Pearson Edexcel A Level Chemistry (9CH0)

Topic 1 — Atomic Structure and the Periodic Table

1. Specification Coverage

Pearson Edexcel Topic 1: Atomic Structure and the Periodic Table contains 26 specification statements.

This topic is assessed directly in Paper 1: Advanced Inorganic and Physical Chemistry, and may also appear synoptically in Paper 3.

SPEC: 1 — Structure of the atom

Know the structure of an atom in terms of:

• electrons;
• protons;
• neutrons.

SPEC: 2 — Relative mass and charge

Know the relative:

• mass;
• charge;

of:

• protons;
• neutrons;
• electrons.

SPEC: 3 — Atomic number and mass number

Know what is meant by:

• atomic (proton) number;
• mass number.

SPEC: 4 — Numbers of subatomic particles

Determine numbers of:

• protons;
• neutrons;
• electrons;

in an:

• atom;
• molecule;
• ion;

from atomic number and mass number.

SPEC: 5 — Isotopes

Understand the term isotopes.

SPEC: 6 — Relative isotopic and atomic mass

Define:

• relative isotopic mass;
• relative atomic mass;

using the carbon-12 scale.

SPEC: 7 — Relative molecular and formula mass

Understand:

• relative molecular mass;
• relative formula mass;

and calculate them from relative atomic masses.

Use relative formula mass for giant structures.

Formal definitions of these two terms are not required by Pearson.

SPEC: 8 — Mass spectrometry and isotopic abundance

Analyse and interpret mass-spectrometry data to calculate:

• relative atomic mass from isotopic abundance;
• isotopic abundance from relative atomic mass.

SPEC: 9 — Mass spectra of diatomic molecules

Predict:

• mass spectra;
• relative peak heights;

for diatomic molecules, including chlorine.

SPEC: 10 — Molecular-ion peak

Understand how mass spectrometry determines relative molecular mass using the:

[ M^+ ]

molecular-ion peak.

SPEC: 11 — Ionisation energy definitions

Define:

• first ionisation energy;
• successive ionisation energies.

SPEC: 12 — Factors affecting ionisation energy

Understand effects of:

• number of protons;
• electron shielding;
• electron subshell from which the electron is removed.

SPEC: 13 — Across a period

Understand why first ionisation energy generally increases across a period.

SPEC: 14 — Down a group

Understand why first ionisation energy decreases down a group.

SPEC: 15 — Evidence for electronic configuration

Understand how electronic-configuration ideas developed from:

• atomic emission spectra as evidence for quantum shells;
• successive ionisation energies as evidence for quantum shells and group;
• first ionisation energies of successive elements as evidence for subshells.

SPEC: 16 — Capacity of quantum shells

Know the number of electrons that can fill the first four quantum shells.

SPEC: 17 — Orbital definition

Know that an orbital is a region within an atom that can hold up to two electrons with opposite spins.

SPEC: 18 — Orbital shapes

Know the shapes of:

• (s)-orbitals;
• (p)-orbitals.

SPEC: 19 — Subshell capacities

Know the number of electrons that occupy:

• (s);
• (p);
• (d);

subshells.

SPEC: 20 — Filling orbitals

Know that:

• electrons fill orbitals singly before pairing;
• two electrons in the same orbital have opposite spins.

SPEC: 21 — Electronic configurations

Predict electronic configurations using:

• (1s) notation;
• electrons-in-boxes notation;

for:

• atoms up to (Z=36);
• s- and p-block ions only up to (Z=36).

SPEC: 22 — Periodic-table blocks

Know that elements are classified as:

• (s)-block;
• (p)-block;
• (d)-block.

SPEC: 23 — Configuration and chemical properties

Understand that electronic configuration determines the chemical properties of an element.

SPEC: 24 — Periodicity

Understand periodicity as a repeating pattern across different periods.

SPEC: 25 — Period 2 and 3 trends

Understand reasons for trends in Period 2 and Period 3:

• melting and boiling temperatures, using structure and bonding;
• first ionisation energy, using data or recalled plots.

SPEC: 26 — Illustrating periodicity using data

Use data including:

• electronic configurations;
• atomic radii;
• melting temperatures;
• boiling temperatures;
• first ionisation energies;

to illustrate periodicity.

2. Core Content — Extreme Detail

2.1 SPEC: 1–4 — Atomic Structure and Subatomic Particles

2.1.1 Structure of an atom

An atom consists of:

• a small, dense nucleus;
• electrons outside the nucleus.

The nucleus contains:

• protons;
• neutrons.

Most of an atom's mass is concentrated in the nucleus.

2.1.2 Proton

DEFINITION — Proton
A positively charged subatomic particle found in the nucleus, with relative charge (+1) and relative mass approximately (1).

2.1.3 Neutron

DEFINITION — Neutron
An uncharged subatomic particle found in the nucleus, with relative charge (0) and relative mass approximately (1).

2.1.4 Electron

DEFINITION — Electron
A negatively charged subatomic particle found outside the nucleus, with relative charge (-1) and very small relative mass.

The relative mass of an electron is approximately:

[ \frac{1}{1836} ]

of the mass of a proton.

2.1.5 Required particle table

2.2 Atomic Number and Mass Number

2.2.1 Atomic number

DEFINITION — Atomic (proton) number, (Z)
The number of protons in the nucleus of an atom.

The atomic number identifies the element.

For a neutral atom:

[ \boxed{ \text{number of electrons}=Z } ]

2.2.2 Mass number

DEFINITION — Mass number, (A)
The total number of protons and neutrons in the nucleus.

Therefore:

[ \boxed{ A=p+n } ]

and:

[ \boxed{ n=A-Z } ]

2.2.3 Nuclide notation

[ {}^{A}_{Z}X ]

Example:

[ {}^{27}_{13}\mathrm{Al} ]

contains:

• 13 protons;
• (27-13=14) neutrons;
• 13 electrons if neutral.

2.3 Particle Numbers in Ions and Molecules

2.3.1 Positive ions

A positive ion has lost electrons.

Example:

[ {}^{24}_{12}\mathrm{Mg}^{2+} ]

contains:

• 12 protons;
• 12 neutrons;
• 10 electrons.

2.3.2 Negative ions

A negative ion has gained electrons.

Example:

[ {}^{35}_{17}\mathrm{Cl}^{-} ]

contains:

• 17 protons;
• 18 neutrons;
• 18 electrons.

2.3.3 General relationship

[ \boxed{ \text{charge}=\text{protons}-\text{electrons} } ]

when charge is counted in elementary-charge units.

2.3.4 Particles in molecules

For a molecule, total numbers are obtained by adding contributions from every atom.

Example:

[ {}^{1}\mathrm{H}_2{}^{16}\mathrm{O} ]

Total protons:

[ 2(1)+8=10 ]

Total neutrons:

[ 2(0)+8=8 ]

Total electrons in the neutral molecule:

[ 10 ]

2.4 SPEC: 5 — Isotopes

DEFINITION — Isotopes
Atoms of the same element with the same number of protons but different numbers of neutrons.

Therefore isotopes have:

• same atomic number;
• different mass numbers;
• different masses.

Example:

[ {}^{35}_{17}\mathrm{Cl} ]

and:

[ {}^{37}_{17}\mathrm{Cl} ]

Both have 17 protons.

Their neutron numbers are:

[ 35-17=18 ]

and:

[ 37-17=20 ]

respectively.

2.4.1 Chemical similarity

Neutral isotopes of the same element have the same electron configuration.

Chemical behaviour is determined mainly by electrons, especially the outer-shell configuration.

Therefore isotopes have essentially the same chemical properties.

Examiner language

“The isotopes have the same number and arrangement of electrons, so they show the same chemical behaviour.”

2.5 SPEC: 6 — Relative Isotopic Mass

DEFINITION — Relative isotopic mass
The mass of an atom of an isotope relative to one-twelfth of the mass of an atom of carbon-12.

It has no unit because it is a ratio.

Crucial distinction

Mass number is an integer count of protons and neutrons.

Relative isotopic mass is an experimentally determined relative mass and need not be an exact integer.

2.6 Relative Atomic Mass

DEFINITION — Relative atomic mass, (A_r)
The weighted mean mass of an atom of an element relative to one-twelfth of the mass of an atom of carbon-12.

The phrase weighted mean is essential.

[ \boxed{ A_r= \frac{\sum(\text{relative isotopic mass}\times\text{abundance})} {\sum\text{abundance}} } ]

2.6.1 Example

An element consists of:

• isotope 24: 79.0%;
• isotope 25: 10.0%;
• isotope 26: 11.0%.

[ A_r= \frac{24(79)+25(10)+26(11)}{100} ]

[ =\frac{2432}{100} ]

[ \boxed{ A_r=24.32 } ]

No unit.

2.7 SPEC: 7 — Relative Molecular and Formula Mass

2.7.1 Relative molecular mass

For a discrete molecule:

[ \boxed{ M_r=\sum A_r } ]

Example:

[ \mathrm{CO_2} ]

[ M_r=12.0+2(16.0) ]

[ \boxed{ M_r=44.0 } ]

Pearson does not require a formal definition of (M_r) in Topic 1, but you must understand and calculate it correctly.

2.7.2 Relative formula mass

Use relative formula mass for compounds with giant structures.

Examples:

• (\mathrm{NaCl})
• (\mathrm{MgO})
• (\mathrm{SiO_2})

For:

[ \mathrm{NaCl} ]

[ 23.0+35.5 ]

[ \boxed{ 58.5 } ]

Lost-mark distinction

Do not call NaCl's value a relative molecular mass because NaCl does not exist as discrete NaCl molecules in its giant ionic lattice.

2.8 SPEC: 8 — Mass Spectrometry and Isotopic Abundance

A mass spectrum gives information about ions according to their:

[ \boxed{ m/z } ]

ratio.

For singly charged isotopic ions:

[ z=1 ]

so the numerical (m/z) value corresponds to relative isotopic mass.

2.8.1 Interpreting peaks

A simple elemental mass spectrum plots:

• x-axis: [ m/z ]
• y-axis: relative abundance / relative intensity.

Peak position gives:

[ m/z ]

Peak height or area represents relative abundance.

2.8.2 Calculating (A_r) from a spectrum

Suppose peaks occur at:

[ 63,\ 65 ]

with relative intensities:

[ 69,\ 31 ]

Then:

[ A_r= \frac{63(69)+65(31)}{100} ]

[

63.62 ]

[ \boxed{ A_r=63.62 } ]

2.8.3 Calculating abundance from (A_r)

Suppose an element has two isotopes:

[ {}^{35}X ]

and:

[ {}^{37}X ]

and:

[ A_r=35.5 ]

Let fraction of isotope 35 be:

[ x ]

Then fraction of isotope 37 is:

[ 1-x ]

Weighted mean:

[ 35x+37(1-x)=35.5 ]

[ 35x+37-37x=35.5 ]

[ -2x=-1.5 ]

[ x=0.75 ]

Therefore:

[ \boxed{ 75%\ {}^{35}X } ]

and:

[ \boxed{ 25%\ {}^{37}X } ]

2.9 SPEC: 9 — Mass Spectra of Diatomic Molecules

This is an important Edexcel-specific requirement.

A diatomic molecule such as:

[ \mathrm{Cl_2} ]

can contain different isotope combinations.

If chlorine has isotopes:

[ {}^{35}\mathrm{Cl} ]

and:

[ {}^{37}\mathrm{Cl} ]

then the possible chlorine molecules are:

[ {}^{35}\mathrm{Cl}-{}^{35}\mathrm{Cl} ]

[ {}^{35}\mathrm{Cl}-{}^{37}\mathrm{Cl} ]

[ {}^{37}\mathrm{Cl}-{}^{37}\mathrm{Cl} ]

Their molecular-ion (m/z) values are:

[ 70,\ 72,\ 74 ]

respectively.

2.9.1 Relative peak heights using probability

Natural chlorine is approximately:

[ 75%{}^{35}\mathrm{Cl} ]

and:

[ 25%{}^{37}\mathrm{Cl} ]

Represent as:

[ p({}^{35}\mathrm{Cl})=\frac34 ]

[ p({}^{37}\mathrm{Cl})=\frac14 ]

(m/z=70)

[ P(35,35)

\frac34\times\frac34

\frac9{16} ]

(m/z=72)

There are two arrangements:

[ 35-37 ]

and:

[ 37-35 ]

Thus:

[ P

2\left(\frac34\times\frac14\right)

\frac6{16} ]

(m/z=74)

[ P(37,37)

\frac14\times\frac14

\frac1{16} ]

Therefore peak-height ratio:

[ \boxed{ 9:6:1 } ]

for:

[ 70:72:74 ]

This is a classic Edexcel calculation.

2.10 SPEC: 10 — Molecular Ion, (M^+)

In mass spectrometry, a molecule can lose one electron while remaining intact:

[ M(g)\rightarrow M^+(g)+e^- ]

The resulting intact ion is the:

DEFINITION — Molecular ion, (M^+)
The positively charged intact molecular species formed by removal of one electron from a molecule.

For a singly charged molecular ion:

[ \boxed{ m/z=M_r } ]

Thus the molecular-ion peak can determine relative molecular mass.

Example

If the molecular-ion peak is:

[ m/z=86 ]

then:

[ \boxed{ M_r=86 } ]

assuming the peak corresponds to singly charged (M^+).

2.11 SPEC: 11 — First Ionisation Energy

DEFINITION — First ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous (1+) ions.

Equation:

[ \boxed{ X(g)\rightarrow X^+(g)+e^- } ]

Essential definition components:

• one electron;
• one mole;
• gaseous atoms;
• gaseous (1+) ions.

2.12 Successive Ionisation Energies

DEFINITION — Successive ionisation energies
The energies required to remove electrons one at a time from one mole of gaseous species, forming ions with successively higher positive charges.

Second ionisation:

[ \boxed{ X^+(g)\rightarrow X^{2+}(g)+e^- } ]

Third ionisation:

[ \boxed{ X^{2+}(g)\rightarrow X^{3+}(g)+e^- } ]

Do not define the second ionisation energy as removal of two electrons from a neutral atom.

2.13 SPEC: 12 — Factors Affecting Ionisation Energy

Pearson specifies three central factors:

1. number of protons;
2. electron shielding;
3. subshell from which the electron is removed.

Distance from the nucleus is also naturally involved in explaining shell effects.

2.13.1 Number of protons

More protons give a greater nuclear charge.

All else equal, this increases electrostatic attraction between:

• the positive nucleus;
• the electron.

Therefore more energy is needed to remove the electron.

2.13.2 Shielding

DEFINITION — Electron shielding
Reduction in the attraction between the nucleus and an outer electron due to repulsion by electrons between that electron and the nucleus.

Greater shielding lowers the effective nuclear attraction experienced by the outer electron.

Therefore ionisation energy tends to decrease.

2.13.3 Subshell

Electrons in different subshells have different energies and penetration/shielding environments.

For example, a (3p) electron is higher in energy and more shielded than a (3s) electron in the same principal shell.

This explains some departures from the general period trend.

2.14 SPEC: 13 — First Ionisation Energy Across a Period

General trend:

[ \boxed{ IE_1\text{ generally increases across a period} } ]

Reason:

• proton number increases;
• nuclear charge increases;
• electrons are added to the same principal shell;
• shielding increases only slightly;
• nuclear attraction to the outer electron becomes stronger;
• atomic radius generally decreases.

Examiner-safe wording

“Across the period, nuclear charge increases while shielding changes relatively little because electrons are added to the same principal shell. The outer electron is therefore more strongly attracted to the nucleus, so more energy is required to remove it.”

2.15 Period 3 Anomaly — Mg to Al

Mg:

[ [\mathrm{Ne}],3s^2 ]

Al:

[ [\mathrm{Ne}],3s^2,3p^1 ]

The first ionisation energy decreases from Mg to Al.

The electron removed from Al is in a:

[ 3p ]

subshell.

This is higher in energy and more shielded than the (3s) electron removed from Mg.

Therefore the Al electron is easier to remove.

High-mark wording

“Al loses a 3p electron whereas Mg loses a 3s electron. The 3p electron is higher in energy and more shielded, so it experiences weaker attraction to the nucleus and requires less energy to remove.”

2.16 Period 3 Anomaly — P to S

P:

[ [\mathrm{Ne}],3s^2,3p^3 ]

S:

[ [\mathrm{Ne}],3s^2,3p^4 ]

In P, the three (3p) electrons occupy separate orbitals.

In S, one (3p) orbital contains a pair.

The paired electrons repel one another.

Therefore one electron in S is easier to remove.

So:

[ \boxed{ IE_1(\mathrm S)<IE_1(\mathrm P) } ]

2.17 SPEC: 14 — First Ionisation Energy Down a Group

General trend:

[ \boxed{ IE_1\text{ decreases down a group} } ]

Down a group:

• proton number increases;
• number of occupied electron shells increases;
• outer electron is further from nucleus;
• shielding increases significantly.

The increased distance and shielding outweigh increased nuclear charge.

Therefore attraction to the outer electron is weaker.

Examiner-safe wording

“Down the group, the outer electron occupies a higher shell and experiences greater shielding. These effects outweigh the increased nuclear charge, so less energy is required to remove the outer electron.”

2.18 SPEC: 15(i) — Atomic Emission Spectra as Evidence for Quantum Shells

Atoms can absorb energy, promoting electrons to higher-energy states.

When electrons fall to lower-energy states, they emit photons.

Photon energy is:

[ \boxed{ \Delta E=hf } ]

An atomic emission spectrum consists of discrete lines rather than a continuous range.

DEFINITION — Atomic emission spectrum
A spectrum of discrete wavelengths/frequencies emitted when excited atoms release photons as electrons move between allowed energy levels.

The discrete lines show that electrons can occupy only specific permitted energy levels.

Therefore atomic emission spectra provide evidence for:

[ \boxed{ \text{quantised electron energy levels / quantum shells} } ]

Examiner language

“Only specific photon energies are emitted, showing that electron energies are quantised rather than continuous.”

2.19 SPEC: 15(ii) — Successive Ionisation Energies as Evidence for Shells and Groups

Successive ionisation energies normally increase because each electron is removed from an increasingly positive ion.

A very large jump indicates that the next electron is being removed from a lower principal shell.

Example: magnesium

[ 1s^2,2s^2,2p^6,3s^2 ]

The first two ionisations remove (3s) electrons.

The third removes an electron from the (n=2) shell.

Therefore there is a large jump between:

[ IE_2 ]

and:

[ IE_3 ]

This indicates:

[ 2 ]

outer-shell electrons.

Therefore the element belongs to Group 2.

Deduction rule

If the large jump occurs after removal of (n) outer electrons, the main-group element has approximately (n) outer-shell electrons.

2.20 SPEC: 15(iii) — First Ionisation Energies as Evidence for Subshells

The detailed pattern of first ionisation energies across successive elements is not perfectly smooth.

Drops such as:

• Be → B;
• Mg → Al;

provide evidence that:

[ p ]

subshells are higher in energy than corresponding:

[ s ]

subshells.

Drops such as:

• N → O;
• P → S;

provide evidence for pairing effects in (p) orbitals.

Thus experimental ionisation-energy data support the existence of:

• shells;
• subshells;
• orbitals.

2.21 SPEC: 16 — First Four Quantum Shell Capacities

The theoretical maximum number of electrons in shell (n) is:

[ 2n^2 ]

Thus:

Therefore:

[ \boxed{ 2,\ 8,\ 18,\ 32 } ]

for the first four shells.

Note that actual filling order depends on subshell energies, so (4s) fills before (3d).

2.22 SPEC: 17 — Orbitals

DEFINITION — Orbital
A region within an atom that can hold up to two electrons with opposite spins.

This is the definition Pearson expects.

Maximum occupancy:

[ \boxed{ 2\text{ electrons per orbital} } ]

The two electrons must have opposite spins.

2.23 SPEC: 18 — Shapes of (s)- and (p)-Orbitals

2.23.1 (s)-orbital

An (s)-orbital is:

[ \boxed{ \text{spherical} } ]

around the nucleus.

Diagram description

Draw:

• nucleus at centre;
• spherical boundary around it;
• label (s)-orbital.

A two-dimensional page representation is commonly shown as a circle, but it represents a three-dimensional sphere.

2.23.2 (p)-orbital

A (p)-orbital has a:

[ \boxed{ \text{two-lobed / dumbbell shape} } ]

with the nucleus between the lobes.

There are three (p)-orbitals:

[ p_x,\ p_y,\ p_z ]

oriented at right angles to one another.

Diagram description

Draw:

• two lobes on opposite sides of the nucleus;
• nucleus at the centre/node;
• label as a (p)-orbital.

2.24 SPEC: 19 — Subshell Capacities

(s)-subshell

Contains:

[ 1\text{ orbital} ]

Maximum:

[ \boxed{ 2e^- } ]

(p)-subshell

Contains:

[ 3\text{ orbitals} ]

Maximum:

[ \boxed{ 6e^- } ]

(d)-subshell

Contains:

[ 5\text{ orbitals} ]

Maximum:

[ \boxed{ 10e^- } ]

2.25 SPEC: 20 — Filling Orbitals Singly Before Pairing

Within orbitals of equal energy in the same subshell, electrons occupy orbitals singly before pairing.

For a (p^3) arrangement:

[ [\uparrow]\ [\uparrow]\ [\uparrow] ]

not:

[ [\uparrow\downarrow]\ [\uparrow]\ [\ ] ]

This reduces electron-electron repulsion.

When two electrons occupy the same orbital, they must have opposite spins:

[ \boxed{ \uparrow\downarrow } ]

not:

[ \uparrow\uparrow ]

2.26 SPEC: 21 — Electronic Configurations to (Z=36)

Relevant filling order:

[ \boxed{ 1s\rightarrow2s\rightarrow2p\rightarrow3s\rightarrow3p\rightarrow4s\rightarrow3d\rightarrow4p } ]

2.26.1 Examples

Hydrogen:

[ 1s^1 ]

Carbon:

[ 1s^2,2s^2,2p^2 ]

Oxygen:

[ 1s^2,2s^2,2p^4 ]

Sodium:

[ 1s^2,2s^2,2p^6,3s^1 ]

Chlorine:

[ 1s^2,2s^2,2p^6,3s^2,3p^5 ]

Potassium:

[ [\mathrm{Ar}],4s^1 ]

Calcium:

[ [\mathrm{Ar}],4s^2 ]

Scandium:

[ [\mathrm{Ar}],4s^2,3d^1 ]

Zinc:

[ [\mathrm{Ar}],4s^2,3d^{10} ]

Bromine:

[ [\mathrm{Ar}],4s^2,3d^{10},4p^5 ]

Krypton:

[ [\mathrm{Ar}],4s^2,3d^{10},4p^6 ]

2.26.2 Chromium and copper

Common configurations:

[ \boxed{ \mathrm{Cr}=[\mathrm{Ar}],3d^5,4s^1 } ]

[ \boxed{ \mathrm{Cu}=[\mathrm{Ar}],3d^{10},4s^1 } ]

These differ from simplistic filling predictions and are associated with particularly stable half-filled or fully filled (d)-subshell arrangements.

2.27 Electron-in-Boxes Notation

Each box represents one orbital.

Each arrow represents one electron.

Example:

Carbon:

[ 1s^2,2s^2,2p^2 ]

Boxes:

Oxygen:

[ 1s^2,2s^2,2p^4 ]

Electrons occupy (p)-orbitals singly before pairing.

2.28 Ions — Edexcel Topic 1 Scope

Pearson explicitly limits Topic 1 ion configurations to:

[ \boxed{ s\text{- and }p\text{-block ions only, up to }Z=36 } ]

Examples:

Na:

[ [\mathrm{Ne}]3s^1 ]

[ \mathrm{Na}^+=[\mathrm{Ne}] ]

Mg:

[ [\mathrm{Ne}]3s^2 ]

[ \mathrm{Mg}^{2+}=[\mathrm{Ne}] ]

Cl:

[ [\mathrm{Ne}]3s^23p^5 ]

[ \mathrm{Cl}^-=[\mathrm{Ar}] ]

O:

[ 1s^22s^22p^4 ]

[ \mathrm{O}^{2-}=1s^22s^22p^6 ]

Do not turn Topic 1 into an advanced transition-metal-ion configuration chapter; that content is developed later.

2.29 SPEC: 22 — (s)-, (p)- and (d)-Blocks

The block of an element is determined by the subshell into which its highest-energy / differentiating electron is added.

(s)-block

Outer configuration ends in:

[ s^1 ]

or:

[ s^2 ]

Examples:

• Group 1;
• Group 2.

(p)-block

Outer configuration ends in:

[ p^1\text{ to }p^6 ]

Examples include Groups 13–18, excluding helium by electron configuration.

(d)-block

Electrons are filling a:

[ d ]

subshell.

These elements occupy the central region of the Periodic Table.

2.30 SPEC: 23 — Electronic Configuration Determines Chemical Properties

Chemical reactions involve electron rearrangements.

The outer-electron configuration determines:

• tendency to lose/gain/share electrons;
• typical ionic charges;
• bonding behaviour;
• reactivity patterns.

Elements in the same group have similar outer-shell configurations.

Therefore they show similar chemistry.

Example:

Group 1 elements have:

[ ns^1 ]

They commonly lose one electron to form:

[ M^+ ]

ions.

2.31 SPEC: 24 — Periodicity

DEFINITION — Periodicity
Repeating trends or patterns in physical and chemical properties across successive periods of the Periodic Table.

Periodicity arises because outer-electron configurations repeat in a regular pattern.

For example:

• Group 1 elements repeatedly have (ns^1);
• Group 17 repeatedly have (ns^2np^5);
• noble gases have filled outer-shell configurations.

2.32 SPEC: 25(i) — Melting and Boiling Trends in Periods 2 and 3

Melting and boiling temperatures depend on:

1. structure;
2. bonding;
3. strength/number of forces or bonds that must be overcome.

Do not explain these trends simply in terms of “more electrons” without first identifying the structure.

2.33 Period 3 Structures

A useful Period 3 structure sequence is:

• Na — giant metallic lattice;
• Mg — giant metallic lattice;
• Al — giant metallic lattice;
• Si — giant covalent structure;
• P — simple molecular, mainly (\mathrm{P_4});
• S — simple molecular, mainly (\mathrm{S_8});
• Cl — simple molecular, (\mathrm{Cl_2});
• Ar — monatomic.

2.34 Na → Mg → Al: Metallic Bonding

Metallic bonding is the electrostatic attraction between:

• positive metal ions;
• delocalised electrons.

Across Na → Mg → Al:

• ionic charge increases;
• number of delocalised electrons per atom increases;
• ionic radius generally decreases;
• charge density increases.

Metallic bonding generally becomes stronger.

Therefore melting temperature tends to increase.

Strong explanation

“Al forms (3+) ions and supplies three delocalised electrons per atom, giving stronger electrostatic attraction between the positive ions and delocalised electrons than in sodium.”

2.35 Silicon

Silicon has a giant covalent structure.

Many strong covalent bonds extend throughout the structure.

Melting requires breaking many strong covalent bonds.

Therefore silicon has a very high melting temperature.

Common lost mark

Do not write:

“Intermolecular forces in silicon are strong.”

There are no discrete silicon molecules in giant covalent silicon.

2.36 Phosphorus, Sulfur and Chlorine

These are simple molecular substances in the elemental state relevant to the trend.

Intermolecular attractions are London dispersion forces.

Strength generally increases with:

• number of electrons;
• polarizability;
• molecular size/contact.

Important molecular forms:

[ \mathrm{P_4} ]

[ \mathrm{S_8} ]

[ \mathrm{Cl_2} ]

Sulfur's (\mathrm{S_8}) molecules are larger and contain more electrons than (\mathrm{P_4}) or (\mathrm{Cl_2}), so stronger London forces act between them.

Hence sulfur has a comparatively higher melting/boiling temperature than neighbouring simple molecular elements.

2.37 Argon

Argon is monatomic.

Only weak London forces act between Ar atoms.

Therefore its melting and boiling temperatures are low.

2.38 Period 2 Structure Overview

Period 2 contains:

• Li — metallic;
• Be — metallic;
• B — giant covalent/network character;
• C — giant covalent allotropes such as diamond/graphite;
• N — (\mathrm{N_2}), simple molecular;
• O — (\mathrm{O_2}), simple molecular;
• F — (\mathrm{F_2}), simple molecular;
• Ne — monatomic.

Again, melting/boiling trends are interpreted using:

• structure;
• bonding;
• intermolecular forces.

2.39 SPEC: 25(ii) — First Ionisation Energies in Periods 2 and 3

General trend:

[ \boxed{ IE_1\text{ increases across Periods 2 and 3} } ]

because:

• nuclear charge increases;
• shielding changes relatively little within the period;
• outer electrons are in the same principal shell;
• atomic radius generally decreases.

Key dips:

Period 2:

• Be → B;
• N → O.

Period 3:

• Mg → Al;
• P → S.

These provide evidence for:

• subshell energy differences;
• electron pairing.

2.40 SPEC: 26 — Illustrating Periodicity from Data

You may be given tables or graphs containing:

• atomic radius;
• melting temperature;
• boiling temperature;
• first ionisation energy;
• electronic configuration.

You must identify repeating patterns.

2.40.1 Atomic radius across a period

Atomic radius generally:

[ \boxed{ \text{decreases across a period} } ]

because:

• proton number increases;
• shielding changes relatively little;
• electrons remain in the same principal shell;
• stronger nuclear attraction pulls electron density closer.

2.40.2 Atomic radius down a group

Atomic radius generally:

[ \boxed{ \text{increases down a group} } ]

because:

• an additional principal shell is occupied at each step;
• shielding increases;
• outer electrons lie further from the nucleus.

2.40.3 Linking data to configuration

Example:

A sharp drop in first ionisation energy after a noble gas indicates the start of a new shell.

Example:

[ \mathrm{Ne}\rightarrow\mathrm{Na} ]

The outer electron in Na occupies a new:

[ n=3 ]

shell, further from the nucleus and more shielded.

Therefore first ionisation energy falls sharply.

This repeating noble-gas → Group 1 drop is strong evidence for periodicity.

3. Exact Edexcel Exam Language

These are original examiner-safe phrasings aligned to Pearson command-word expectations.

3.1 Define

Isotope

“Atoms of the same element with the same number of protons but different numbers of neutrons.”

Relative isotopic mass

“The mass of an atom of an isotope relative to one-twelfth of the mass of an atom of carbon-12.”

Relative atomic mass

“The weighted mean mass of an atom of an element relative to one-twelfth of the mass of an atom of carbon-12.”

First ionisation energy

“The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous (1+) ions.”

3.2 State

Give the fact only.

Example:

“A (p)-subshell contains three orbitals and can hold six electrons.”

3.3 Explain

Use:

[ \boxed{ \text{cause}\rightarrow\text{atomic mechanism}\rightarrow\text{observed effect} } ]

Example:

“Shielding increases down the group, reducing the attraction between the nucleus and the outer electron, so less energy is required to remove it.”

3.4 Calculate

For weighted means:

1. multiply mass by abundance;
2. sum products;
3. divide by total abundance;
4. give a sensible numerical answer;
5. no unit for (A_r).

3.5 Deduce

Use the evidence and state the implication.

“The large jump between the second and third ionisation energies shows that two electrons occupy the outer shell, so the element is in Group 2.”

3.6 Explain a trend

Never give only the trend.

Use factors such as:

• nuclear charge;
• shielding;
• distance;
• subshell;
• electron pairing;
• structure and bonding.

3.7 Compare

Explicitly mention both species.

“Both phosphorus and sulfur lose a 3p electron, but sulfur has a paired 3p orbital, so electron-electron repulsion makes an electron easier to remove.”

4. Common Misconceptions & Lost Marks

4.1 Atomic number = protons + neutrons

❌ Incorrect.

[ Z=\text{protons} ]

[ A=\text{protons}+\text{neutrons} ]

4.2 Ion formation changes proton number

❌ Incorrect.

Chemical ion formation involves electron loss/gain.

4.3 Isotopes have different protons

❌ Incorrect.

Same protons, different neutrons.

4.4 Relative mass has units

(A_r), relative isotopic mass and (M_r) are dimensionless ratios.

4.5 Relative formula mass called (M_r) of an ionic molecule

For giant ionic/covalent structures, Pearson explicitly expects relative formula mass terminology.

4.6 Peak position = abundance

❌ Incorrect.

• x-position → (m/z);
• peak height/area → abundance.

4.7 Dividing isotope calculation by 100 regardless of data

Divide by total abundance.

Only use 100 automatically when percentages sum to 100.

4.8 Forgetting mixed-isotope permutations in (\mathrm{Cl_2})

For (35/37):

[ 35-37 ]

and:

[ 37-35 ]

are two equally probable arrangements.

Hence:

[ 9:6:1 ]

rather than:

[ 9:3:1 ]

4.9 Molecular-ion peak confused with base peak

The molecular-ion peak corresponds to intact:

[ M^+ ]

The base peak is simply the most intense peak and need not be (M^+).

4.10 Ionisation-energy definition missing “gaseous”

This can lose the definition mark.

4.11 Second ionisation equation written from neutral atom

❌:

[ X(g)\rightarrow X^{2+}(g)+2e^- ]

✅:

[ X^+(g)\rightarrow X^{2+}(g)+e^- ]

4.12 Across-period trend explained only by “more protons”

Incomplete.

Also state:

• same principal shell;
• similar shielding;
• stronger attraction.

4.13 Mg → Al anomaly explained only by shielding

Key point:

[ 3p ]

electron in Al is higher in energy/more shielded than Mg's (3s) electron.

4.14 P → S anomaly missing pairing

State:

• S has a paired (3p) orbital;
• electron-electron repulsion makes removal easier.

4.15 Down-group trend ignores extra shells

State:

• outer electron further away;
• greater shielding;
• these outweigh increased nuclear charge.

4.16 Atomic emission spectrum described as continuous

Atomic emission spectra have discrete lines.

That discreteness is evidence for quantised energy levels.

4.17 Orbital defined as “where an electron is”

Too vague.

Use:

“A region within an atom that can hold up to two electrons with opposite spins.”

4.18 p subshell has three electrons

❌ It has three orbitals and can hold six electrons.

4.19 d subshell has five electrons

❌ It has five orbitals and can hold ten electrons.

4.20 Pairing before singly filling

Incorrect for equal-energy orbitals.

Electrons fill singly first.

4.21 (p)-orbital drawn spherical

(s) is spherical.

(p) is dumbbell/two-lobed.

4.22 Treating (4s) and (3d) ordering as purely shell-number order

For neutral atoms in this region:

[ 4s ]

fills before:

[ 3d ]

4.23 Adding advanced transition-metal ion configurations to Topic 1

Pearson Topic 1 specifies ions only for:

[ s\text{- and }p\text{-block ions} ]

up to (Z=36).

4.24 Silicon melting described using intermolecular forces

Silicon has a giant covalent structure.

Strong covalent bonds must be broken.

4.25 Molecular substances described as breaking covalent bonds when boiling

Boiling simple molecular substances overcomes intermolecular forces, not covalent bonds within molecules.

5. Worked Exam-Style Questions

Question 1 — Atomic Structure [4 marks]

For:

[ {}^{56}_{26}\mathrm{Fe}^{3+} ]

determine the numbers of protons, neutrons and electrons.

Model answer

Protons:

[ \boxed{26} ]

Neutrons:

[ 56-26=30 ]

[ \boxed{30} ]

Electrons:

[ 26-3=23 ]

[ \boxed{23} ]

Indicative marks:

• proton number correct;
• neutron method;
• neutron answer;
• electron answer.

Question 2 — Isotopic Abundance [6 marks]

An element has isotopes of relative isotopic mass 35 and 37.

Its relative atomic mass is 35.40.

Calculate the percentage abundance of each isotope.

Model answer

Let fraction of isotope 35 be:

[ x ]

Then isotope 37 fraction is:

[ 1-x ]

[ 35x+37(1-x)=35.40 ]

[ 35x+37-37x=35.40 ]

[ -2x=-1.60 ]

[ x=0.800 ]

Therefore:

[ \boxed{ 80.0%\text{ isotope 35} } ]

and:

[ \boxed{ 20.0%\text{ isotope 37} } ]

Question 3 — Diatomic Mass Spectrum [6 marks]

An element (X) has isotopes:

• (^{10}X): 20%;
• (^{11}X): 80%.

Predict the molecular-ion peaks and relative probabilities for (X_2^+).

Model answer

Possible combinations:

[ 10+10=20 ]

[ 10+11=21 ]

[ 11+11=22 ]

Probabilities:

[ P(20)=0.20^2=0.04 ]

[ P(21)=2(0.20)(0.80)=0.32 ]

[ P(22)=0.80^2=0.64 ]

Ratio:

[ 0.04:0.32:0.64 ]

divide by 0.04:

[ \boxed{ 1:8:16 } ]

Peaks:

[ \boxed{ m/z=20,\ 21,\ 22 } ]

with relative heights:

[ \boxed{ 1:8:16 } ]

Question 4 — Successive Ionisation Energies [6 marks]

Successive ionisation energies of an element are:

Deduce the group of the element and explain.

Model answer

The large jump occurs between:

[ IE_3 ]

and:

[ IE_4 ]

Therefore three electrons are removed relatively easily before an inner-shell electron must be removed.

The atom has three outer-shell electrons.

Hence it is in:

[ \boxed{ \text{Group 13 / Group 3 in older main-group numbering} } ]

Within modern Periodic Table group numbering, the appropriate group is:

[ \boxed{ 13 } ]

Explanation:

“The fourth electron is in an inner shell, closer to the nucleus and less shielded, so much more energy is required to remove it.”

Question 5 — Period 3 Ionisation Energies [6 marks]

Explain why first ionisation energy generally increases from Na to Ar and why the value for sulfur is lower than phosphorus.

Model answer

“Across Period 3, proton number and nuclear charge increase. Electrons are added to the same principal shell, so shielding changes relatively little. The outer electron is therefore increasingly strongly attracted to the nucleus and first ionisation energy generally increases. Phosphorus has three 3p electrons occupying separate orbitals, whereas sulfur has four 3p electrons so one orbital contains a pair. Repulsion between the paired electrons makes one sulfur electron easier to remove, so sulfur has a lower first ionisation energy than phosphorus.”

Question 6 — Period 3 Melting Trend [8 marks]

Explain the major changes in melting temperature from Na to Ar.

Full-mark model answer

“Na, Mg and Al form giant metallic lattices. Across these elements the metal-ion charge and number of delocalised electrons increase and ionic radius generally decreases, so electrostatic attraction between the positive ions and delocalised electrons becomes stronger. Silicon has a giant covalent structure and many strong covalent bonds must be broken to melt it, giving a very high melting temperature. Phosphorus, sulfur and chlorine are simple molecular substances, so melting mainly overcomes London forces between molecules rather than covalent bonds within molecules. Sulfur exists mainly as larger (S_8) molecules and therefore has stronger London forces than (P_4) or (Cl_2). Argon is monatomic and has only weak London forces between atoms, giving a very low melting temperature.”

Question 7 — Emission Spectra and Quantum Shells [4 marks]

Explain how an atomic emission spectrum supports the existence of quantum shells.

Model answer

“The emission spectrum contains discrete lines rather than a continuous range. Each line corresponds to a photon of a specific energy released when an electron moves between two allowed energy levels. Because only particular photon energies are observed, the electron energies must be quantised.”

Indicative marks:

• discrete lines;
• photons emitted in transitions;
• specific photon energies;
• allowed/quantised energy levels.

6. Links to Other Topics

6.1 Topic 2 — Bonding and Structure

Electron configuration directly explains:

• ion formation;
• ionic bonding;
• covalent bonding;
• valence electrons;
• Periodic Table block position.

6.2 Topic 4 — Inorganic Chemistry and the Periodic Table

Topic 1 provides the basis for:

• Group 2 reactivity;
• halogen trends;
• ionic radii;
• periodic behaviour.

6.3 Topic 5 — Formulae, Equations and Amounts of Substance

Relative atomic/molecular/formula masses feed directly into:

[ n=\frac{m}{M} ]

and empirical/molecular-formula calculations.

6.4 Topic 7 — Modern Analytical Techniques I

The molecular-ion idea develops into fuller mass-spectral interpretation of organic compounds.

6.5 Topic 15 — Transition Metals

(d)-subshell ideas and (d)-block classification become central to:

• variable oxidation states;
• coloured ions;
• complex formation;
• catalysis.

6.6 Topic 19 — Modern Analytical Techniques II

Mass spectrometry combines with:

• NMR;
• IR;
• chromatography;

for structural elucidation.

7. Summary Notes

7.1 Subatomic particles

7.2 Atomic and mass number

[ \boxed{ Z=\text{number of protons} } ]

[ \boxed{ A=p+n } ]

[ \boxed{ n=A-Z } ]

Neutral atom:

[ e=p ]

Positive ion: electrons lost.

Negative ion: electrons gained.

7.3 Isotopes

Same number of protons, different numbers of neutrons.

Same chemistry because neutral isotopes have the same electron configuration.

7.4 Relative isotopic mass

Mass of an atom of an isotope relative to (1/12) of the mass of a carbon-12 atom.

No unit.

7.5 Relative atomic mass

Weighted mean mass of an atom of an element relative to (1/12) of carbon-12.

[ \boxed{ A_r= \frac{\sum(mass\times abundance)} {\sum abundance} } ]

No unit.

7.6 Relative molecular/formula mass

Discrete molecules:

[ M_r=\sum A_r ]

Giant structures:

use relative formula mass.

No unit.

7.7 Mass spectra

• x-axis → (m/z);
• y-axis → relative intensity/abundance;
• peak position → mass-to-charge ratio;
• peak height/area → abundance.

Molecular ion:

[ \boxed{ M^+ } ]

For singly charged (M^+):

[ \boxed{ m/z=M_r } ]

7.8 Diatomic isotope pattern

For chlorine approximately (75:25):

[ {}^{35}\mathrm{Cl}:{}^{37}\mathrm{Cl}=3:1 ]

[ \mathrm{Cl_2} ]

peaks:

[ 70,\ 72,\ 74 ]

ratio:

[ \boxed{ 9:6:1 } ]

7.9 First ionisation energy

Energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous (1+) ions.

[ \boxed{ X(g)\rightarrow X^+(g)+e^- } ]

Second:

[ \boxed{ X^+(g)\rightarrow X^{2+}(g)+e^- } ]

7.10 Factors affecting ionisation energy

• proton number / nuclear charge;
• shielding;
• distance/shell;
• subshell;
• electron pairing.

Across a period:

[ \boxed{ IE_1\text{ generally increases} } ]

Down a group:

[ \boxed{ IE_1\text{ generally decreases} } ]

7.11 Evidence for shells and subshells

Emission spectrum:

[ \boxed{ \text{discrete lines}\Rightarrow\text{quantised energy levels} } ]

Successive IE:

[ \boxed{ \text{large jump}\Rightarrow\text{new inner shell} } ]

First IE irregularities:

[ \boxed{ \text{evidence for subshells and pairing} } ]

7.12 Shell capacities

First four quantum shells:

[ \boxed{ 2,\ 8,\ 18,\ 32 } ]

7.13 Orbital

Region within an atom that can hold up to two electrons with opposite spins.

(s)-orbital:

[ \boxed{ \text{spherical} } ]

(p)-orbital:

[ \boxed{ \text{dumbbell / two-lobed} } ]

7.14 Subshell capacities

[ s:2 ]

[ p:6 ]

[ d:10 ]

Electrons fill equal-energy orbitals singly before pairing.

7.15 Filling order to (Z=36)

[ \boxed{ 1s,2s,2p,3s,3p,4s,3d,4p } ]

Pearson Topic 1 ion configurations:

[ \boxed{ s\text{- and }p\text{-block ions only} } ]

7.16 Blocks

(s)-block → differentiating electron enters (s).

(p)-block → enters (p).

(d)-block → enters (d).

7.17 Periodicity

Repeating pattern of properties across successive periods.

Caused by repeating outer-electron configurations.

Atomic radius generally decreases across a period.

First ionisation energy generally increases across a period.

7.18 Period 3 structures

• Na — metallic;
• Mg — metallic;
• Al — metallic;
• Si — giant covalent;
• P — (P_4), simple molecular;
• S — (S_8), simple molecular;
• Cl — (Cl_2), simple molecular;
• Ar — monatomic.

Melting/boiling explanations must use:

[ \boxed{ \text{structure + bonding/forces} } ]