What you'll learn
This revision guide covers how the periodic table is organised into groups and periods, and how this arrangement reflects the electronic structure of elements. You'll learn to identify patterns in properties across periods and down groups, understand why elements in the same group behave similarly, and predict properties of elements based on their position in the periodic table.
Key terms and definitions
Group — a vertical column in the periodic table containing elements with the same number of electrons in their outer shell, giving them similar chemical properties.
Period — a horizontal row in the periodic table; the period number indicates how many electron shells an atom of that element has.
Electronic structure — the arrangement of electrons in shells (or energy levels) around the nucleus of an atom.
Noble gases — elements in Group 0 (or Group 8) that have full outer electron shells, making them unreactive.
Alkali metals — highly reactive metals in Group 1 that have one electron in their outer shell.
Halogens — reactive non-metals in Group 7 that have seven electrons in their outer shell.
Transition metals — elements found in the central block of the periodic table between Groups 2 and 3, which form coloured compounds and can act as catalysts.
Outer shell — the highest energy level containing electrons in an atom; the number of electrons in this shell determines an element's chemical properties.
Core concepts
Organisation of the periodic table
The modern periodic table arranges all known elements in order of increasing atomic number (number of protons). This arrangement was developed by Dmitri Mendeleev in 1869, though he ordered elements by atomic mass and left gaps for undiscovered elements.
Elements are organised into:
- Vertical columns called groups, numbered 1–7 and 0 (or 1–8 depending on convention)
- Horizontal rows called periods, numbered 1–7
- A central block of transition metals between Groups 2 and 3
The periodic table is divided into metals (found on the left and centre) and non-metals (found on the right). A stepped line separates these two categories.
Groups and outer shell electrons
All elements within the same group share the same number of electrons in their outer shell:
- Group 1 elements have 1 outer electron
- Group 2 elements have 2 outer electrons
- Group 7 elements have 7 outer electrons
- Group 0 elements have 8 outer electrons (except helium with 2)
This shared outer shell configuration explains why elements in the same group have similar chemical properties. Chemical reactions involve the loss, gain or sharing of outer shell electrons, so elements with the same number of outer electrons react in similar ways.
Example electronic structures:
- Lithium (Li): 2,1 — Group 1
- Sodium (Na): 2,8,1 — Group 1
- Potassium (K): 2,8,8,1 — Group 1
All three have 1 outer electron, placing them in Group 1.
Periods and electron shells
The period number tells you how many electron shells an atom contains:
- Period 1 elements (hydrogen and helium) have 1 electron shell
- Period 2 elements (lithium to neon) have 2 electron shells
- Period 3 elements (sodium to argon) have 3 electron shells
- Period 4 elements (potassium to krypton) have 4 electron shells
Example:
- Carbon is in Period 2, so it has 2 electron shells: 2,4
- Silicon is in Period 3, so it has 3 electron shells: 2,8,4
- Both are in Group 4 (4 outer electrons)
Group 1: The alkali metals
The alkali metals include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr).
Properties:
- Soft metals that can be cut with a knife
- Low melting and boiling points compared to other metals
- Low density (lithium, sodium and potassium float on water)
- Shiny when freshly cut but tarnish quickly in air
- One electron in their outer shell
Reactions: All alkali metals react vigorously with water to produce a metal hydroxide and hydrogen gas:
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
The reaction produces an alkaline solution (hence the name "alkali metals"), with pH 12–14.
Trends down Group 1:
- Reactivity increases as you go down the group
- Melting and boiling points decrease
- Density increases
The increase in reactivity occurs because the outer electron is further from the nucleus in larger atoms. The electrostatic attraction between the positive nucleus and the outer electron is weaker, making it easier to lose the electron during reactions.
Group 7: The halogens
The halogens include fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂) and astatine (At₂).
Properties:
- Non-metals that exist as diatomic molecules (X₂)
- Seven electrons in their outer shell
- Coloured elements: chlorine is pale green gas, bromine is red-brown liquid, iodine is dark grey solid
- Poisonous and reactive
Reactions: Halogens react with metals to form ionic compounds called salts:
2Na(s) + Cl₂(g) → 2NaCl(s)
They gain one electron to form 1− ions (halide ions):
- Chlorine forms Cl⁻ (chloride)
- Bromine forms Br⁻ (bromide)
- Iodine forms I⁻ (iodide)
Trends down Group 7:
- Reactivity decreases as you go down the group
- Melting and boiling points increase
- Colour becomes darker
Reactivity decreases because larger atoms have the outer shell further from the nucleus. The electrostatic attraction for an additional electron is weaker, making it harder to gain an electron.
Displacement reactions: A more reactive halogen will displace a less reactive halogen from an aqueous solution of its salt:
Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
Chlorine displaces bromine because chlorine is more reactive. Bromine cannot displace chlorine from chloride solutions.
Group 0: The noble gases
The noble gases include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn).
Properties:
- Unreactive (inert) gases
- Colourless at room temperature
- Full outer shell of electrons (8 electrons, except helium with 2)
- Exist as single atoms (monatomic)
- Low boiling points
The noble gases are unreactive because they have a stable electronic structure with a full outer shell. They do not need to lose, gain or share electrons.
Uses based on their inertness:
- Helium: party balloons, airships (less dense than air, non-flammable)
- Neon: advertising signs (glows when electricity passes through it)
- Argon: light bulbs, welding (prevents oxidation of hot metals)
Trends down Group 0:
- Boiling points increase as you go down the group
- Density increases
Transition metals
The transition metals occupy the central block of the periodic table between Groups 2 and 3. Examples include iron (Fe), copper (Cu), chromium (Cr), nickel (Ni), zinc (Zn), silver (Ag) and gold (Au).
Properties:
- Hard, strong metals with high densities
- High melting and boiling points (except mercury)
- Good conductors of heat and electricity
- Can form ions with different charges (e.g., Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺)
- Form coloured compounds (e.g., copper compounds are often blue/green, iron compounds are often orange/brown)
- Many act as catalysts (e.g., iron in the Haber process, nickel in margarine production)
Transition metals differ from Group 1 metals in being:
- Harder and stronger
- Higher melting points
- Less reactive (don't react vigorously with water or oxygen)
- Form ions with different charges
- Form coloured compounds
Worked examples
Example 1: Predicting electronic structure from position
Question: An element is in Group 6 and Period 3 of the periodic table. Write its electronic structure and identify the element. [3 marks]
Answer:
- Period 3 means the element has 3 electron shells [1 mark]
- Group 6 means the element has 6 electrons in its outer shell [1 mark]
- Electronic structure: 2,8,6
- The element is sulfur (S) [1 mark]
Examiner's note: Always check the group number matches the number of outer electrons and the period number matches the number of shells.
Example 2: Explaining trends in Group 1
Question: Potassium reacts more vigorously with water than sodium. Explain why, in terms of electronic structure. [3 marks]
Answer:
- Potassium has more electron shells than sodium / potassium's outer electron is in the 4th shell while sodium's is in the 3rd shell [1 mark]
- The outer electron in potassium is further from the nucleus [1 mark]
- The electrostatic attraction between the nucleus and outer electron is weaker, so the electron is lost more easily, making potassium more reactive [1 mark]
Examiner's note: Link electronic structure to distance from nucleus, then to attraction, then to reactivity.
Example 3: Displacement reaction prediction
Question: Predict what you would observe when chlorine solution is added to potassium iodide solution. Write a balanced symbol equation for the reaction. [4 marks]
Answer:
- The colourless solution turns brown/orange [1 mark]
- Because iodine is produced [1 mark]
- Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq) [2 marks: 1 for correct formulae, 1 for balancing]
Alternative: Chlorine is more reactive than iodine / chlorine is higher in Group 7 than iodine [1 mark], so chlorine displaces iodine [1 mark].
Common mistakes and how to avoid them
Confusing groups and periods — Remember: groups are vertical columns, periods are horizontal rows. The group number tells you outer electrons; the period number tells you electron shells.
Thinking all Group 1 elements react the same way — While they react similarly (same type of reaction), reactivity increases down the group. Potassium reacts more vigorously than lithium.
Getting halogen displacement backwards — More reactive halogens displace less reactive ones. Chlorine (top of group, most reactive) displaces bromine and iodine. Iodine (bottom of group, least reactive) cannot displace chlorine or bromine.
Writing incorrect electronic structures — The first shell holds maximum 2 electrons, subsequent shells hold maximum 8. For Group 0 elements, remember helium has 2 outer electrons (full first shell), others have 8.
Forgetting noble gases are unreactive — Noble gases have full outer shells so don't react. Don't confuse them with Group 7 elements which are very reactive.
Applying Group 1 trends to Group 7 — Reactivity increases down Group 1 but decreases down Group 7. Make sure you learn which direction each trend goes.
Exam technique for "Atomic structure and the periodic table: groups and periods in the periodic table"
"Explain" questions require you to give reasons. For reactivity trends, mention electron shells, distance from nucleus, and strength of electrostatic attraction. Link electronic structure to the property being explained.
Predicting observations — State both the colour change and what substance causes it. For displacement reactions: "solution turns from colourless to brown/orange due to formation of iodine/bromine."
Extended response questions (6 marks) on periodic table trends require detailed comparison. Structure your answer: state the trend, compare at least two elements with data/observations, explain using electronic structure. Use scientific terminology precisely.
Working out group/period from electronic structure — Count total shells for period number, count electrons in the outer (highest) shell for group number. Check your answer makes sense (Groups 1-7, Period 1-7).
Quick revision summary
The periodic table organises elements by atomic number into groups (vertical) and periods (horizontal). Group number equals outer shell electrons; period number equals total electron shells. Group 1 metals become more reactive down the group as outer electrons are further from the nucleus. Group 7 halogens become less reactive down the group for the opposite reason. Group 0 noble gases are unreactive due to full outer shells. Transition metals form coloured compounds and act as catalysts.