IGCSE Chemistry Topic 8: The Periodic Table

8.1 Arrangement 8.2 Group I 8.3 Group VII 8.4 Transition & Noble Gases

Hey Tara! Welcome to Topic 8 — the Periodic Table. Here is the thing that most students never quite realise: this topic is almost entirely Topic 2 in disguise. Every trend, every reaction, every prediction comes back to electronic configuration — how many electrons sit in the outer shell, and how far that shell is from the nucleus. Learn to read those two numbers off the table and you can answer questions about elements that were not even discovered when the table was drawn. In this guide you will meet the alkali metals that get more violent as you go down, the halogens that get gentler as you go down (yes, the opposite way — that catches out thousands of candidates every session), the transition elements that behave nothing like the metals on the far left, and the noble gases that refuse to react with anything at all. Four subtopics, eighty practice questions, and one very powerful map. Let us get started!

8.1 Arrangement of Elements ▼

The Big Idea: The Table Is Sorted by Electrons

In 1869 the Russian chemist Dmitri Mendeleev laid the known elements out in order of atomic mass and noticed that similar properties repeated at regular intervals — periodically. He was so confident in the pattern that he left gaps for elements nobody had found yet, and predicted their properties. When gallium and germanium were discovered a few years later and matched his predictions almost exactly, the table became the most powerful single idea in chemistry.

The modern table improves on his in one crucial way. Elements are arranged in order of increasing proton number (also called atomic number), not atomic mass. Proton number is what actually defines an element, and it fixes the number of electrons in a neutral atom — which is what really controls chemical behaviour.

Elements are arranged in order of increasing proton number

Period — a horizontal row. There are seven of them. Group — a vertical column. Numbered I, II, III, IV, V, VI, VII and VIII (or 0) in Roman numerals at IGCSE. Elements in the same group have similar chemical properties, because they have the same number of outer-shell electrons.

The Two Rules That Unlock Everything

These two sentences are worth more marks in Topic 8 than anything else on this page. Learn them word for word.

Group number = number of outer-shell electrons

Group I ⇒ 1 outer electron. Group VII ⇒ 7 outer electrons. Group IV ⇒ 4 outer electrons. The exception: Group VIII (the noble gases) have a full outer shell — 8 electrons, except helium which has only 2. That is why the group is often labelled 0 instead.

Period number = number of occupied electron shells

Period 2 ⇒ 2 occupied shells. Period 3 ⇒ 3 occupied shells. Period 4 ⇒ 4 occupied shells. Sodium is 2,8,1 — three shells occupied and one electron in the outer one ⇒ Period 3, Group I.

Memory Trick

"Groups go down, periods go across." The letters help: Period is like a sentence — you read it across the Page. Groups hang down like a bunch of grapes. And for the numbers: Group = outer electrons (the last number in the configuration), Period = number of shells (how many numbers there are). Configuration 2,8,7 has three numbers (Period 3) ending in 7 (Group VII) ⇒ chlorine.

A simplified Periodic Table. The white staircase is the metal / non-metal divide — elements touching it (B, Si, Ge, As, Sb, Te) are metalloids with in-between properties.

Metals on the Left, Non-Metals on the Right

Draw a rough staircase starting between boron and carbon and stepping down to the right. Everything to the left of it is a metal; everything to the right is a non-metal. Roughly 80% of the elements are metals, which is why the metallic block is so much bigger.

Property	Metals (left)	Non-metals (right)
Appearance	Shiny (lustrous) when fresh	Dull; many are gases
Conduction of heat and electricity	Good conductors	Poor conductors (graphite is the famous exception)
Malleable / ductile?	Yes — bend and stretch without breaking	Brittle when solid
Melting and boiling points	Usually high	Usually low
Density	Usually high	Usually low
Behaviour of the atoms	Lose electrons → positive ions	Gain or share electrons → negative ions
Oxide formed	Basic (e.g. Na₂O, MgO)	Acidic (e.g. SO₂, CO₂)

That last row is a direct bridge back to Topic 7. If an oxide dissolves to give a solution of pH 12, the element is a metal on the left-hand side; if it gives pH 3, the element is a non-metal on the right. Cambridge sets this crossover question regularly.

Exam Tip

Metallic character increases going down a group and decreases going across a period from left to right. So in Group IV, carbon at the top is a non-metal, silicon and germanium are metalloids, and tin and lead at the bottom are proper metals. If a question asks "which element in Group IV is most metallic?", the answer is always the lowest one.

Why Elements in a Group Behave Alike

Chemical reactions involve only the outer-shell electrons. Two atoms with the same number of outer electrons will therefore react in the same way, even if their inner shells are completely different sizes. Lithium (2,1), sodium (2,8,1) and potassium (2,8,8,1) all have exactly one outer electron, so all three lose that one electron to form a 1+ ion and all three react with water to give a hydroxide and hydrogen.

The properties are similar, not identical. Going down a group the atoms get larger, so the outer electron behaves slightly differently — and that gradual change is what we call a trend. Sections 8.2 and 8.3 are entirely about explaining two of those trends.

Supplement

Predicting Properties from Position

Extended candidates are expected to predict the properties of an unfamiliar element from its position. The method is always the same three steps:

Locate it. Which group and which period? That gives you outer electrons and number of shells.
Compare it with its neighbours above and below. Its properties should sit between theirs, following the trend.
Quote the trend as your reason. A prediction without a reason usually scores half marks.

Example: francium sits below caesium in Group I. Predict its reaction with water. Answer: extremely violent — an explosion — because reactivity increases down Group I; its outer electron is furthest from the nucleus with the most shielding, so it is lost most easily of all.

You can also work backwards. If an element is a soft, silvery solid that floats on water and fizzes to give an alkaline solution, it must be in Group I. If a solid sublimes to a violet vapour and forms a compound of formula NaX, it is in Group VII.

Ion Charges Straight from the Group Number

Because the group number gives the outer electrons, it also gives the charge on the ion — and therefore the formula of the compound. This is the fastest formula-writing shortcut in the whole syllabus.

Group	Outer electrons	Electrons lost or gained	Ion charge	Example compound with the opposite type
I	1	Loses 1	1+	NaCl, K₂O
II	2	Loses 2	2+	MgCl₂, CaO
III	3	Loses 3	3+	AlCl₃, Al₂O₃
IV	4	Usually shares (covalent)	—	CO₂, SiO₂, CH₄
V	5	Gains 3	3−	Mg₃N₂, NH₃
VI	6	Gains 2	2−	Na₂O, MgS
VII	7	Gains 1	1−	NaCl, MgBr₂
VIII / 0	8 (He: 2)	Neither — already full	No ions	No common compounds

Memory Trick

For Groups V, VI and VII, the charge is 8 minus the group number, made negative. Group VII: 8 − 7 = 1, so 1−. Group VI: 8 − 6 = 2, so 2−. Group V: 8 − 5 = 3, so 3−. For Groups I, II and III the charge is simply the group number, made positive.

Worked Example 1 An element X has the electronic configuration 2,8,7. State its group and period, say whether it is a metal or a non-metal, give the charge on its ion, and write the formula of the compound it forms with calcium.

Step 1: Count the shells for the period

The configuration has three numbers, so three shells are occupied ⇒ Period 3.

Step 2: Read the last number for the group

The outer shell contains 7 electrons ⇒ Group VII. Period 3, Group VII is chlorine.

Step 3: Metal or non-metal?

Group VII sits on the far right, beyond the staircase ⇒ non-metal. It is close to a full shell, so it gains one electron rather than losing seven.

Step 4: Charge and formula

Gaining one electron gives a 1− ion, X⁻. Calcium is in Group II, so it forms Ca²⁺. To balance 2+ against 1− you need two X ions: CaX₂ (in real terms, CaCl₂).

Period 3, Group VII; a non-metal; forms a 1− ion; the compound with calcium is CaX₂. The element is chlorine.

Worked Example 2 Element Q has proton number 20. Without using a data book, deduce its electronic configuration, its position in the Periodic Table, the formula of its oxide, and predict whether that oxide will be acidic or basic.

Step 1: Build the configuration

Twenty protons means twenty electrons in the neutral atom. Fill the shells 2, then 8, then 8, then the rest: 2 + 8 + 8 = 18, leaving 2 over. Configuration = 2,8,8,2.

Step 2: Read off the position

Four shells ⇒ Period 4. Two outer electrons ⇒ Group II. That is calcium.

Step 3: The oxide

Group II forms 2+ ions; oxygen is Group VI and forms O²⁻. Charges cancel one-to-one, so the formula is QO (CaO).

Step 4: Acidic or basic?

Q is on the far left of the table, so it is a metal, and metal oxides are basic. Calcium oxide is indeed basic — it dissolves slightly in water to give limewater, an alkali of about pH 12.

2,8,8,2; Period 4, Group II; oxide formula QO; the oxide is basic because Q is a metal.

Worked Example 3 Mendeleev left a gap below silicon and predicted the properties of the missing element, which he called "eka-silicon". It was later found and named germanium. Using only the Periodic Table, predict two physical properties of germanium and the formula of its chloride, explaining your reasoning.

Step 1: Locate the element and its neighbours

Germanium is in Group IV, Period 4. Directly above it is silicon (a metalloid, melting point 1410 °C, density 2.3 g/cm³); directly below is tin (a metal, melting point 232 °C, density 7.3 g/cm³).

Step 2: Interpolate between the neighbours

Germanium's properties should lie between those of silicon and tin. Predicted density: somewhere around 5 g/cm³ (the true value is 5.3). Predicted melting point: between 232 and 1410 °C (the true value is 938 °C). Predicted appearance: a hard, greyish, brittle solid with some metallic lustre — a metalloid, because it sits on the staircase.

Step 3: The chloride

Group IV has four outer electrons, so it shares four of them in covalent bonds — exactly as carbon does in CCl₄ and silicon in SiCl₄. Formula: GeCl₄.

Germanium is predicted to be a hard grey metalloid of density about 5 g/cm³ and melting point several hundred degrees, forming GeCl₄ — because a group is a family whose properties change gradually, so an element's properties sit between those of the elements above and below it.

Exam Tips for 8.1

1. Say "proton number", not "atomic mass". The modern table is ordered by increasing proton number. Writing "in order of atomic mass" describes Mendeleev's table, not today's, and loses the mark.

2. Never mix up group and period. Group = column = outer electrons. Period = row = number of shells. Configuration 2,8,3 is Period 3, Group III — not Period 3, Group 3 electrons in some other sense. Write the two words out in full in your answer.

3. Answer "why do they have similar properties?" with electrons. The full-mark sentence is: "They have the same number of electrons in the outer shell." "They are in the same group" simply restates the question and earns nothing.

4. Use Roman numerals for groups at IGCSE (I, II, VII, VIII), not 1, 2, 17, 18. Both are usually accepted, but the paper uses Roman numerals and matching them avoids confusion.

5. The staircase is not a straight line. Aluminium is a metal even though it is quite far right; silicon and germanium are metalloids. Learn the six metalloids: B, Si, Ge, As, Sb, Te.

6. Metal oxides are basic, non-metal oxides are acidic. This links 8.1 straight to Topic 7 and turns a pH value into a Periodic Table position.

7. Predictions need reasons. "Rubidium is more reactive than potassium" is one mark at best. Add "because reactivity increases down Group I" and then the electron explanation to get the rest.

8. Transition elements sit between Groups II and III. They are not given Roman-numeral group numbers at IGCSE, so never say "iron is in Group VIII".

🌎 Apply It: Real-World Chemistry

The table on your classroom wall decides which metal goes in a battery, which gas fills a light bulb, and which element a mining company spends a billion rupees looking for.

In 1871 Mendeleev left a gap in his table between silicon and tin, and published a detailed prediction of the missing element's density, melting point and oxide formula. Fifteen years later, Clemens Winkler in Freiberg isolated germanium from a silver ore and found the density was 5.35 g/cm³ against a prediction of 5.5.

Why could Mendeleev make such an accurate prediction about an element that nobody had ever seen?

▼

The Reasoning

The table is not a list — it is a set of families whose properties change gradually and predictably. Any element in the gap had to be in Group IV, so it must have four outer electrons and form a chloride of formula XCl₄ and an oxide XO₂. Its density and melting point had to lie roughly halfway between those of silicon above and tin below.

Why It Convinced Everyone

Predicting something that already exists proves nothing; predicting something nobody has found is a real test. When gallium (1875), scandium (1879) and germanium (1886) all appeared with the predicted properties, chemists accepted that the table reflected a genuine physical pattern rather than an accident of arrangement.

Chemistry Connection

Mendeleev did not know about electrons — they were not discovered until 1897. He was seeing the consequences of electron configuration without knowing the cause. Today you have the better tool: you can explain why Group IV forms XCl₄ (four outer electrons, four shared pairs), which is exactly what Extended papers reward.

A materials engineer at a solar panel plant near Bengaluru works with silicon, and adds tiny traces of boron (Group III) or phosphorus (Group V) to it. Adding boron creates a shortage of electrons in the crystal; adding phosphorus creates a surplus.

Use group numbers to explain why boron creates a shortage and phosphorus a surplus.

▼

Counting the Outer Electrons

Silicon is Group IV: 4 outer electrons, each shared with a neighbouring silicon atom, so every bond in the crystal is complete. Boron is Group III: only 3 outer electrons, so one bond is left one electron short. Phosphorus is Group V: 5 outer electrons, one more than the crystal has room for.

The Consequence

The missing electron behaves as a mobile positive "hole"; the spare electron is free to move. Either way the silicon can now conduct, which is what makes the solar cell work. The whole semiconductor industry runs on the difference between Group III, IV and V.

Chemistry Connection

Nothing in this scenario requires you to know anything about boron or phosphorus except which column they are in. That is exactly what the Periodic Table is for — the group number is the number of outer electrons, and the number of outer electrons is the behaviour.

A geologist surveying a lithium prospect in the Salar de Atacama, Chile, is asked by a journalist why lithium and sodium turn up in the same kinds of salty brine deposits, and why nobody ever finds a lump of native lithium metal lying on the ground the way they find native gold.

Explain both observations using the element's position in the Periodic Table.

▼

Why They Occur Together

Lithium and sodium are both in Group I, so both form 1+ ions of similar chemistry. Their compounds are almost all soluble, so both stay dissolved in water and become concentrated together wherever a lake evaporates. Chemically similar elements end up in chemically similar places.

Why Never Found as the Metal

Group I metals are extremely reactive — they lose that single outer electron to water and to oxygen within seconds. Any native lithium would have reacted away long ago. Gold sits far to the right of the transition block and is famously unreactive, so it survives as the element.

Chemistry Connection

Where an element is found in nature is decided by its Periodic Table position. Unreactive elements (gold, platinum, the noble gases) are found uncombined; reactive ones (Group I and II metals) are only ever found as compounds, which is why extracting them needs electrolysis rather than a furnace.

A chemistry teacher in Manchester shows two unlabelled oxides. Oxide A is a white solid that dissolves to give a solution of pH 12 which conducts electricity well. Oxide B is a colourless gas that dissolves to give pH 3. She asks the class to say which side of the Periodic Table each element comes from.

Deduce the answer and say what else you can conclude about each element.

▼

Oxide A

pH 12 means the oxide is basic, and basic oxides come from metals, which sit on the left of the staircase. The good conductivity of the solution confirms an ionic compound — another metal signature. A likely candidate is sodium oxide or calcium oxide.

Oxide B

pH 3 means an acidic oxide, which comes from a non-metal on the right. Being a gas at room temperature points to small simple molecules with weak forces between them — sulfur dioxide or carbon dioxide.

Chemistry Connection

This one experiment ties together three topics: bonding (ionic versus simple molecular), acids and bases (Topic 7), and Periodic Table position (Topic 8). Examiners love this style of question precisely because it cannot be answered by memorising one fact.

In 2016 four new elements were officially named: nihonium (113), moscovium (115), tennessine (117) and oganesson (118). None of them exists for more than a fraction of a second, and no chemist has ever performed a normal test-tube reaction with any of them. Yet the naming committee placed tennessine in Group VII and oganesson in Group VIII without hesitation.

How can chemists be confident about the position of an element that has barely existed?

▼

Position Comes from Proton Number Alone

The table is ordered by proton number, and the number of protons is measured directly from the nuclear reaction that created the atom. Element 117 must follow element 116 and precede 118, which places it in Group VII — no chemistry required.

What Follows from the Position

Once it is in Group VII, the configuration must end in 7 outer electrons, so tennessine is predicted to form 1− ions and compounds such as NaTs. And because reactivity decreases down Group VII, it should be the least reactive halogen of all — so unreactive that some chemists argue it will behave more like a metal.

Chemistry Connection

This is Mendeleev's trick, still working 150 years later. Notice how the two group trends point in opposite directions: going down, Group I elements get more reactive and Group VII elements get less reactive. Getting that pair the right way round is the single most valuable thing you can take from Topic 8.

Practice Questions: 8.1

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

In the modern Periodic Table, the elements are arranged in order of increasing

A relative atomic mass

B proton number

C number of neutrons

D density

Mendeleev used atomic mass; the modern table uses proton number, which is what actually defines an element and fixes its electron arrangement.

Question 2

An element has the electronic configuration 2,8,5. In which group and period is it?

A Group III, Period 5

B Group V, Period 2

C Group V, Period 3

D Group VIII, Period 3

Three shells ⇒ Period 3. Five outer electrons ⇒ Group V. The element is phosphorus. Option A swaps the two numbers over — the classic error.

Question 3

Why do all the elements in Group II have similar chemical properties?

A They have the same number of shells

B They have the same number of electrons in the outer shell

C They have the same number of protons

D They have the same relative atomic mass

Chemical reactions involve only outer-shell electrons, so elements with the same number of them behave the same way. Same shells would mean the same period, not group.

Question 4

Which element is a non-metal?

A Aluminium

B Calcium

C Sulfur

D Iron

Sulfur is in Group VI, on the right of the staircase. Aluminium is a metal despite being in Group III, calcium is a Group II metal and iron is a transition metal.

Question 5

An element forms an ion with a charge of 3−. In which group is the element?

A Group III

B Group IV

C Group V

D Group VII

A negative ion means electrons were gained. 8 − 5 = 3, so Group V gains three electrons to form a 3− ion (for example N³⁻). Group III forms a 3+ ion.

Question 6

Which statement about the Periodic Table is correct?

A A period is a vertical column

B A group is a vertical column of elements with similar properties

C Metals are found on the right-hand side

D The noble gases have one outer-shell electron

Groups are columns and periods are rows. Metals are on the left, and noble gases have a full outer shell (8, or 2 for helium).

Question 7

Element Z is in Period 4 and Group VI. What is its electronic configuration?

A 2,8,6

B 2,8,8,4

C 2,8,18,6

D 2,6,8,8

Period 4 means four occupied shells and Group VI means six outer electrons. Selenium is 2,8,18,6. Option A is only three shells (that is sulfur, Period 3).

Question 8

The oxide of element J dissolves in water to give a solution of pH 2. Element J is most likely to be

A a Group I metal

B a transition element

C a non-metal on the right of the table

D a noble gas

Acidic oxides come from non-metals. Group I and transition metals give basic oxides, and noble gases do not form oxides at all.

Question 9

Which pair of elements would you expect to have the most similar chemical properties?

A Sodium and magnesium

B Sodium and potassium

C Sodium and chlorine

D Sodium and argon

Sodium and potassium are both in Group I, so both have one outer electron. Magnesium is next door in Group II, chlorine is a non-metal and argon is inert.

Question 10

Element X is in Group II and element Y is in Group VII. What is the formula of the compound they form?

A XY

B XY₂

C X₂Y

D X₂Y₃

X forms X²⁺ and Y forms Y⁻. Two 1− ions are needed to balance one 2+ ion, giving XY₂ — the pattern of MgCl₂ and CaBr₂.

Question 11

Which property is typical of a non-metal?

A Good conductor of electricity

B Malleable

C Brittle when solid

D Forms basic oxides

Non-metal solids shatter rather than bend. The other three options are all metallic properties.

Question 12

Where are the transition elements found in the Periodic Table?

A To the right of Group VII

B Between Group II and Group III

C In Period 2 only

D Below the main body of the table

The transition block sits between Groups II and III, starting in Period 4 with scandium. The two rows printed below the table are the lanthanoids and actinoids, not the transition block.

Question 13

An atom has 15 protons. Which statement is correct?

A It is in Period 5

B It has 5 outer-shell electrons

C It forms a 5+ ion

D It is a metal

15 electrons fill as 2,8,5 ⇒ Period 3, Group V, five outer electrons. It is phosphorus, a non-metal that gains three electrons to form P³⁻ rather than losing five.

Question 14

Going from left to right across Period 3, the elements change from

A metallic to non-metallic

B non-metallic to metallic

C gases to metals

D acidic to basic in their oxides

Period 3 runs Na, Mg, Al (metals) → Si (metalloid) → P, S, Cl, Ar (non-metals). The oxides change the other way round from option D: basic → amphoteric → acidic.

Question 15

Mendeleev left gaps in his Periodic Table. Why?

A He had run out of paper

B He predicted that elements with those properties existed but had not yet been discovered

C He knew about the noble gases but could not place them

D He wanted to leave space for the transition elements

Rather than force a known element into the wrong family, Mendeleev left the space empty and predicted the missing element's properties. Gallium, scandium and germanium later filled three of those gaps.

Question 16

Which element is a metalloid, showing properties of both metals and non-metals?

A Sodium

B Silicon

C Sulfur

D Argon

Silicon sits on the staircase. It looks metallic and conducts a little, but it is brittle and forms an acidic oxide — hence "semiconductor".

Question 17

Element R is directly below element S in the same group. Which prediction is most reliable?

A R and S have the same number of shells

B R and S form ions with the same charge

C R and S have the same melting point

D R has fewer protons than S

Same group means the same number of outer electrons, so the same ion charge. R has one more shell and more protons than S, and physical properties change gradually rather than staying equal.

Question 18

An element forms an oxide of formula X₂O₃ and a chloride of formula XCl₃. In which group is X?

A Group II

B Group III

C Group V

D Group VI

One X balances three Cl⁻ ions, so X is 3+ ⇒ three outer electrons lost ⇒ Group III. Aluminium fits exactly: Al₂O₃ and AlCl₃.

Question 19

Which statement about metallic character is correct?

A It increases from left to right across a period

B It increases going down a group

C All elements in Group IV are non-metals

D Metalloids are found only in Group I

Group IV runs from carbon (non-metal) through silicon and germanium (metalloids) to tin and lead (metals) — metallic character rises going down. Across a period it falls.

Question 20

A newly made element has proton number 119. What can be predicted with most confidence?

A It will be a gas at room temperature

B It will be in Group I and form a 1+ ion

C It will be a non-metal with 7 outer electrons

D It will be chemically unreactive

Element 118 (oganesson) completes a noble-gas row, so element 119 must start the next period in Group I — one outer electron, a 1+ ion, and predicted to be extremely reactive.

8.2 Group I Properties ▼

Meet the Alkali Metals

Group I is the first column of the table: lithium, sodium, potassium, then rubidium, caesium and francium. They are called the alkali metals because when they react with water they produce a hydroxide solution that is alkaline — pH 13 or 14. The syllabus names three of them for detailed study, and expects you to predict the rest.

Every alkali metal atom has exactly one electron in its outer shell: lithium 2,1; sodium 2,8,1; potassium 2,8,8,1. That single electron is the whole story of this section. Losing it leaves a stable full shell underneath and a 1+ ion, so every Group I compound contains M⁺: NaCl, KBr, Li₂O, K₂CO₃.

Physical Properties — Nothing Like "Normal" Metals

If your mental picture of a metal is iron or copper, Group I will surprise you. These metals are:

Soft — you can cut them with a knife like cold butter. Potassium is softer than sodium, which is softer than lithium.
Low density — lithium (0.53 g/cm³), sodium (0.97) and potassium (0.86) are all less dense than water, so they float. Lithium is the least dense metal there is.
Low melting points for metals — lithium 181 °C, sodium 98 °C, potassium 63 °C. Caesium melts at 29 °C, so it would melt in your hand (if it did not react with your skin first).
Shiny when freshly cut, but the bright silvery surface tarnishes within seconds as it reacts with oxygen and water vapour in the air, turning dull grey.
Good conductors of heat and electricity, like all metals.
Stored under oil, precisely because they react with air and water.

Element	Symbol	Melting point / °C	Density / g cm⁻³	Reaction with water
Lithium	Li	181	0.53	Steady fizzing; floats and moves slowly; does not melt
Sodium	Na	98	0.97	Melts into a ball; darts about; rapid fizzing; may ignite (orange flame)
Potassium	K	63	0.86	Ignites immediately with a lilac flame; may spit or explode
Rubidium	Rb	39 (predicted trend)	1.53	Violent — ignites instantly, small explosion
Caesium	Cs	29 (predicted trend)	1.88	Explosive; shatters the container

Exam Tip

Density is the one Group I trend that does not behave neatly: it broadly increases down the group, but potassium (0.86) is slightly less dense than sodium (0.97). Melting point and reactivity are the safe trends to quote. If a question shows you a data table with that potassium anomaly in it, the examiner is checking whether you read the data or just recited a rule.

The Reaction with Water

This is the defining reaction of the group and the source of its name. The pattern never changes:

alkali metal + water → metal hydroxide + hydrogen

2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g) 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) The hydroxide is soluble, so the solution left behind is an alkali — universal indicator turns purple, pH 13–14.

Memory Trick

The equation always needs a 2 in front of the metal and a 2 in front of the water: "two metal, two water, two hydroxide, one hydrogen". The reason is that hydrogen gas is diatomic (H₂), so you need two water molecules to supply the two hydrogen atoms. Forgetting this is the most common balancing error in the whole topic.

What You Actually See

Cambridge asks for observations, which means what your eyes, ears and nose report — not what you conclude. Learn these:

All three: the metal floats (density less than water), fizzes / effervesces (hydrogen gas), gets smaller and disappears, and the resulting solution turns universal indicator purple.
Lithium: steady fizzing, moves gently on the surface, stays solid throughout, takes the longest to disappear.
Sodium: the heat of the reaction melts it into a shiny ball which darts rapidly across the surface; vigorous fizzing; sometimes an orange flame.
Potassium: reacts so fast that the hydrogen ignites, burning with a lilac flame; the metal may spit and jump out of the trough.

The flame colour is your flame-test evidence from Topic 11 showing up early: lithium red, sodium yellow-orange, potassium lilac.

The two headline Group I trends point in opposite directions: reactivity increases downwards while melting point decreases.

Explaining the Reactivity Trend

This is the four-mark question that appears again and again. A Group I metal reacts by losing its single outer electron. So whichever metal loses that electron most easily is the most reactive. Three things change as you go down the group:

Each element has one more occupied electron shell than the one above, so the atomic radius increases — the outer electron is further from the nucleus.
There are more inner shells of electrons between the nucleus and the outer electron, so the shielding (screening) of the nuclear charge increases.
Together these mean the electrostatic attraction between the nucleus and the outer electron is weaker.

Therefore the outer electron is lost more easily, and the element is more reactive. Caesium loses its outer electron more readily than lithium, so caesium is the more violent of the two.

Down Group I: bigger atom + more shielding ⇒ weaker attraction ⇒ electron lost more easily ⇒ more reactive

Note what is not in that chain: "because it has more shells" on its own is not an explanation, it is an observation. The mark scheme wants: distance from the nucleus, shielding, weaker attraction, electron lost more easily.

Nuclear charge does increase down the group (3+ to 19+), but the extra distance and extra shielding more than cancel it out — so the net pull on the outer electron gets weaker.

Exam Tip

A very common wrong answer is: "potassium is more reactive because it has more electrons to lose." It does not — every Group I atom loses exactly one electron. The difference is how easily that one electron is lost. Another loser: "because it has a bigger nuclear charge". A bigger nuclear charge on its own would make it less reactive; the point is that distance and shielding outweigh it.

Supplement

Predicting Rubidium and Caesium

Extended papers regularly hand you data for Li, Na and K and ask you to predict a property of Rb, Cs or Fr. Work like this:

Melting point: the values fall 181 → 98 → 63, so the gaps are shrinking. A sensible prediction for rubidium is about 35–45 °C (true value 39 °C) and for caesium about 25–30 °C (true value 29 °C). Always give a number, and always make it lower than the one above.
Density: increases down the group overall — caesium is 1.88 g/cm³ and would sink. Predict a value above potassium's.
Reactivity with water: more violent than potassium — ignites instantly, may explode, producing RbOH or CsOH and hydrogen: 2Cs + 2H₂O → 2CsOH + H₂.
Reason: always the same — larger atom, outer electron further from the nucleus, more shielding, weaker attraction, so the electron is lost more easily.

Softness, tarnishing speed and hydroxide alkalinity also increase downwards, and every one of these predictions is worth a mark with its reason.

Two More Reactions Worth Knowing

alkali metal + oxygen → metal oxide

4Na(s) + O₂(g) → 2Na₂O(s) This is why a freshly cut surface tarnishes in seconds and why the metals are stored under oil.

alkali metal + chlorine → metal chloride

2Na(s) + Cl₂(g) → 2NaCl(s) — a vigorous reaction with a bright yellow flame and clouds of white smoke. Note the Cl₂: chlorine is diatomic, which forces the 2 in front of the sodium.

Worked Example 1 A small piece of potassium is dropped into a trough of water containing universal indicator. (a) Describe three observations. (b) Write the balanced equation with state symbols. (c) Explain why the reaction is more vigorous than the same experiment with sodium.

(a) Observations — what you can see

The potassium floats on the surface; it melts into a ball and moves rapidly; the hydrogen ignites with a lilac flame; there is vigorous fizzing; the metal disappears; and the indicator turns purple, showing the solution has become alkaline.

(b) The equation

Products are the hydroxide and hydrogen. Balance by remembering H₂ is diatomic: 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g). Check: 2 K each side; 4 H each side; 2 O each side.

A potassium atom has one more occupied shell than a sodium atom, so its outer electron is further from the nucleus and there is more shielding by the inner shells. The attraction between the nucleus and the outer electron is therefore weaker, so that electron is lost more easily and the reaction is faster and more vigorous.

(a) Floats, melts to a ball, lilac flame, fizzing, disappears, indicator turns purple. (b) 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g). (c) Bigger atom, outer electron further from nucleus with more shielding ⇒ weaker attraction ⇒ electron lost more easily ⇒ more reactive.

Worked Example 2 Use the data (Li mp 181 °C, density 0.53; Na mp 98 °C, density 0.97; K mp 63 °C, density 0.86) to predict the melting point and density of rubidium, and state whether rubidium would float on water. Explain your answers.

Step 1: Identify the melting point trend

181 → 98 is a drop of 83; 98 → 63 is a drop of 35. The melting point is decreasing down the group, and by a smaller amount each time. Predict a further drop of roughly 20–25 °C: about 40 °C. (The true value is 39 °C.)

Step 2: Handle the density trend carefully

0.53 → 0.97 → 0.86 is not a smooth rise, but the overall direction down the group is upwards. Predict a density greater than 1 g/cm³ — around 1.5 is sensible. (The true value is 1.53.)

Step 3: Answer the float question from the number

Water has a density of 1.0 g/cm³. A predicted density of 1.5 is greater, so rubidium would sink — unlike Li, Na and K. In practice the reaction is so violent that it explodes on contact.

Melting point about 40 °C, density about 1.5 g/cm³. Rubidium would sink, because its density exceeds that of water — density increases down the group even though the Na/K values are slightly out of order.

Worked Example 3 A student writes: "Sodium is more reactive than lithium because sodium atoms have more electrons, so there are more electrons available to be lost in the reaction." Identify the errors and write a correct explanation.

Error 1: The number of electrons lost

Every Group I atom loses exactly one electron, whichever element it is, because each has one outer-shell electron and losing it gives a stable full shell underneath. Sodium does not lose more electrons than lithium — it loses the same single electron.

Error 2: Confusing "how many" with "how easily"

Reactivity here is entirely about how easily that one electron is removed. The student has answered a quantity question when the examiner asked an energy question.

Step 3: Build the correct chain

Sodium has three occupied shells to lithium's two, so the outer electron is further from the nucleus and is shielded by more inner shells. The electrostatic attraction holding it is therefore weaker, so it is lost more easily and sodium reacts more vigorously.

Corrected: "Both atoms lose one outer electron. In sodium that electron is further from the nucleus and shielded by more inner shells, so the attraction to the nucleus is weaker and it is lost more easily — making sodium the more reactive metal."

Exam Tips for 8.2

1. Reactivity INCREASES down Group I. Write it on your hand if you must. The mistake that costs most marks in this topic is applying the Group VII direction here (or vice versa).

2. Melting point DECREASES down Group I. Reactivity up, melting point down — two trends, opposite directions, same group.

3. Three words earn the explanation marks: distance, shielding, attraction. "It has more shells" alone is a description, not an explanation, and typically scores zero out of three.

4. Never say "more electrons to lose". Every alkali metal loses exactly one.

5. Balance with the diatomic in mind. 2M + 2H₂O → 2MOH + H₂. The lone H₂ on the right is what forces both 2s on the left.

6. Observations must be observable. "Hydrogen is produced" is a deduction; "effervescence / bubbles / fizzing" is an observation. Write both if you have room, but never only the deduction.

7. Learn the flame colours: lithium red, sodium yellow-orange, potassium lilac. The lilac flame is a favourite one-mark answer.

8. The solution is alkaline, pH 13–14. Say "the universal indicator turns purple" or "the solution is alkaline because a soluble hydroxide is formed" — that reason is often the second mark.

🌎 Apply It: Real-World Chemistry

Group I metals are too reactive to be useful as metals — which is exactly why they are so useful as ions, in batteries, in fertiliser, in street lamps and in the human body.

A lithium-ion battery plant in Gujarat receives its raw lithium as lithium carbonate powder, never as lithium metal, and the shipping documents forbid the metal entirely. Meanwhile a school laboratory in Leeds stores its small lump of lithium in a jar of oil, and a technician who once left the lid off found the surface had turned dull grey overnight.

Explain both the shipping rule and the oil storage, and say what the grey coating is.

▼

Why Not the Metal

Lithium metal reacts with water — including water vapour in humid air — giving lithium hydroxide and hydrogen, which is flammable: 2Li + 2H₂O → 2LiOH + H₂. Shipping a cargo that generates hydrogen if it meets moisture is a serious fire risk. The carbonate is a stable, unreactive ionic compound in which the lithium has already lost its outer electron.

Why Oil

Oil is a hydrocarbon — no oxygen, no water — so it forms a barrier that keeps air away from the metal surface. The grey coating on the exposed sample is lithium oxide (with some hydroxide and carbonate): 4Li + O₂ → 2Li₂O.

Chemistry Connection

Once lithium is an ion (Li⁺) it has already lost the electron that made it dangerous, so it is perfectly stable. In a working battery the lithium ions shuttle back and forth without ever becoming metal atoms — the reactivity that makes the element hazardous is precisely what makes the ion such a good charge carrier.

Motorway lighting across Britain used to glow a distinctive deep orange. Those lamps contained sodium vapour. A rival design using potassium was never built, and lamps using lithium would glow red.

What links these colours to the Periodic Table, and why does the colour identify the metal so reliably?

▼

The Chemistry of the Colour

Heating (or an electric discharge) pushes the outer electron into a higher energy level. When it drops back, the energy is released as light of a fixed colour. Because the electron arrangement is unique to each element, so is the colour: lithium red, sodium yellow-orange, potassium lilac.

Why It Is a Reliable Test

The colour comes from the metal ion, not the compound, so sodium chloride, sodium carbonate and sodium nitrate all give the same yellow-orange flame. That is exactly what makes the flame test in Topic 11 work — you are identifying the cation regardless of what it is paired with.

Chemistry Connection

Sodium's yellow is so intense that it drowns out other colours — which is why laboratory glassware must be cleaned with acid before a flame test, and why astronomers use sodium lasers to create artificial guide stars. One electron, one energy jump, one very famous colour.

A sugarcane farmer in Uttar Pradesh buys muriate of potash (potassium chloride) as fertiliser. His neighbour asks why he does not simply spread potassium metal, which after all contains far more potassium per kilogram, and would surely be better value.

Explain, chemically, why the neighbour's idea is both useless and extremely dangerous.

▼

What Happens on Damp Soil

Potassium metal reacts violently with water: 2K + 2H₂O → 2KOH + H₂. The reaction is fast enough to ignite the hydrogen, so scattering it on damp soil would start fires. Potassium is the third member of Group I, and reactivity increases down the group, so it is far more violent than lithium or sodium.

Why It Would Not Even Work

The product, potassium hydroxide, is a strongly alkaline solution of pH 14. It would raise the soil pH far outside the range crops tolerate and would damage roots directly. Plants take up potassium as the K⁺ ion anyway, which is exactly what potassium chloride already supplies — pH-neutral and ready to use.

Chemistry Connection

This is the same principle as the lithium battery: the useful thing is the ion, not the element. Group I metals are never used as metals in agriculture, medicine or food — the sodium in your food, the potassium in a banana and the lithium in a mood stabiliser are all ions in compounds.

A YouTube science channel drops caesium into a bath of water and the tank shatters. In a school demonstration the teacher uses only a rice-grain-sized piece of sodium behind a safety screen, and refuses point-blank to demonstrate caesium at all. A student argues that caesium atoms have a much bigger positive nuclear charge (55+ against sodium's 11+), so surely caesium should hold its electrons more tightly and be less reactive.

Resolve the student's paradox.

▼

The Student Is Right About One Thing

Nuclear charge genuinely does increase down the group, and on its own a larger positive charge would indeed hold electrons more tightly. This is a good, thoughtful objection — and it is exactly what Extended mark schemes are probing.

The Two Effects That Win

Caesium's outer electron sits in the sixth shell, far further from the nucleus than sodium's third-shell electron, and electrostatic attraction falls off sharply with distance. On top of that, five inner shells of electrons shield the outer electron from the nuclear charge, so the charge it actually experiences is much smaller than 55+. Distance and shielding together outweigh the extra protons.

Chemistry Connection

Never argue a group trend from nuclear charge alone — it points the wrong way. The winning three-part chain is always distance + shielding ⇒ weaker net attraction. Notice too that the same three-part chain explains the Group VII trend in the opposite direction, because there the atom is trying to gain an electron rather than lose one.

In 1807 Humphry Davy at the Royal Institution in London passed electricity through molten potash and saw tiny globules of a soft, bright metal appear that instantly burst into lilac flame. He had isolated potassium — an element that had sat inside a common compound for centuries without anyone suspecting it. Days later he isolated sodium the same way.

Why had nobody extracted these metals before, and why did electricity succeed where furnaces had failed?

▼

Why Heating with Carbon Fails

The traditional method — heating an ore with carbon — works only if carbon is more reactive than the metal, so it can take the oxygen away. Potassium sits at the very top of the reactivity series, far above carbon, so carbon cannot displace it. The compound simply refuses to give the metal up.

Why Electrolysis Works

Electrolysis forces the reaction using electrical energy instead of a chemical reducing agent. K⁺ ions migrate to the negative electrode and are forced to gain the electron they lost: K⁺ + e⁻ → K. It is the reverse of the reaction the element wants to do, driven by an external supply.

Chemistry Connection

The link runs straight into Topic 5: the more reactive the metal, the harder it is to extract, and the more energy the extraction costs. Group I metals and aluminium all need electrolysis; iron and zinc can be smelted with carbon; gold needs no extraction at all. Davy's lilac flame was Group I reactivity announcing itself the instant the metal existed.

Practice Questions: 8.2

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

Which property is characteristic of the Group I metals?

A High density

B High melting point

C Soft enough to cut with a knife

D Unreactive with water

Alkali metals are soft, low-density, low-melting and very reactive — the opposite of the typical "strong metal" picture people carry from iron and copper.

Question 2

Which is the correct equation for the reaction of sodium with water?

A Na + H₂O → NaOH + H

B 2Na + 2H₂O → 2NaOH + H₂

C Na + 2H₂O → NaOH₂ + H₂

D 2Na + H₂O → Na₂O + H₂

Hydrogen is diatomic, so two water molecules are needed, forcing 2Na and 2NaOH. Option A leaves a lone H atom, and NaOH₂ does not exist — sodium forms only 1+ ions.

Question 3

Going down Group I, which pair of changes is correct?

A Reactivity decreases, melting point increases

B Reactivity increases, melting point decreases

C Both increase

D Both decrease

The two headline trends point in opposite directions: more violent reactions but lower melting points as you descend. Caesium melts at 29 °C and explodes in water.

Question 4

Why does potassium react more vigorously with water than lithium?

A Potassium has more outer-shell electrons to lose

B Potassium has a smaller nuclear charge

C Its outer electron is further from the nucleus and more shielded, so it is lost more easily

D Potassium has a lower melting point

Both metals lose exactly one electron; the difference is how easily. Distance plus shielding weaken the attraction. The low melting point is a separate consequence of the same bigger atom, not the cause.

Question 5

What is the pH of the solution formed when potassium reacts with water?

A About 1

B Exactly 7

C About 13

D About 5

Potassium hydroxide is a soluble, strongly alkaline hydroxide — hence the group name. Universal indicator turns purple.

Question 6

Why are Group I metals stored under oil?

A To keep them cool

B To stop them reacting with oxygen and water vapour in the air

C To stop them evaporating

D Because they dissolve in water

Oil is a hydrocarbon barrier containing no oxygen or water. Note the wording of D: the metals react with water, they do not dissolve in it.

Question 7

Lithium, sodium and potassium all float on water. This is because they

A are held up by the hydrogen bubbles only

B have densities less than that of water

C are gases

D have low melting points

Li 0.53, Na 0.97 and K 0.86 g/cm³ are all below water's 1.00. Caesium (1.88) would sink. The hydrogen produced does help push them about, but density is the reason they float.

Question 8

What flame colour is seen when potassium reacts with water?

A Red

B Yellow-orange

C Lilac

D Green

Lithium red, sodium yellow-orange, potassium lilac. Green is a copper flame colour — a transition element, not Group I.

Question 9

Rubidium is below potassium in Group I. Which prediction about rubidium is correct?

A It reacts less vigorously with water than potassium

B It has a higher melting point than potassium

C It forms an ion with a charge of 1+

D It is harder than potassium

Ion charge is fixed by the group, so every alkali metal forms M⁺. Rubidium is more reactive, lower melting and softer than potassium.

Question 10

A freshly cut piece of sodium is shiny but goes dull within seconds. Why?

A It melts at room temperature

B It reacts with oxygen in the air to form a layer of sodium oxide

C It absorbs light

D It turns into sodium chloride

4Na + O₂ → 2Na₂O. The dull grey coating is the oxide (plus some hydroxide and carbonate from water vapour and CO₂).

Question 11

Which observation is seen with sodium but not with lithium?

A It floats on the water

B Bubbles of gas are produced

C The metal melts into a ball

D The solution becomes alkaline

Sodium melts at 98 °C and the reaction releases enough heat to melt it; lithium melts at 181 °C and stays solid. The other three happen with both.

Question 12

What is the formula of the compound formed when potassium burns in oxygen (simple oxide)?

A KO

B K₂O

C KO₂

D K₂O₃

K⁺ is 1+ and O²⁻ is 2−, so two potassium ions are needed per oxide ion: K₂O. Group I oxides always follow the M₂O pattern.

Question 13

Which statement about the electronic structure of Group I atoms is correct?

A They all have the same number of shells

B They all have one electron in the outer shell

C They all have eight electrons in the outer shell

D They all have the same number of protons

Li 2,1; Na 2,8,1; K 2,8,8,1 — the outer 1 is what they share. The number of shells and protons differs for each element.

Question 14

A piece of an unknown Group I metal is added to water. It sinks, ignites instantly and shatters the beaker. Which metal is it most likely to be?

A Lithium

B Sodium

C Potassium

D Caesium

Two clues: an explosive reaction points to the bottom of the group, and sinking means a density above 1 g/cm³ — caesium is 1.88 while Li, Na and K all float.

Question 15

Which statement explains why Group I elements have similar chemical reactions?

A They all have low melting points

B They are all stored under oil

C They all lose one outer electron to form a 1+ ion

D They all have the same atomic radius

Identical outer-shell behaviour means identical reaction patterns. The physical similarities in A and B are consequences, not causes.

Question 16

Sodium reacts with chlorine. What is the correct balanced equation?

A Na + Cl → NaCl

B 2Na + Cl₂ → 2NaCl

C Na + Cl₂ → NaCl₂

D Na₂ + Cl₂ → 2NaCl

Chlorine exists as Cl₂ molecules, which forces 2Na and 2NaCl. Sodium is a metal, so it is never written as Na₂, and NaCl₂ would require a 2+ sodium ion.

Question 17

The melting points of Li, Na and K are 181, 98 and 63 °C. The best prediction for caesium is

A 210 °C

B 120 °C

C 29 °C

D −40 °C

Caesium is two places below potassium, so the melting point must fall below 63 °C but stay above room temperature (it is a solid in the bottle). 29 °C fits the shrinking-gap pattern.

Question 18

Which is not a valid reason for the increase in reactivity down Group I?

A The outer electron is further from the nucleus

B There are more inner shells shielding the outer electron

C The attraction between nucleus and outer electron is weaker

D There are more outer electrons available to be lost

Every Group I atom has exactly one outer electron and loses exactly one. A, B and C are the three linked points the mark scheme wants.

Question 19

50 cm³ of water is used in three separate experiments with equal masses of Li, Na and K. Which statement about the resulting solutions is correct?

A Only the potassium solution is alkaline

B All three solutions are alkaline and turn universal indicator purple

C Only the lithium solution is alkaline because lithium is the least reactive

D All three solutions are neutral

All three produce a soluble metal hydroxide, so all three are alkaline. Reactivity changes how fast and how violently, not what is formed.

Question 20

Which explains why Group I metals are extracted by electrolysis rather than by heating with carbon?

A Their ores are liquids

B Carbon is too expensive

C They are more reactive than carbon, so carbon cannot displace them from their compounds

D They have low melting points

Reduction with carbon works only for metals below carbon in the reactivity series. Group I metals sit at the very top, so electricity must supply the electrons: Na⁺ + e⁻ → Na.

8.3 Group VII Properties ▼

Meet the Halogens

Group VII is the second-from-last column: fluorine, chlorine, bromine, iodine and astatine. They are called the halogens, from the Greek for "salt-formers", because they react with metals to make salts — sodium chloride being the most famous salt on Earth. The syllabus focuses on chlorine, bromine and iodine.

Every halogen atom has seven electrons in its outer shell: fluorine 2,7; chlorine 2,8,7; bromine 2,8,18,7. Being one short of a full shell, a halogen atom gains one electron to form a 1− ion — the halide ion: Cl⁻, Br⁻, I⁻. Everything in this section follows from that single fact.

Halogens exist as diatomic molecules: F₂, Cl₂, Br₂, I₂

Two atoms share one pair of electrons, so each reaches a full outer shell — a single covalent bond. Always write the element as Cl₂, never Cl, in an equation. Dropping the 2 is one of the most common equation errors on the paper. The ion, by contrast, is written singly: Cl⁻ not Cl₂²⁻.

Physical Properties — A Colour Chart You Must Memorise

The halogens are the one group where colour and state at room temperature are examinable facts in their own right. Learn this table cold.

Halogen	Formula	State at r.t.p.	Colour	Melting point / °C
Fluorine	F₂	Gas	Pale yellow	−220
Chlorine	Cl₂	Gas	Pale yellow-green	−101
Bromine	Br₂	Liquid	Red-brown	−7
Iodine	I₂	Solid	Grey-black (shiny, sublimes to a purple vapour)	114
Astatine	At₂	Solid (predicted)	Black (predicted)	~302 (predicted)

Memory Trick

"Green Gas, Brown Bromine liquid, Black Iodine solid." Read down the group and two things get bigger together: the colour gets darker (pale yellow → yellow-green → red-brown → grey-black) and the state gets more condensed (gas → gas → liquid → solid). Both are the same underlying trend: bigger molecules, stronger forces between them.

Exam Tip

Be precise with iodine. As a solid it is grey-black and shiny; the famous purple / violet colour belongs to its vapour (formed on gentle warming — it sublimes) and to its solution in an organic solvent. Writing "iodine is purple" when asked for the appearance of the solid can lose the mark. In water, iodine gives a pale brown solution.

The Trends Down Group VII

Here is the sentence to tattoo on your brain: reactivity DECREASES down Group VII. This is the exact opposite of Group I, and swapping the two directions is the single biggest source of lost marks in Topic 8.

Property	Trend going DOWN Group VII	Evidence
Reactivity	Decreases	Chlorine displaces bromine and iodine; iodine displaces neither
Melting and boiling point	Increases	Gas → liquid → solid at room temperature
Colour	Gets darker	Pale yellow-green → red-brown → grey-black
Density	Increases	Chlorine gas → dense liquid bromine → solid iodine
Atomic radius	Increases	One more occupied electron shell each step

Supplement

Why Reactivity Decreases Down the Group

A halogen reacts by gaining one electron to complete its outer shell. So the most reactive halogen is the one that attracts an extra electron most strongly. Going down the group:

Each atom has one more occupied electron shell, so the atomic radius increases and the outer shell — where the incoming electron must go — is further from the nucleus.
There is more shielding by the inner shells of electrons.
So the attraction between the nucleus and the incoming electron is weaker.
The electron is therefore gained less easily, and the element is less reactive.

Look closely and this is the same physical reasoning as Group I — bigger atom, more shielding, weaker attraction. The direction of the reactivity trend flips only because Group I atoms want to lose an electron (weaker attraction helps them) while Group VII atoms want to gain one (weaker attraction hinders them).

Weaker attraction ⇒ easier to LOSE (Group I, more reactive) but harder to GAIN (Group VII, less reactive)

One physical cause, two opposite consequences. If you can write this sentence in the exam you will never mix the two groups up again.

The melting point trend has a different cause and must be explained differently: going down the group the molecules get larger, so the intermolecular forces between molecules are stronger and more energy is needed to separate them. Note carefully — you are breaking forces between molecules, not the covalent bonds inside them.

Displacement Reactions of the Halogens

This is the classic laboratory demonstration of the reactivity trend, and it comes up on almost every paper.

A more reactive halogen displaces a less reactive halogen from a solution of its halide

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq) — solution turns orange / red-brown Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq) — solution turns brown, black solid may appear Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq) — solution turns brown I₂(aq) + KBr(aq) → no reaction — iodine is less reactive than bromine, so it cannot displace it

Read those equations carefully and notice two things. First, the halogen is always diatomic on both sides. Second, the 2 in front of the potassium halide is there because one Cl₂ molecule contains two chlorine atoms, each needing its own K⁺ partner.

Halogen added ↓ / halide solution →	Potassium chloride	Potassium bromide	Potassium iodide
Chlorine water	No reaction	Reaction — turns orange (Br₂ formed)	Reaction — turns brown (I₂ formed)
Bromine water	No reaction	No reaction	Reaction — turns brown (I₂ formed)
Iodine solution	No reaction	No reaction	No reaction

The pattern is a neat triangle: reactions happen only in the top-right corner, where a halogen meets the halide of an element below it in the group. Chlorine displaces both the others; bromine displaces only iodine; iodine displaces nothing. That order — Cl > Br > I — is the reactivity trend, demonstrated in a test tube.

Every tick is a colour change from colourless to orange or brown — the displaced halogen is what you are seeing.

Supplement

Displacement as Electron Transfer

Extended candidates should be able to describe displacement in terms of ionic equations and electron transfer. Take chlorine added to potassium bromide. The potassium ions are spectators, so cancel them:

Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)

Each chlorine atom gains an electron: Cl₂ + 2e⁻ → 2Cl⁻ — chlorine is reduced, so it is the oxidising agent. Each bromide ion loses an electron: 2Br⁻ → Br₂ + 2e⁻ — the bromide is oxidised. Chlorine is a stronger oxidising agent than bromine, which is why it wins the electron.

So "chlorine is more reactive than bromine" and "chlorine is a stronger oxidising agent than bromine" are two ways of saying the same thing. Oxidising power, like reactivity, decreases down Group VII.

Two Practical Points

Chlorine bleaches damp litmus paper. Damp blue litmus first turns red (chlorine dissolves to give an acidic solution) and then is bleached white. This is the standard test for chlorine gas, and the bleaching is what distinguishes it from other acidic gases.

Halogens are toxic. Chlorine and bromine are used in a fume cupboard; bromine liquid causes severe burns and its vapour is dangerous even in small amounts. Any practical description that mentions handling them should include working in a fume cupboard.

Worked Example 1 Chlorine water is added to a colourless solution of potassium iodide. (a) State what you would see. (b) Write the balanced symbol equation and the ionic equation. (c) Explain what the result tells you about the reactivity of chlorine and iodine.

(a) The observation

The colourless potassium iodide solution turns brown (a dark brown or even black solid of iodine may appear if the solution is concentrated). The brown colour is the displaced iodine.

(b) The equations

Full: Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq). Both halogens are diatomic, and the 2 in front of KI supplies two K⁺ for the two Cl⁻ formed. Ionic, with the spectator K⁺ cancelled: Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq).

Chlorine has displaced iodine from its compound, so chlorine is more reactive than iodine. Chlorine is higher in Group VII, and reactivity decreases down the group. In electron terms, chlorine attracts an electron more strongly, so it takes the electron from the iodide ion: chlorine is reduced (the oxidising agent) and iodide is oxidised.

(a) The solution turns brown. (b) Cl₂ + 2KI → 2KCl + I₂; ionically Cl₂ + 2I⁻ → 2Cl⁻ + I₂. (c) Chlorine is more reactive than iodine, consistent with reactivity decreasing down Group VII.

Worked Example 2 A student adds bromine water to three colourless solutions: potassium chloride, potassium bromide and potassium iodide. Predict the result in each tube and explain the pattern.

Step 1: Place bromine in the reactivity order

Reactivity decreases down Group VII, so the order is Cl₂ > Br₂ > I₂. Bromine is in the middle: it can displace only what is below it, which is iodine.

Step 2: Tube by tube

Potassium chloride: chlorine is above bromine and more reactive, so bromine cannot displace it — no reaction; the solution just keeps the orange colour of the added bromine water. Potassium bromide: same element — no reaction. Potassium iodide: iodine is below bromine and less reactive, so it is displaced — the solution turns brown.

Step 3: Write the one reaction that happens

Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq), or ionically Br₂ + 2I⁻ → 2Br⁻ + I₂.

Step 4: Watch the trap

In the first tube the mixture is still coloured, because bromine water is orange to begin with. "It went orange, so a reaction happened" is wrong — you must compare with the original colour of what was added. Only a change to a new colour indicates displacement.

KCl: no reaction. KBr: no reaction. KI: reaction — turns brown as iodine is displaced, Br₂ + 2KI → 2KBr + I₂. Bromine displaces only the halogen below it.

Worked Example 3 Astatine lies below iodine in Group VII. Predict (a) its colour and physical state at room temperature, (b) whether it would react with potassium iodide solution, and (c) the formula of the compound it forms with sodium. Explain each prediction.

(a) Colour and state

Colour gets darker and melting point increases down the group: pale yellow gas → yellow-green gas → red-brown liquid → grey-black solid. Continuing the pattern, astatine should be a black solid with a melting point above iodine's 114 °C. The reason: larger molecules mean stronger intermolecular forces, so more energy is needed to separate them.

(b) Reaction with potassium iodide

No reaction. Astatine is below iodine, so it is less reactive than iodine and cannot displace it. A halogen can only displace a halide of an element below itself in the group, and there is nothing below astatine on the syllabus.

Group VII means seven outer electrons, so astatine gains one to form At⁻. Sodium forms Na⁺. Charges cancel one-to-one: NaAt — exactly the pattern of NaCl, NaBr and NaI.

(a) A black solid, melting above 114 °C. (b) No reaction — astatine is less reactive than iodine. (c) NaAt, because the group number fixes the 1− charge whatever the element.

Exam Tips for 8.3

1. Reactivity DECREASES down Group VII. Opposite to Group I. Before writing any trend answer, check which group you are in.

2. Halogens are diatomic. Cl₂, Br₂, I₂ in every equation. The ion is single: Cl⁻. Losing the 2 costs the equation mark even when the chemistry is right.

3. Learn the colours precisely. Chlorine pale yellow-green gas; bromine red-brown liquid; iodine grey-black solid that sublimes to a purple vapour. Vague answers such as "dark" or "coloured" score nothing.

4. Explain the reactivity trend with electron GAIN. Bigger atom, more shielding, weaker attraction for the incoming electron, so it is gained less easily. Never explain Group VII with "loses an electron more easily" — that is Group I.

5. Explain melting points with intermolecular forces. Larger molecules ⇒ stronger forces between molecules ⇒ more energy needed. Do not say the covalent bonds get stronger.

6. Predict displacement by position. A halogen displaces the halide of any element below it. Higher displaces lower. Never the other way round.

7. Describe colour changes as changes. "Colourless to brown" scores; "brown" alone may not. And remember the added halogen solution already has a colour of its own.

8. Chlorine bleaches damp litmus. Red first, then white — the bleaching is the identifying step.

🌎 Apply It: Real-World Chemistry

Halogens sterilise your drinking water, protect your teeth, disinfect a swimming pool in Dubai and turn up in the salt on every table on Earth.

A municipal water treatment plant in Chennai adds a carefully controlled trace of chlorine to drinking water. A resident complains that chlorine is a poison gas used in the First World War and demands that iodine be used instead, since iodine is added to table salt and must therefore be safer.

Explain why chlorine is chosen, and evaluate the suggestion of using iodine.

▼

Why Chlorine Works

Chlorine is the most reactive halogen in normal use, a powerful oxidising agent that destroys bacteria and viruses by ripping electrons out of their cell chemistry. It is also cheap, and at the concentrations used (about 1 part per million) it kills microorganisms without harming people.

Why Iodine Is a Poor Substitute

Iodine is at the bottom of the group, so it is much less reactive — a far weaker oxidising agent, so it disinfects much more slowly and less thoroughly. It is also barely soluble in water, far more expensive, and it stains. Iodine tablets are used for emergency field purification of small volumes, but never for a city supply.

Chemistry Connection

The resident's confusion is a Topic 8 classic: the element chlorine (Cl₂, toxic gas) is a completely different substance from the ion chloride (Cl⁻, in table salt). Iodine in iodised salt is present as the iodide or iodate ion, not as I₂. Whenever a question compares an element with its ion, name the difference explicitly — it is almost always where the mark lives.

A forensic technician in Birmingham tests an unknown colourless solution for halide ions. She adds chlorine water: the solution turns brown. She repeats with a second sample, adds bromine water, and the solution turns brown again. A third sample plus iodine solution gives no change at all.

Which halide is present, and how do you know?

▼

Reading the First Two Results

A brown colour means iodine has been displaced. Both chlorine and bromine displaced it, and both chlorine and bromine are above iodine in the group — so the solution must contain iodide ions, I⁻.

Confirming with the Third Result

Iodine cannot displace itself, and it cannot displace anything above it, so "no change" is exactly what iodide should give. If the solution had contained bromide, chlorine would have turned it orange and bromine would have given no change. The pattern of results is unique to iodide.

Chemistry Connection

Notice the logic: a negative result carries as much information as a positive one. Cambridge often gives you a table with several "no change" entries, and identifying the halide requires you to use them. In practice the silver nitrate test from Topic 11 (white / cream / yellow precipitates) is quicker — but displacement is what proves the reactivity order.

A photography museum in Rochester displays 19th-century glass plates coated with silver bromide, and explains that silver iodide plates were used for the earliest daguerreotypes. Meanwhile a modern cloud-seeding programme in Maharashtra fires silver iodide crystals into monsoon clouds.

Both technologies use silver halides. What do the compounds have in common chemically, and what does "silver bromide" tell you about the bonding?

▼

The Common Chemistry

Both are silver halides: AgBr and AgI, formed from Ag⁺ and a 1− halide ion. Because all halogens form 1− ions, all silver halides have the same 1 : 1 formula — AgCl, AgBr, AgI. All three are insoluble, which is exactly why they can be used as solid crystals and why silver nitrate is the standard test for halide ions.

What the Name Tells You

The ending -ide signals a compound of two elements in which the halogen is present as the negative ion, not as the diatomic element. Silver is a metal and bromine a non-metal, so the bonding is ionic: an electron transfers from silver to bromine.

Chemistry Connection

Photographic film works because light breaks the silver halide back down to silver metal — the reverse of the electron transfer that formed it. The lighter, more reactive halides respond fastest, which is why film speed depends on which halide is used. Group VII trends decide the exposure time of a photograph.

A chemical store in Singapore keeps chlorine in a pressurised cylinder, bromine in a thick-walled glass bottle standing in a tray of sand inside a fume cupboard, and iodine as loose crystals in a screw-top jar on an open shelf. A new technician asks why three elements from the same group need three completely different storage arrangements.

Explain the difference using the physical properties of the group.

▼

Three States, Three Problems

At room temperature chlorine is a gas, so it must be compressed into a cylinder to store any useful quantity. Bromine is a dense, volatile liquid whose red-brown vapour is highly toxic, so it needs a sealed bottle, a fume cupboard, and sand to absorb any spill. Iodine is a solid with a much lower vapour pressure, so a simple jar is enough.

Why the States Differ

Going down the group the molecules get larger, so the intermolecular forces between molecules get stronger and more energy is needed to separate them. That raises the melting and boiling points steadily, taking the elements from gas to liquid to solid over three steps of a single column.

Chemistry Connection

Do not confuse this with the reactivity trend, which runs the other way. Chlorine is the most reactive and the most volatile; iodine is the least reactive and the least volatile. Physical trend and chemical trend have different causes — intermolecular forces for one, electron gain for the other — and examiners give separate marks for each explanation.

Bromine is extracted commercially from seawater and from brine wells such as those in Arkansas and at the Dead Sea. The industrial process bubbles chlorine gas through the concentrated brine, and the bromine is then blown out of the liquid with air. A plant manager is asked why they do not simply extract the bromine directly, or use iodine to do the job instead of chlorine.

Explain the chemistry of the process and why chlorine is the only sensible choice.

▼

What Is Actually Happening

Seawater contains bromine as bromide ions, Br⁻, not as the element. Chlorine is more reactive, so it displaces bromine: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq). This is exactly the test-tube displacement reaction, scaled up to thousands of tonnes.

Why Not Iodine, and Why Not "Directly"

Iodine is below bromine in Group VII and therefore less reactive, so it cannot displace bromide — nothing would happen. And there is no free bromine in seawater to extract directly; a chemical reaction is the only way to convert the ion into the element.

Chemistry Connection

In electron terms, chlorine takes an electron from each bromide ion: chlorine is reduced (it is the oxidising agent) and bromide is oxidised. An industry worth billions rests on one line of the Periodic Table — chlorine sits above bromine, therefore chlorine wins the electron. That is Topic 8 earning its keep.

Practice Questions: 8.3

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

What is the colour and state of chlorine at room temperature?

A Pale yellow-green gas

B Red-brown liquid

C Grey-black solid

D Colourless gas

B describes bromine and C describes iodine. Chlorine is definitely coloured, so "colourless" is wrong.

Question 2

Going down Group VII, reactivity

A increases, as it does in Group I

B decreases

C stays the same

D increases then decreases

Group VII runs the opposite way to Group I. Halogens react by gaining an electron, and a bigger, better-shielded atom attracts that electron less strongly.

Question 3

Which formula correctly represents the element bromine?

A Br

B Br₂

C Br⁻

D Br₃

All halogens are diatomic molecules. Br⁻ is the bromide ion, found in compounds such as KBr, not the element.

Question 4

Chlorine water is added to potassium bromide solution. What is observed?

A No change

B The solution turns orange as bromine is displaced

C A white precipitate forms

D The solution turns purple

Cl₂ + 2KBr → 2KCl + Br₂. Chlorine is above bromine, so it displaces it, and the free bromine colours the solution orange / red-brown.

Question 5

Which reaction will not take place?

A Chlorine + potassium iodide

B Chlorine + potassium bromide

C Bromine + potassium iodide

D Iodine + potassium bromide

Iodine is below bromine, so it is less reactive and cannot displace it. Displacement only works downwards: a halogen displaces the halide of an element below it.

Question 6

Why does reactivity decrease down Group VII?

A The atoms have more electrons to lose

B The outer shell is further from the nucleus and more shielded, so an electron is gained less easily

C The molecules get heavier so they move more slowly

D The nuclear charge decreases

Halogens react by gaining an electron, so a weaker attraction makes them less reactive — the mirror image of Group I, where a weaker attraction makes losing easier.

Question 7

Which balanced equation is correct?

A Cl + KBr → KCl + Br

B Cl₂ + 2KBr → 2KCl + Br₂

C Cl₂ + KBr → KCl₂ + Br

D 2Cl₂ + KBr → 2KCl + Br₂

Both halogens must be diatomic, and one Cl₂ molecule liberates two Cl⁻ ions, so two KBr are needed. KCl₂ would require a 2+ potassium ion, which does not exist.

Question 8

Solid iodine is warmed gently in a test tube. What is seen?

A It melts to a red-brown liquid

B It sublimes to a purple vapour

C It burns with a lilac flame

D It turns into a colourless gas

Iodine sublimes — solid straight to vapour — giving the distinctive purple colour. The red-brown liquid is bromine, and lilac flames belong to potassium.

Question 9

Why do the melting points of the halogens increase down the group?

A The covalent bonds inside the molecules get stronger

B The molecules get larger, so the forces between molecules are stronger

C The elements become more reactive

D The atoms gain more electrons

Melting a molecular substance separates whole molecules, so it is the intermolecular forces that matter. The covalent bonds inside stay intact.

Question 10

What charge does a halide ion carry?

A 7+

B 1+

C 1−

D 7−

Seven outer electrons, one short of eight, so one electron is gained: 8 − 7 = 1, negative. Gaining seven electrons or losing seven would take far too much energy.

Question 11

A colourless solution turns brown when chlorine water is added, but shows no change with iodine solution. Which ion is present?

A Chloride

B Bromide

C Iodide

D Fluoride

Brown means iodine was displaced, so iodide was present. Bromide would have gone orange with chlorine, and chloride would show no change with any of them.

Question 12

Astatine is below iodine in Group VII. Which prediction is correct?

A It is a pale yellow gas

B It displaces iodine from potassium iodide

C It is a dark solid with a melting point above that of iodine

D It is the most reactive halogen

Colour darkens and melting point rises down the group, while reactivity falls — so astatine is the darkest, highest-melting and least reactive of the halogens listed.

Question 13

In the reaction Cl₂ + 2I⁻ → 2Cl⁻ + I₂, the chlorine has been

A oxidised, because it lost electrons

B reduced, because it gained electrons

C neither oxidised nor reduced

D both oxidised and reduced

Cl₂ goes to 2Cl⁻, so each atom gained an electron — reduction. Chlorine is therefore the oxidising agent, and the iodide ions are oxidised to I₂.

Question 14

Which is the correct test result for chlorine gas?

A It relights a glowing splint

B It gives a squeaky pop with a lighted splint

C It bleaches damp litmus paper white

D It turns limewater milky

Damp blue litmus turns red then white. Option A is oxygen, B is hydrogen and D is carbon dioxide. The paper must be damp for any gas test using litmus.

Question 15

Which statement about the halogens is correct?

A They form 1+ ions

B They are all gases at room temperature

C They react with metals to form salts containing halide ions

D They have full outer electron shells

"Halogen" means salt-former: 2Na + Cl₂ → 2NaCl. Bromine is a liquid and iodine a solid, halide ions are 1−, and a full outer shell belongs to Group VIII.

Question 16

Bromine water is added to potassium chloride solution. What happens?

A The solution turns brown

B No reaction — the mixture keeps the orange colour of the bromine water

C Chlorine gas is given off

D A yellow precipitate forms

Chlorine is above bromine and more reactive, so bromine cannot displace it. Careful: the mixture is still coloured, but only because the bromine water you added was orange to begin with.

Question 17

Which list places the halogens in order of increasing reactivity?

A Chlorine, bromine, iodine

B Iodine, bromine, chlorine

C Bromine, iodine, chlorine

D Chlorine, iodine, bromine

Reactivity decreases down the group, so going up the group means increasing reactivity: I < Br < Cl. Option A is the order of decreasing reactivity — read the question wording carefully.

Question 18

Bromine is extracted industrially by bubbling chlorine through seawater. Why does this work?

A Chlorine dissolves the bromine out of the water

B Chlorine is more reactive and displaces bromine from bromide ions

C Chlorine is denser than bromine

D Bromine evaporates when chlorine is added

Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂. Seawater contains bromide ions, not bromine molecules, so a displacement reaction is needed to release the element.

Question 19

The electronic configuration of a halogen atom in Period 3 is

A 2,7

B 2,8,7

C 2,8,8

D 2,8,1

Period 3 means three shells and Group VII means seven outer electrons — chlorine. Option A is fluorine (Period 2), C is argon and D is sodium.

Question 20

Equal volumes of chlorine water are added to solutions of sodium bromide and sodium iodide. Which statement is correct?

A Neither tube reacts, because sodium salts are stable

B Only the sodium bromide reacts

C Both react; the bromide turns orange and the iodide turns brown

D Both react and both turn colourless

Chlorine is above both, so it displaces both. The metal is irrelevant — sodium and potassium ions are spectators, so NaBr behaves exactly like KBr.

8.4 Transition Elements and Noble Gases ▼

The Transition Elements

The block sitting between Group II and Group III is the transition elements: scandium, titanium, vanadium, chromium, manganese, iron, cobalt, nickel, copper, zinc and the rows below them. These are the metals you actually meet — the ones in bridges, wiring, coins, cutlery and catalytic converters.

At IGCSE they are not given a Roman-numeral group number, and you are not expected to explain why they behave as they do. You are expected to know five characteristic properties and to contrast them with Group I.

The five properties of transition elements

1. High density — iron 7.9, copper 8.9 g/cm³ (compare sodium at 0.97). 2. High melting points — iron 1538 °C, copper 1085 °C (compare sodium at 98 °C). 3. Variable oxidation number — iron forms Fe²⁺ and Fe³⁺; copper forms Cu⁺ and Cu²⁺. 4. Coloured compounds — copper(II) sulfate blue, iron(II) sulfate pale green, iron(III) chloride yellow-brown, potassium manganate(VII) purple. 5. Catalytic behaviour — the elements and their compounds are used as catalysts.

Memory Trick

"Very Colourful Cats Have Density" — Variable oxidation number, Coloured compounds, Catalysts, High melting point, Density high. Five properties, five words. If a question asks for "two properties of transition elements not shown by Group I metals", pick from the first three — they are the most distinctive and the hardest to confuse.

Variable Oxidation Number in Practice

This is the property that shows up most often. Because a transition element can form more than one ion, its compounds need a Roman numeral in the name to say which one you mean:

Iron(II) chloride is FeCl₂ — the iron is Fe²⁺, and the solution is pale green.
Iron(III) chloride is FeCl₃ — the iron is Fe³⁺, and the solution is yellow-brown.
Copper(I) oxide is Cu₂O (red); copper(II) oxide is CuO (black).

Compare that with sodium, which is always Na⁺. Nobody ever writes "sodium(I) chloride" because there is no alternative to distinguish it from — and that is exactly why Group I compounds never carry Roman numerals.

Catalysts You Should Be Able to Name

Catalyst	Process	Reaction
Iron	Haber process	N₂ + 3H₂ ⇌ 2NH₃
Vanadium(V) oxide, V₂O₅	Contact process	2SO₂ + O₂ ⇌ 2SO₃
Nickel	Hydrogenation of alkenes (making margarine)	C₂H₄ + H₂ → C₂H₆
Manganese(IV) oxide, MnO₂	Laboratory preparation of oxygen	2H₂O₂ → 2H₂O + O₂
Platinum / rhodium	Catalytic converter in a car exhaust	2CO + 2NO → 2CO₂ + N₂

Exam Tip

A catalyst speeds up a reaction without being used up and is chemically unchanged at the end. Two more marks live in the detail: a catalyst provides an alternative reaction pathway of lower activation energy, and it does not change the yield or the position of an equilibrium — only the speed at which it is reached. That last point is a favourite Topic 6 crossover.

Transition Elements versus Group I — The Comparison Table

This table is the single most examinable item in 8.4. Cambridge asks for the comparison directly, and expects you to be specific.

Property	Group I (alkali metals)	Transition elements
Density	Low — Li 0.53, Na 0.97 (float on water)	High — Fe 7.9, Cu 8.9
Melting point	Low — Na 98 °C, K 63 °C	High — Fe 1538 °C, Cu 1085 °C
Hardness	Very soft — cut with a knife	Hard, strong, tough
Reactivity	Very reactive — react vigorously with cold water	Much less reactive — iron rusts only slowly; copper does not react with water at all
Oxidation number in compounds	Always +1 only	Variable — Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺
Colour of compounds	White solids giving colourless solutions	Coloured — CuSO₄ blue, FeSO₄ green, Fe(NO₃)₃ yellow-brown
Catalytic behaviour	Not used as catalysts	Widely used as catalysts (Fe, Ni, V₂O₅, MnO₂)
Storage	Under oil — too reactive to leave in air	On an open shelf

Two blocks of metals, almost opposite in every property. When a question says "compare", give a property AND both values.

Exam Tip

A "compare" question needs both sides of each point. "Transition metals have high melting points" is half an answer. "Transition metals have much higher melting points than Group I metals — iron melts at 1538 °C compared with sodium at 98 °C" is a full one. Comparative words (higher, denser, less reactive) are what earn the marks.

The Noble Gases

Group VIII — also written Group 0 — is the last column: helium, neon, argon, krypton, xenon and radon. They were discovered late (mostly by William Ramsay in London in the 1890s) precisely because they do nothing: they have no reactions to give them away.

Noble gases have a full outer electron shell

Helium 2 — a full first shell. Neon 2,8. Argon 2,8,8. A full outer shell is a stable arrangement, so the atom has no tendency to lose, gain or share electrons. Therefore noble gases are unreactive (inert) — they form almost no compounds at all.

Two more consequences follow, and both are examinable:

They are monatomic — they exist as single atoms (He, Ne, Ar), not as molecules. Every other gaseous element on the syllabus is diatomic (H₂, O₂, N₂, Cl₂), because those atoms still need to share electrons to complete their shells. A noble gas atom is already complete, so it has no reason to bond even to another atom of itself.
They have very low boiling points and are all gases at room temperature, because there are only weak forces between the separate atoms.

Exam Tip

The most common wrong answer in the whole of 8.4 is: "noble gases are unreactive because they are gases". That is backwards — being a gas is a consequence of their structure, not the cause of their inertness. Chlorine is a gas and is extremely reactive. The mark scheme wants: full outer shell ⇒ no tendency to lose, gain or share electrons ⇒ unreactive.

Uses of the Noble Gases

Noble gas	Use	Property that makes it suitable
Helium	Filling balloons and airships	Very low density so it provides lift, and non-flammable — unlike hydrogen, which is lighter but exploded on the Hindenburg
Helium	Cooling MRI scanner magnets	Liquefies at only 4 K and is chemically inert, so it cannot attack the equipment
Argon	Filling filament light bulbs	Unreactive, so the hot tungsten filament cannot burn away in oxygen — the bulb lasts far longer
Argon	Shielding gas in arc welding	Blankets the weld, keeping oxygen and nitrogen away from the hot metal so no brittle oxides form
Neon	Advertising signs and lasers	Glows bright red-orange when a current is passed through it at low pressure
Argon / krypton	Filling the gap in double glazing	Inert and a poorer conductor of heat than air

Memory Trick

For any "why is this noble gas used here?" question, the answer is nearly always "because it is unreactive" — plus one extra physical property specific to the job: helium's low density for lifting, neon's red glow for signs, argon's cheapness (it is 0.9% of the air) for bulk uses such as welding and light bulbs. Two marks, two reasons.

Worked Example 1 Give three ways in which the physical and chemical properties of iron differ from those of sodium, quoting evidence for each, and explain why iron compounds need a Roman numeral in their name but sodium compounds do not.

Difference 1: Melting point and density

Iron melts at 1538 °C and has a density of 7.9 g/cm³; sodium melts at 98 °C and has a density of 0.97 g/cm³. Iron is much higher melting and much denser — and it sinks in water while sodium floats.

Difference 2: Reactivity

Sodium reacts violently with cold water, melting into a ball and fizzing. Iron does not react with cold water at all; it only rusts slowly when both water and oxygen are present over days or weeks.

Difference 3: Compounds

Sodium compounds are white solids giving colourless solutions (NaCl, Na₂CO₃). Iron compounds are coloured — iron(II) sulfate is pale green, iron(III) chloride yellow-brown. Iron is also used as a catalyst in the Haber process, which sodium never is.

The Roman numeral

Iron shows variable oxidation number: it can form Fe²⁺ or Fe³⁺, so "iron chloride" is ambiguous — FeCl₂ and FeCl₃ are different compounds. The numeral removes the ambiguity. Sodium forms only Na⁺, so there is nothing to distinguish and no numeral is needed.

Iron has a far higher melting point and density, is far less reactive, and forms coloured compounds with variable oxidation numbers (and acts as a catalyst). The Roman numeral is needed only because iron has more than one possible ion charge.

Worked Example 2 Explain, in terms of electronic structure, why argon does not form compounds, and explain why argon is used both inside filament light bulbs and as a shielding gas in welding.

Step 1: The electronic structure

Argon is 2,8,8 — its outer shell holds eight electrons, which is full.

Step 2: Why that means no reactions

Chemical reactions happen when atoms lose, gain or share electrons to reach a full outer shell. Argon already has one, so it has no tendency to do any of the three. There is nothing to be gained energetically, so it forms essentially no compounds.

Step 3: The light bulb

A tungsten filament glowing at over 2000 °C would react with oxygen and burn away within seconds in air. Argon fills the bulb and, being unreactive, will not attack the filament even at that temperature. Argon is chosen over the other noble gases because it makes up nearly 1% of the atmosphere and is therefore cheap.

Step 4: The welding torch

Molten metal reacts readily with oxygen and nitrogen, forming brittle oxides and nitrides that weaken the joint. A stream of argon blankets the weld pool, physically excluding air, and because argon is inert it takes no part in the process itself.

Argon has a full outer shell (2,8,8), so it has no tendency to lose, gain or share electrons and is unreactive. That inertness protects the hot tungsten filament from oxidation in a bulb, and shields molten metal from oxygen and nitrogen during welding.

Worked Example 3 Element E is a hard, dense metal with a melting point of 1495 °C. It forms two chlorides, ECl₂ and ECl₃, and its aqueous compounds are pink. It is used as a catalyst. (a) Where in the Periodic Table is E? (b) Give the oxidation numbers of E in the two chlorides. (c) State two properties of E that a Group I metal would not share.

(a) Locating E

Four clues all point the same way: high melting point, high density, two different chlorides (variable oxidation number), coloured compounds and catalytic behaviour. That combination is the definition of a transition element, so E lies in the block between Groups II and III. (The pink colour suggests cobalt.)

(b) Oxidation numbers

Chlorine is always 1− in a chloride. In ECl₂ two chlorides means E must be +2. In ECl₃ three chlorides means E must be +3. The two chlorides would be named E(II) chloride and E(III) chloride.

A Group I metal has only one oxidation state (+1), so it could never form two different chlorides; and Group I compounds are white / colourless, never pink. You could equally cite the high melting point (Group I melts below 200 °C), the high density (Group I floats on water), or the catalytic use.

(a) E is a transition element, between Groups II and III. (b) +2 in ECl₂ and +3 in ECl₃. (c) Variable oxidation number and coloured compounds — neither of which any Group I metal shows.

Exam Tips for 8.4

1. Learn the five transition properties as a list — high density, high melting point, variable oxidation number, coloured compounds, catalytic behaviour. Questions ask for two or three, and vague answers such as "they are strong metals" earn nothing.

2. Give both sides in a comparison. "Iron is denser than sodium" scores; "iron is dense" may not.

3. Never call a transition element a Group I metal. Iron and copper are not reactive, not soft, not low-melting, and are not stored under oil.

4. Explain noble gas inertness with the full outer shell. "Because they are gases" is the classic zero-mark answer. Say: full outer shell, so no tendency to lose, gain or share electrons.

5. Monatomic means single atoms. Write He, Ne, Ar — never He₂. This is one of the few facts examiners can test in a single word.

6. Match the use to the property. Helium in balloons: unreactive and low density. Argon in lamps: unreactive and cheap. Always pair inertness with the second reason.

7. Roman numerals signal variable oxidation number. If a compound is called copper(II) sulfate, the numeral is telling you the metal has other possible states — a transition-element fingerprint.

8. Catalysts do not change yield. They speed up the reaction by providing a pathway of lower activation energy and are unchanged at the end. A Topic 6 sentence that scores marks in Topic 8.

🌎 Apply It: Real-World Chemistry

Transition elements build the physical world and speed up its chemistry; the noble gases protect it by refusing to join in.

A steel plant in Jamshedpur produces girders that hold up buildings and bridges. Its chief engineer is asked, half-jokingly, whether the company could save weight by making girders out of sodium instead of iron, since sodium is a metal too and is far lighter.

Give three chemical reasons why sodium girders are impossible.

▼

Reason 1: It Is Not Strong

Sodium is soft enough to cut with a knife. A Group I metal has nothing like the hardness or tensile strength needed to carry a load; a sodium girder would sag under its own weight.

Reason 2: It Melts in a Warm Room

Sodium melts at 98 °C against iron's 1538 °C. Even a modest fire — or a hot afternoon on a black-painted roof beam — would soften it. High melting point is exactly why transition metals are the structural metals.

Reason 3: It Reacts with Rain

Sodium reacts violently with water: 2Na + 2H₂O → 2NaOH + H₂. The first shower would destroy the structure and release flammable hydrogen. Iron does corrode, but slowly, and it can be painted or galvanised.

Chemistry Connection

This is the Group I versus transition table brought to life. Every property in that table matters somewhere: density decides what floats, melting point decides what survives a fire, reactivity decides what survives the rain. When a question asks you to "suggest why a transition metal is chosen", pick the property that matches the job.

A fertiliser plant in Kanpur runs the Haber process over a bed of iron. The catalyst has been in the reactor for three years and the plant chemist reports that essentially none of it has been consumed. A trainee asks why they do not simply add more iron to increase the yield of ammonia.

Explain why the iron is not used up, and evaluate the trainee's suggestion.

▼

Why the Catalyst Survives

A catalyst provides an alternative reaction pathway with a lower activation energy. It takes part in the mechanism — the nitrogen and hydrogen bond briefly to the iron surface — but it is regenerated at the end of each cycle, so it is chemically unchanged and can catalyse molecule after molecule for years.

Why More Iron Will Not Help the Yield

A catalyst changes only the rate, never the position of equilibrium. It speeds up the forward and reverse reactions equally, so the plant reaches the same equilibrium yield — just sooner. To raise the yield you must change the conditions: higher pressure, or a lower temperature (at the cost of rate). More catalyst is a rate lever, not a yield lever.

Chemistry Connection

Catalytic behaviour is one of the five transition-element properties, and this is why it matters commercially: an unreactive, high-melting metal can sit in a 450 °C reactor for years without degrading. A Group I metal in that reactor would have reacted with everything in sight before the first tonne of ammonia came out.

A conservator at the British Museum stores a fragile 3000-year-old bronze artefact in a sealed case filled with argon rather than air. A colleague suggests nitrogen would be cheaper, since air is 78% nitrogen and it is famously unreactive too.

Explain the choice of argon and comment fairly on the nitrogen suggestion.

▼

Why Not Air

Air contains oxygen and water vapour, which slowly corrode metals and degrade organic materials. Removing them stops the corrosion reactions entirely — you cannot oxidise something in the absence of an oxidising agent.

Why Argon Wins

Argon has a full outer shell (2,8,8), so it has genuinely no tendency to react with anything under any museum conditions. It is also denser than air, so it stays put in an open-topped case, and it makes up 0.9% of the atmosphere, so it is cheap to obtain from liquefied air.

Judging the Nitrogen Idea

The colleague has a reasonable point — nitrogen is cheap and is used for exactly this purpose in many cases. But nitrogen is N₂, a molecule held together by a strong triple bond; it is unreactive, not inert, and at high temperatures it does react (it forms nitrides with lithium and magnesium, and nitrogen oxides in engines). Argon needs no such caveat.

Chemistry Connection

Notice the sharp distinction the examiner is testing: nitrogen is unreactive because its triple bond is hard to break; argon is unreactive because there is nothing to break — it is not even bonded to itself, being monatomic. Same practical outcome, completely different reasons.

A party balloon supplier in Dubai advertises helium balloons at four times the price of the hydrogen ones a competitor sells. Hydrogen is lighter than helium and much cheaper to make. A physics student points out that hydrogen would give more lift per litre and argues the helium is a rip-off.

Settle the argument chemically.

▼

The Student Is Right About Lift

Hydrogen (Mₕ = 2) is genuinely lighter than helium (Mₕ = 4), so litre for litre it does give slightly more lift. On physics alone, hydrogen wins.

Why Chemistry Overrules Physics

Hydrogen is extremely flammable: 2H₂ + O₂ → 2H₂O is violently exothermic and a hydrogen-air mixture explodes from the smallest spark. Helium has a full outer shell of two electrons, so it has no tendency to react with anything at all — it cannot burn under any circumstances. The Hindenburg airship disaster of 1937 ended hydrogen's use in passenger craft for good.

Chemistry Connection

This is the standard exam answer for "why is helium used in balloons?" and it needs two reasons: low density (so it floats) and unreactive / non-flammable (so it is safe). Giving only the density earns one mark of two. Notice also that helium is the one noble gas whose full shell holds 2 electrons, not 8.

A hospital in Leeds runs an MRI scanner whose superconducting magnet is bathed in liquid helium at 4 K (−269 °C). Every few months a tanker delivers more. The purchasing manager notes that helium is the second most abundant element in the universe and cannot understand why it costs so much or why any is ever lost.

Explain why helium is scarce on Earth despite being abundant in the universe, and why its inertness matters here.

▼

Why It Escapes

Helium is monatomic and has the second-lowest relative atomic mass of any element, so its atoms move extremely fast. Once released into the atmosphere, they reach escape velocity and are lost into space permanently. And because helium is chemically inert, it never gets locked into any compound — there is no helium ore. The only supply is what accumulates underground from radioactive decay, trapped alongside natural gas.

Why the Inertness Is Essential Here

The magnet is a delicate assembly of metals and superconducting alloy. A coolant that reacted with any of it — corroding, oxidising or forming deposits — would destroy the scanner. Helium has a full outer shell, so it cools without reacting. It is also the only substance that stays liquid down to 4 K, which is precisely because the forces between its separate atoms are so weak.

Chemistry Connection

Every property in this scenario traces back to one fact: helium's outer shell is full. That makes it monatomic (so light and fast-moving), inert (so it forms no compounds and no ores), and weakly attracted to its own atoms (so it liquefies only at 4 K). One line of electronic structure explains a global supply crisis.

Practice Questions: 8.4

20 multiple choice questions. Click an option to check your answer.

Your Score 0 / 20

Question 1

Which property is characteristic of the transition elements but not of Group I metals?

A They conduct electricity

B They form coloured compounds

C They are solids at room temperature

D They form positive ions

All metals conduct, are solid (except mercury) and form positive ions. Coloured compounds — blue CuSO₄, green FeSO₄ — are a transition-only signature.

Question 2

Why are the noble gases unreactive?

A Because they are gases

B Because they have no electrons

C Because they have a full outer electron shell

D Because they have low densities

A full outer shell means no tendency to lose, gain or share electrons. Being a gas is a consequence, not a cause — chlorine is a gas and is extremely reactive.

Question 3

Which statement about the noble gases is correct?

A They exist as diatomic molecules

B They are monatomic

C They form 1− ions

D They are good conductors of electricity

A noble gas atom already has a full outer shell, so it has no reason to bond even to another atom of itself. Write He and Ar, never He₂ or Ar₂.

Question 4

Iron forms both FeCl₂ and FeCl₃. This illustrates which transition element property?

A Catalytic behaviour

B High density

C Variable oxidation number

D High melting point

Fe is +2 in FeCl₂ and +3 in FeCl₃. Sodium can only ever be +1, which is why sodium chloride needs no Roman numeral.

Question 5

Why is helium used in balloons rather than hydrogen?

A Helium is lighter than hydrogen

B Helium has a low density and, unlike hydrogen, is not flammable

C Helium is cheaper

D Helium dissolves in air

Hydrogen is actually lighter and cheaper, but it burns explosively. Helium's full outer shell makes it completely non-flammable — the safety reason wins.

Question 6

Which is used as the catalyst in the Haber process?

A Nickel

B Iron

C Vanadium(V) oxide

D Sodium

Iron for the Haber process, V₂O₅ for the Contact process, nickel for hydrogenating alkenes. All three are transition elements or their compounds — sodium is never a catalyst.

Question 7

The electronic configuration of argon is 2,8,8. What does this tell you?

A It readily gains electrons

B It forms a 2+ ion

C Its outer shell is full, so it does not need to react

D It is in Group II

Eight outer electrons is a full shell. Argon is in Group VIII (0), Period 3, and forms essentially no compounds.

Question 8

Which comparison between sodium and copper is correct?

A Sodium has the higher melting point

B Sodium is the denser metal

C Sodium is much more reactive with water

D Sodium forms coloured compounds

Sodium reacts violently with cold water; copper does not react with water at all. Copper is denser (8.9 vs 0.97), higher melting (1085 vs 98 °C) and forms the coloured compounds.

Question 9

Why is argon used inside filament light bulbs?

A It conducts electricity to the filament

B It is unreactive, so the hot filament cannot burn away

C It makes the light brighter

D It cools the filament to room temperature

In air the tungsten would oxidise and break within seconds. Argon excludes oxygen and, being inert, will not attack the filament itself even at 2000 °C.

Question 10

Copper(II) sulfate solution is blue and iron(II) sulfate solution is pale green. What does this indicate?

A Both are Group I compounds

B Both metals are transition elements, which form coloured compounds

C Both solutions are alkaline

D The sulfate ion is coloured

The colour comes from the transition metal ion, not the sulfate — sodium sulfate and magnesium sulfate solutions are both colourless.

Question 11

Which statement about a catalyst is correct?

A It increases the yield of product

B It speeds up the reaction and is chemically unchanged at the end

C It is used up during the reaction

D It raises the activation energy

A catalyst offers an alternative pathway of lower activation energy. It changes only the rate, never the equilibrium yield, and it is regenerated each cycle.

Question 12

Which element would you expect to be stored on an open laboratory shelf rather than under oil?

A Sodium

B Potassium

C Copper

D Lithium

Copper is a transition element and far less reactive — it does not react with water and tarnishes only slowly in air. All three alkali metals need oil.

Question 13

An element is a hard, dense metal that forms two different sulfates, one green and one yellow-brown. It is best described as

A a Group I metal

B a Group II metal

C a transition element

D a noble gas

Two different sulfates means variable oxidation number, and the colours confirm it — those two are iron(II) and iron(III). Groups I and II give one oxidation state and white compounds.

Question 14

What is the oxidation number of manganese in MnO₂?

A +2

B +3

C +4

D +7

Each oxygen is −2, so two oxygens total −4; the compound is neutral, so manganese must be +4. Hence the name manganese(IV) oxide. It is +7 in potassium manganate(VII).

Question 15

Neon is used in advertising signs. Which pair of properties explains this?

A It is reactive and cheap

B It is unreactive and glows red-orange when a current passes through it

C It is a good conductor of heat and is dense

D It burns with a bright flame

Every noble-gas use question wants inertness plus one specific physical property. Neon does not burn — nothing in Group VIII does.

Question 16

Which statement about Group I metals compared with transition elements is incorrect?

A Group I metals have lower melting points

B Group I metals are more reactive

C Group I metals show variable oxidation numbers

D Group I metals have lower densities

Group I metals are always +1 — that fixed single oxidation state is one of the clearest contrasts with the transition block.

Question 17

Why is argon used as a shielding gas in welding?

A It reacts with the metal to strengthen the joint

B It excludes oxygen and nitrogen from the hot metal without reacting itself

C It burns to provide the heat

D It dissolves impurities in the weld

Hot metal would otherwise form brittle oxides and nitrides. Argon blankets the weld pool and, having a full outer shell, takes no part in any reaction.

Question 18

Which set of formulae shows a transition element with variable oxidation number?

A NaCl and Na₂O

B MgCl₂ and MgO

C Cu₂O and CuO

D CaCl₂ and CaO

In Cu₂O copper is +1; in CuO it is +2 — the same element in two different states. In A, B and D the metal keeps the same charge in both compounds.

Question 19

The noble gases have very low boiling points. Why?

A Their covalent bonds are weak

B They exist as separate atoms with only weak forces between them

C They have no protons

D They are ionic

Being monatomic there are no covalent bonds at all — only very weak attractions between individual atoms, so very little energy is needed to separate them.

Question 20

A student writes: "Iron is a Group I metal because it is a metal that forms positive ions." The best correction is

A Iron is a non-metal

B Iron is in Group II because it forms Fe²⁺

C Iron is a transition element — it forms Fe²⁺ and Fe³⁺, has a high melting point and forms coloured compounds

D Iron is a noble gas

Forming positive ions is common to all metals and proves nothing about the group. Option B is a tempting trap: an element is placed by its position in the table, not by the charge of one of several possible ions.