What is the Periodic Table and why is it important in Chemistry?

The Periodic Table is a tabular arrangement of chemical elements, organized by increasing atomic number, electron configuration, and recurring chemical properties. It is crucial in Chemistry, especially for the Edexcel IGCSE curriculum, as it helps predict the properties of elements, understand chemical reactions, and provides a framework for analyzing chemical behavior.

How are elements arranged in the Periodic Table?

Elements in the Periodic Table are arranged in order of increasing atomic number, which is the number of protons in an atom's nucleus. The table is divided into periods (rows) and groups (columns). Elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell, which is a key concept in the Principles of Chemistry.

What are the main groups in the Periodic Table and their characteristics?

The main groups in the Periodic Table include Group 1 (alkali metals), Group 2 (alkaline earth metals), Group 17 (halogens), and Group 18 (noble gases). Alkali metals are highly reactive, especially with water. Alkaline earth metals are less reactive than alkali metals. Halogens are very reactive non-metals, and noble gases are inert due to their full outer electron shells. Understanding these groups is essential for Edexcel IGCSE Chemistry students.

What is the significance of periods in the Periodic Table?

Periods in the Periodic Table represent the horizontal rows. Each period corresponds to the number of electron shells an atom has. As you move across a period from left to right, the atomic number increases, and elements become less metallic and more non-metallic. This trend is important for understanding the chemical properties and reactivity of elements, a key topic in the Principles of Chemistry.

How does the Periodic Table help in predicting chemical reactions?

The Periodic Table helps predict chemical reactions by allowing chemists to understand the reactivity of elements based on their position. Elements in the same group often react similarly due to having the same number of valence electrons. For example, alkali metals react vigorously with water. This predictive power is a fundamental aspect of Chemistry and is emphasized in the Edexcel IGCSE curriculum.

Why is "Confusing 'group' (vertical) and 'period' (horizontal)" a common mistake in The Periodic Table ?

Why it happens: Both terms refer to rows/columns and are easily muddled. How to avoid it: **Groups go DOWN (vertical columns)**, periods go ACROSS (horizontal rows). Mnemonic: 'group' has the letters in 'going down'. Period number = number of shells. Source: 4CH1 Examiner Reports 2022-2024

Why is "Explaining the Group 7 trend with 'easier to lose an electron' reasoning" a common mistake in The Periodic Table ?

Why it happens: Reusing the Group 1 explanation without adapting. How to avoid it: Group 1 LOSES electrons; Group 7 GAINS electrons. For Group 7 the reasoning is: outer shell further from nucleus + more shielding → harder to ATTRACT a new electron → less reactive. Source: 4CH1 Examiner Reports 2022-2024

Why is "Saying 'reactivity increases because the metals get bigger' without explaining WHY" a common mistake in The Periodic Table ?

Why it happens: Stopping the chain of reasoning too early. How to avoid it: Spell out: more shells → outer electron FURTHER from nucleus + MORE SHIELDING → weaker attraction → easier to lose (Group 1) or harder to gain (Group 7). The keywords 'shielding' and 'attraction' are mark-bearing.

Why is "Forgetting that Group 0 (or Group 8) elements have FULL outer shells" a common mistake in The Periodic Table ?

Why it happens: Reading 'Group 0' literally as zero outer electrons. How to avoid it: Group 0 (also written Group 8) means a FULL outer shell — 2 electrons for He; 8 for Ne, Ar, Kr… The '0' refers to the number of 'extra' electrons beyond the full shell.

Why is "Listing 'transition metals are different colours' as a property" a common mistake in The Periodic Table ?

Why it happens: Confusing the metal with its compounds. How to avoid it: Transition metals themselves are usually silvery (Cu is reddish, Au is yellow — exceptions). The key property is that their COMPOUNDS / IONS in solution are coloured (Cu²⁺ blue, Fe²⁺ pale green, Fe³⁺ orange-brown, etc.).

Why is "Saying Mendeleev arranged elements by atomic NUMBER" a common mistake in The Periodic Table ?

Why it happens: Confusing the historical and modern versions. How to avoid it: Mendeleev (1869) arranged by atomic MASS. The atomic number wasn't understood until Moseley's work in 1913. The modern table uses atomic number.

Principles of ChemistryPearson Edexcel IGCSE Chemistry 4CH1 Higher

The Periodic Table

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The Periodic Table — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

How elements are arranged by atomic number into Groups (vertical, same outer electrons → similar properties) and Periods (horizontal, same number of shells). Mendeleev's early version and how he predicted undiscovered elements. Reactivity trends DOWN Group 1 (increases) and DOWN Group 7 (decreases) explained by atomic radius and shielding. Special properties of transition metals and noble gases.

At a glance

Modern PT: elements in order of increasing atomic (proton) number.
Groups = vertical columns; same outer-shell electrons → similar chemical properties.
Periods = horizontal rows; period number = number of electron shells.
Metals on the LEFT + centre; non-metals on the RIGHT, separated by a staircase line.
Group 1 reactivity INCREASES down: outer e⁻ further from nucleus + more shielding → easier to LOSE.
Group 7 reactivity DECREASES down: outer shell further + more shielding → harder to GAIN an e⁻.
Group 0 noble gases: full outer shell → very unreactive (He 2, Ne 8, Ar 8…).
Transition metals (d-block): hard, high mp, less reactive, variable charges, COLOURED compounds, catalysts.
Mendeleev (1869): arranged by atomic MASS, grouped by properties, left gaps for predicted elements (e.g. germanium).

What you’ll learn

Mapped to the Cambridge IGCSE 4CH1 syllabus (2026 onwards).

1.84 — Understand that the modern periodic table is arranged in order of increasing atomic (proton) number.
1.85 — Recall that elements in the same group have similar chemical properties because they have the same number of outer-shell electrons.
1.86 — Recall the position of metals and non-metals in the periodic table (metals on the left and centre; non-metals on the right, separated by a 'staircase' line).
1.87 — Use the position of an element in the periodic table to predict its electron configuration, ion charge, and likely properties.
1.88 — Describe and explain the reactivity trend going DOWN Group 1 (reactivity INCREASES) in terms of atomic structure and electron loss.
1.89 — Describe and explain the reactivity trend going DOWN Group 7 (reactivity DECREASES) in terms of atomic structure and electron gain. Recall the displacement reactions of halogens with halide solutions.
1.90 — Describe the unreactivity of Group 0 (noble gases) in terms of full outer electron shells.
1.91 — Identify the key properties of transition metals: high density, high melting points, less reactive than Group 1 metals, form ions with variable charges, form coloured compounds, often act as catalysts.

How the periodic table is arranged (spec 1.84-1.86)

Atomic number order; groups = same outer e⁻; periods = same number of shells.

The modern periodic table lists all known elements in order of increasing atomic (proton) number. The table is organised into a 2D grid where the position of each element communicates two pieces of information at once.

Vertical columns = GROUPS.

All elements in the same group have the SAME NUMBER OF OUTER-SHELL ELECTRONS.
Outer-shell electrons control chemical behaviour, so elements in the same group have SIMILAR CHEMICAL PROPERTIES.
For Groups 1, 2, 3 and 5, 6, 7: group number = number of outer-shell electrons.
For Group 0 (noble gases): outer shell is FULL (2 for He; 8 for Ne, Ar, Kr…).
Group 4 elements (C, Si, Ge…) bridge metallic and non-metallic behaviour.

Group	Outer e⁻	Family name	Typical ion
1	1	Alkali metals	+1
2	2	Alkaline earth metals	+2
3	3	(boron group)	+3
4	4	(carbon group)	usually share 4
5	5	(nitrogen group)	−3
6	6	(oxygen group, chalcogens)	−2
7	7	Halogens	−1
0	8 (full)	Noble gases	none

Horizontal rows = PERIODS.

The period number = the number of electron SHELLS in atoms of elements in that row.
Across a period, atomic number increases by 1 with each element — every successive element has one more proton AND one more electron, but the SAME number of shells (the outer shell fills up across the period).

Period	Number of shells	Example: outer shell
1	1	H, He (1st shell filling)
2	2	Li → Ne (2nd shell filling)
3	3	Na → Ar (3rd shell filling)
4	4	K → Kr (4th shell + transition metals)

Metal vs non-metal regions.

Metals occupy the LEFT and CENTRE of the table (Groups 1, 2, 3, transition metals).
Non-metals are on the RIGHT (most of Group 4-7 + Group 0 noble gases), above the 'staircase' line that runs from B-Si-As-Te.
Hydrogen is placed above Group 1 but behaves more like a non-metal — a special case.

The blocks (extension only). The table also has an s-block (Groups 1-2), p-block (Groups 3-0), d-block (transition metals) and f-block (lanthanides + actinides — usually shown separately at the bottom).

Groups are columns, periods are rows; metals sit left and centre, non-metals right of the staircase line.

Order by atomic (proton) number.
Groups = same outer-shell electrons = similar properties.
Periods = same number of electron shells.
Metals on left/centre; non-metals on right (staircase line).

See the full worked example for the periodic table →

Mendeleev's contribution (spec 1.84)

Arranged by atomic MASS in 1869; predicted unknown elements.

Mendeleev's 1869 periodic table is the most famous early version. Working with about 60 known elements, he made three key innovations:

1. Arranged by atomic mass. Mendeleev listed elements in order of increasing relative atomic mass (he didn't know about protons or the structure of the atom). When the mass-order brought him to an element with the wrong chemistry for that column, he KNEW something was up.

2. Left gaps for undiscovered elements. Rather than force elements into ill-fitting columns, Mendeleev left GAPS where the patterns suggested an unknown element should sit. He even predicted the properties of these missing elements based on the trends within each column:

"Eka-silicon" — atomic mass ~72, density ~5.5, forms oxide XO₂ and chloride XCl₄. Discovered 1886 → germanium (Ge). Properties matched Mendeleev's prediction within a few percent.
"Eka-aluminium" — discovered 1875 → gallium (Ga).
"Eka-boron" — discovered 1879 → scandium (Sc).

3. Switched mass-order when chemistry demanded. Mendeleev placed a few pairs of elements 'out of order' by mass when their chemistry pointed to a different position. Famous example: tellurium (Te, mass 128) placed BEFORE iodine (I, mass 127) so Te aligned with the oxygen group and I with the chlorine group.

Why the modern table differs.

Atomic NUMBER instead of mass. Henry Moseley (1913) used X-ray spectroscopy to determine the number of protons in each element. Arranging by proton number resolves the Te/I mass anomaly (Te = 52, I = 53; Te before I is correct by proton number).
Noble gases added. Mendeleev's table had no noble gases — Group 0 was discovered later (Ramsay & Rayleigh, 1894 onwards). They slotted in neatly without disrupting the structure.
More elements known. The table now has 118 elements (up to oganesson, Z = 118), with all natural elements identified and many synthetic.

Why Mendeleev's table is considered a scientific milestone. He didn't just CATALOGUE the elements — he made FALSIFIABLE PREDICTIONS that were later confirmed. The discovery of germanium with the predicted properties was one of the most striking confirmations in the history of chemistry, establishing periodicity as a real natural pattern rather than a coincidence.

Mendeleev (1869) arranged by atomic MASS.
Grouped elements with similar properties in same columns.
Left GAPS for undiscovered elements and predicted their properties.
Germanium (eka-silicon), gallium (eka-aluminium), scandium (eka-boron) confirmed predictions.
Modern PT uses atomic NUMBER (Moseley 1913) and includes noble gases.

Group 1 — the alkali metals: reactivity INCREASES down (spec 1.88)

Bigger atom + more shielding → outer e⁻ lost more easily.

Group 1 elements (alkali metals): Li, Na, K, Rb, Cs, Fr. Each has 1 outer-shell electron and reacts by LOSING that electron to form +1 ions:

$\text{M} \rightarrow \text{M}^+ + e^-$

Common Group 1 reactions.

With water: produces a metal hydroxide solution + hydrogen gas. e.g. $2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2$ . Vigour: Li fizzes; Na melts into a silvery ball and zooms across the surface; K ignites and burns with a lilac flame; Rb and Cs react explosively.
With oxygen: forms a metal oxide. e.g. $4\text{Na} + \text{O}_2 \rightarrow 2\text{Na}_2\text{O}$ .
With chlorine: forms a metal chloride. e.g. $2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$ .

Reactivity trend: increases DOWN the group. Li (least reactive) < Na < K < Rb < Cs (most reactive).

Explanation in terms of atomic structure.

Going down the group:

Each successive element has ONE MORE electron SHELL than the one above. Li has 2 shells, Na 3, K 4, etc.
So the ATOM IS LARGER going down the group (atomic radius increases).
The outer-shell electron is FURTHER from the positive nucleus → the electrostatic attraction between the nucleus and the outer electron is WEAKER.
Additionally, the outer electron is MORE SHIELDED by the increasing number of inner electrons — these inner-shell electrons repel the outer electron and screen it from the nucleus.
Both effects reduce the effective nuclear attraction on the outer electron → less energy is needed to remove it → atom loses its electron more easily → the metal is MORE REACTIVE.

Down Group 1, more shells push the outer electron away from the nucleus and shield it, so it is lost more easily.

Two competing factors (extension). Going down the group, nuclear CHARGE also increases (more protons), which on its own would tighten the grip on the outer electron. However, the effects of (a) increased distance and (b) increased shielding outweigh the (c) increased nuclear charge — so net effect = weaker attraction = greater reactivity.

Other properties of Group 1 metals.

Soft (can be cut with a knife).
Low density — Li, Na and K all float on water.
Low melting points (Cs melts at 28 °C — almost room temperature!).
Tarnish in air rapidly (must be stored under oil).
Form ionic compounds with non-metals (e.g. NaCl, K₂O).
Solutions of their hydroxides are ALKALINE (hence 'alkali metals').

Group 1: 1 outer electron; reacts by LOSING that electron → +1 ions.
Reactivity INCREASES down the group.
Reason: more shells → outer e⁻ further from nucleus + more shielding → easier to lose.
Cs reacts explosively with water; Li only fizzes.

See the full worked example for the periodic table →

Group 7 — the halogens: reactivity DECREASES down (spec 1.89)

Bigger atom + more shielding → harder to attract a new electron.

Group 7 elements (halogens): F, Cl, Br, I, At. Each has 7 outer-shell electrons and reacts by GAINING 1 electron to form −1 ions:

$\text{X} + e^- \rightarrow \text{X}^-$

Halogens also exist as diatomic molecules (F₂, Cl₂, Br₂, I₂) by sharing 1 pair of electrons between two atoms.

Physical state at room temperature (very useful to know).

Halogen	State at rt	Colour
F₂	Gas	Pale yellow
Cl₂	Gas	Pale green
Br₂	Liquid	Orange / red-brown
I₂	Solid	Dark grey / purple-black (gives purple vapour)

Reactivity trend: DECREASES down the group. F (most reactive) > Cl > Br > I.

Explanation in terms of atomic structure.

Going down the group:

Each successive halogen has ONE MORE electron SHELL than the one above. F has 2 shells, Cl 3, Br 4, I 5.
So the ATOM IS LARGER; the outer shell (where a new electron must arrive) is FURTHER from the nucleus.
The electrostatic attraction on an INCOMING electron is WEAKER (greater distance from nucleus).
Additionally, more inner-shell electrons SHIELD the outer shell from the full positive nuclear charge.
Both effects make it HARDER to attract an incoming electron → the halogen is LESS REACTIVE.

So F (smallest, least shielding) attracts an electron most strongly → most reactive. I (largest, most shielding) attracts least strongly → least reactive.

A more reactive halogen displaces a less reactive halide from solution; the reverse never works.

Halogen displacement reactions. A more reactive halogen DISPLACES a less reactive halide from its salt:

$\text{Cl}_2 + 2\text{KBr}(aq) \rightarrow 2\text{KCl}(aq) + \text{Br}_2(aq)$ (colourless → orange-brown)

$\text{Cl}_2 + 2\text{KI}(aq) \rightarrow 2\text{KCl}(aq) + \text{I}_2(aq)$ (colourless → dark brown / black)

$\text{Br}_2 + 2\text{KI}(aq) \rightarrow 2\text{KBr}(aq) + \text{I}_2(aq)$

LESS reactive halogen CANNOT displace more reactive halide:

Br₂ + KCl(aq) → NO reaction
I₂ + KCl(aq) → NO reaction
I₂ + KBr(aq) → NO reaction

Reaction with hydrogen. Halogens react with hydrogen to form hydrogen halides (HX):

F₂ + H₂ → 2HF (explosive even in dark)
Cl₂ + H₂ → 2HCl (needs light or heat)
Br₂ + H₂ → 2HBr (needs heat and catalyst)
I₂ + H₂ → 2HI (slow, reversible, incomplete)

Vigour DECREASES down the group, mirroring the reactivity trend.

Group 7: 7 outer electrons; reacts by GAINING 1 electron → −1 ions.
Reactivity DECREASES down the group.
Reason: more shells → outer shell further from nucleus + more shielding → harder to gain an e⁻.
More reactive halogen displaces less reactive halide from solution.

See the full worked example for the periodic table →

Group 0 — the noble gases (spec 1.90)

Full outer shells → very unreactive.

Group 0 elements (noble gases): He, Ne, Ar, Kr, Xe, Rn. All exist as monatomic gases at room temperature (single, unbonded atoms — NOT diatomic).

Why they are unreactive. Noble gases have a FULL OUTER ELECTRON SHELL:

Element	Electron config	Outer e⁻
He	2	2 (full 1st shell)
Ne	2, 8	8 (full 2nd shell)
Ar	2, 8, 8	8 (full 3rd shell)
Kr	2, 8, 18, 8	8 (full outer)

A full outer shell is energetically the most stable configuration. Other elements REACT in order to ATTAIN a noble-gas configuration (Group 1 metals lose, halogens gain, etc.). Noble gases already HAVE this configuration, so they have no driving force to form bonds:

They do NOT lose electrons easily (very high ionisation energy — the next electron is in a lower-energy full shell).
They do NOT gain electrons (their outer shell is already full; an extra e⁻ would have to occupy a high-energy new shell).
They do NOT share electrons (no incomplete shell to fill).

So under normal conditions noble gases exist as separate atoms and do not form compounds. (A few noble-gas compounds — XeF₂, XeF₄, KrF₂ — have been synthesised in extreme conditions, but this is beyond IGCSE.)

Noble gases already have full outer shells, so they have no tendency to react.

Uses (taking advantage of unreactivity / inertness).

Helium: weather balloons and airships (low density + non-flammable, unlike H₂); MRI scanners (very low boiling point — used as a coolant for superconductors); cylinder mixtures for deep-sea diving.
Neon: advertising signs (glows red-orange when an electric current is passed); some lasers.
Argon: filling incandescent light bulbs (prevents the tungsten filament from oxidising); shielding gas for arc welding (excludes O₂ from the hot weld).
Krypton, Xenon: specialised lasers, high-intensity lamps, ion-engine propellants for satellites.

Trends.

mp / bp INCREASE down the group (larger molecules → stronger intermolecular forces, even though they're individual atoms).
All are colourless and odourless.
All are extremely low-density gases.

Group 0 / Noble gases: He, Ne, Ar, Kr, Xe, Rn.
Full outer shell (2 for He; 8 for the rest) → very stable.
No tendency to gain, lose or share electrons → unreactive.
Uses: He in balloons/MRI; Ar in light bulbs/welding; Ne in signs.

Transition metals (spec 1.91)

Hard, high mp, less reactive, variable charges, coloured compounds, catalysts.

Transition metals occupy the central d-block of the periodic table — the row from scandium (Sc) to zinc (Zn) is the first transition series, with iron (Fe), copper (Cu) and chromium (Cr) as the most-mentioned examples in 4CH1.

Key properties (always contrasted with Group 1 metals).

Property	Transition metals	Group 1 metals
Hardness	Hard	Soft (can cut with knife)
Density	High (Fe = 7.9, Cu = 8.9 g/cm³)	Low (Na = 0.97 g/cm³)
Melting point	High (Fe 1538 °C, Cu 1085 °C)	Low (Na 98 °C, K 63 °C)
Reactivity	Much LESS reactive	Very reactive
Ion charges	VARIABLE (Fe²⁺ + Fe³⁺; Cu⁺ + Cu²⁺)	Only +1
Compounds	Often COLOURED	Usually white / colourless
Catalysts	Often yes	No

1. High density and hardness. Transition-metal atoms are smaller and pack closely, giving high density and hardness. Their metallic bonding involves not just the outer s-electrons but also some d-electrons, giving more delocalised electrons per atom → stronger metallic bond.

2. High melting points. The stronger metallic bonding (more delocalised electrons + closer packing) means more energy is needed to break the lattice. Iron melts at 1538 °C; tungsten at 3422 °C (highest of any metal — used for light-bulb filaments).

3. Less reactive than Group 1. Transition metals don't react vigorously with water (iron rusts slowly; copper doesn't react with water at all) and many react only slowly with acids (e.g. Cu won't displace H₂ from dilute HCl). This makes them suitable as structural materials, unlike Group 1 metals.

4. Variable ion charges. Iron forms BOTH Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)). Copper forms BOTH Cu⁺ (copper(I)) and Cu²⁺ (copper(II)). Manganese forms Mn²⁺, Mn³⁺, Mn⁴⁺ and Mn⁷⁺! The Roman numeral in a name (e.g. "iron(II) chloride") tells you the charge of the metal ion in that compound.

5. Coloured compounds.

Ion	Colour
Cu²⁺(aq)	Blue
Cu⁺	(often colourless or white)
Fe²⁺(aq)	Pale green
Fe³⁺(aq)	Orange / brown
Cr³⁺(aq)	Green
Mn²⁺(aq)	Pale pink
MnO₄⁻ (manganate(VII))	Deep purple
Cr₂O₇²⁻ (dichromate)	Orange

The colours come from d-electron transitions (extension content).

6. Catalysts. Many transition metals (and their compounds) act as catalysts:

Fe in the Haber process ( $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ ).
V₂O₅ (vanadium(V) oxide) in the Contact process (making H₂SO₄).
Pt and Pd in catalytic converters (in car exhausts — convert CO + NO to CO₂ + N₂).
Ni in hydrogenation of vegetable oils (margarine).

Catalysts speed up reactions WITHOUT being used up — the transition metal's ability to vary its oxidation state plays a key role.

Transition metals in central d-block (Sc-Zn first row; Fe, Cu, Cr most-tested).
Hard, dense, high mp — much more than Group 1 metals.
Much LESS reactive than Group 1; useful as structural materials.
Form ions with VARIABLE charges (e.g. Fe²⁺ + Fe³⁺).
COLOURED compounds (Cu²⁺ blue, Fe²⁺ pale green, Fe³⁺ orange-brown).
Often catalysts (Fe in Haber, V₂O₅ in Contact, Pt in catalytic converters).

See the full worked example for the periodic table →

How it’s examined

The Periodic Table is a guaranteed Paper 1 (4CH1/1C) topic — typically 6-12 marks per session. Most-common items: (1) describe how the modern PT is arranged (3-4 marks); (2) explain Group 1 or Group 7 reactivity trend in terms of atomic structure (5-6 marks); (3) properties of transition metals vs Group 1 (4 marks); (4) noble gases unreactivity + uses (4-5 marks); (5) halogen displacement reactions with observations (4-5 marks). Paper 2 (4CH1/2C) sometimes extends with Mendeleev's predictions and historical questions. Worth 8-15 marks combined across both papers per session.

Sources: Pearson Edexcel International GCSE Chemistry 4CH1 specification — Issue 4, September 2024 (valid for 2026 onwards); Pearson Edexcel 4CH1 Examiner Reports 2022-2024; 4CH1/1C May/Jun 2024 question paper and mark scheme; 4CH1/1C January 2024 question paper and mark scheme; 4CH1/1CR January 2024 question paper and mark scheme; 4CH1/2C May/Jun 2024 question paper and mark scheme. Last reviewed 2026-05-12.

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Step-by-step worked examples — The Periodic Table

Step-by-step solutions to past-paper-style questions on the periodic table , written exactly the way a tutor would explain them at the board.

1How the modern periodic table is arranged (4 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q1• periodic table, spec-1.84, spec-1.85
Question
Describe how the modern periodic table is arranged. Refer to GROUPS, PERIODS, and what determines an element's position in the table. (4 marks)
Step-by-step solution
1. Step 1
  Order of elements (1 M). Elements are arranged in order of INCREASING ATOMIC (PROTON) NUMBER from left to right and top to bottom.
2. Step 2
  Groups (1 A). A GROUP is a VERTICAL column. Elements in the same group have the same number of electrons in their outer shell (= the group number, for groups 1, 2, 3 and 5, 6, 7; group 0 = full outer shell). Same outer electrons → similar chemical properties.
3. Step 3
  Periods (1 B). A PERIOD is a HORIZONTAL row. The period number = the number of electron SHELLS the atom has. Across a period, each successive element has one more proton and one more electron, but the same number of shells.
4. Step 4
  Metals vs non-metals (1 B). Metals are on the LEFT and CENTRE of the table; non-metals are on the RIGHT, separated by a 'staircase' line. Transition metals occupy the central block between Groups 2 and 3.
Answer
Elements arranged by increasing atomic (proton) number. GROUPS = vertical columns; elements in the same group have the same number of outer-shell electrons (= group number) → similar properties. PERIODS = horizontal rows; period number = number of electron shells. Metals are on the left/centre; non-metals on the right, separated by the staircase.
Examiner tip
Edexcel 4-mark mark scheme: 1 for 'atomic / proton number'; 1 for groups defined correctly; 1 for periods defined correctly; 1 for metal / non-metal regions. Common error: confusing groups and periods. Memorise: GROUPS go DOWN (vertical), PERIODS go ACROSS (horizontal).
2Trend in reactivity DOWN Group 1 (5 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q7• group 1, reactivity trend, spec-1.88
Question
Reactivity INCREASES as you go DOWN Group 1 of the periodic table — from Li (least reactive) through Na, K, Rb, Cs (most reactive). Explain this trend in terms of atomic structure. (5 marks)
Step-by-step solution
1. Step 1
  Group 1 reaction step (1 M). Group 1 metals react by LOSING their 1 outer-shell electron to form +1 ions: $\text{M} \rightarrow \text{M}^+ + e^-$ . The more easily this electron is lost, the more reactive the metal.
2. Step 2
  Going down the group — more shells (1 A). Each successive Group 1 element has ONE MORE electron shell than the one above. Li has 2 shells, Na has 3, K has 4, etc. So the atom gets LARGER going down the group.
3. Step 3
  Outer electron further from nucleus (1 A). Because the atom is larger, the outer-shell electron is further from the positive nucleus → the electrostatic attraction between the nucleus and the outer electron is WEAKER.
4. Step 4
  More inner-shell shielding (1 B). With more inner shells, there are more inner-shell electrons SHIELDING the outer electron from the full positive charge of the nucleus. This further weakens the attraction felt by the outer electron.
5. Step 5
  Easier to lose → more reactive (1 B). Because the attraction holding the outer electron is weaker, LESS energy is needed to remove it. The atom loses its electron more easily → the metal is MORE REACTIVE. Hence reactivity increases down the group.
Answer
Group 1 metals react by losing 1 outer electron. Going down the group: more electron shells → outer electron is FURTHER from the nucleus AND more shielded by inner shells. Both effects WEAKEN the attraction between the nucleus and the outer electron → LESS energy needed to remove it → the metal loses its electron more easily → more reactive.
Examiner tip
Edexcel 5-mark mark scheme: 1 for the reaction step (losing 1 outer e⁻); 1 for more shells / larger atom; 1 for outer electron further from nucleus; 1 for more shielding; 1 for 'easier to lose / weaker attraction → more reactive'. Mark-bearing keywords: 'shielding' and 'attraction between nucleus and outer electron'. Vague answers like 'they get more reactive because they have more electrons' score 0.
3Trend in reactivity DOWN Group 7 (5 marks, Paper 1)
Core• Adapted from 4CH1/1CR Jan 2024 Q9• group 7, halogens, reactivity trend, spec-1.89
Question
Reactivity DECREASES going DOWN Group 7 (the halogens) — F (most reactive) > Cl > Br > I (least reactive). Explain this trend in terms of atomic structure. (5 marks)
Step-by-step solution
1. Step 1
  Group 7 reaction step (1 M). Group 7 elements (halogens) react by GAINING 1 electron to form −1 ions: $\text{X} + e^- \rightarrow \text{X}^-$ . The more easily an atom can attract an electron, the more reactive it is.
2. Step 2
  Going down — more shells (1 A). Each successive halogen has ONE MORE electron shell than the one above. F has 2 shells, Cl has 3, Br has 4, I has 5. So the atom gets LARGER.
3. Step 3
  Outer shell further from nucleus (1 A). The outer shell (where the new electron must be added) is FURTHER from the positive nucleus → the electrostatic attraction on an incoming electron is WEAKER.
4. Step 4
  More shielding from inner shells (1 B). More inner-shell electrons mean more shielding of the outer-shell region from the nucleus. The effective positive charge felt by the incoming electron is further reduced.
5. Step 5
  Harder to attract → less reactive (1 B). Because the attraction for an incoming electron is weaker, MORE difficult to capture a new electron → the halogen reacts less readily → LESS REACTIVE. So reactivity decreases down the group.
Answer
Halogens react by gaining 1 electron. Going down the group: more electron shells → outer shell FURTHER from the nucleus and more shielded by inner electrons. Both weaken the attraction the nucleus exerts on an incoming electron → harder to gain an electron → less reactive.
Examiner tip
Edexcel 5-mark mark scheme: 1 for the reaction step (gaining 1 electron); 1 for more shells / larger atom; 1 for outer shell further from nucleus; 1 for more shielding; 1 for 'harder to attract → less reactive'. Examiner reports note many candidates use the WRONG reasoning here — confusing the gain-an-electron mechanism with the lose-an-electron Group 1 mechanism. Make sure your explanation matches the actual reaction.
4Properties of transition metals (4 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q2• transition metals, spec-1.91
Question
Iron, copper and chromium are examples of TRANSITION METALS. Compare the properties of transition metals with the properties of Group 1 metals. State FOUR differences. (4 marks)
Step-by-step solution
1. Step 1
  Density and hardness (1 M). Transition metals are HARDER and DENSER than Group 1 metals. For example, iron has a density ~7.9 g/cm³ vs sodium ~0.97 g/cm³; you can cut sodium with a knife but not iron.
2. Step 2
  Melting points (1 A). Transition metals have MUCH HIGHER melting points. Fe melts at 1538 °C; Cu at 1085 °C; Cr at 1907 °C. Group 1 metals melt at low temperatures: Na 98 °C, K 63 °C, Cs 28 °C (almost room temperature!).
3. Step 3
  Reactivity (1 A). Transition metals are MUCH LESS REACTIVE than Group 1 metals. Group 1 metals react vigorously with water (releasing H₂ + heat) and explosively with chlorine. Iron rusts slowly; copper barely reacts at all. Transition metals can be used as structural materials, unlike Group 1 metals.
4. Step 4
  Ion charges and colour (1 B). Transition metals form ions with VARIABLE charges (e.g. Fe²⁺ AND Fe³⁺; Cu⁺ AND Cu²⁺) and often form COLOURED COMPOUNDS (Cu²⁺ = blue; Fe²⁺ = pale green; Fe³⁺ = orange/brown; Cr³⁺ = green). Group 1 metals form ONLY +1 ions and give COLOURLESS / WHITE compounds.
Answer
Differences: (1) Transition metals are HARDER and DENSER. (2) Transition metals have MUCH HIGHER melting points. (3) Transition metals are MUCH LESS REACTIVE. (4) Transition metals form ions with VARIABLE charges (e.g. Fe²⁺/Fe³⁺) and COLOURED compounds; Group 1 form only +1 ions and white/colourless compounds.
Examiner tip
Edexcel 4-mark mark scheme: 1 mark for each of FOUR distinct differences. Candidates often list 'transition metals can act as catalysts' as another property — also credited (Fe in Haber, V₂O₅ in contact, Pt in catalytic converters). Avoid lazy answers like 'transition metals are different colours' — be specific (e.g. Cu²⁺ blue, Fe³⁺ orange-brown).

Model Answers — The Periodic Table

High-scoring sample answers for the periodic table on the Pearson Edexcel IGCSE 4CH1 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question 1
4CH1/2C-style — Mendeleev and history5 marks
Dmitri Mendeleev proposed an early form of the periodic table in 1869. Describe how he arranged the elements, explain how he predicted the existence of undiscovered elements, and state how the modern periodic table differs from his. (5 marks)
Model answer
Mendeleev's arrangement. In 1869, Mendeleev arranged the known elements (about 60 at the time) in order of increasing atomic MASS (relative atomic mass), grouping together elements with similar chemical properties in the same vertical column. He noticed that the properties of elements REPEATED at regular intervals — the periodicity that gives the table its name.

Leaving gaps for undiscovered elements. Crucially, Mendeleev did NOT force every element into a slot. When the next element in the mass order had properties that did not match the column it would fall into, he left a GAP for an as-yet-undiscovered element. He even predicted the properties of these missing elements based on the trends within their column. For example:

He predicted "eka-silicon" — an element with $A_r$ ~72, density ~5.5 g/cm³, and the ability to form oxides and chlorides like silicon.

This element was later discovered (1886) and named germanium (Ge). Its actual properties matched Mendeleev's predictions extremely closely — strong evidence that his arrangement was correct.

He similarly predicted "eka-aluminium" (later gallium) and "eka-boron" (later scandium).

Switching pairs of elements. Mendeleev occasionally placed an element 'out of order' by mass when the chemistry demanded it. For example, he put tellurium (Te, $A_r$ 128) BEFORE iodine (I, $A_r$ 127) so that Te aligned with the oxygen group and I aligned with chlorine, fluorine, etc. — even though by atomic mass it should be the other way round.

The modern periodic table.

Elements are now arranged by atomic (proton) NUMBER, not by atomic mass. This was discovered in 1913 by Henry Moseley using X-ray spectroscopy and resolves the Te/I "anomaly" — by proton number, Te (52) DOES come before I (53), matching its chemistry.

The noble gases (Group 0) had not yet been discovered when Mendeleev published in 1869. They were added later (1894 onwards) without disrupting his structure.

Many more elements are known today (118 vs Mendeleev's 60-ish).

Why he is considered the founder. Mendeleev's foresight — predicting elements and their properties decades before they were discovered — established his table as a powerful scientific tool, not just a catalogue. The modern table refines but retains his core insight: elements show periodicity in their properties when arranged by an appropriate ordering principle.
Why this scores
Edexcel 5-mark mark scheme: 1 for arrangement by atomic mass; 1 for grouping similar properties; 1 for leaving gaps; 1 for predicting properties (with a specific example like Ge / eka-silicon); 1 for the modern difference (atomic number instead of mass). Naming a specific predicted element (germanium, gallium, scandium) earns the prediction mark with confidence.
Question 2
4CH1/2C-style — noble gases5 marks
Group 0 (the noble gases) are very unreactive. (a) Explain why noble gases are unreactive in terms of electron configuration. (b) Give TWO uses of noble gases that take advantage of their unreactivity. (5 marks)
Model answer
(a) Electron configuration explanation.

Noble gases have a full outer electron shell:

Helium (He, Z = 2): 2 outer electrons — full first shell.

Neon (Ne, Z = 10): 2, 8 — full second shell.

Argon (Ar, Z = 18): 2, 8, 8 — full outer shell.

Krypton (Kr, Z = 36): 2, 8, 18, 8 — full outer shell of 8.

A full outer shell is the most stable electron configuration — energetically there is no driving force for the atom to GAIN, LOSE or SHARE electrons. Other elements react in order to ATTAIN this stable noble-gas configuration; noble gases already have it.

Specifically:

They do NOT lose electrons (their ionisation energies are extremely high; the outer shell is full and the next one to lose is in a much lower-energy shell).

They do NOT gain electrons (their outer shell is full; an extra electron would have to go into a new, much-higher-energy shell).

They do NOT share electrons (no incomplete-shell driving force).

So under normal conditions, noble gases exist as single, unbonded atoms (monatomic) and do not form compounds with other elements. (Some noble-gas compounds — e.g. XeF₂, XeO₃, KrF₂ — have been synthesised under extreme conditions, but this is outside the IGCSE syllabus.)

(b) Two uses based on unreactivity.

Helium in airships and weather balloons. Helium has a very low density (lighter than air) AND is non-flammable (unlike hydrogen, which is also low-density but burns explosively in O₂). So helium is the safe choice for lifting balloons. The 1937 Hindenburg disaster — which used H₂ — is a famous lesson in why unreactive gases are preferred.

Argon in light bulbs. Argon (and sometimes other noble gases) fills the inside of incandescent light bulbs. The hot tungsten filament would react quickly with oxygen if air were inside the bulb — burning out almost instantly. Argon's inertness lets the filament glow without reacting, extending the bulb's life.

Other valid uses (any one credited): neon in advertising signs (glows orange when an electric current is passed); helium in MRI scanners and superconductors (used as a coolant — very low boiling point); argon shielding gas in arc welding to prevent the hot weld metal from oxidising.
Why this scores
Edexcel 5-mark mark scheme: 3 marks for the unreactivity explanation (full outer shell / stable / no tendency to gain, lose or share electrons); 2 marks for two uses with justification (1 each). Most-frequent error: candidates saying noble gases are 'inert because they don't bond' — that's just restating the question. The explanation requires the full-outer-shell reasoning.

Question 3

4CH1/2C-style — halogen displacement6 marks

Chlorine is added to a solution of potassium bromide. A reaction occurs: $\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$ (a) Explain why this reaction occurs. (b) Predict what would happen if BROMINE was added to a solution of potassium chloride instead, justifying your answer. (c) State and explain what is observed if CHLORINE is added to a solution of potassium iodide. (6 marks)

Model answer

(a) Why chlorine displaces bromide. Chlorine is HIGHER in Group 7 than bromine, so chlorine is MORE REACTIVE than bromine. More reactive halogen DISPLACES less reactive halogen from its salt. Mechanism:

Chlorine atoms attract electrons more strongly than bromine atoms (smaller atom, less shielding, stronger attraction for an incoming electron).
Each Cl₂ molecule gains 2 electrons (one per atom) from 2 Br⁻ ions: $\text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^-$ .
The Br⁻ ions lose their electrons (oxidation): $2\text{Br}^- \rightarrow \text{Br}_2 + 2e^-$ .

Net: Cl₂ is REDUCED to Cl⁻ (it gains electrons); 2Br⁻ are OXIDISED to Br₂. K⁺ is a spectator ion (unchanged).

Observation. The colourless KBr(aq) turns orange-brown as Br₂ is liberated.

(b) Bromine + potassium chloride. Bromine is LESS reactive than chlorine, so Br₂ CANNOT displace Cl⁻ from KCl. Therefore NO reaction would occur — the solution would just remain a mixture of bromine water and KCl(aq). The colour would be the colour of the bromine (orange) but no NEW species would form.

Justification. Displacement reactions follow the rule "MORE reactive halogen displaces LESS reactive halogen". Since Br is less reactive than Cl, it cannot displace Cl⁻ from KCl.

(c) Chlorine + potassium iodide. Chlorine is MORE reactive than iodine, so chlorine DISPLACES iodide:

$\text{Cl}_2 + 2\text{KI} \rightarrow 2\text{KCl} + \text{I}_2$

Observation. The colourless KI(aq) turns to a dark BROWN solution (or a black/brown solid I₂ may precipitate out if there is little water) as iodine is liberated. The reaction is more vigorous than the chlorine + bromide case because the reactivity gap between Cl and I is bigger than between Cl and Br.

Summary table.

Reactants	Reaction?	Observation
Cl₂ + KBr(aq)	YES	colourless → orange (Br₂ produced)
Br₂ + KCl(aq)	NO	no change in colour (no reaction)
Cl₂ + KI(aq)	YES	colourless → dark brown (I₂ produced)
Br₂ + KI(aq)	YES	orange → dark brown
I₂ + KCl(aq)	NO	no change
I₂ + KBr(aq)	NO	no change

Why this scores

Edexcel 6-mark mark scheme: 2 marks for part (a) — reactivity-based explanation + electron-transfer detail; 2 marks for part (b) — 'no reaction' + correct justification (Br less reactive than Cl); 2 marks for part (c) — predicting I₂ formed and stating the colour change. Specific colours expected: Cl₂ pale green; Br₂ orange/brown; I₂ dark brown / black solid. Examiner reports flag that candidates often confuse Br₂ orange-brown with I₂ dark-brown — both are similar shades.

Question 4
4CH1/2C-style — using the periodic table to predict properties5 marks
Element X has the electron configuration 2, 8, 7. Predict and justify (a) the GROUP and PERIOD it is in; (b) whether it is a metal or non-metal; (c) the formula of its ion; (d) ONE typical chemical property; (e) the formula of the compound formed between X and sodium. (5 marks)
Model answer
Element analysis.

Electron configuration 2, 8, 7 — outer-shell electrons = 7. Total electrons = 17 → atomic number Z = 17 → the element is chlorine (Cl).

(a) Group and period.

Group: 7 outer-shell electrons → Group 7 (the halogens).

Period: 3 electron shells → Period 3.

(b) Metal or non-metal. Cl is in Group 7, on the RIGHT-HAND side of the periodic table (above the staircase line). Therefore it is a NON-METAL.

(c) Formula of its ion. Cl has 7 outer electrons and needs to GAIN 1 to reach the argon configuration (2, 8, 8). Gain of 1 electron gives a charge of −1. So the ion is Cl⁻ (chloride).

(d) One typical chemical property. Chlorine reacts with metals to form metal chlorides (ionic compounds), e.g. $2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$ . It is a strong oxidising agent (gains electrons from other species). It also reacts with hydrogen to form HCl gas (covalent). Any one of: forms acidic compounds when dissolved in water (HCl), bleaches dyes, displaces less reactive halogens (Br₂, I₂) from their salts.

(e) Formula with sodium. Sodium is in Group 1 and forms Na⁺ ions. To balance charges:

1 × Na⁺ (+1) and 1 × Cl⁻ (−1) → NaCl (sodium chloride).
Why this scores
Edexcel 5-mark mark scheme: 1 each for (a) group AND period correctly stated; (b) non-metal (with justification — right side of table or 'gains electrons'); (c) Cl⁻; (d) any one valid chemical property; (e) NaCl formula. Tests the ability to extract multiple types of information from electron configuration / position in the table.

Key Definitions and Keywords — The Periodic Table

Definitions to memorise and the exact keywords mark schemes credit for the periodic table answers — sharpened from recent examiner reports for the 2026 Pearson Edexcel IGCSE 4CH1 sitting.

Periodic Table
Examiner keyword
A tabular arrangement of all known elements in order of increasing atomic (proton) number, grouped vertically by similar chemical properties and arranged horizontally so that the period number equals the number of electron shells.
Atomic number (proton number)
Examiner keyword
The number of protons in the nucleus of an atom of an element. The atomic number is unique to each element and is the basis for the modern periodic table's ordering.
Group
Examiner keyword
A VERTICAL column in the periodic table. Elements in the same group have the same number of OUTER-SHELL electrons (equal to the group number for Groups 1, 2, 3 and 5, 6, 7), giving them similar chemical properties.
Period
Examiner keyword
A HORIZONTAL row in the periodic table. The period number equals the number of electron SHELLS in atoms of elements in that row. Across a period, atomic number increases by 1 and an outer electron is added.
Alkali metals (Group 1)
Examiner keyword
The Group 1 elements (Li, Na, K, Rb, Cs, Fr). Soft, low-density, low-melting-point metals. Each has 1 outer electron → forms +1 ions. Very reactive; reactivity INCREASES down the group.
Halogens (Group 7)
Examiner keyword
The Group 7 elements (F, Cl, Br, I, At). Diatomic non-metals (F₂, Cl₂, Br₂, I₂). Each has 7 outer electrons → forms −1 ions. Reactivity DECREASES down the group.
Noble gases (Group 0)
Examiner keyword
The Group 0 elements (He, Ne, Ar, Kr, Xe, Rn). Monatomic, very unreactive gases. Have a full outer shell (2 for He, 8 for the others) → no tendency to gain, lose or share electrons.
Transition metals
Examiner keyword
The metals in the central d-block of the periodic table (Sc through Zn in the first row). Hard, high-density, high-melting metals. Form ions with VARIABLE charges (e.g. Fe²⁺ and Fe³⁺) and COLOURED compounds. Often act as catalysts.
Electron shielding
Examiner keyword
The reduction in attraction felt by an outer-shell electron from the positive nucleus, caused by the repulsion / screening effect of inner-shell electrons. More inner shells = more shielding = weaker effective attraction.
Mendeleev's periodic table
The 1869 forerunner of the modern table. Mendeleev arranged elements by atomic MASS (not atomic number), grouped them by similar properties, and left GAPS for undiscovered elements whose properties he successfully predicted (e.g. germanium = eka-silicon).

Common Mistakes and Misconceptions — The Periodic Table

The traps other students keep falling into on the periodic table questions — taken from recent Pearson Edexcel IGCSE 4CH1 examiner reports and mark schemes — and how to avoid them.

✕Confusing 'group' (vertical) and 'period' (horizontal)
4CH1 Examiner Reports 2022-2024
Why it happens
Both terms refer to rows/columns and are easily muddled.
How to avoid it
Groups go DOWN (vertical columns), periods go ACROSS (horizontal rows). Mnemonic: 'group' has the letters in 'going down'. Period number = number of shells.
✕Explaining the Group 7 trend with 'easier to lose an electron' reasoning
4CH1 Examiner Reports 2022-2024
Why it happens
Reusing the Group 1 explanation without adapting.
How to avoid it
Group 1 LOSES electrons; Group 7 GAINS electrons. For Group 7 the reasoning is: outer shell further from nucleus + more shielding → harder to ATTRACT a new electron → less reactive.
✕Saying 'reactivity increases because the metals get bigger' without explaining WHY
Why it happens
Stopping the chain of reasoning too early.
How to avoid it
Spell out: more shells → outer electron FURTHER from nucleus + MORE SHIELDING → weaker attraction → easier to lose (Group 1) or harder to gain (Group 7). The keywords 'shielding' and 'attraction' are mark-bearing.
✕Forgetting that Group 0 (or Group 8) elements have FULL outer shells
Why it happens
Reading 'Group 0' literally as zero outer electrons.
How to avoid it
Group 0 (also written Group 8) means a FULL outer shell — 2 electrons for He; 8 for Ne, Ar, Kr… The '0' refers to the number of 'extra' electrons beyond the full shell.
✕Listing 'transition metals are different colours' as a property
Why it happens
Confusing the metal with its compounds.
How to avoid it
Transition metals themselves are usually silvery (Cu is reddish, Au is yellow — exceptions). The key property is that their COMPOUNDS / IONS in solution are coloured (Cu²⁺ blue, Fe²⁺ pale green, Fe³⁺ orange-brown, etc.).
✕Saying Mendeleev arranged elements by atomic NUMBER
Why it happens
Confusing the historical and modern versions.
How to avoid it
Mendeleev (1869) arranged by atomic MASS. The atomic number wasn't understood until Moseley's work in 1913. The modern table uses atomic number.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

Get a report showing which sub-topics you've nailed and which ones still need work.

Video lesson

Short walkthrough of the concepts students most often get stuck on.

The Periodic Table — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

The Periodic Table

Short Study Notes in the form of Slides

Short Study Notes — The Periodic Table

Detailed Notes

What you’ll learn

How it’s examined

1How the modern periodic table is arranged (4 marks, Paper 1)

2Trend in reactivity DOWN Group 1 (5 marks, Paper 1)

3Trend in reactivity DOWN Group 7 (5 marks, Paper 1)

4Properties of transition metals (4 marks, Paper 1)

Periodic Table

Atomic number (proton number)

Group

Period

Alkali metals (Group 1)

Halogens (Group 7)

Noble gases (Group 0)

Transition metals

Electron shielding

Mendeleev's periodic table

✕Confusing 'group' (vertical) and 'period' (horizontal)

✕Explaining the Group 7 trend with 'easier to lose an electron' reasoning

✕Saying 'reactivity increases because the metals get bigger' without explaining WHY

✕Forgetting that Group 0 (or Group 8) elements have FULL outer shells

✕Listing 'transition metals are different colours' as a property

✕Saying Mendeleev arranged elements by atomic NUMBER

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

The Periodic Table — frequently asked questions