AQA A-Level Chemistry: Atomic Structure and Bonding -- The Complete Guide

Atomic structure and bonding sit at the very foundation of A-Level Chemistry. Every reaction you study, every property you explain, and every calculation you perform ultimately traces back to the arrangement of electrons in atoms and the forces that hold atoms together. These topics appear early in the AQA specification for good reason -- they provide the conceptual framework on which the rest of the course is built.

At GCSE you learned a simplified model of the atom and a basic treatment of bonding types. A-Level Chemistry takes this much further. You will encounter sub-shells and orbitals, quantitative treatments of ionisation energy, mass spectrometry as an analytical tool, and a far more nuanced picture of how and why atoms bond. This guide covers the key content across both topics, explains the connections between them, and highlights the areas that examiners test most frequently.

Atomic Structure

The Development of the Atomic Model

The AQA specification expects you to understand how the model of the atom evolved through experimental evidence. Key milestones include:

• Dalton's atomic theory -- atoms are indivisible, identical within an element, and combine in fixed ratios
• Thomson's plum pudding model -- the discovery of the electron showed atoms contain negatively charged particles embedded in a positive "pudding"
• Rutherford's nuclear model -- alpha particle scattering demonstrated that most of the atom is empty space, with mass concentrated in a tiny, dense, positively charged nucleus
• Bohr's model -- electrons orbit the nucleus in fixed energy levels, explaining line spectra
• The quantum mechanical model -- electrons occupy orbitals (regions of probability), described by quantum numbers

You will not be asked to derive quantum mechanics, but you must understand that electrons exist in defined energy levels, sub-levels, and orbitals, and that this model explains the properties of elements across the periodic table.

Sub-shells and Orbitals

At A-Level, the simple electron shells (1, 2, 3...) are divided into sub-shells, each containing a specific number of orbitals:

• s sub-shell -- one orbital, holds a maximum of 2 electrons
• p sub-shell -- three orbitals, holds a maximum of 6 electrons
• d sub-shell -- five orbitals, holds a maximum of 10 electrons
• f sub-shell -- seven orbitals, holds a maximum of 14 electrons

Each orbital holds a maximum of two electrons, which must have opposite spins (the Pauli exclusion principle). Orbitals within the same sub-shell are filled singly before pairing begins (Hund's rule). Electrons fill sub-shells in order of increasing energy, following the Aufbau principle. For the first four periods, the filling order is:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p

Note that the 4s sub-shell fills before the 3d sub-shell because it has a lower energy. However, when transition metal atoms form ions, the 4s electrons are lost first -- this is a common source of confusion and a frequent exam question.

Writing Electron Configurations

You need to write electron configurations using sub-shell notation. For example:

• Sodium (Na, Z = 11): 1s2 2s2 2p6 3s1
• Iron (Fe, Z = 26): 1s2 2s2 2p6 3s2 3p6 3d6 4s2
• Iron(III) ion (Fe3+): 1s2 2s2 2p6 3s2 3p6 3d5

There are two important exceptions to learn: chromium (Cr) has the configuration [Ar] 3d5 4s1 rather than [Ar] 3d4 4s2, and copper (Cu) has [Ar] 3d10 4s1 rather than [Ar] 3d9 4s2. In both cases, the stability of a half-filled or fully filled d sub-shell provides the driving force for the unexpected arrangement.

Ionisation Energies

The first ionisation energy is the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. It is always endothermic and is measured in kJ mol-1.

Ionisation energies provide direct evidence for electron shells and sub-shells. The key trends are:

Across a period (left to right): First ionisation energy generally increases because nuclear charge increases while shielding remains roughly constant, so the effective nuclear charge experienced by the outer electron increases. However, there are two important drops within each period:

• Group 2 to Group 3 -- the Group 3 element has its outer electron in a p sub-shell, which is slightly higher in energy and further from the nucleus than the s sub-shell. This electron is easier to remove despite the increased nuclear charge. For example, the first ionisation energy of aluminium is lower than that of magnesium.
• Group 5 to Group 6 -- the Group 6 element has a pair of electrons in one of its p orbitals. The repulsion between these paired electrons makes one of them easier to remove. For example, the first ionisation energy of oxygen is lower than that of nitrogen.

Down a group: First ionisation energy decreases because additional electron shells increase the distance from the nucleus and the shielding effect, outweighing the increase in nuclear charge.

Successive ionisation energies of a single element reveal the shell structure. Large jumps in ionisation energy occur when an electron is removed from a shell closer to the nucleus. By analysing a graph of successive ionisation energies, you can deduce how many electrons are in each shell and therefore identify the element or confirm its group.

Mass Spectrometry

Mass spectrometry is a key analytical technique at A-Level. For atomic structure, it allows the determination of relative atomic mass and the identification of isotopes.

The basic process involves four stages:

1. Ionisation -- atoms are bombarded with high-energy electrons to knock out an electron, forming positive ions (e.g. by electron impact ionisation)
2. Acceleration -- the positive ions are accelerated by an electric field so that they all have the same kinetic energy
3. Deflection -- the ions pass through a magnetic field that deflects lighter ions more than heavier ones (ions are separated by their mass-to-charge ratio, m/z)
4. Detection -- the ions reach a detector that records the m/z ratio and the relative abundance of each ion

From a mass spectrum, you can:

• Identify the isotopes of an element from the m/z values of the peaks

• Use the relative abundances to calculate the relative atomic mass using the weighted mean formula:

  Ar = sum of (isotopic mass x relative abundance) / total relative abundance

For example, chlorine shows peaks at m/z = 35 and m/z = 37 with relative abundances of approximately 75% and 25%, giving a relative atomic mass of approximately 35.5.

At A-Level, mass spectrometry is also applied to molecular compounds (covered in organic chemistry), where fragmentation patterns help identify molecular structures.

Bonding

Ionic Bonding

Ionic bonding occurs between metals and non-metals. Metals lose electrons to form positive ions (cations), and non-metals gain electrons to form negative ions (anions). The resulting electrostatic attraction between oppositely charged ions is the ionic bond.

Key points for the AQA specification:

• Ionic compounds form giant ionic lattices, not discrete molecules. Each ion is surrounded by ions of the opposite charge in a regular three-dimensional arrangement.
• The physical properties of ionic compounds -- high melting points, brittleness, conduction only when molten or dissolved -- are all explained by the strength of the electrostatic attractions and the structure of the lattice.
• Ionic bonding is strongest between small, highly charged ions. Lattice energy increases as ionic charge increases and as ionic radius decreases.
• You should be able to draw dot-and-cross diagrams showing the transfer of electrons and the resulting electron configurations of the ions formed.

Covalent Bonding

Covalent bonding involves the sharing of one or more pairs of electrons between two non-metal atoms. The shared pair of electrons is attracted to both nuclei, holding the atoms together.

At A-Level, you need to distinguish between several types:

• Single covalent bonds -- one shared pair (e.g. H-H, Cl-Cl)
• Double covalent bonds -- two shared pairs (e.g. O=O, C=O)
• Triple covalent bonds -- three shared pairs (e.g. the triple bond in nitrogen gas, N2)
• Dative (coordinate) covalent bonds -- both electrons in the shared pair come from one atom. This occurs in the ammonium ion (NH4+), where the nitrogen lone pair forms a bond with H+, and in many transition metal complexes.

Covalent bonds can exist in simple molecular structures (e.g. water, carbon dioxide) or in giant covalent structures (e.g. diamond, silicon dioxide, graphite). The distinction is crucial for explaining physical properties:

• Simple molecular substances have low melting and boiling points because the weak intermolecular forces (not the covalent bonds) are broken during changes of state.
• Giant covalent substances have very high melting points because many strong covalent bonds must be broken.

Metallic Bonding

Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons. The outer electrons of metal atoms are released into a shared "cloud" that extends throughout the structure.

This model explains the characteristic properties of metals:

• High melting points -- strong electrostatic attraction between positive ions and the electron sea requires significant energy to overcome
• Electrical conductivity -- delocalised electrons are free to move through the structure when a potential difference is applied
• Malleability and ductility -- layers of ions can slide over each other without disrupting the bonding, because the electron sea adjusts to the new arrangement
• Thermal conductivity -- delocalised electrons transfer kinetic energy rapidly through the structure

The strength of metallic bonding increases with the number of delocalised electrons and with decreasing ionic radius. This explains why metals such as aluminium (3 delocalised electrons per atom, small ionic radius) have higher melting points than metals such as sodium (1 delocalised electron per atom, larger ionic radius).

Electronegativity and Bond Polarity

Electronegativity is a measure of the ability of an atom to attract the bonding pair of electrons in a covalent bond. The Pauling scale is used to assign electronegativity values, with fluorine being the most electronegative element (4.0) and francium the least (0.7).

Electronegativity trends:

• Across a period: electronegativity increases because nuclear charge increases while atomic radius decreases, so the nucleus exerts a stronger pull on bonding electrons
• Down a group: electronegativity decreases because atomic radius increases and additional electron shells provide greater shielding, reducing the pull on bonding electrons

When two atoms with different electronegativities form a covalent bond, the bonding pair is drawn more towards the more electronegative atom, creating a polar bond. The more electronegative atom carries a partial negative charge (delta-), and the less electronegative atom carries a partial positive charge (delta+).

If the electronegativity difference is very large, the bond is considered ionic rather than covalent. In reality, most bonds fall on a spectrum between purely covalent and purely ionic -- the distinction is not absolute.

A molecule with polar bonds may or may not be a polar molecule overall. Molecular polarity depends on the shape of the molecule. Symmetrical molecules such as carbon dioxide (O=C=O, linear) and tetrachloromethane (CCl4, tetrahedral) contain polar bonds but are non-polar overall because the individual bond dipoles cancel out. Asymmetrical molecules such as water (H-O-H, bent) and trichloromethane (CHCl3) are polar because the bond dipoles do not cancel.

Shapes of Molecules and Ions

The shapes of molecules and ions are predicted using Valence Shell Electron Pair Repulsion (VSEPR) theory. The principle is straightforward: electron pairs around a central atom repel each other and arrange themselves as far apart as possible to minimise repulsion.

There are two types of electron pair:

• Bonding pairs -- shared between two atoms
• Lone pairs -- not shared, held closer to the central atom

Lone pairs repel more strongly than bonding pairs (because they are held closer to the nucleus and occupy more space), so they compress the bond angles slightly. The order of repulsion is:

lone pair--lone pair > lone pair--bonding pair > bonding pair--bonding pair

The key shapes you must know:

Bonding pairs	Lone pairs	Shape	Bond angle	Example
2	0	Linear	180 degrees	CO2, BeCl2
3	0	Trigonal planar	120 degrees	BF3, AlCl3
3	1	Pyramidal	107 degrees	NH3, PCl3
2	2	Bent (non-linear)	104.5 degrees	H2O
4	0	Tetrahedral	109.5 degrees	CH4, NH4+
5	0	Trigonal bipyramidal	90 and 120 degrees	PCl5
6	0	Octahedral	90 degrees	SF6

You should be able to draw these shapes, state the bond angles, and explain why lone pairs reduce the bond angle from the ideal value. For ions such as NH4+ and H3O+, you should also show the dative covalent bond and apply VSEPR theory in the same way.

Intermolecular Forces

Intermolecular forces are the weak forces of attraction between molecules. They determine the physical properties of simple molecular substances -- boiling point, solubility, and viscosity. There are three types, in order of increasing strength:

London (dispersion) forces are present in all molecules, whether polar or non-polar. They arise from temporary fluctuations in the electron cloud that create instantaneous dipoles, which then induce dipoles in neighbouring molecules. The strength of London forces increases with:

• The number of electrons in the molecule (larger molecules have stronger London forces)
• The surface area of the molecule (long-chain molecules have stronger London forces than branched isomers because of greater contact area)

London forces are the only intermolecular forces present in non-polar molecules such as alkanes, noble gases, and symmetrical molecules like methane.

Permanent dipole--dipole forces exist between molecules that have a permanent dipole (i.e. polar molecules). The partial positive end of one molecule is attracted to the partial negative end of another. These forces are stronger than London forces alone, which is why polar molecules often have higher boiling points than non-polar molecules of similar size.

Hydrogen bonding is a special case of permanent dipole--dipole interaction that occurs when hydrogen is bonded directly to one of the three most electronegative elements: fluorine, oxygen, or nitrogen. The hydrogen atom carries a significant partial positive charge and is small enough to get very close to a lone pair on an F, O, or N atom of a neighbouring molecule, forming a strong directional interaction.

Hydrogen bonding explains several anomalous properties:

• The unusually high boiling points of water, hydrogen fluoride, and ammonia compared to other hydrides in their respective groups
• The relatively high surface tension and viscosity of water
• The low density of ice -- in ice, hydrogen bonds hold water molecules in an open hexagonal lattice that is less dense than liquid water, which is why ice floats
• The solubility of alcohols and carboxylic acids in water -- these molecules can form hydrogen bonds with water molecules

For the AQA exam, you should be able to:

• Identify which type of intermolecular force is present in a given substance
• Explain trends in boiling points within a homologous series (increasing London forces with increasing chain length)
• Explain the anomalously high boiling point of water and similar compounds
• Draw diagrams showing hydrogen bonds between molecules, including the lone pairs involved and the delta charges

Connecting Atomic Structure to Bonding

One of the strengths of A-Level Chemistry is the way topics interconnect. The electron configuration of an atom determines its bonding behaviour:

• Elements with one, two, or three outer electrons tend to lose them and form ionic bonds with non-metals
• Elements with high electronegativities tend to share electrons and form covalent bonds
• The number of bonding pairs and lone pairs (determined by electron configuration) dictates molecular shape
• Electronegativity differences (which arise from nuclear charge and shielding) determine bond polarity, which in turn determines intermolecular forces, which in turn determine physical properties

When answering exam questions, examiners reward students who make these connections explicit. If asked to explain why water has a higher boiling point than hydrogen sulfide, the best answers will link electronegativity to bond polarity, bond polarity to hydrogen bonding, and hydrogen bonding to the energy required for vaporisation.

Exam Tips for These Topics

• Ionisation energy questions almost always require reference to nuclear charge, shielding, and distance from the nucleus. If you leave any of these out when explaining a trend, you will lose marks.
• Electron configuration questions may ask you to write configurations for atoms or ions. Remember that transition metal ions lose 4s electrons before 3d electrons.
• Mass spectrometry calculations require careful use of the weighted mean formula. Show your working clearly and give your final answer to an appropriate number of significant figures.
• Bonding questions frequently ask you to explain physical properties in terms of structure and bonding. Always be specific about what is being broken -- do not write "bonds are broken" without saying whether you mean ionic bonds, covalent bonds, or intermolecular forces.
• Shape questions require both the name of the shape and the bond angle. If lone pairs are involved, you must explain how they reduce the bond angle from the ideal value.
• Hydrogen bonding questions require you to name the electronegative atom involved, show the lone pair, and mark the delta charges. A labelled diagram is often the clearest way to answer.

Prepare with LearningBro

These topics form the essential groundwork for A-Level Chemistry. Practise applying your understanding with targeted questions on LearningBro:

• AQA A-Level Chemistry: Atomic Structure in Depth -- thorough coverage of electron configurations, ionisation energies, and mass spectrometry
• AQA A-Level Chemistry: Bonding in Depth -- detailed practice on ionic, covalent, and metallic bonding, along with electronegativity and bond polarity
• AQA A-Level Chemistry: Structure and Bonding -- questions linking structure type to physical properties, molecular shapes, and intermolecular forces

Build confidence in these foundational topics and the rest of A-Level Chemistry becomes significantly more manageable.