What Are the Halogens?

The halogens are a family of reactive non-metals found in Group 17 of the modern periodic table.

In some older school systems and textbooks, this group is called Group 7. Both names may therefore be encountered.

The main halogens are:

• fluorine (F)
• chlorine (Cl)
• bromine (Br)
• iodine (I)
• astatine (At)
• tennessine (Ts)

For most introductory chemistry, the main focus is on:

fluorine, chlorine, bromine, and iodine.

They have similar chemical properties because they all have seven valence electrons.

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Finding the Halogens on the Periodic Table

Group 17 is the second column from the right side of the periodic table, immediately beside the noble gases.

Moving down the group:

F → Cl → Br → I → At → Ts

Like other groups in the periodic table, the halogens show predictable patterns in their physical and chemical properties.

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Electron Arrangements of the Halogens

All halogens have seven electrons in their outer shell.

For example:

Fluorine:

2,7

Chlorine:

2,8,7

Although the number of occupied electron shells increases down the group, every halogen has the same basic outer-shell arrangement:

7 valence electrons

This explains their similar chemical behaviour.

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Halogens Are Non-Metals

Unlike the Group 1 alkali metals, halogens are non-metals.

They generally:

• do not conduct electricity well
• have relatively low melting and boiling points compared with many metals
• are not malleable
• are not ductile
• tend to gain electrons during chemical reactions

Their physical properties change noticeably as we move down the group.

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Halogens Exist as Diatomic Molecules

In their elemental form, halogens normally exist as diatomic molecules.

Diatomic means that each molecule contains two atoms.

Therefore:

Fluorine = F₂

Chlorine = Cl₂

Bromine = Br₂

Iodine = I₂

When writing equations involving elemental halogens, it is important to use these diatomic formulas.

For example, write:

Cl₂

not simply:

Cl

for elemental chlorine.

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Physical States of the Halogens

One of the clearest trends down Group 17 is the change in physical state at room temperature.

Fluorine – gas

Chlorine – gas

Bromine – liquid

Iodine – solid

Moving down the group, melting points and boiling points generally increase.

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Why the Physical State Changes

Halogen molecules become larger as we move down the group.

For example:

F₂ < Cl₂ < Br₂ < I₂

Larger molecules contain more electrons and have more easily distorted electron clouds.

This strengthens London dispersion forces between molecules.

More energy is therefore required to separate the molecules.

As a result:

melting point increases down the group

and:

boiling point increases down the group.

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Colours of the Halogens

The halogens also show distinctive colours.

Common descriptions include:

Fluorine: pale yellow gas

Chlorine: greenish-yellow gas

Bromine: red-brown liquid with brown-orange vapour

Iodine: grey-black solid that produces purple/violet vapour when heated

These visible differences can help distinguish the halogens during demonstrations and investigations.

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A Physical Trend Down the Group

As we move down Group 17:

• atomic size increases
• molecular size increases
• melting point generally increases
• boiling point generally increases
• physical state changes from gases to liquid to solid

This is an important reminder that elements in the same group have similar chemical behaviour but not identical physical properties.

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Halogens Gain Electrons

The chemical behaviour of halogens is strongly connected to their seven valence electrons.

A halogen atom needs only one additional electron to complete its outer shell.

Therefore, halogens tend to:

gain one electron

during reactions.

For example:

Cl + e⁻ → Cl⁻

The resulting ion is called a:

chloride ion

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Why Halogens Form −1 Ions

Consider chlorine.

Its electron arrangement is:

2,8,7

Chlorine gains one electron.

Its new arrangement becomes:

2,8,8

The chlorine atom originally has equal numbers of protons and electrons.

After gaining one negatively charged electron, it has one more electron than protons.

Therefore, the ion has a charge of:

−1

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Halide Ions

When halogen atoms gain electrons, they form halide ions.

Fluorine forms:

F⁻ – fluoride

Chlorine forms:

Cl⁻ – chloride

Bromine forms:

Br⁻ – bromide

Iodine forms:

I⁻ – iodide

The general ion-formation equation is:

X + e⁻ → X⁻

where X represents a halogen atom.

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Halogen vs Halide

These two terms are easy to confuse.

A halogen is an element in Group 17.

Examples:

Cl₂, Br₂, I₂

A halide is a −1 ion formed when a halogen atom gains an electron.

Examples:

Cl⁻, Br⁻, I⁻

Therefore:

chlorine → chloride

bromine → bromide

iodine → iodide

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Halogens React with Metals

Halogens readily react with metals.

For example, sodium reacts with chlorine:

2Na + Cl₂ → 2NaCl

Sodium loses an electron:

Na → Na⁺ + e⁻

Chlorine gains an electron:

Cl + e⁻ → Cl⁻

The resulting oppositely charged ions attract each other and form an ionic compound.

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Metal Halides

Compounds formed between metals and halogens are called metal halides.

Examples include:

NaCl – sodium chloride

KBr – potassium bromide

LiF – lithium fluoride

MgCl₂ – magnesium chloride

CaBr₂ – calcium bromide

The formula depends on the charges of the ions involved.

For example:

Mg²⁺ requires two Cl⁻ ions:

MgCl₂

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Reactivity of the Halogens

The most important chemical trend in Group 17 is:

reactivity decreases down the group.

Therefore:

F > Cl > Br > I

Fluorine is more reactive than chlorine.

Chlorine is more reactive than bromine.

Bromine is more reactive than iodine.

This is the opposite of the Group 1 reactivity trend.

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Why Halogen Reactivity Decreases

Halogens react by gaining an electron.

Moving down Group 17:

• the number of occupied electron shells increases
• atomic radius increases
• shielding increases
• an incoming electron is farther from the nucleus
• the effective attraction between the nucleus and an incoming electron becomes weaker

Therefore, the atom gains an electron less readily.

As a result:

reactivity decreases down Group 17.

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The Reactivity Explanation Chain

A strong explanation can be written as:

Down Group 17

↓

More occupied electron shells

↓

Greater atomic radius

↓

Greater shielding

↓

Weaker effective attraction for an incoming electron

↓

Electron is gained less readily

↓

−1 ion forms less readily

↓

Reactivity decreases

This explains the trend rather than simply memorizing it.

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Electronegativity and the Halogens

The electronegativity trend supports this explanation.

Electronegativity generally decreases down Group 17.

Fluorine has an extremely strong attraction for bonding electrons.

Farther down the group, increased atomic radius and shielding reduce this attraction.

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Group 1 and Group 17 Compared

The two groups show opposite reactivity trends.

Group 1

Atoms react by losing one electron.

Reactivity:

increases down the group.

Group 17

Atoms react by gaining one electron.

Reactivity:

decreases down the group.

The difference makes sense once we consider whether the atom needs to lose or gain an electron.

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What Is a Displacement Reaction?

A displacement reaction occurs when a more reactive element replaces a less reactive element in a compound.

For halogens:

a more reactive halogen can displace a less reactive halogen from a halide compound.

This provides an experimental way to compare halogen reactivity.

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The Rule for Halogen Displacement

Remember:

A more reactive halogen displaces a less reactive halogen.

Since:

Cl > Br > I

chlorine can displace:

• bromide ions
• iodide ions

Bromine can displace:

• iodide ions

Iodine cannot displace chlorine or bromine from their halides.

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Chlorine and Potassium Bromide

Consider:

chlorine + potassium bromide

Chlorine is more reactive than bromine.

Therefore, chlorine displaces bromine.

Word equation:

chlorine + potassium bromide → potassium chloride + bromine

Symbol equation:

Cl₂ + 2KBr → 2KCl + Br₂

Bromine is released from the bromide compound.

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Chlorine and Potassium Iodide

Chlorine is also more reactive than iodine.

Therefore:

chlorine + potassium iodide → potassium chloride + iodine

Symbol equation:

Cl₂ + 2KI → 2KCl + I₂

The chlorine has displaced iodine from the iodide compound.

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Bromine and Potassium Iodide

Bromine is more reactive than iodine.

Therefore:

bromine + potassium iodide → potassium bromide + iodine

Symbol equation:

Br₂ + 2KI → 2KBr + I₂

Again, the more reactive halogen replaces the less reactive halogen.

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When No Reaction Occurs

Now consider:

bromine + potassium chloride

Would bromine displace chlorine?

No.

Chlorine is more reactive than bromine.

A less reactive halogen cannot displace a more reactive halogen.

Therefore:

Br₂ + KCl → no reaction

Similarly:

I₂ + KBr → no reaction

and:

I₂ + KCl → no reaction

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Displacement Reaction Pattern

For chlorine, bromine, and iodine:

Chlorine

Displaces bromide and iodide.

Bromine

Displaces iodide but not chloride.

Iodine

Displaces neither chloride nor bromide.

The results provide evidence for:

Cl > Br > I

in reactivity.

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Understanding What Is Actually Displaced

Consider:

Cl₂ + 2KBr → 2KCl + Br₂

It is useful to focus on the halogens.

Chlorine begins as:

Cl₂

Bromine begins as:

Br⁻

After the reaction:

chlorine becomes:

Cl⁻

and bromide becomes:

Br₂

The potassium ions do not change.

They are spectator ions.

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Ionic Equation for Halogen Displacement

The ionic equation for chlorine reacting with bromide ions is:

Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

This equation shows the important chemical change more clearly.

Chlorine gains electrons.

Bromide ions lose electrons.

Therefore, halogen displacement reactions are also examples of redox reactions.

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Electron Transfer in Displacement

Consider:

Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

Chlorine gains electrons:

Cl₂ + 2e⁻ → 2Cl⁻

Bromide loses electrons:

2Br⁻ → Br₂ + 2e⁻

The more reactive halogen has the greater tendency to gain electrons.

Therefore, chlorine can take electrons from bromide ions.

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Explaining Observations

Halogen displacement reactions often produce visible colour changes because the halogens have distinctive colours.

For example, if chlorine reacts with bromide ions and bromine forms, the appearance of bromine's characteristic colour provides evidence that a displacement reaction has occurred.

If iodine forms, its characteristic colour can provide evidence for iodine production.

The exact observed colour can depend on the solvent and experimental conditions.

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Observation vs Explanation

As with other chemical reactions, distinguish between what is observed and what is explained.

Observation:

The solution changes colour.

Explanation:

A different halogen has formed.

Observation:

The characteristic colour of bromine appears.

Explanation:

Chlorine has displaced bromine from bromide ions because chlorine is more reactive.

This distinction is important when writing experimental conclusions.

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Worked Example 1

Will chlorine react with sodium bromide?

Compare the halogens:

Cl > Br

Chlorine is more reactive.

Therefore:

Yes, a displacement reaction occurs.

Equation:

Cl₂ + 2NaBr → 2NaCl + Br₂

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Worked Example 2

Will iodine react with potassium chloride?

Compare:

Cl > I

Iodine is less reactive than chlorine.

Therefore:

No displacement reaction occurs.

Iodine cannot displace chloride ions.

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Worked Example 3

Will bromine react with sodium iodide?

Compare:

Br > I

Bromine is more reactive.

Therefore:

Yes.

Equation:

Br₂ + 2NaI → 2NaBr + I₂

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Worked Example 4

An unknown halogen X displaces bromide ions but does not displace chloride ions.

Where is X likely to be in the reactivity series?

X is more reactive than bromine because it displaces bromide.

X is less reactive than chlorine because it cannot displace chloride.

Therefore:

Cl > X > Br

This shows how displacement experiments can be used to determine the relative position of an unfamiliar halogen.

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Worked Example 5

Why is chlorine more reactive than iodine?

Chlorine has fewer occupied electron shells than iodine.

Therefore:

• chlorine has a smaller atomic radius
• chlorine has less electron shielding
• an incoming electron experiences stronger effective attraction to chlorine's nucleus

Chlorine therefore gains an electron more readily.

As a result:

chlorine is more reactive than iodine.

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Predicting the Reactivity of an Unfamiliar Halogen

Suppose element X is a halogen located below bromine.

We can predict that X:

• has seven valence electrons
• tends to gain one electron
• forms X⁻ ions
• has more occupied electron shells than bromine
• experiences greater shielding
• attracts an incoming electron less strongly
• is less reactive than bromine

We could also predict that bromine should be able to displace X⁻ ions from a compound.

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Predicting Physical Properties

Periodic trends also allow us to predict physical properties.

Suppose an unfamiliar halogen is below iodine.

We would expect it to have:

• a larger atomic and molecular size
• stronger intermolecular attractions
• a higher melting point
• a higher boiling point

It would therefore be expected to be less volatile than iodine under comparable conditions.

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Halogens in Everyday Life

Halogens and halogen compounds have many important applications.

Examples include:

Fluoride compounds

Used in many toothpastes and dental treatments.

Chlorine compounds

Used in water treatment and many chemical products.

Bromine compounds

Used in a variety of industrial and chemical applications.

Iodine compounds

Important in nutrition and used in some antiseptic applications.

Many applications use halogen compounds, rather than the elemental halogens themselves.

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Halogens and Biological Importance

Some halide ions are important in living organisms.

For example:

chloride ions, Cl⁻

are important in body fluids.

Iodine is required in small amounts for the production of thyroid hormones.

Fluoride compounds can help strengthen tooth enamel.

This illustrates how reactive elements can have very different properties once they are present as ions or compounds.

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Safety and Halogens

Elemental halogens can be hazardous substances.

For example, chlorine and bromine require appropriate controlled laboratory handling.

Halogen experiments should therefore be performed using appropriate laboratory procedures and supervision.

Observations from demonstrations or controlled investigations can still provide valuable evidence for periodic trends.

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A Strategy for Halogen Displacement Questions

When given a displacement problem:

Step 1: Identify the free halogen.

Example:

Cl₂

Step 2: Identify the halide ion in the compound.

Example:

Br⁻

Step 3: Compare their positions in Group 17.

Step 4: Remember:

reactivity decreases down the group.

Step 5: Ask:

Is the free halogen more reactive?

If yes:

displacement occurs.

If no:

no reaction occurs.

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A Quick Example

Question:

Will chlorine displace iodide ions?

Reactivity order:

Cl > I

Therefore:

yes

Equation:

Cl₂ + 2I⁻ → 2Cl⁻ + I₂

Chlorine gains electrons more readily than iodine and therefore converts iodide ions into iodine.

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Common Mistakes

Mistake 1: Saying halogens have one valence electron

Halogens have:

seven valence electrons.

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Mistake 2: Saying halogens lose one electron

Halogens tend to:

gain one electron.

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Mistake 3: Saying halogens form +1 ions

Gaining one electron produces:

−1 ions.

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Mistake 4: Saying reactivity increases down Group 17

The correct trend is:

reactivity decreases down Group 17.

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Mistake 5: Forgetting that elemental halogens are diatomic

Write:

F₂, Cl₂, Br₂, I₂

when referring to elemental halogens.

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Mistake 6: Reversing the displacement rule

The correct rule is:

a more reactive halogen displaces a less reactive halogen from its halide compound.

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Mistake 7: Confusing halogen and halide

Chlorine, Cl₂ is a halogen.

Chloride, Cl⁻ is a halide ion.

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Did You Know?

The word halogen comes from Greek roots meaning approximately "salt-former."

This is appropriate because halogens readily react with metals to produce salts.

For example:

2Na + Cl₂ → 2NaCl

The product, sodium chloride, is a familiar example of a metal halide.

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Key Terms

Halogen: An element in Group 17 of the periodic table.

Halide: A −1 ion formed when a halogen atom gains an electron.

Valence electron: An electron in the outermost occupied shell.

Diatomic: Existing as molecules containing two atoms.

Reactivity: How readily a substance undergoes a chemical reaction.

Shielding: The reduction in effective nuclear attraction caused by inner electrons.

Atomic radius: A measure of atomic size.

Electronegativity: An atom's tendency to attract bonding electrons.

Displacement reaction: A reaction in which a more reactive element replaces a less reactive element in a compound.

Spectator ion: An ion that is present but does not undergo chemical change in the reaction.

Redox reaction: A reaction involving electron transfer.

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Key Equations and Patterns

Halide ion formation:

X + e⁻ → X⁻

Examples:

Cl + e⁻ → Cl⁻

Br + e⁻ → Br⁻

Metal + halogen:

2Na + Cl₂ → 2NaCl

Halogen displacement:

Cl₂ + 2KBr → 2KCl + Br₂

Ionic equation:

Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

Reactivity:

F > Cl > Br > I

Reactivity:

decreases down Group 17 ↓

Melting and boiling points:

generally increase down Group 17 ↓

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Key Takeaways

• The halogens are found in Group 17, traditionally called Group 7 in some school systems.
• Important halogens include fluorine, chlorine, bromine, and iodine.
• Halogens are non-metals.
• All halogens have seven valence electrons.
• Elemental halogens exist as diatomic molecules, such as F₂, Cl₂, Br₂, and I₂.
• Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid.
• Melting and boiling points generally increase down Group 17.
• Halogen atoms tend to gain one electron.
• Gaining one electron produces a −1 halide ion.
• Fluorine forms F⁻, chlorine forms Cl⁻, bromine forms Br⁻, and iodine forms I⁻.
• Halogens react with metals to form ionic compounds called metal halides.
• Halogen reactivity decreases down Group 17.
• Down the group, atomic radius and electron shielding increase.
• Greater distance and shielding reduce the effective attraction for an incoming electron.
• This makes electron gain less favourable and explains the decreasing reactivity.
• A more reactive halogen can displace a less reactive halogen from a halide compound.
• Chlorine can displace bromide and iodide ions.
• Bromine can displace iodide ions but cannot displace chloride ions.
• Iodine cannot displace chloride or bromide ions.
• Halogen displacement reactions involve electron transfer and are therefore redox reactions.
• Displacement experiments can be used to establish the relative reactivity of unfamiliar halogens.
• The strongest explanations connect position in Group 17 → electron arrangement → atomic radius and shielding → electron attraction → ion formation → reactivity → displacement behaviour.