.

Condensed formula

• Condensed formula shows enough to make the structure clear but omits most bonds, e.g. CH₃CH(OH)CH₃.

Brackets

• Brackets identical groups can be bracketed together.

Double and triple bonds

• Double and triple bonds need not always be shown, and can be deduced from the number of hydrogens on carbon.

Moving between formulas

• Moving between formulas butane is C₄H₁₀ with the condensed formula CH₃CH₂CH₂CH₃.

Section 27 of 45Skeletal formulae

C–C bonds

• C–C bonds drawn as straight lines.

Carbon atoms

• Carbon atoms each end or junction of a line is a carbon atom.

Hydrogen atoms

• Hydrogen atoms on carbon are omitted, but those in functional groups such as –OH are shown.

Other atoms

• Other atoms and functional groups are shown explicitly.

Methane

• Methane has no skeletal formula because skeletal formulae rely on C–C bonds;
  • a dot could be mistaken for a radical.

Section 28 of 45Stereochemical formulae

Purpose

• Purpose show the three-dimensional arrangement of atoms or groups around a carbon.

Solid line

• Solid line the bond lies in the plane of the paper.

Solid wedge

• Solid wedge the bond comes out towards the viewer.

Dashed wedge

• Dashed wedge the bond goes behind the plane, away from the viewer.

Chiral carbon

• Chiral carbon bonded to four different groups in a tetrahedral arrangement of about 109.5°.

Section 29 of 45Homologous series

Definition

Definition a family of compounds with the same functional group and general formula, in which successive members differ by a CH₂ unit (the homologous increment).

Every added carbon

• Every added carbon adds two hydrogen atoms as well.

Chemical properties

• Chemical properties similar, because they share a functional group.

Physical properties

• Physical properties change gradually, such as boiling point, melting point and density.

Series	General formula	Example
Alkanes	CnH2n+2	Propane, C3H8
Alkenes (n ≥ 2)	CnH2n	Propene, C3H6
Alkynes	CnH2n−2	Propyne, C3H4
Halogenoalkanes	CnH2n+1X	Chloropropane, C3H7Cl
Alcohols	CnH2n+1OH	Propanol, C3H7OH
Aldehydes and ketones	CnH2nO	Propanal and propanone, C3H6O
Carboxylic acids	CnH2n+1COOH	Propanoic acid, C2H5COOH
Amines	CnH2n+1NH2	Propylamine, C3H7NH2

Section 30 of 45Physical trends in a homologous series

Physical trends in a homologous series

Boiling and melting points

• Boiling and melting points increase as the number of carbons in the chain increases.

Cause

• Cause a larger molecule has more electrons and a greater surface area,
• so London dispersion forces are stronger and more energy is needed to separate molecules.

Series showing the trend

• Series showing the trend alkanes, primary alcohols, halogenoalkanes and carboxylic acids.

Alkane states

• Alkane states methane and ethane are gases, while longer alkanes are liquids and then solids.
  • (ethane boils at −89 °C)

Other properties

• Other properties viscosity and density also change gradually.

Section 31 of 45Classes of compound and functional groups

Classes of compound and functional groups

Alcohol

• –OH, suffix –ol, e.g. ethanol.

Aldehyde

• –CHO, suffix –al, e.g. propanal.

Carboxylic acid

• –COOH, suffix –oic acid, e.g. ethanoic acid.

Halogenoalkane

• –X, prefix chloro– or bromo–, e.g. chloromethane.

Haloalkane definition

• Haloalkane definition an alkane in which one or more hydrogens are replaced by halogens.

Example of naming

• CH₃CHBrCH₃ is 2-bromopropane.

Section 32 of 45More from the specification

More from the specification

IUPAC Nomenclature Principles

• A systematic nomenclature system where the longest unbranched carbon chain gives the stem name,
• and attached alkyl substituents are identified by numbering.

Structural Isomerism in Alkanes

• Saturated hydrocarbons can exist as straight-chain or branched isomers that share the same molecular formula but have different structural connectivities.

Primary, Secondary, and Tertiary Carbons

• Carbon atoms are classified as primary, secondary,
• or tertiary according to whether they are bonded to one, two,
• or three other alkyl groups.

Aromatic Hydrocarbon Nomenclature

• Derivatives of benzene are named systematically by prefixing substituent group names or treating phenyl as an attached radical group.

IUPAC Nomenclature Principles

• A systematic nomenclature system where the longest unbranched carbon chain gives the stem name,
• and attached alkyl substituents are identified by numbering.

Structural Isomerism in Alkanes

• Saturated hydrocarbons can exist as straight-chain or branched isomers that share the same molecular formula but have different structural connectivities.

Primary, Secondary, and Tertiary Carbons

• Carbon atoms are classified as primary, secondary,
• or tertiary according to whether they are bonded to one, two,
• or three other alkyl groups.

Aromatic Hydrocarbon Nomenclature

• Derivatives of benzene are named systematically by prefixing substituent group names or treating phenyl as an attached radical group.

Isomer

• Same formula, different structure

Branched chain

• A carbon chain with a side group attached.

Structural isomers

• Compounds with the same molecular formula but different structural formulae.

Section 33 of 45Three definitions of oxidation and reduction

Definition	Oxidation	Reduction
Oxygen	Gain of oxygen, e.g. 2Mg + O2 → 2MgO	Loss of oxygen, e.g. 2CuO + C → 2Cu + CO2
Hydrogen	Loss of hydrogen, e.g. CH3OH → CH2O + H2O	Gain of hydrogen, e.g. C2H4 + H2 → C2H6
Electrons	Loss of electrons, e.g. Al → Al3+ + 3e−	Gain of electrons, e.g. F2 + 2e− → 2F−
Oxidation number	Increase	Decrease

• Memory aid
  • OIL RIGOxidation Is Loss, Reduction Is Gain of electrons.

Section 34 of 45What an oxidation number is

Definition

Definition the charge an atom would have if all its bonds were completely ionic.

Purpose

• Purpose shows whether oxidation or reduction has occurred, identifies what is oxidised or reduced and helps build half-equations.

Writing

• Writing the sign comes before the number, as +2, and the terms oxidation number and oxidation state are interchangeable.

Atoms and ions

• Na is 0, Na⁺ is +1 and Cl⁻ is −1.

Electronegativity

• Electronegativity the more electronegative atom is given the negative oxidation number.

Section 35 of 45Rules for oxidation numbers

Rule	Example
Uncombined element is 0	H2, Zn, O2
Monatomic ion equals its charge	Zn2+ = +2, Cl− = −1
Group 1 is +1, group 2 is +2, fluorine is −1	NaCl: Na +1, Cl −1
Hydrogen is +1, but −1 in metal hydrides	NaH: H −1
Oxygen is −2, but −1 in peroxides and +2 in OF2	H2O2: O −1
Sum is 0 in a compound and equals the charge in an ion	SO42−: S +6, O −2
Averages may be fractional	S4O62−: average S +2.5

• Electronegativity trend
  • it increases across a period and decreases down a group, which decides which atom is negative.

Section 36 of 45Calculating an oxidation number

1

Write the known oxidation numbers under each element.

2

Multiply each by the number of atoms.

3

Let the unknown be x and write an equation equal to the total charge.

4

Solve for x.

Examples

• Examples in KMnO₄, +1 + x + 4(−2) = 0 gives Mn = +7; in Cr₂O₇²⁻, 2x + 7(−2) = −2 gives Cr = +6.

Whole numbers

• Whole numbers a single atom always has a whole-number oxidation state, and fractions are averages over atoms in different environments.

Section 37 of 45Identifying redox reactions

Redox reaction

• Redox reaction any reaction in which an element changes oxidation number.

Oxidation

• Oxidation an increase in oxidation number.

Reduction

• Reduction a decrease in oxidation number.

Reducing agent

• Reducing agent is oxidised and the oxidising agent is reduced, e.g. in Zn + Cu²⁺ → Zn²⁺ + Cu zinc (0 to +2) is oxidised and copper(II) (+2 to 0) is reduced.

Electron balance

• Electron balance in 2Mg + O₂ → 2MgO the four electrons lost equal the four gained, which is why the equation balances.

Disproportionation

• Disproportionation the same element is both oxidised and reduced, as chlorine in cold NaOH forms Cl⁻ (−1) and ClO⁻ (+1).

Section 38 of 45More from the specification

More from the specification

Functional organic substituent groups

• Functional organic substituent groups characteristic reactive arrangements of atoms form functional groups that dictate chemical reactivity across homologous series.

Atomic Radius and Nuclear Charge

• Moving across a period, atomic radius decreases because increasing positive charge in the nucleus pulls outer electron shells closer with similar shielding.

Reactivity of Group 1 Alkali Metals

• As atomic radius increases down the group from sodium to potassium, less energy is needed to remove the outer valence electron,
• increasing reactivity with water.

Electronic Configuration Trends

• The periodic table arranges metals and non-metals according to ground-state electron configuration, reflecting progressive filling of outer s, p,
• and d subshells.

Alkali Metals in Aqueous Solution

• When alkali metals dissolve and react in water, they form strongly alkaline hydroxide solutions with an overall release of heat and hydrogen gas.

Alkali metals

• The Group I metals, which form alkaline solutions with water.

Outer electrons

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

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.

Electron configuration

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

Electronic configuration

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

Ionic radius

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

Electron shell

• An energy level for electrons.

Valence electrons

• Electrons in the outer shell.

Nucleus

• Central part of an atom with protons and neutrons.

Shell

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

Ionisation

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

Section 39 of 45Why Roman numerals are used

Variable states

• Variable states many transition metals form more than one ion, so the Roman numeral in the name gives the oxidation state.

Stock notation

• Stock notation the system of Roman numerals is named after the German chemist Alfred Stock.

Examples

• FeO contains iron(II) and Fe₂O₃ contains iron(III).

Non-metals

• Non-metals the notation is not normally used for non-metals, e.g. SO₂ is sulfur dioxide.

Section 40 of 45Method for finding the oxidation state

1

Assign the known oxidation states, e.g. oxygen −2 and group 1 +1.

2

Multiply each by its subscript.

3

Find the value of the unknown that makes the total zero (or the ion charge).

4

Write the Roman numeral after the metal name.

Neutral compound

• Neutral compound the oxidation numbers must sum to zero.

Read the subscripts

• Read the subscripts the formula must be read carefully before assigning states.

Section 41 of 45Copper and manganese compounds

Cu₂O

• Cu₂O oxygen is −2, so the two copper atoms total +2 and each is +1, giving copper(I) oxide.

MnSO₄

• MnSO₄ sulfate has a charge of −2, so manganese is +2, giving manganese(II) sulfate.

Naming rule

• Naming rule only the transition metal is given Stock notation in this context.

Section 42 of 45Chromium and manganese oxoanions

Na₂CrO₄

• Na₂CrO₄ two Na give +2 and four O give −8, so Cr is +6, giving sodium chromate(VI).

KMnO₄

• K gives +1 and four O give −8, so Mn is +7, giving potassium manganate(VII).

Na₂Cr₂O₇

• Na gives +2 and seven O give −14, so the two Cr total +12 and each is +6, giving sodium dichromate(VI).

Prefix di

• Prefix di shows the ion contains two chromium atoms, distinguishing it from chromate.

Section 43 of 45More from the specification

1

Aliphatic alkyl substituent groups

• Aliphatic alkyl substituent groups branched hydrocarbon side chains act as alkyl groups, altering boiling points and steric hindrance around reactive centres.

2

Periodic Trends Across Period 3

• Across Period 3 from sodium to argon, nuclear charge increases by one proton per element while inner shielding remains constant,
• causing atomic radius to decrease across the period.

3

First Ionisation Energy Across a Period

• First ionisation energy generally increases across Period 3 due to higher effective nuclear charge, with slight dips at aluminium and sulfur (paired 3p electrons).
  • (electron in higher-energy 3p subshell)

4

Acid-Base Nature of Period 3 Oxides

• Across Period 3, bonding in oxides changes from ionic to covalent: Na₂O and MgO are basic, Al₂O₃ is amphoteric,
• and SiO₂, P₄O₁₀,
• and SO₃ form acidic oxides in water.

5

Successive Ionisation Energy Patterns

• Removing successive electrons from an atom requires increasing amounts of energy;
  • dramatic jumps in ionisation energy reveal the boundary between valence electrons and inner noble gas shells.

6

Ionic radius

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

7

Alkali metals

• The Group I metals, which form alkaline solutions with water.

8

Ionisation

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

9

Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

10

Outer electrons

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

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

Electron configuration

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

14

Ionisation energy

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

15

Electronic configuration

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

16

Electron

• Negative particle moving around the nucleus.

17

Period

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

18

Element

• A substance made of only one type of atom.

19

Nucleus

• Central part of an atom with protons and neutrons.

20

Electronegativity

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

Section 44 of 45Key terms

Term	Meaning
Period	A horizontal row of the periodic table; the number equals the occupied main energy levels.
Group	A vertical column of elements with the same number of valence electrons.
Block	A region of the table named after the sublevel being filled (s, p, d or f).
Electron affinity	The energy change when one mole of gaseous atoms gains electrons to form 1− ions.
Shielding	The reduction in nuclear attraction caused by inner-shell electrons.
Oxidation state	The charge an atom would have if all its bonds were ionic.
Amphoteric	Able to react as both an acid and a base, e.g. Al2O3.

Term	Meaning
Periodic trend	A predictable change in a property across a period or down a group.
Atomic radius	The distance from the nucleus to the outermost shell.
Ionisation energy	The energy to remove one mole of electrons from one mole of gaseous atoms.
Electron affinity	The enthalpy change when gaseous atoms gain electrons.
Shielding	The repulsion of outer electrons by inner electrons.
Effective nuclear charge	The net attraction felt by outer electrons from the nucleus.

Term	Meaning
Period	A horizontal row of the periodic table.
Group	A vertical column of elements with similar properties.
Block	A region of the table defined by the subshell being filled (s, p, d or f).
Metalloid	An element with properties between metals and non-metals.
Valence electrons	Electrons in the outer shell.
Atomic number	The number of protons in the nucleus.

Term	Meaning
Empirical formula	The simplest whole-number ratio of atoms in a compound.
Molecular formula	The actual number of atoms of each element in one molecule.
Skeletal formula	A simplified diagram in which each line end or junction is a carbon atom.
Homologous series	A family of compounds with the same functional group and general formula, differing by CH2.
Functional group	The atom or group that gives a compound its characteristic reactions.
Stereochemical formula	A formula using wedges to show the three-dimensional arrangement of bonds.

Term	Meaning
Oxidation	Gain of oxygen, loss of hydrogen or electrons, or an increase in oxidation number.
Reduction	Loss of oxygen, gain of hydrogen or electrons, or a decrease in oxidation number.
Oxidation number	The charge an atom would have if bonding were completely ionic.
Peroxide	A compound containing O–O in which oxygen is −1.
Redox reaction	A reaction involving both oxidation and reduction.
Disproportionation	A reaction in which the same element is both oxidised and reduced.

Term	Meaning
Stock notation	Roman numerals in a name that show the oxidation state of a metal.
Oxidation state	The charge an atom would have if all bonds were ionic.
Chromate(VI)	The CrO42− ion containing chromium in the +6 state.
Dichromate(VI)	The Cr2O72− ion containing chromium in the +6 state.
Manganate(VII)	The MnO4− ion containing manganese in the +7 state.

Section 45 of 45Exam technique

1

Saying atomic radius increases across a period — nuclear charge increases and the radius decreases.

2

Assuming all halogens become more reactive down the group — reactivity decreases down group 17 but increases down group 1.

3

Forgetting oxygen is −1 in peroxides — it is not always −2.

4

Confusing electronegativity with electron affinity — electronegativity concerns bonded electrons and electron affinity is a measured enthalpy change.

1

Giving trends without nuclear charge and shielding — explain with nuclear charge, shells and shielding.

2

Claiming nuclear charge decreases across a period — it increases.

3

Forgetting that anions are larger than their atoms — extra electrons increase repulsion.

4

Omitting the gaseous state in the ionisation energy definition — atoms must be gaseous.

1

Mixing up groups (columns) and periods (rows) — groups are columns and periods are rows.

2

Using 4s to decide the block of a transition metal — the block is set by the last subshell filled in the order of increasing energy (d).

3

Saying elements in a period have similar properties — it is elements in a group that do.

4

Forgetting that the period equals the highest occupied shell.

1

Confusing empirical and molecular formulae — the empirical formula is the simplest ratio and the molecular formula is the actual count.

2

Drawing methane as a skeletal formula — skeletal formulae need C–C bonds.

3

Thinking members of a series have identical chemical and physical properties — chemistry is similar and physical properties change gradually.

4

Counting a hydrogen on carbon in a skeletal formula — these are omitted but those in functional groups are shown.

1

Writing oxidation numbers as 2+ — write +2, with the sign first.

2

Assuming oxygen is always −2 — it is −1 in peroxides and +2 in OF₂.

3

Forgetting the ion charge when the sum is not zero — the sum equals the charge.

4

Thinking oxidation numbers must be whole for the average — an average over several atoms can be a fraction.

1

Forgetting to multiply by the subscript — oxidation numbers are totalled over all atoms of each element.

2

Using the wrong charge for the oxoanion — find the ion charge from the metal and other ions in the formula.

3

Omitting the di prefix — dichromate(VI) and chromate(VI) are different ions.

4

Giving the Roman numeral as the ion charge for covalent oxoanions — it is the oxidation state of the metal.