S3.1Structure 3 — Classification of matter
The periodic table: classification of elements — IB Chemistry HL Revision Notes & Topical Questions
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Lesson summary
S3.1 The periodic table: classification of elements — IB CHEMISTRY HL notes: Use the layout of the periodic table, periodic trends, group chemistry, oxides, oxidation numbers and transition element properties.Order — elements are arranged by increasing atomic number, 1 to 118.Per…
- Periodicity
- Non-metals
- Metalloids
- Know without the table
- Shorthand configuration
- Reading a configuration
Part 1 of 11Layout 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.
Part 2 of 11Metals, 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.
Part 3 of 11Configuration 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.
Part 4 of 11Atomic 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.
Part 5 of 11Ionisation 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.
Part 6 of 11Electronegativity
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.
Part 7 of 11Group 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.
Part 8 of 11Period 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.
Part 9 of 11Oxidation 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.
Part 10 of 11Transition 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.
- and absorbing light of energy equal to the gap promotes an electron;
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.
Part 11 of 11More 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)
Key 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. |
Avoid these mistakes
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.
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