Atomic Structure and the Periodic Table

Chemistry (8CH0) · Pearson Edexcel AS Level · 8 min read · Updated: 2 Oct 2026

Welcome to Atomic Structure and the Periodic Table!

Welcome to the first step of your AS Level Chemistry journey! This chapter is the foundation of everything else you’ll learn. We are going to explore what atoms are made of, how we can "weigh" them, and why the Periodic Table is organized the way it is. Don't worry if some of the math or the patterns seem a bit strange at first—once you see the logic behind them, it all starts to click like a puzzle!

1. The Heart of the Atom: Sub-atomic Particles

Atoms are the tiny building blocks of everything around you. Even though they are incredibly small, they are made of three even smaller pieces: protons, neutrons, and electrons.

Mass and Charge

Think of the nucleus (the center) as a heavy suitcase and the electrons as tiny flies buzzing around it.

Protons: Found in the nucleus. Relative mass = 1. Relative charge = +1.
Neutrons: Found in the nucleus. Relative mass = 1. Relative charge = 0 (they are neutral!).
Electrons: Found in shells orbiting the nucleus. Relative mass = 1/1840 (so small we often say "negligible"). Relative charge = -1.

Atomic Number and Mass Number

To identify an atom, we look at its numbers on the Periodic Table:

1. Atomic (proton) number ( \(Z\)): This is the number of protons in the nucleus. It defines the element! If you change the proton number, you change the element.
2. Mass number ( \(A\)): This is the total number of protons + neutrons.

Quick Calculation Trick:
To find the number of neutrons, just subtract the bottom number from the top number: \( \text{Neutrons} = \text{Mass Number} - \text{Atomic Number} \).
In a neutral atom, the number of electrons is always equal to the number of protons.

Key Takeaway: Protons and neutrons live in the center (nucleus) and provide the mass; electrons orbit on the outside and provide the charge.

2. Isotopes and Atomic Mass

Nature isn't always identical. Sometimes, atoms of the same element have different numbers of neutrons. We call these isotopes.

Isotopes are atoms of the same element with the same number of protons but a different number of neutrons. This means they have the same atomic number but different mass numbers.

The Carbon-12 Scale

Because atoms are too small to weigh on a normal scale, we compare them to a standard: the Carbon-12 isotope.

Relative isotopic mass: The mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon-12.
Relative atomic mass (RAM or \(A_r\)): The weighted mean mass of an atom of an element compared with 1/12th of the mass of an atom of carbon-12.

Analogy: If you have a bag of marbles where 75% weigh 10g and 25% weigh 12g, the "weighted average" is the RAM.

Key Takeaway: Isotopes react chemically the same way because they have the same number of electrons, but they have different physical masses.

3. Mass Spectrometry: Weighing Atoms

How do we actually know how many isotopes an element has? We use a machine called a Mass Spectrometer.

Calculating Relative Atomic Mass

You might be asked to calculate the RAM from a graph (a mass spectrum). The x-axis is \(m/z\) (mass/charge ratio) and the y-axis is abundance (percentage).
\( RAM = \frac{\sum (\text{isotopic mass} \times \text{relative abundance})}{\text{total abundance}} \)

Diatomic Molecules (like Chlorine, \(\text{Cl}_2\))

Chlorine has two main isotopes: \(^{35}\text{Cl}\) and \(^{37}\text{Cl}\). When they form a molecule (\(\text{Cl}_2\)), you can get three different combinations: 35+35 (mass 70), 35+37 (mass 72), and 37+37 (mass 74). Chemistry exams love asking you to predict the height of these peaks based on probability!

The Molecular Ion Peak

For a whole molecule, the mass spectrometer gives a peak called the molecular ion peak (\(M^+\)). This peak tells you the relative molecular mass of the substance.

Key Takeaway: The mass spectrometer is like a very sensitive set of scales that separates particles by their mass.

4. Ionisation Energy: The Cost of an Electron

If you want to pull an electron away from an atom, you have to "pay" with energy. This is Ionisation Energy (IE).

First Ionisation Energy: The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Equation: \( X(g) \rightarrow X^+(g) + e^- \)

Factors affecting IE:

1. Nuclear Charge: More protons = more "pull" on electrons (IE increases).
2. Shielding: More inner shells of electrons "block" the pull of the nucleus (IE decreases).
3. Distance (Atomic Radius): Further away the electron is, the easier it is to remove (IE decreases).

General Trends:

Down a Group: IE decreases. Even though nuclear charge increases, the extra electron shells increase shielding and distance, making the outer electron easier to remove.
Across a Period: IE generally increases. Nuclear charge increases while shielding remains roughly constant, pulling electrons closer to the nucleus.

Explaining the Specific Dips Across a Period:

There are two characteristic dips across Period 2 (and Period 3) that provide experimental evidence for sub-shells:

Group 2 to Group 3 Dip (e.g., \(\text{Be} \rightarrow \text{B}\) or \(\text{Mg} \rightarrow \text{Al}\)): The outer electron in Group 3 is in a higher-energy \(p\)-subshell (e.g., \(2p\) instead of \(2s\)). This \(p\)-subshell is further from the nucleus and shielded by the inner \(s\)-electrons, making it easier to remove.
Group 5 to Group 6 Dip (e.g., \(\text{N} \rightarrow \text{O}\) or \(\text{P} \rightarrow \text{S}\)): In Group 5, the \(p\)-orbitals are all singly occupied (\(p_x^1 p_y^1 p_z^1\)). In Group 6, the fourth \(p\)-electron must pair up in an orbital (\(p_x^2 p_y^1 p_z^1\)). The repulsion between the two paired electrons in the same orbital makes that electron easier to remove.

Successive Ionisation Energies

Successive ionisation energies (removing the 1st, 2nd, 3rd, etc., electron) always increase as electrons are removed from an increasingly positive ion.
When plotted as a graph of \(\log_{10}(\text{IE})\) against the number of electrons removed, large jumps in energy indicate the removal of an electron from a new inner shell closer to the nucleus. The number of electrons removed before the first big jump tells you the group number of the element.

5. Electronic Configuration: Where are the electrons?

Electrons don't just fly around randomly. They live in quantum shells, sub-shells, and orbitals.

Shells and Sub-shells

Quantum Shells: Numbered \(n = 1, 2, 3, 4\). The maximum electrons they can hold is \( 2n^2 \) (Shell 1 = 2, Shell 2 = 8, Shell 3 = 18, Shell 4 = 32).
Sub-shells: Shells are split into \(s, p,\) and \(d\) sub-shells:

\(s\)-subshell: 1 orbital (max 2 electrons)
\(p\)-subshell: 3 orbitals (max 6 electrons)
\(d\)-subshell: 5 orbitals (max 10 electrons)

Orbitals

An orbital is a region of space where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons with opposite spins.

\(s\)-orbital: Spherical shape.
\(p\)-orbital: Dumbbell shape (three degenerate orbitals along the axes: \(p_x, p_y, p_z\)).

Filling Rules

1. Aufbau Principle: Orbitals are filled in order of increasing energy (\(1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p\)).
2. Hund's Rule: Electrons fill degenerate orbitals singly before pairing up.
3. Pauli Exclusion Principle: Two electrons in the same orbital must have opposite spins.

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

Example for Magnesium (12 electrons): \( 1s^2 2s^2 2p^6 3s^2 \)

Important Anomalies to Know:

Exam boards frequently test two elements that do not follow the standard Aufbau order because a half-filled or fully filled \(d\)-subshell provides extra stability:

Chromium (\(Z=24\)): \( 1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^1 \) (or \([\text{Ar}] 3d^5 4s^1\)), NOT \(3d^4 4s^2\).
Copper (\(Z=29\)): \( 1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^1 \) (or \([\text{Ar}] 3d^{10} 4s^1\)), NOT \(3d^9 4s^2\).

Key Takeaway: Electronic configuration determines how an atom reacts. Elements in the same group have the same outer shell configuration!

6. Periodicity: Patterns in the Table

Periodicity is the repeating pattern of physical and chemical properties across periods.

Blocks

The table is divided into blocks based on which sub-shell contains the highest-energy electron:

\(s\)-block: Groups 1 and 2.
\(p\)-block: Groups 3 to 0 (Groups 13 to 18).
\(d\)-block: The transition metals.

Melting and Boiling Points

The trends in melting and boiling points across Periods 2 and 3 depend on bonding and structure:

Metals (Li, Be / Na, Mg, Al): Giant metallic structures. Melting points increase as the charge on the metal ion increases and more delocalised electrons strengthen metallic bonding.
Giant Covalent (C / Si): Very high melting points because many strong covalent bonds throughout the macromolecular lattice must be broken.
Simple Molecular (N, O, F, Ne / P, S, Cl, Ar): Low melting points because only weak London (intermolecular) forces between molecules need to be overcome.

Quick Review Box:
- Atomic Radii: Decrease across a period (higher nuclear charge pulls shells closer).
- First Ionisation Energy: Generally increases across a period (with dips at Groups 3 and 6).
- Chemical Properties: Dictated by outer shell electronic configurations.

Key Takeaway: The Periodic Table isn't just a list; it's a map. If you know where an element is, you can predict how it will behave!

Practice writing out full electronic configurations and explaining the IE dips, and you'll master Topic 1!

References

Pearson Edexcel Level 3 Advanced Subsidiary GCE in Chemistry Specification (8CH0) — Pearson Edexcel
Topic 1: Atomic Structure and the Periodic Table Revision Guide — Chemrevise
Detailed Notes - Topic 1 Atomic Structure and the Periodic Table — Physics & Maths Tutor

Quick check

Can you answer these now?

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

What are the relative masses and charges 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 \(1/1840\) (negligible) and charge of -1.

What do the atomic number (\(Z\)) and mass number (\(A\)) of an atom represent?

Atomic number (\(Z\)) is the number of protons in the nucleus, while mass number (\(A\)) is the total number of protons plus neutrons.

What are isotopes?

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons (same atomic number, different mass numbers).

How is relative atomic mass (\(A_r\)) defined on the Carbon-12 scale?

Relative atomic mass (\(A_r\)) is the weighted mean mass of an atom of an element compared with \(1/12\)th of the mass of an atom of carbon-12.

How is relative atomic mass calculated from mass spectrometry data?

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

What does the molecular ion peak (\(M^+\)) indicate in a mass spectrum of a compound?

The molecular ion peak (\(M^+\)) represents the unfragmented ion of the complete molecule and gives its relative molecular mass (\(M_r\)).

What is the definition of first ionisation energy, and what is its general equation?

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions: \( X(g) \rightarrow X^+(g) + e^- \).

What are the general periodic trends in first ionisation energy across a period and down a group?

First ionisation energy generally increases across a period (greater nuclear charge, similar shielding) and decreases down a group (increased atomic radius and shielding).

Why does first ionisation energy dip from Group 2 to Group 3 and from Group 5 to Group 6 across a period?

The Group 2 to 3 dip occurs because the outer electron is in a higher-energy, shielded \(p\)-subshell; the Group 5 to 6 dip occurs due to electron-pair repulsion in a doubly occupied \(p\)-orbital.

How can successive ionisation energies be used to determine the group number of an element?

The group number corresponds to the number of electrons removed before the first large jump in ionisation energy on a \(\log_{10}(\text{IE})\) plot.

What are the maximum electron capacities and shapes of atomic orbitals and sub-shells?

An orbital holds up to 2 electrons with opposite spins; \(s\)-subshells have 1 spherical orbital (2 electrons), \(p\)-subshells have 3 dumbbell orbitals (6 electrons), and \(d\)-subshells have 5 orbitals (10 electrons).

What are the anomalous ground-state electron configurations of Chromium (\(Z=24\)) and Copper (\(Z=29\))?

Chromium is \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^1\) and Copper is \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^1\).

How does structure and bonding explain the variation in melting points across Periods 2 and 3?

Melting points increase through giant metallic structures, peak at giant covalent structures due to extensive strong covalent bonding, and drop sharply for simple molecular structures held by weak London forces.

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