Q: What does H2 Chemistry Notes: Topic 1 - Atomic Structure cover? A: Review quantum numbers, electron configurations, ionisation trends, and spectroscopy evidence for the Core Idea 1 Atomic Structure topic in the 2026 H2 Chemistry syllabus.
TL;DR Atomic structure explains where electrons are, how they are arranged, and why spectra and ionisation energy data support that model.
Electrons sit in shells, subshells, and orbitals: Identify the electron configuration.
Trends depend on nuclear charge, distance, and shielding: Use those three ideas in every ionisation-energy explanation.
Spectroscopy turns electron structure into evidence: Link peaks or lines to subshells and energy changes.
Concrete example: Aluminium has a lower first ionisation energy than magnesium because its outer electron is in a higher-energy 3p subshell, so it is easier to remove than magnesium's 3s electron.
Route map: choose the atomic-structure tool first
Before calculating or explaining, decide what kind of evidence the question gives you. Most mistakes in this topic come from using a trend explanation when the question actually needs configuration, or using configuration when the graph is showing shell changes.
Question cue
First tool to use
What your answer must name
Common trap
"Write the configuration" or "which electron is removed"
Fill subshells, then handle ions.
Aufbau order, Hund's rule, Pauli exclusion, and 4s before 3d removal for transition metal cations.
Removing 3d before 4s because 3d appears earlier in the written configuration.
Nuclear charge, distance from nucleus, shielding, and subshell energy.
Saying only "more protons" without saying whether shielding or subshell energy changes.
"Successive ionisation energy graph"
Find the large jump first.
The jump shows a move to an inner shell with lower n and stronger attraction.
Treating every small increase as a new shell.
"PES spectrum"
Match peak position and peak size separately.
Binding energy identifies subshell depth; peak area or height gives electron count.
Reading a taller peak as higher binding energy instead of more electrons.
"Line spectrum"
Link energy changes to electron transitions.
Electrons absorb or emit photons when moving between quantised energy levels.
Describing spectral lines as continuous colours instead of discrete energy changes.
Atomic structure underpins every other area of H2 Chemistry. You need to know how quantum numbers define electron positions, how orbitals fill, and how evidence from spectroscopy backs the quantum model. These notes focus on the ideas that repeatedly appear in structured and data-based questions.
Status: SEAB's current H2 Chemistry (9476) syllabus PDF is labelled for 2026, and the current Chemistry Data Booklet is labelled 8873/9476/9813 for use from 2026 in non-practical papers. Core Idea 1: Atomic Structure is assessed across Papers 1-3 (with data booklet support where relevant). For the full topic map and paper weightings, see our H2 Chemistry Syllabus 2026-27 overview.
Quick revision box
What this topic tests: Quantum numbers, electron configurations, ionisation trends, spectroscopy evidence.
Top mistakes to avoid: Mixing up subshell filling order; weak PES justification; missing shielding language in trend explanations.
20-minute sprint plan: 5 min quantum numbers + filling order; 10 min trend explanation drill; 5 min PES interpretation.
1 Quantum Numbers at a Glance
Symbol
Name
Typical values
What it tells you
n
Principal quantum number
1,2,3,...
Shell index; larger n means electrons are, on average, further from the nucleus and higher in energy.
l
Azimuthal (angular momentum) quantum number
0 to n−1
Subshell shape: l=0(s),1(p),2(d).
ml
Magnetic quantum number
−l to +l
Orientation of the orbital in space.
ms
Spin quantum number
+1/2 or −1/2
Electron spin; explains paired and unpaired electrons.
Key reminders
The Pauli exclusion principle: no two electrons share the same four quantum numbers.
Hund's rule: electrons occupy degenerate orbitals singly with parallel spins before pairing.
Aufbau order for multi-electron atoms: 1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p→6s. Use this sequence unless exceptions (Cr,Cu) are specified.
2 Electron Configurations
Ground-state examples
Magnesium: 1s22s22p63s2.
Chlorine: [Ne]3s23p5 (one unpaired electron explains high electron affinity).
Copper: [Ar]3d104s1 (fully filled 3d¹⁰ subshell stabilises the atom).
Cations of transition metals
Remove 4s electrons before 3d electrons even though 4s fills first.
Reason: in the ion, the 3d orbitals drop below 4s in energy because of electron-electron repulsion changes.
Configuration anomalies
Chromium [Ar]3d54s1 and copper [Ar]3d104s1
3 Ionisation Energy Trends
3.1 Successive ionisation energies
Large jumps indicate removal from an inner shell with lower n.
Link explanations to effective nuclear charge: higher proton number and lower shielding raise IE.
Successive ionisation jump checkpoint
When a graph shows successive ionisation energies, do not start by naming the element. First count how many electrons are removed before the large jump.
Graph clue
What it means
Answer move
Big jump after IE1
One outer electron was removed before the jump.
The element is in Group 1 because the second electron is from an inner shell.
Big jump after IE2
Two outer electrons were removed before the jump.
The element is in Group 2 because the third electron is much closer to the nucleus.
Big jump after IE3
Three outer electrons were removed before the jump.
The element is in Group 13 because three valence electrons are removed first.
No single dramatic jump in the first few removals
The listed electrons may still be from the same shell.
Compare the size of increases, then wait for the clear shell-change jump before assigning group.
Misconception check: the large jump does not mean that electron is the valence electron. It means the previous electron was the last valence electron removed, so the next removal starts an inner shell.
3.2 Factors affecting first ionisation energy
Nuclear charge: more protons mean stronger attraction.
Atomic radius: larger atoms have valence electrons further out and more weakly held.
Shielding: inner electrons reduce the pull felt by valence electrons.
Subshell energy: electrons removed from p subshells are at higher energy than those from s subshells in the same shell.
3.3 Evidence from photoelectron spectroscopy (PES)
Each peak represents electrons from a subshell; peak height is proportional to electron count.
Use PES data to justify electron configurations or identify anomalies in question stem.
PES reading checkpoint
Read a PES spectrum in two passes. First use peak position to identify how tightly the electrons are held, then use relative peak size to count how many electrons are in that subshell.
Spectrum feature
What it tells you
How to use it
Common trap
Peak at higher binding energy
Electrons are held more strongly and are usually closer to the nucleus.
Assign inner-shell subshells before outer-shell subshells.
Treating the rightmost or tallest peak as the outer electron without checking the axis.
Peak at lower binding energy
Electrons are easier to remove and are usually in a higher-energy outer subshell.
Match these peaks to valence subshells such as 3s or 3p.
Calling the lowest-binding-energy peak the most stable electrons.
Taller or larger-area peak
More electrons contribute to that peak.
Use peak size to infer counts such as 2 for an s subshell or 6 for a full p subshell.
Reading peak height as stronger binding.
Two nearby valence peaks
Two different subshells in the same shell may be present.
Separate, for example, 3s and 3p electrons before writing the configuration.
Combining them into one shell count and losing subshell evidence.
Worked check: a PES spectrum with one small low-binding-energy peak and one larger nearby peak in the valence region suggests different valence subshells, not just "two outer electrons". Identify the lower-binding-energy valence peak by position, then use relative area to decide how many electrons it represents.
Misconception check: position answers "how strongly held"; peak size answers "how many electrons". Do not swap those two readings.
4 Worked Example
Question: Explain why aluminium has a lower first ionisation energy than magnesium, while silicon's first ionisation energy is higher than aluminium's.
Answer outline:
Aluminium vs magnesium: aluminium's outer electron is in the 3p subshell, which is higher in energy and more shielded than magnesium's 3s electron. Reduced effective nuclear charge outweighs the extra proton, so IE₁ decreases (Data Booklet: Mg 736, Al 577 kJ⋅mol−1).
Silicon vs aluminium: silicon still removes a 3p electron, but the increased proton number raises effective nuclear charge more than the slight increase in shielding, so IE₁ rises (Data Booklet: Si 786 kJ⋅mol−1).
State your reasoning in terms of effective nuclear charge, subshell type, and energy level to earn explanation marks.
5 Exam Skills Check
Paper 1 MCQ: practise assigning quantum numbers to electrons quickly. Multi-statement questions often mix configurations with ionisation trends.
Paper 2 structured: expect graph interpretation (successive ionisation data) or sketching PES spectra. Quote units clearly (kJ⋅mol−1).
Paper 3 free-response: link electron configurations to bonding and periodic trends; you may have to justify shapes or polarities from configuration data.
Paper 4 practical: atomic-structure ideas show up indirectly through evidence-based thinking (clear variables, units, and calibration), but the data booklet is not used for the practical paper.
6 Misconceptions to Fix
Shell vs subshell confusion: always clarify whether you mean n or l.
Forgetting shielding: effective nuclear charge explanations must include both proton number and shielding.
Removing 3d electrons first: correct sequence is 4s before 3d when forming cations.
Misapplying Hund's rule: it applies only within degenerate orbitals, not across subshells of different energy.
7 Quick Practice Prompts
Write the electronic configuration for FeX2+ and explain the order in which electrons are removed.
Interpret a PES spectrum for phosphorus by labelling peak positions and electron counts.
Compare IE₁ values of oxygen and nitrogen; include half-filled subshell stability in your reasoning.
Check your answers against lecture notes or textbook examples and refine the wording until each explanation references effective nuclear charge, shielding, and energy level.
Common exam mistakes
Wrong removal order for transition metal cations: Students write FeX2+ by removing 3d electrons first. Always remove 4s electrons before 3d when forming cations.
Omitting shielding in trend explanations: Stating "nuclear charge increases" without mentioning shielding gives an incomplete explanation and loses marks; always pair proton number with shielding and effective nuclear charge.
Confusing period vs group ionisation energy trends: IE₁ generally increases across a period and decreases down a group - mixing up the direction is a common slip under time pressure.
Missing the subshell dip anomalies: Forgetting that IE₁ drops from Be to B (2s → 2p) and from N to O (half-filled 2p stability) leads to wrong predictions in MCQ sequences.
Misidentifying PES peaks: Assuming peak height represents energy rather than electron count; the area/height of a PES peak is proportional to the number of electrons in that subshell.
Applying Hund's rule across subshells: Hund's rule only governs filling within degenerate orbitals of the same subshell, not across different subshells.
Forgetting to justify Cr and Cu anomalies: Simply stating the configuration without explaining extra stability from half-filled or fully filled d subshells will not earn explanation marks.
Frequently asked questions
Is Atomic Structure tested in Paper 1? Yes. Paper 1 MCQs regularly include quantum number assignments, electron configurations of ions, and ionisation energy comparisons. Multi-statement questions often combine configuration knowledge with periodic trend reasoning.
Do I need to memorise ionisation energy values? No. The SEAB Chemistry Data Booklet provides ionisation energies and other atomic data. You are expected to interpret and explain trends using the values given, not reproduce them from memory.
What is the most common question type for this topic in Paper 2? Interpret or sketch a photoelectron spectrum, and explain successive ionisation energy graphs. Both require you to link peak positions or jump patterns to electron configurations and effective nuclear charge.
How should I explain an ionisation energy anomaly in an exam answer? Name the specific subshells involved (e.g. 2s vs 2p), state which is higher in energy or why it is more shielded, then relate this to why less energy is needed to remove the electron - typically two to three concise sentences earns full marks.
Struggling with Atomic Structure? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
