Group Properties: Groups 1, 17 and 18

Last updated: 8 October 2026

Diagnose → Learn → Master: Every study-note section includes 15 free MCQs — start with 5 Diagnostic MCQs to identify gaps in your understanding, then complete 10 Mastery MCQs near the end to check what you have learned.

The Periodic Table allows us to predict chemical behaviour. We do not need to memorise every element separately. Once we understand why elements in the same group behave similarly, the position of an element becomes a useful clue to its reactions and properties.

Group 1, Group 17 and Group 18 illustrate this particularly well. Group 1 metals readily lose electrons, Group 17 halogens tend to gain electrons, while Group 18 noble gases already have filled outer electron shells and are therefore very unreactive.

The important skill is not simply remembering that one group becomes more reactive while another becomes less reactive. We want to understand what the elements actually do, recognise their characteristic properties and use the group trends to make predictions.

Diagnostic Quiz

What do you already know about groups?

Try these five questions before reading the chapter. They test the electron ideas that help explain the behaviour of Groups 1, 17 and 18.

5 diagnostic questions
Important misconceptions tested
Instant explanation for every answer option

Why do elements in the same group usually have similar chemical properties?

They have the same number of occupied electron shells.

Elements in the same period have the same number of occupied shells. Moving down a group adds additional occupied shells.

They have the same number of outer-shell electrons.

Correct. Chemical reactions involve the outer electrons, so similar outer-shell arrangements lead to similar chemical behaviour.

They have the same proton number.

Different elements have different proton numbers. The proton number identifies the element.

They contain the same total number of electrons.

Different elements in the same group contain different total numbers of electrons.

Elements in the same group have similar chemical properties because they have similar outer-shell electronic configurations.

Which electronic configuration belongs to a Group 17 element?

2,1

This atom has one outer-shell electron and belongs to Group 1.

2,8,2

This atom has two outer-shell electrons rather than seven.

2,8,7

Correct. Seven outer-shell electrons identify a Group 17 element.

2,8,8

This arrangement has a filled outer shell and represents a Group 18 noble gas.

Group 17 atoms have seven electrons in their outermost occupied shell.

What normally happens when a Group 1 atom forms a simple ion?

It loses one electron.

Correct. A Group 1 atom loses its single outer-shell electron and forms an ion with a 1+ charge.

It gains one electron.

Gaining an electron would produce a negative ion, which is not the characteristic behaviour of Group 1 metals.

It loses seven electrons.

Only the single outer-shell electron needs to be lost.

It gains seven electrons.

Losing one outer electron is much more favourable than gaining seven.

Group 1 atoms contain one outer-shell electron and normally lose it to form 1+ ions.

Which formula correctly represents elemental chlorine?

Cl represents one chlorine atom, but elemental chlorine normally exists as diatomic molecules.

Cl₂

Correct. Chlorine exists as diatomic Cl₂ molecules.

Cl⁻

Cl⁻ is a chloride ion, not elemental chlorine.

Cl₂⁻

This is not the formula used for elemental chlorine.

Halogens exist as diatomic molecules, so elemental chlorine is written Cl₂.

Why are noble gases very unreactive?

They have no electrons.

Noble-gas atoms contain electrons. Their important feature is a filled outer shell.

They have only one outer-shell electron.

One outer-shell electron is characteristic of Group 1.

They have filled outer electron shells.

Correct. Their stable electronic configurations give them little tendency to lose, gain or share electrons.

Their nuclei contain no neutrons.

Their low reactivity is explained by their electron arrangement, not neutron number.

Noble gases are very unreactive because their atoms already have filled outer electron shells.

Why group position predicts chemical behaviour

A group is a vertical column in the Periodic Table. For the main-group elements studied here, the group tells us something very important about the electrons in the outermost occupied shell.

Group 1 atoms have one outer-shell electron. Group 17 atoms have seven. Group 18 atoms have filled outer electron shells.

This matters because chemical reactions usually involve changes in the outer electrons. An atom may lose electrons, gain electrons or share electrons with another atom. The arrangement of the outer electrons therefore helps determine how an element reacts.

The group-property principle

Elements in the same group have similar chemical properties because their atoms have the same number of outer-shell electrons.

Group	Outer-shell arrangement	Typical electron behaviour	Typical result
Group 1	1 outer-shell electron	Loses 1 electron	Forms a 1+ ion
Group 17	7 outer-shell electrons	Gains 1 electron	Forms a 1− ion
Group 18	Filled outer shell	Little tendency to lose, gain or share electrons	Very unreactive

Elements in the same group are similar, but they are not identical. Moving down a group adds another occupied electron shell. The atoms therefore become larger, and the outer electrons experience more shielding from inner-shell electrons. These changes produce the trends that we observe.

Group 1: the alkali metals

Group 1 contains metals such as lithium, sodium and potassium. They are called the alkali metals because they react with water to produce soluble metal hydroxides, giving alkaline solutions.

Electronic structure and ions

Every Group 1 atom has one electron in its outer shell. During a chemical reaction, it can obtain a stable electron arrangement by losing this electron.

Na → Na⁺ + e⁻

Sodium therefore forms Na⁺ ions. In the same way, lithium forms Li⁺ and potassium forms K⁺.

Atom versus ion

Group 1 atom: has one outer-shell electron.

Group 1 ion: has lost that electron and carries a 1+ charge.

Physical properties of Group 1 metals

Lithium, sodium and potassium are metals, so they conduct electricity and have a shiny metallic surface when freshly cut. However, they are unusual compared with familiar structural metals such as iron.

The first three Group 1 metals are:

soft and can be cut with a knife;
relatively low in density;
relatively low in melting point compared with many other metals; and
very reactive, so freshly cut surfaces quickly become dull in air.

Lithium, sodium and potassium are all less dense than water, so pieces of these metals float when placed on water.

Physical trends down Group 1

Property	Trend down Group 1
Atomic size	Increases
Softness	Generally increases
Melting point	Decreases
Reactivity	Increases
Density	Generally increases, although the trend is not perfectly regular

Be careful with density. The general pattern is an increase down the group, but potassium is less dense than sodium. Density should therefore not be treated as a perfectly smooth trend.

Why does the melting point decrease down Group 1?

Down the group, the metal ions become larger while each still carries a 1+ charge. The attraction between the positive metal ions and the delocalised electrons becomes weaker.

Less energy is therefore needed to overcome the metallic bonding, so the melting point decreases.

Melting-point trend in Group 1

Down the group, the metal ions become larger, metallic bonding becomes weaker, and the melting point decreases.

Reaction of Group 1 metals with water

Group 1 metals react with cold water to form a metal hydroxide and hydrogen gas.

Group 1 metal + water → metal hydroxide + hydrogen

If M represents a Group 1 metal:

2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)

For sodium:

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Sodium hydroxide dissolves in the water. The resulting solution is alkaline because it contains OH⁻ ions.

Hydrogen can be confirmed using a lighted splint. The gas burns with a squeaky pop.

Comparing lithium, sodium and potassium with water

Metal	Typical observations	Relative reactivity
Lithium	Floats, fizzes and moves across the water. It gradually becomes smaller and disappears.	Least reactive of the three
Sodium	Floats, fizzes more vigorously and moves rapidly. The heat released can melt the sodium into a small silvery ball.	More reactive than lithium
Potassium	Reacts very vigorously and usually ignites, producing a characteristic lilac flame.	Most reactive of the three

Comparison of lithium, sodium and potassium reacting with water, showing increasingly vigorous reactions down Group 1 — Lithium, sodium and potassium all react with water to form a metal hydroxide and hydrogen, but the reaction becomes increasingly vigorous down Group 1.

Notice that all three metals undergo the same type of reaction. What changes is the vigour of the reaction. This is exactly what we expect from elements in the same group: similar chemical behaviour, but a systematic change in reactivity.

Why does reactivity increase down Group 1?

A Group 1 metal reacts by losing its single outer-shell electron. Therefore, the easier it is to remove this electron, the more reactive the metal is.

As we move down Group 1:

the number of occupied electron shells increases;
the outer electron is farther from the nucleus;
shielding by inner-shell electrons increases; and
the attraction between the positive nucleus and the outer electron becomes weaker.

The outer electron is therefore lost more easily. This is why potassium is more reactive than sodium, and sodium is more reactive than lithium.

Reactivity trend in Group 1

More electron shells → greater distance and shielding → weaker attraction for the outer electron → electron lost more easily → greater reactivity.

Group 17: the halogens

Group 17 contains the halogens. Familiar examples include fluorine, chlorine, bromine and iodine.

Unlike Group 1, the halogens are non-metals. Their atoms have seven outer-shell electrons, so their chemical behaviour commonly involves gaining one electron rather than losing one.

Halogen molecules and halide ions

The halogens exist as diatomic molecules. Each molecule contains two atoms joined together.

F₂ Cl₂ Br₂ I₂

During many reactions, a halogen atom gains one electron and forms a halide ion with a 1− charge.

Cl + e⁻ → Cl⁻

It is important to distinguish the element from its ion.

Halogen versus halide

Chlorine: Cl₂, a diatomic element.

Chloride: Cl⁻, a negatively charged ion.

Similarly, bromine is Br₂ while bromide is Br⁻, and iodine is I₂ while iodide is I⁻.

Colours and physical states of the halogens

The physical properties of the halogens change steadily down Group 17. The elements become darker in colour, while their melting points and boiling points increase.

Halogen	State at room temperature	Appearance of the element
Chlorine, Cl₂	Gas	Yellow-green
Bromine, Br₂	Liquid	Red-brown
Iodine, I₂	Solid	Grey-black; forms a purple vapour when heated

The change in physical state at room temperature can be summarised as:

chlorine: gas → bromine: liquid → iodine: solid

Moving down the group, the halogen molecules become larger and contain more electrons. The attractive forces between neighbouring molecules become stronger.

More energy is therefore required to separate the molecules. The melting points and boiling points increase down the group.

Physical trend in Group 17

Down Group 17, the molecules become larger and intermolecular attractions become stronger. More energy is needed to separate the molecules, so melting points and boiling points increase.

The covalent bond inside each Cl₂, Br₂ or I₂ molecule is not broken when the substance melts or boils. It is the attraction between separate molecules that is overcome.

Why does reactivity decrease down Group 17?

A halogen reacts by gaining an electron. The more strongly a halogen atom attracts an incoming electron, the more reactive it is.

As we move down Group 17:

the number of occupied electron shells increases;
the outer shell becomes farther from the nucleus;
shielding by inner-shell electrons increases; and
the attraction between the nucleus and an incoming electron becomes weaker.

The atom therefore gains an electron less readily. Reactivity decreases down Group 17.

F₂ > Cl₂ > Br₂ > I₂

Reactivity trend in Group 17

More electron shells → greater distance and shielding → weaker attraction for an incoming electron → electron gained less readily → lower reactivity.

Why are the Group 1 and Group 17 trends opposite?

Down both groups, atomic size and shielding increase. However, the atoms are trying to do different things.

Group	What the atom does	Effect of weaker attraction down the group	Reactivity trend
Group 1	Loses an outer electron	The electron becomes easier to lose	Increases
Group 17	Gains an incoming electron	The incoming electron is attracted less strongly	Decreases

Therefore, simply memorising “more shielding down a group” is not enough. You must ask whether the element reacts by losing an electron or by gaining one.

Halogen displacement reactions

The reactivity order of the halogens can be demonstrated using displacement reactions.

A more reactive halogen can displace the halide ions of a less reactive halogen from solution.

The halogen displacement rule

A halogen can displace halide ions belonging to a halogen below it in Group 17.

For example, chlorine is more reactive than bromine. Chlorine therefore displaces bromide ions from potassium bromide solution.

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)

Potassium ions do not take part in the chemical change. The ionic equation is:

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

Bromine is formed, so an orange solution is observed.

Chlorine can also displace iodide ions:

Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)

Iodine in water gives a brown colour.

Halogen displacement reactions showing chlorine displacing bromide and iodide ions while bromine cannot displace chloride ions — A more reactive halogen displaces the halide ions of a less reactive halogen. Chlorine can displace bromide and iodide ions, while bromine cannot displace chloride ions.

Predicting halogen displacement reactions

Instead of memorising every possible reaction, begin with the reactivity order:

Cl₂ > Br₂ > I₂

Then ask: Is the added halogen more reactive than the halogen represented by the halide ion?

Halogen added	Chloride ions, Cl⁻	Bromide ions, Br⁻	Iodide ions, I⁻
Cl₂	No displacement	Br₂ formed	I₂ formed
Br₂	No displacement	No displacement	I₂ formed
I₂	No displacement	No displacement	No displacement

Use reactivity, not physical state

Correct reasoning: chlorine displaces Br⁻ because chlorine is more reactive than bromine.

Incorrect reasoning: chlorine displaces Br⁻ because chlorine is a gas and bromine is a liquid. Physical state does not determine which halogen is more reactive.

Group 18: the noble gases

Group 18 contains helium, neon, argon and the other noble gases. Their chemical behaviour is very different from Groups 1 and 17 because their atoms already have filled outer electron shells.

Why are noble gases very unreactive?

Helium has two electrons in its first and only shell, so that shell is filled. Neon, argon and the noble gases below them have eight electrons in their outer shells.

Because the outer shell is already filled, a noble-gas atom has little tendency to lose electrons, gain electrons or share electrons with another atom. Noble gases are therefore very unreactive.

A stable outer electron shell

Noble gases are very unreactive because their atoms already have filled outer electron shells.

At O-Level, noble gases are normally treated as very unreactive under ordinary conditions. Some heavier noble gases can form compounds under unusual conditions, but this is not needed for the standard Group 18 chemistry here.

Noble gases are monatomic

Noble gases exist as separate atoms rather than as molecules. They are described as monatomic.

He Ne Ar

This contrasts with the Group 17 halogens, which exist as diatomic molecules such as Cl₂ and Br₂.

Physical trends down Group 18

Noble gases have very low melting points and boiling points. Moving down the group, however, the atoms become larger and contain more electrons.

Attractions between neighbouring atoms become stronger, so more energy is needed to separate them. Their melting points and boiling points therefore increase down the group.

Their densities also increase down the group as the atoms become progressively more massive.

Uses of noble gases

A good Chemistry answer should connect a use to the property that makes the substance suitable. For noble gases, their very low reactivity is particularly important.

Noble gas	Use	Why it is suitable
Helium	Balloons	Helium has a low density and is non-flammable, making it safer than hydrogen.
Neon	Advertising signs and discharge lamps	Neon emits characteristic light when an electric current passes through it at low pressure.
Argon	Filament lamps	Argon is very unreactive, so it reduces unwanted reaction of the hot metal filament.
Argon	Protective atmosphere during welding	Argon provides an unreactive atmosphere around hot metals.

Always connect a use to a property

Incomplete: Argon is used in filament lamps because it is a noble gas.

Better: Argon is used because it is very unreactive, so it does not readily react with the hot filament.

Comparing Groups 1, 17 and 18

The three groups become much easier to remember when their main features are placed side by side.

Feature	Group 1	Group 17	Group 18
Type of element	Metal	Non-metal	Non-metal
Outer-shell electrons	1	7	Filled outer shell
Usual electron behaviour	Loses 1 electron	Gains 1 electron	Little tendency to lose or gain electrons
Typical ion	1+	1−	No common simple ion
Structure of the element	Giant metallic structure	Diatomic molecules	Separate atoms
Reactivity down the group	Increases	Decreases	Remains very low

Predicting the properties of unfamiliar elements

One of the most useful features of the Periodic Table is prediction. If you know the position of an unfamiliar element and understand the trend in its group, you can make sensible predictions without having studied that element directly.

A reliable method is:

Identify the group. This tells you the outer electron arrangement and general chemical behaviour.
Identify the relevant trend. Decide whether the property increases or decreases down the group.
Locate the unfamiliar element. Decide whether it lies above or below an element whose properties you know.
Apply the trend and justify the prediction.

Worked prediction: rubidium

Rubidium lies below potassium in Group 1.

Since Group 1 metals become more reactive down the group, rubidium should be more reactive than potassium. Its outer electron is farther from the nucleus and more strongly shielded, so it is lost more easily.

Rubidium should therefore react very vigorously with water to form rubidium hydroxide and hydrogen.

2Rb(s) + 2H₂O(l) → 2RbOH(aq) + H₂(g)

Worked prediction: an element below iodine

Suppose an unfamiliar element lies below iodine in Group 17. Applying the Group 17 trends, we would predict that it:

is less reactive than iodine;
has a higher melting point and boiling point than iodine;
has a darker appearance; and
is likely to be a solid at room temperature.

The important skill is the reasoning. We know the direction of the trend and apply it to an element lower in the same group.

Common mistakes in group-property questions

Common mistake	Why it is wrong	Better idea
Using the same reactivity trend for Groups 1 and 17.	The two groups react by different electron changes.	Group 1 loses an electron more easily down the group, while Group 17 gains an electron less easily.
Saying Group 1 becomes more reactive simply because the atoms are larger.	Atomic size alone does not complete the explanation.	Explain greater distance and shielding, weaker attraction and easier loss of the outer electron.
Saying Group 17 becomes more reactive because shielding increases.	Greater shielding weakens attraction for an incoming electron.	State that Group 17 becomes less reactive down the group.
Writing Cl⁻ when elemental chlorine is required.	Cl⁻ is a chloride ion rather than chlorine.	Write chlorine as Cl₂ and chloride as Cl⁻.
Using physical state to decide whether a halogen displacement occurs.	Displacement depends on relative reactivity rather than whether a halogen is a gas, liquid or solid.	Compare the positions of the halogens in Group 17.
Saying boiling iodine breaks the covalent I–I bond.	The iodine molecules remain intact during a physical change.	State that intermolecular attractions between I₂ molecules are overcome.
Saying noble gases have empty outer shells.	Their low reactivity comes from a filled outer shell.	State that helium has a filled first shell and the other noble gases have filled outer shells.
Giving a noble-gas use without explaining the relevant property.	The chemical reason for the use is missing.	Link the use to a property, such as argon's very low reactivity.
Calling the density trend in Group 1 perfectly regular.	Potassium is less dense than sodium.	Describe density as generally increasing down the group, with exceptions.

Group properties checklist

Understanding groups

Can I explain why elements in the same group have similar chemical properties?
Can I link group position to outer-shell electrons?
Can I distinguish a shared property from a trend down a group?
Can I use a known trend to predict an unfamiliar element?

Group 1

Can I explain why Group 1 metals form 1+ ions?
Can I write the reaction of a Group 1 metal with water?
Can I compare lithium, sodium and potassium reacting with water?
Can I explain why reactivity increases down the group?

Group 17

Can I write halogens as F₂, Cl₂, Br₂ and I₂?
Can I state the colour and physical-state trends?
Can I explain why reactivity decreases down the group?
Can I predict halogen displacement reactions?

Group 18

Can I explain why noble gases are very unreactive?
Can I explain what monatomic means?
Can I describe physical trends down Group 18?
Can I link helium, neon and argon uses to suitable properties?

Mastery Quiz

Can you apply the group trends?

These ten questions go beyond simple recall. Use electronic configuration, observations and group trends to reason your way to each answer.

10 mastery questions
Application and misconception testing
Instant explanation for every answer option

Element X is a Group 1 metal below potassium. Which prediction about X is most reasonable?

X reacts less vigorously with water because its atom contains more electrons.

More electrons do not make a Group 1 metal less reactive. Increased distance and shielding make the outer electron easier to lose.

X reacts more vigorously with water because its outer electron is lost more easily.

Correct. X is below potassium, so greater distance and shielding weaken attraction for its outer electron.

X is unreactive because it has more occupied electron shells.

Additional shells increase shielding; they do not give a Group 1 atom a filled outer shell.

X gains one electron more easily than potassium.

Group 1 metals characteristically react by losing their single outer-shell electron.

Group 1 reactivity increases down the group because the outer electron becomes farther from the nucleus and more shielded, so it is lost more easily.

A student compares sodium and potassium. Which combination of properties is correct?

Potassium is less reactive and has a higher melting point.

Both trends are reversed. Reactivity increases and melting point decreases down Group 1.

Potassium is less reactive and has a lower melting point.

The melting-point comparison is correct, but potassium is more reactive than sodium.

Potassium is more reactive and has a lower melting point.

Correct. Potassium loses its outer electron more readily and has weaker metallic bonding than sodium.

Potassium is more reactive and has a higher melting point.

The reactivity comparison is correct, but melting point decreases down Group 1.

Potassium lies below sodium in Group 1. It is more reactive but has a lower melting point.

A Group 1 metal floats on water, reacts very vigorously and burns with a lilac flame. Which metal is it most likely to be?

Lithium

Lithium floats and fizzes, but its reaction is much less vigorous.

Sodium

Sodium reacts vigorously and often melts into a ball, but the lilac flame is characteristic of potassium.

Potassium

Correct. Potassium reacts very vigorously with water and commonly burns with a lilac flame.

Magnesium

Magnesium is not a Group 1 metal and does not show this characteristic reaction with cold water.

Potassium is the most reactive of lithium, sodium and potassium and characteristically produces a lilac flame.

Halogen Y is below bromine in Group 17. Which statement is the best prediction?

Y is more reactive than bromine and has a lower melting point.

Both predictions oppose the Group 17 trends.

Y is more reactive than bromine and has a higher melting point.

The melting-point prediction is reasonable, but reactivity decreases down Group 17.

Y is less reactive than bromine and has a higher melting point.

Correct. Reactivity decreases while melting point increases down Group 17.

Y has the same melting point and reactivity as bromine.

Members of a group behave similarly, but their physical properties and reactivities change down the group.

Down Group 17, reactivity decreases while melting point and boiling point increase.

Chlorine water is added to potassium bromide solution. Which observation and explanation are correct?

No reaction occurs because bromine is a liquid.

Physical state does not determine displacement. Relative halogen reactivity does.

An orange solution forms because Br₂ is produced.

Correct. Chlorine is more reactive than bromine and displaces bromide ions.

A colourless solution forms because chlorine changes into Cl₂.

Chlorine is already present as Cl₂. It gains electrons and forms Cl⁻ ions.

A purple gas forms because potassium is displaced.

Potassium ions are spectator ions. The displacement involves chlorine and bromide ions.

Chlorine is more reactive than bromine and displaces Br⁻ ions to form Br₂.

Which ionic equation correctly represents chlorine displacing bromide ions?

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

Correct. The atoms and charge are balanced and chlorine displaces bromide.

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

This is the reverse reaction. Bromine is less reactive and cannot displace chloride ions.

Cl₂(aq) + Br⁻(aq) → Cl⁻(aq) + Br₂(aq)

The direction is sensible, but the equation is not balanced.

Cl⁻(aq) + Br⁻(aq) → ClBr(aq)

Halogen displacement is not represented by combining two halide ions.

Chlorine gains electrons to form Cl⁻ while bromide ions lose electrons to form Br₂.

Iodine has a higher boiling point than chlorine. Which explanation is most appropriate?

The I–I covalent bond is ionic.

The bond within an iodine molecule is covalent.

Iodine atoms have fewer electrons than chlorine atoms.

Iodine atoms contain more electrons than chlorine atoms.

I₂ molecules experience stronger intermolecular attractions.

Correct. I₂ molecules are larger and contain more electrons, so the attractions between neighbouring molecules are stronger.

Boiling iodine requires breaking the covalent bonds inside I₂.

Boiling overcomes attractions between molecules rather than breaking the I–I covalent bonds.

Larger halogen molecules have stronger intermolecular attractions, so more energy is required to separate them.

Why is argon used around a hot metal filament in a lamp?

Argon reacts rapidly with the filament and cools it.

Argon's usefulness comes from its very low reactivity.

Argon supplies oxygen for the filament.

Oxygen would encourage unwanted reaction of the hot filament.

Argon is very unreactive and provides a protective atmosphere.

Correct. Its filled outer shell makes argon very unreactive.

Argon conducts electricity better than metals.

Electrical conductivity is not the reason argon is used around the filament.

Argon is very unreactive, so it provides an atmosphere that reduces unwanted reaction of the hot filament.

An unfamiliar element lies below iodine in Group 17. Which set of predictions follows the group trends?

More reactive, lower melting point and lighter colour.

All three predictions run opposite to the Group 17 trends.

More reactive, higher melting point and darker colour.

The physical predictions fit the trend, but reactivity should decrease.

Less reactive, higher melting point and darker colour.

Correct. These are the expected trends on moving below iodine in Group 17.

Less reactive, lower melting point and lighter colour.

The reactivity prediction is correct, but melting point increases and the elements become darker down the group.

Down Group 17, reactivity decreases, melting and boiling points increase, and colour becomes darker.

Three atoms have electronic configurations P: 2,8,1; Q: 2,8,7; R: 2,8,8. Which statement is correct?

P and Q are both noble gases because they have three occupied shells.

Group membership depends on the outer-shell arrangement here, not simply the number of occupied shells.

P tends to lose one electron, Q tends to gain one electron, and R is very unreactive.

Correct. This follows directly from their outer electron arrangements.

P tends to gain one electron while Q tends to lose one electron.

The electron changes have been reversed. Group 1 loses one electron, while Group 17 gains one.

R is highly reactive because it has the greatest number of outer-shell electrons.

Eight outer-shell electrons give R a filled and stable outer shell, making it very unreactive.

P represents a Group 1 atom, Q represents a Group 17 atom and R has the filled outer shell characteristic of Group 18.

Frequently asked questions

Questions students often ask about group properties

Once the main trends are clear, most remaining difficulties come from explaining why those trends occur or applying them to unfamiliar situations.

Do all elements in the same group have exactly the same properties?

No. They have similar chemical properties because they have the same number of outer-shell electrons. However, properties such as atomic size, melting point, boiling point, density and reactivity change as you move down the group.

Lithium and potassium, for example, both react with water to form a hydroxide and hydrogen, but potassium reacts much more vigorously.

Why are Group 1 metals stored under oil?

Group 1 metals react readily with substances in the air, including oxygen and water vapour. Storing them under oil reduces contact with air and moisture.

Why does sodium form a ball when it reacts with water?

The reaction between sodium and water releases heat. Sodium has a relatively low melting point, so this heat can melt it. The molten sodium forms a small approximately spherical droplet on the water surface.

Why do Group 1 and Group 17 show opposite reactivity trends?

Down both groups, atomic size and electron shielding increase. For Group 1, this makes the outer electron easier to lose, so reactivity increases.

For Group 17, the nucleus attracts an incoming electron less strongly, so the electron is harder to gain and reactivity decreases.

Why are halogens written as Cl₂ and Br₂ instead of Cl and Br?

Elemental halogens exist as diatomic molecules. Each molecule contains two halogen atoms joined by a covalent bond, so chlorine is Cl₂ and bromine is Br₂.

Do not confuse these molecules with the halide ions Cl⁻ and Br⁻.

Why can chlorine displace bromide ions but bromine cannot displace chloride ions?

Chlorine is more reactive than bromine, so Cl₂ can gain electrons from Br⁻ ions and form Cl⁻ ions. Bromine is less reactive than chlorine, so Br₂ cannot remove electrons from Cl⁻ ions.

The direction of the reaction follows the Group 17 reactivity order.

Why may iodine be described as grey-black, brown or purple?

The observed colour depends on the physical form and surroundings. Solid iodine is grey-black, iodine vapour is purple, and iodine dissolved in water appears brown.

Always read the experimental conditions before stating the expected colour.

Are noble gases completely unable to react?

At O-Level, noble gases are treated as very unreactive because their atoms have filled outer electron shells.

Some heavier noble gases can form compounds under unusual conditions, but this does not change the explanation required for ordinary O-Level Chemistry.

Why is helium safer than hydrogen for filling balloons?

Both gases have low densities, but helium is very unreactive and non-flammable. Hydrogen is flammable and can form explosive mixtures with air, so helium is safer.

How should I answer a question asking me to predict the property of an unfamiliar element?

First identify the group. Then state the known trend and locate the unfamiliar element relative to elements whose properties are known. Finally, apply the trend and, where required, explain it using electron shells, shielding and nuclear attraction.

A prediction is strongest when it is justified rather than presented as an unsupported guess.

Group properties show why the Periodic Table is such a powerful tool. Group 1 metals become more reactive because their outer electron becomes easier to lose, while Group 17 halogens become less reactive because an incoming electron is attracted less strongly. Group 18 atoms already have filled outer electron shells and are therefore very unreactive. Once you connect these patterns to electronic configuration, you can move beyond memorising individual facts and begin predicting the chemistry of unfamiliar elements.