A-Level Chemistry Periodic Table: How to Read and Use It in Exams

The periodic table is the single most important reference tool in A-Level chemistry. Every exam provides you with one, and countless questions require you to extract information from it or explain trends within it. Yet many students treat it as just a list of elements rather than the powerful analytical tool it actually is.

Understanding how to read the periodic table properly – and recognising the trends and patterns it reveals – can transform your performance across organic, inorganic, and physical chemistry questions.

What Information Does the Periodic Table Provide?

At first glance, the periodic table looks like a grid of element symbols with numbers. But each element’s entry contains multiple pieces of information:

Basic Information for Each Element

Element symbol (e.g., Na for sodium, Cl for chlorine)
Atomic number (number of protons in the nucleus)
Relative atomic mass (average mass of an atom, accounting for isotopes)

Some periodic tables also include:

Electron configuration (though you’re expected to work this out yourself)
Oxidation states (common charge states of the element)
Electronegativity values (in some versions)

For A-Level exams, the AQA periodic table is relatively minimal. Edexcel and OCR versions vary slightly but all provide atomic number and relative atomic mass at minimum.

How the Periodic Table is Organised

The genius of the periodic table lies in its organisation. Elements are arranged by:

1. Atomic Number (Left to Right)

Elements are ordered by increasing atomic number (number of protons). Each element has one more proton than the element before it.

2. Periods (Horizontal Rows)

There are seven periods. Each period represents elements with the same number of electron shells:

Period 1: Elements have 1 electron shell (H, He)
Period 2: Elements have 2 electron shells (Li to Ne)
Period 3: Elements have 3 electron shells (Na to Ar)
Period 4: Elements have 4 electron shells (K to Kr)

And so on.

3. Groups (Vertical Columns)

There are 18 groups (though Groups 3-12 are often treated as one block for transition metals). Elements in the same group have:

The same number of electrons in their outer shell
Similar chemical properties
Similar reactivity patterns

Key groups you need to know:

Group 1: Alkali metals (Li, Na, K, etc.) – 1 outer electron
Group 2: Alkaline earth metals (Mg, Ca, etc.) – 2 outer electrons
Group 7 (17): Halogens (F, Cl, Br, I) – 7 outer electrons
Group 0 (18): Noble gases (He, Ne, Ar, Kr) – Full outer shell

4. Blocks (s, p, d, f)

The periodic table is divided into blocks based on which orbital is being filled:

s-block: Groups 1 and 2 (left side)
p-block: Groups 3-8 (right side)
d-block: Transition metals (middle section)
f-block: Lanthanides and actinides (bottom two rows)

Understanding blocks helps you work out electron configurations quickly.

Key Periodic Trends You Must Know

The periodic table isn’t just a reference – it reveals patterns in chemical and physical properties. Here are the trends examiners expect you to know.

1. Atomic Radius

Across a period (left to right): Atomic radius decreases

Why? More protons in the nucleus pull electrons closer, despite having the same number of shells

Down a group (top to bottom): Atomic radius increases

Why? More electron shells, so the outer electrons are further from the nucleus

Exam application: Explains why sodium is larger than chlorine, and why potassium is larger than sodium.

2. Ionisation Energy

Ionisation energy is the energy required to remove an electron from an atom.

Across a period (left to right): First ionisation energy increases

Why? Stronger nuclear attraction (more protons) makes it harder to remove an electron
Exception: Small drops between Groups 2 and 3 (new subshell) and Groups 5 and 6 (paired electrons)

Down a group (top to bottom): First ionisation energy decreases

Why? Outer electrons are further from the nucleus and more shielded, so they’re easier to remove

Exam application: Questions often ask you to explain trends in ionisation energy data or predict values for unfamiliar elements.

3. Electronegativity

Electronegativity measures an atom’s ability to attract electrons in a covalent bond.

Across a period (left to right): Electronegativity increases

Why? More protons means stronger pull on bonding electrons

Down a group (top to bottom): Electronegativity decreases

Why? Bonding electrons are further from the nucleus, so the attraction is weaker

Exam application: Explains bond polarity (e.g., why H-Cl is polar but Cl-Cl is not) and predicts which element is more electronegative in a bond.

4. Melting and Boiling Points

This is more complex because it depends on structure and bonding:

Across Period 3 (Na to Ar):

Na, Mg, Al: Metallic bonding – increases as more delocalised electrons
Si: Giant covalent – very high melting point
P, S, Cl, Ar: Simple molecular – low melting points

Down Group 7 (halogens):

Melting/boiling points increase (F₂ < Cl₂ < Br₂ < I₂)
Why? Larger molecules have stronger van der Waals forces

Exam application: Explaining anomalies in melting point data or predicting physical properties of elements.

5. Reactivity

Group 1 (alkali metals): Reactivity increases down the group

Why? Outer electron is further from nucleus, easier to lose

Group 7 (halogens): Reactivity decreases down the group

Why? Outer electrons are further from nucleus, harder to gain an electron

Exam application: Displacement reactions (Cl₂ displaces Br⁻ because chlorine is more reactive) and predicting reaction rates.

How Exam Boards Differ in Their Periodic Tables

AQA Periodic Table

AQA provides a clean, minimal periodic table showing:

Element symbol
Atomic number (top)
Relative atomic mass (bottom)

AQA does NOT include oxidation states or electronegativity values. You’re expected to recall or deduce these.

Edexcel Periodic Table

Similar to AQA but sometimes includes additional information like common oxidation states for transition metals. Check your exam board’s specimen papers to see exactly what’s provided.

OCR Periodic Table

OCR’s periodic table is comprehensive, often including:

Electronegativity values
Common oxidation states
More detailed formatting for transition metals

This means OCR students have slightly more information readily available, but the same knowledge is still tested.

Key point: Always use your exam board’s periodic table during revision so you’re familiar with its exact layout and what information it does (or doesn’t) provide.

How Exams Test Your Periodic Table Knowledge

Questions fall into several categories:

1. Extracting Information

“Using the periodic table, state the number of protons, neutrons, and electrons in an atom of chlorine-35.”

How to answer:

Atomic number = 17 (protons and electrons)
Mass number = 35
Neutrons = 35 - 17 = 18

2. Predicting Trends

“Explain why the first ionisation energy of magnesium is greater than the first ionisation energy of sodium.”

How to answer:

Magnesium has more protons (12 vs 11)
Both have the same number of shells (same period)
Greater nuclear charge in magnesium pulls electrons more strongly
Therefore more energy is needed to remove an electron from magnesium

3. Explaining Anomalies

“Explain why the first ionisation energy of aluminium is lower than the first ionisation energy of magnesium, despite aluminium having a higher atomic number.”

How to answer:

Magnesium’s outer electrons are in the 3s subshell (paired)
Aluminium’s outer electron is in the 3p subshell (unpaired, higher energy)
The 3p electron is slightly further from the nucleus and more shielded
Therefore less energy is needed to remove it

4. Applying Trends to Reactions

“Explain why chlorine is able to displace bromide ions from solution.”

How to answer:

Chlorine is more reactive than bromine (higher up Group 7)
Chlorine has a stronger attraction for electrons (higher electronegativity)
Chlorine atoms gain electrons more easily than bromine atoms
Therefore chlorine displaces bromide ions: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

5. Determining Electron Configurations

“State the electron configuration of a chlorine atom.”

How to answer:

Chlorine is in Period 3, Group 7
Atomic number = 17 (17 electrons)
Configuration: 1s² 2s² 2p⁶ 3s² 3p⁵

Or in shorthand: [Ne] 3s² 3p⁵

You can work this out using the periodic table’s structure (s-block, p-block, d-block).

Common Mistakes Students Make

Mistake 1: Confusing Atomic Number with Mass Number

Atomic number (number of protons) is always the smaller number on the periodic table. Mass number (protons + neutrons) is sometimes given in questions (e.g., carbon-12, carbon-14) but is NOT always shown on the periodic table.

Mistake 2: Misunderstanding Relative Atomic Mass

Relative atomic mass is often not a whole number (e.g., chlorine is 35.5) because it’s the weighted average of all isotopes. Don’t assume all atoms of an element have the same mass.

Mistake 3: Forgetting About Electron Shielding

When explaining ionisation energy trends down a group, students often say “the outer electrons are further away” but forget to mention shielding. Both distance and shielding are important.

Mistake 4: Applying Group Trends Across Periods

Reactivity trends are specific to each group. Don’t assume that because Group 1 gets more reactive down the group, all groups behave the same way. Group 7 gets less reactive down the group.

Mistake 5: Not Using Data to Support Answers

If a question provides data (e.g., ionisation energies or electronegativities), you must reference the data in your answer. Simply stating a general trend without linking it to the specific numbers given won’t earn full marks.

How to Use the Periodic Table Effectively in Your Exam

1. Locate Elements Quickly

Practise finding elements fast. If a question mentions “an element in Period 3 with 5 outer electrons,” you should immediately locate phosphorus (P).

2. Identify Patterns Instantly

See an unfamiliar element? Look at its position:

Same group as a known element? It will have similar properties
Same period? It will have the same number of electron shells

3. Cross-Check Your Electron Configurations

If you write an electron configuration, cross-check it against the periodic table. The period number tells you the highest shell, and the group number (for s and p blocks) tells you the number of outer electrons.

4. Use It for Oxidation States

Though oxidation states aren’t always shown, you can deduce common ones:

Group 1: +1
Group 2: +2
Group 7: -1
Transition metals: variable (often +2 or +3)

5. Don’t Ignore the Bottom Rows

Lanthanides and actinides (f-block) are tested less frequently, but don’t assume they’re irrelevant. Some questions ask about electron configurations or trends involving these elements.

How UpGrades Helps with A-Level Chemistry Periodic Trends

Understanding the periodic table is one thing – applying that knowledge to unfamiliar exam questions is another. UpGrades provides:

Targeted practice questions on periodic trends, ionisation energy, electronegativity, and reactivity
Instant feedback showing where your explanations are incomplete or incorrect
Adaptive difficulty that focuses on the areas you find hardest (e.g., explaining ionisation energy anomalies, predicting reactivity)

With UpGrades, you’ll build confidence interpreting data, spotting patterns, and writing clear, precise explanations that earn full marks. Combine smart periodic table revision with UpGrades’ exam-style practice, and you’ll approach your A-Level chemistry exams fully prepared.

Final Checklist for Periodic Table Mastery

Memorise key groups (Group 1, Group 2, Group 7, Group 0) and their properties
Understand atomic radius, ionisation energy, and electronegativity trends
Practise explaining trends with reference to nuclear charge, shielding, and distance
Learn how to work out electron configurations from the periodic table structure
Apply trends to reactivity and displacement reactions
Familiarise yourself with your exam board’s specific periodic table layout
Practise using the periodic table during timed past papers

The A-Level chemistry periodic table isn’t just a reference sheet – it’s a map of chemical behaviour. Master how to read it, understand the trends it reveals, and you’ll unlock a deeper understanding of chemistry that translates directly into exam success.