Atomic structure

Q: What are the relative charges and relative masses of protons, neutrons, and electrons?

Protons have a relative mass of \(1\) and charge of \(+1\); neutrons have a relative mass of \(1\) and charge of \(0\); electrons have a relative mass of \(\frac{1}{1840}\) (negligible) and charge of \(-1\).

Q: What are isotopes, and why do they exhibit identical chemical properties?

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons; they have identical chemical properties because chemical reactions depend only on the electron configuration.

Q: What is the definition of relative atomic mass (\(A_r\))?

Relative atomic mass (\(A_r\)) is the weighted average mass of an atom of an element compared to \(\frac{1}{12}\text{th}\) of the mass of an atom of carbon-12.

Q: What is an atomic orbital, and what are the shapes and capacities of \(s\) and \(p\) orbitals?

An orbital is a region holding up to two electrons with opposite spins; \(s\)-orbitals are spherical (holding \(2\) electrons), and \(p\)-orbitals are dumbbell-shaped (three per subshell, holding \(6\) electrons total).

Q: What are the three fundamental rules that dictate the electronic configuration of atoms?

The Aufbau Principle (fill lowest energy first), Hund's Rule (occupy degenerate orbitals singly before pairing), and the Pauli Exclusion Principle (max two electrons per orbital with opposite spins).

Q: Why do chromium (\(\text{Cr}\)) and copper (\(\text{Cu}\)) have anomalous electron configurations?

\(\text{Cr}\) is \([\text{Ar}]\, 4s^1\, 3d^5\) and \(\text{Cu}\) is \([\text{Ar}]\, 4s^1\, 3d^{10}\) because half-filled and fully-filled \(d\)-subshells provide extra stability.

Q: Which electrons are removed first when transition metals form positive ions?

Electrons in the \(4s\) subshell are always lost before electrons in the \(3d\) subshell.

Q: What is the definition and general equation for First Ionisation Energy (\(\text{IE}_1\))?

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged gaseous positive ions: \(\text{X}\text{(g)} \rightarrow \text{X}^+\text{(g)} + \text{e}^-\).

Q: Why does first ionisation energy generally decrease down a group?

Extra electron shells increase atomic radius and inner-shell shielding, which outweighs the increase in nuclear charge and weakens the pull on the outer electron.

Chemistry 1110 · CCEA A Level · 9 min read · Updated: 31 Aug 2026

Welcome to Atomic Structure!

Welcome to one of the most fundamental chapters in your AS Chemistry journey! Everything in the universe, from the screen you are reading this on to the tea in your mug, is made of atoms. In this unit (Unit AS 1: Basic Concepts in Physical and Inorganic Chemistry), we will look inside the atom, unpack how electrons arrange themselves, and see how this hidden structure controls the chemical behavior of elements across the Periodic Table.

Don't worry if some of this seems unfamiliar or tricky at first! We will break every concept down step-by-step with clear examples, memory aids, and tips directly tailored to what CCEA examiners look for.

1. Fundamental Subatomic Particles

At the center of every atom lies a tiny, dense central nucleus containing protons and neutrons (collectively called nucleons). Surrounding this nucleus are tiny electrons moving within regions of space called orbitals or energy levels.

Properties of Subatomic Particles

Here are the key properties you must know for your exam:

• Proton: Relative mass = \(1\), Relative charge = \(+1\), Location = Nucleus
• Neutron: Relative mass = \(1\), Relative charge = \(0\) (neutral), Location = Nucleus
• Electron: Relative mass = \(\frac{1}{1840}\) (or \(\approx \frac{1}{1836}\) / negligible), Relative charge = \(-1\), Location = Orbitals surrounding the nucleus

Key Definitions to Master

• Atomic Number (\(Z\)): The number of protons in the nucleus of an atom. (This defines what element it is!)
• Mass Number (\(A\)): The total number of protons and neutrons in the nucleus of an atom.
• Isotopes: Atoms of the same element containing the same number of protons (same atomic number) but different numbers of neutrons (different mass numbers).

Did you know? Because chemical reactions involve only electrons, different isotopes of the same element have identical chemical properties!

Key Takeaway for Section 1

Number of Protons = \(Z\)
Number of Neutrons = \(A - Z\)
Number of Electrons (in a neutral atom) = Number of Protons = \(Z\)

2. Relative Masses and Mass Spectrometry

Atoms are far too light to weigh individually on a standard kitchen balance, so chemists compare their masses to a universal standard.

The Standard Reference: Carbon-12

The standard reference is the carbon-12 isotope (\(^{12}\text{C}\)), which is assigned a mass of exactly \(12.00\) unified atomic mass units.

Essential Definitions

• Relative Isotopic Mass: The mass of an atom of an isotope compared to \(\frac{1}{12}\text{th}\) of the mass of an atom of carbon-12.
• Relative Atomic Mass (\(A_r\)): The weighted average mass of an atom of an element compared to \(\frac{1}{12}\text{th}\) of the mass of an atom of carbon-12.

CCEA Exam Tip: When defining relative atomic mass or relative isotopic mass, always include the phrase "compared to \(\frac{1}{12}\text{th}\) of the mass of an atom of carbon-12" to secure full marks!

Calculating Relative Atomic Mass (\(A_r\))

To find the relative atomic mass of an element from isotopic abundance data, use this formula:

\(A_r = \frac{\sum (\text{isotopic mass} \times \text{relative abundance})}{\text{total abundance}}\)

Worked Example:
A sample of chlorine contains \(75.77\%\) of \(^{35}\text{Cl}\) (mass \(35.0\)) and \(24.23\%\) of \(^{37}\text{Cl}\) (mass \(37.0\)). Calculate the relative atomic mass of chlorine to two decimal places.
\(A_r = \frac{(35.0 \times 75.77) + (37.0 \times 24.23)}{100} = \frac{2651.95 + 896.51}{100} = \frac{3548.46}{100} = 35.48\)

3. Electronic Structure and Orbitals

At GCSE, you learned that electrons sit in simple rings (\(2, 8, 8\)). At AS Level, we discover that electron shells are split into subshells and orbitals.

Quantum Shells, Subshells, and Orbitals

• An orbital is a region around the nucleus that can hold up to two electrons with opposite spins.
• Principal quantum shells (\(n = 1, 2, 3, 4\)) contain subshells designated \(s, p, d, f\):
— \(s\)-subshell: Contains \(1\) orbital \(\rightarrow\) holds a maximum of \(2\) electrons.
— \(p\)-subshell: Contains \(3\) degenerate orbitals (\(p_x, p_y, p_z\)) \(\rightarrow\) holds a maximum of \(6\) electrons.
— \(d\)-subshell: Contains \(5\) degenerate orbitals \(\rightarrow\) holds a maximum of \(10\) electrons.

Orbital Shapes

• \(s\)-orbital: Spherical in shape.
• \(p\)-orbital: Dumbbell-shaped, pointing along three perpendicular axes (\(p_x, p_y, p_z\)).

Governing Rules for Filling Orbitals

1. Aufbau Principle: Electrons fill subshells of lowest available energy first. The energy order fills as:
\(1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p\)
Notice that the \(4s\) subshell fills before the \(3d\) subshell because it is lower in energy!
2. Hund's Rule: Orbitals of equal energy (degenerate orbitals) are each occupied singly before any orbital is doubly occupied, and all singly occupied orbitals have parallel spins.
3. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons, and they must have opposite (paired) spins.

Writing Electron Configurations (up to \(Z = 36\))

• Sodium (\(\text{Na}, Z=11\)): \(1s^2\, 2s^2\, 2p^6\, 3s^1\) or \([\text{Ne}]\, 3s^1\)
• Calcium (\(\text{Ca}, Z=20\)): \(1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^6\, 4s^2\) or \([\text{Ar}]\, 4s^2\)
• Iron (\(\text{Fe}, Z=26\)): \(1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^6\, 4s^2\, 3d^6\) or \([\text{Ar}]\, 4s^2\, 3d^6\)

Two Special Transition Metal Anomalies to Memorise!

Nature loves symmetry and stability. Half-filled and fully-filled \(d\)-subshells offer extra stability:

• Chromium (\(\text{Cr}, Z=24\)): \([\text{Ar}]\, 4s^1\, 3d^5\) (NOT \([\text{Ar}]\, 4s^2\, 3d^4\)) due to the stability of a half-filled \(3d\) subshell.
• Copper (\(\text{Cu}, Z=29\)): \([\text{Ar}]\, 4s^1\, 3d^{10}\) (NOT \([\text{Ar}]\, 4s^2\, 3d^9\)) due to the stability of a fully-filled \(3d\) subshell.

Formation of \(d\)-block Cations (Crucial Rule!)

When transition metals form positive ions, electrons in the \(4s\) subshell are lost first before \(3d\) electrons.
• \(\text{Fe}: [\text{Ar}]\, 4s^2\, 3d^6 \rightarrow \text{Fe}^{2+}: [\text{Ar}]\, 3d^6\)
• \(\text{Fe}^{2+}: [\text{Ar}]\, 3d^6 \rightarrow \text{Fe}^{3+}: [\text{Ar}]\, 3d^5\)

4. Ionisation Energies

Definition of First Ionisation Energy (\(\text{IE}_1\))

The First Ionisation Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged gaseous positive ions.

\(\text{X}\text{(g)} \rightarrow \text{X}^+\text{(g)} + \text{e}^-\)

Must-Remember: Always write state symbols! Both the atom and the ion must be in the gaseous state: \(\text{(g)}\).

Successive Ionisation Energies

The general equation for the \(n^{\text{th}}\) ionisation energy is:
\(\text{X}^{(n-1)+}\text{(g)} \rightarrow \text{X}^{n+}\text{(g)} + \text{e}^-\)

For example, the second ionisation energy of magnesium is:
\(\text{Mg}^+\text{(g)} \rightarrow \text{Mg}^{2+}\text{(g)} + \text{e}^-\)

Why do successive ionisation energies always increase?
Each subsequent electron is removed from an increasingly positive ion. The remaining electrons experience greater electrostatic attraction towards the nucleus, requiring more energy to remove.

Evidence for Shell Structure: When plotting successive ionisation energies on a graph (log scale), large jumps/discontinuities occur when an electron is removed from a new principal shell that is closer to the nucleus with less shielding. This provides direct experimental evidence for quantum electron shells.

Four Factors Affecting Ionisation Energy

1. Nuclear charge: More protons = greater attraction to outer electrons = higher \(\text{IE}\).
2. Atomic radius / Distance: Outer electron further from the nucleus = weaker attraction = lower \(\text{IE}\).
3. Shielding / Screening: More inner electron shells = outer electrons repelled by inner shells = lower \(\text{IE}\).
4. Electron-electron repulsion / Subshell stability: Paired electrons in the same orbital repel each other = easier to remove = lower \(\text{IE}\).

Trends in First Ionisation Energy

1. Trend Down a Group (e.g. Group 1 or Group 2)

• Trend: First ionisation energy decreases down the group.
• Reason: As you go down, extra principal quantum shells are added. The atomic radius increases, and there is increased shielding from inner shells. These factors outweigh the increase in nuclear charge, weakening the pull on the outermost electron.

2. General Trend Across a Period (e.g. Period 2 & Period 3)

• General Trend: First ionisation energy increases across a period.
• Reason: Nuclear charge increases (more protons), atomic radius decreases, and electrons are added to the same shell (similar shielding). Therefore, the outer electrons are pulled more strongly by the nucleus.

3. The Two Notable Dips Across Period 2 and Period 3

Although the general trend increases, there are two distinct dips you must explain in your exam:

• Dip 1: Group 2 to Group 3 (e.g. \(\text{Be} \rightarrow \text{B}\) or \(\text{Mg} \rightarrow \text{Al}\))
— \(\text{Mg}\) configuration: \([\text{Ne}]\, 3s^2\)
— \(\text{Al}\) configuration: \([\text{Ne}]\, 3s^2\, 3p^1\)
— Explanation: The electron removed from aluminium comes from a higher-energy \(3p\) subshell, which is further from the nucleus and shielded by the inner \(3s^2\) electrons. This makes it easier to remove than the \(3s\) electron in magnesium.

• Dip 2: Group 5 to Group 6 (e.g. \(\text{N} \rightarrow \text{O}\) or \(\text{P} \rightarrow \text{S}\))
— \(\text{P}\) configuration: \([\text{Ne}]\, 3s^2\, 3p_x^1\, 3p_y^1\, 3p_z^1\)
— \(\text{S}\) configuration: \([\text{Ne}]\, 3s^2\, 3p_x^2\, 3p_y^1\, 3p_z^1\)
— Explanation: In sulfur, one of the \(3p\) orbitals contains a pair of electrons. Due to spin-pair repulsion between two electrons sharing the same orbital, less energy is needed to remove one of these paired electrons compared to the singly-occupied orbitals in phosphorus.

5. Common Pitfalls & Examiner Warnings

Avoid these frequent mistakes identified in CCEA examiner reports:

1. Forgetting State Symbols: Always write \(\text{X}\text{(g)} \rightarrow \text{X}^+\text{(g)} + \text{e}^-\). Writing \(\text{Na} \rightarrow \text{Na}^+ + \text{e}^-\) will score zero marks for state symbols.
2. Miswriting the 2nd Ionisation Energy: Remember, the 2nd \(\text{IE}\) is the removal of a second electron from a \(+1\) ion: \(\text{X}^+\text{(g)} \rightarrow \text{X}^{2+}\text{(g)} + \text{e}^-\), not \(\text{X}\text{(g)} \rightarrow \text{X}^{2+}\text{(g)} + 2\text{e}^-\).
3. Transition Metal Cations: Always remove \(4s\) electrons before \(3d\) electrons! \(\text{Fe}^{2+}\) is \([\text{Ar}]\, 3d^6\), never \([\text{Ar}]\, 4s^2\, 3d^4\).
4. Vague Explanations for Dips: Do not simply say "Group 5 is more stable". Explicitly state that Group 6 has spin-pair repulsion in a \(p\)-orbital, or that Group 3 has an electron in a higher-energy \(p\)-subshell.

6. Quick Chapter Summary Checklist

Before moving on to the next topic, check that you can:

• State the relative mass and charge of protons, neutrons, and electrons.
• Define atomic number, mass number, isotopes, relative isotopic mass, and relative atomic mass (\(A_r\)).
• Calculate \(A_r\) using isotopic mass and percentage or relative abundance.
• Describe the spherical shape of an \(s\)-orbital and the dumbbell shape of a \(p\)-orbital.
• Write electronic configurations up to \(Z = 36\) using subshell, noble gas, and box notation, including the special cases \(\text{Cr}\) and \(\text{Cu}\).
• Write electron configurations for transition metal ions by removing \(4s\) electrons first.
• Write equations for first and successive ionisation energies with gaseous state symbols.
• Explain the trends in first ionisation energy down a group, across Period 2/3, and account for the dips at Group 3 and Group 6.

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

What are the relative charges and relative masses of protons, neutrons, and electrons?

Protons have a relative mass of \(1\) and charge of \(+1\); neutrons have a relative mass of \(1\) and charge of \(0\); electrons have a relative mass of \(\frac{1}{1840}\) (negligible) and charge of \(-1\).

What are isotopes, and why do they exhibit identical chemical properties?

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons; they have identical chemical properties because chemical reactions depend only on the electron configuration.

What is the definition of relative atomic mass (\(A_r\))?

Relative atomic mass (\(A_r\)) is the weighted average mass of an atom of an element compared to \(\frac{1}{12}\text{th}\) of the mass of an atom of carbon-12.

How is relative atomic mass (\(A_r\)) calculated from isotopic abundances?

\(A_r = \frac{\sum (\text{isotopic mass} \times \text{relative abundance})}{\text{total abundance}}\)

What is an atomic orbital, and what are the shapes and capacities of \(s\) and \(p\) orbitals?

An orbital is a region holding up to two electrons with opposite spins; \(s\)-orbitals are spherical (holding \(2\) electrons), and \(p\)-orbitals are dumbbell-shaped (three per subshell, holding \(6\) electrons total).

What are the three fundamental rules that dictate the electronic configuration of atoms?

The Aufbau Principle (fill lowest energy first), Hund's Rule (occupy degenerate orbitals singly before pairing), and the Pauli Exclusion Principle (max two electrons per orbital with opposite spins).

Why do chromium (\(\text{Cr}\)) and copper (\(\text{Cu}\)) have anomalous electron configurations?

\(\text{Cr}\) is \([\text{Ar}]\, 4s^1\, 3d^5\) and \(\text{Cu}\) is \([\text{Ar}]\, 4s^1\, 3d^{10}\) because half-filled and fully-filled \(d\)-subshells provide extra stability.

Which electrons are removed first when transition metals form positive ions?

Electrons in the \(4s\) subshell are always lost before electrons in the \(3d\) subshell.

What is the definition and general equation for First Ionisation Energy (\(\text{IE}_1\))?

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged gaseous positive ions: \(\text{X}\text{(g)} \rightarrow \text{X}^+\text{(g)} + \text{e}^-\).

Why does first ionisation energy generally decrease down a group?

Extra electron shells increase atomic radius and inner-shell shielding, which outweighs the increase in nuclear charge and weakens the pull on the outer electron.

Why does first ionisation energy generally increase across a period?

Nuclear charge increases and atomic radius decreases while shielding remains similar, causing a stronger attraction to the outermost electron.

Why is the first ionisation energy of aluminium lower than that of magnesium (Group 2 to Group 3 dip)?

The electron removed from aluminium is in a higher-energy \(3p\) subshell, which is further from the nucleus and shielded by the \(3s^2\) electrons.

Why is the first ionisation energy of sulfur lower than that of phosphorus (Group 5 to Group 6 dip)?

Sulfur has a pair of electrons in one of its \(3p\) orbitals; the resulting spin-pair repulsion makes that electron easier to remove.

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