Chemistry: Atoms First 2e · Representative Metals, Metalloids, and Nonmetals

Occurrence, Preparation, and Properties of Halogens

9 min read

Reference-values note: atomic masses are commonly taught reference values based on standard atomic weights; use the periodic table provided in your course for graded work. Physical states at room temperature and the displacement order F₂ > Cl₂ > Br₂ > I₂ are standard textbook facts. Safety note: elemental halogens (especially F₂ and Cl₂) are toxic and corrosive; bromine causes severe burns; never mix bleach with acids (releases Cl₂). General ventilation and containment principles only — no lab procedures provided here.

On this page 9 sections

In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

The halogens — fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) — are Group 17, the elements one electron short of a filled shell. Their name means "salt former," and they live up to it: they are never found free in nature, always paired with metals as salts such as sodium chloride in seawater, fluorite (CaF₂), and the iodides in marine organisms. Because they need only one electron to complete an octet, they are the most reactive family of nonmetals, and reactivity decreases smoothly down the group: fluorine is so aggressive it attacks glass and reacts with almost everything, while iodine is a gentle gray-black solid that sublimes to a violet vapor. This topic covers where halogens occur, how each is prepared (electrolysis for fluorine and chlorine, simple displacement for bromine and iodine), and their signature properties: diatomic molecules, decreasing oxidizing power down the group, the −1 oxidation state in halides, and the positive states they show with oxygen and fluorine.

Why this matters

chemistry is all around you: chlorine disinfects drinking water and pools; fluoride hardens tooth enamel and is in most toothpastes; bromine compounds make flame retardants; iodine is essential for thyroid hormone synthesis, which is why table salt is iodized; and silver halides are the light-sensitive chemistry of traditional photography. The halogen displacement series — F₂ > Cl₂ > Br₂ > I₂ in oxidizing power — is a classic demonstration of periodic trends and redox. Chlorine and fluorine are also safety-critical: chlorine gas was a weapon in World War I, and concentrated fluorine is among the most dangerous substances in chemistry. On exams, the halogens test periodic trends, oxidation states, and displacement and electrolysis stoichiometry.

The college version

Core Concepts

Occurrence: always combined as halides

Halogens are too reactive to exist free, so they occur as halide ions (X^-) in salts:

Chlorine is the most abundant: sodium chloride in seawater (about 3.5% dissolved salts), rock salt (halite), and brines; also KCl (sylvite) and MgCl₂.
Fluorine occurs as fluorite (CaF₂), cryolite (Na₃AlF₆), and fluoroapatite in phosphate rock.
Bromine is found as bromide in seawater and brines, concentrated by marine organisms (seaweed, sponges).
Iodine appears as iodide in seawater, brines, and nitrate deposits, and is concentrated in seaweed and thyroid glands.

Preparation: electrolysis for the top two, displacement for the rest

Fluorine is too powerful an oxidizing agent for any chemical to displace, so it is made by electrolysis of molten potassium hydrogen difluoride, KHF₂:

$2 KHF 2 (l) electrolysis ⟶ 2 KF(l) + H 2 (g) + F 2 (g)$

Chlorine is made industrially by the — electrolysis of brine:

$2 NaCl(aq) + 2 H 2 O(l) electrolysis ⟶ 2 NaOH(aq) + H 2 (g) + Cl 2 (g)$

Bromine and iodine are made by oxidizing their halide ions with a stronger halogen — a . Because chlorine is a stronger oxidizer than bromine, it displaces bromide from brine:

$Cl 2 (g) + 2 Br - (aq) \to 2 Cl - (aq) + Br 2 (l)$

and chlorine (or bromine) displaces iodide:

$Cl 2 (g) + 2 I - (aq) \to 2 Cl - (aq) + I 2 (s)$

This ordering — F₂ > Cl₂ > Br₂ > I₂ — is the halogen displacement series, a direct consequence of decreasing oxidizing power down the group.

Properties: a family with a visible trend

All halogens are diatomic, X₂, with physical states that march down the group: fluorine is a pale-yellow gas, chlorine a greenish-yellow gas, bromine a red-brown liquid (one of only two elements liquid at room temperature, the other being mercury), and iodine a shiny gray-black solid that sublimes to a violet vapor. Electronegativity and reactivity decrease down the group (fluorine is the most electronegative element of all), while atomic radius and boiling point increase. Halogens show a −1 oxidation state in halides (e.g., NaCl, KBr); with more electronegative partners (oxygen and fluorine), they show positive states: +1 (, ClO^-), +3 (chlorite), +5 (chlorate, ClO₃^-), and +7 (perchlorate, ClO₄^-).

Hydrogen halides: a famous exception

The hydrogen halides HF, HCl, HBr, and HI are colorless, fuming gases that dissolve in water to form acids. The trend is a classic exam trap: HCl, HBr, and HI are strong acids, but HF is weak — despite fluorine being the most electronegative element — because the very strong H–F bond resists breaking. is also the only common acid that attacks glass:

$SiO 2 (s) + 4 HF(aq) \to SiF 4 (g) + 2 H 2 O(l)$

Uses and safety

Chlorine compounds disinfect water and bleach (sodium hypochlorite, NaOCl); fluorine chemistry gives toothpaste fluoride, Teflon, and refrigerants; bromine compounds are flame retardants and were once photographic emulsions; iodine tincture is an antiseptic and KI is added to table salt. All elemental halogens are toxic and corrosive — chlorine gas damages the lungs, fluorine reacts violently with moisture and organics, and bromine causes severe burns — so general principles apply: work in ventilated areas with proper containment, and never mix bleach with acids (which releases chlorine gas).

Common Confusions

Do Not Confuse	With	Difference
Halogen (X₂)	Halide (X^-)	Elemental, reactive molecule vs its ion in salts — Cl₂ is toxic, Cl⁻ is essential
Cl₂ (chlorine gas)	Cl^- (chloride ion)	Oxidized, toxic molecule vs reduced, harmless ion; pool "chlorine" is actually hypochlorite
HF (weak acid)	HCl (strong acid)	H–F bond is very strong, so HF ionizes only slightly; HCl, HBr, HI are strong acids
Reactivity order F₂ > Cl₂ > Br₂ > I₂	Size order	Fluorine is the most reactive but the smallest; iodine is largest and least reactive
Sublimation	Evaporation	Sublimation is solid → gas directly (iodine); evaporation is liquid → gas
"Fluoride in toothpaste"	Elemental fluorine	Toothpaste contains F^- ions; F₂ gas is violently reactive and corrosive

Eli explains

The same idea, in plain words

Explain it like I’m 10

The halogens are a family of five elements that are all one step away from being "complete," so they grab electrons from other atoms — that is why they always travel with partners, like salt in the ocean. Fluorine is the bossy big sister who reacts with everything, while iodine is the calm one that turns straight from solid into purple gas. They are the family behind swimming-pool chlorine, toothpaste fluoride, and the iodine your body needs to make thyroid hormones.

Worked example

Example 1: Displacement stoichiometry — making bromine

How many grams of Br₂ are produced when 71.0 g of Cl₂ reacts with excess bromide ions? Balanced equation: Cl₂(g) + 2 Br^-(aq) → 2 Cl^-(aq) + Br₂(l). (Molar masses: Cl₂ = 70.90 g/mol, Br₂ = 159.80 g/mol.)

Step 1 — Grams to moles:

$n(Cl 2) = 71.0 g \times 1 mol Cl 2 70.90 g Cl 2 = 1.00 mol Cl 2$

Step 2 — Mole ratio (1 mol Br₂ per 1 mol Cl₂):

$1.00 mol Cl 2 \times 1 mol Br 2 1 mol Cl 2 = 1.00 mol Br 2$

Step 3 — Moles to grams:

$1.00 mol Br 2 \times 159.80 g Br 2 1 mol Br 2 = 159.8 g Br 2$

Check the sense: bromine atoms are more than twice as heavy as chlorine atoms, so 1 mol of Br₂ (160 g) outweighs 1 mol of Cl₂ (71 g) — consistent.

Example 2: Chlor-alkali electrolysis

What mass of chlorine gas is produced by electrolyzing brine made from 117.0 g of NaCl? Balanced equation: 2 NaCl(aq) + 2 H₂O(l) → 2 NaOH(aq) + H₂(g) + Cl₂(g). (Molar masses: NaCl = 58.44 g/mol, Cl₂ = 70.90 g/mol.)

Step 1 — Grams to moles:

$n(NaCl) = 117.0 g \times 1 mol NaCl 58.44 g NaCl = 2.00 mol NaCl$

Step 2 — Mole ratio (1 mol Cl₂ per 2 mol NaCl):

$2.00 mol NaCl \times 1 mol Cl 2 2 mol NaCl = 1.00 mol Cl 2$

Step 3 — Moles to grams:

$1.00 mol Cl 2 \times 70.90 g Cl 2 1 mol Cl 2 = 70.9 g Cl 2$

Unit check: g NaCl → mol NaCl → mol Cl₂ → g Cl₂. Answer: 70.9 g Cl₂.

Example 3: Percent chlorine in table salt

What is the mass percent of chlorine in NaCl? (Molar masses: Na = 22.99, Cl = 35.45 g/mol.)

Step 1 — Formula for molar mass:

$M(NaCl) = 22.99 + 35.45 = 58.44 g/mol$

Step 2 — Percent chlorine:

$% Cl = 35.45 g Cl 58.44 g NaCl \times 100 % = 60.7 %$

More than half the mass of table salt is chlorine — but as the harmless chloride ion, not the toxic Cl₂ molecule.

Key takeaways

Group 17 = halogens; all are diatomic (F₂, Cl₂, Br₂, I₂) and never found free — always as halide salts.
Physical states down the group: F₂ pale-yellow gas → Cl₂ greenish-yellow gas → Br₂ red-brown liquid → I₂ gray-black solid (sublimes violet).
Reactivity and oxidizing power decrease down the group: F₂ > Cl₂ > Br₂ > I₂ — a stronger halogen displaces a weaker one from its halide.
Preparation: F₂ by electrolysis of molten KHF₂; Cl₂ by chlor-alkali electrolysis of brine (2 NaCl + 2 H₂O → 2 NaOH + H₂ + Cl₂); Br₂ and I₂ by displacement with Cl₂.
Oxidation states: −1 in halides; +1, +3, +5, +7 with oxygen/fluorine (ClO⁻, ClO₃⁻, ClO₄⁻).
HF is a weak acid (strong H–F bond) even though HCl, HBr, and HI are strong; HF etches glass.
Fluoride hardens enamel; iodide is required for thyroid hormones (iodized salt); chlorine disinfects water.
Safety principles: halogens and their vapors are toxic/corrosive; never mix bleach with acid.

Check yourself

5 review questions from the chapter. Try each one, then open the answer.

Why are halogens never found free in nature, and in what form do they occur?
Show answer
They are one electron short of a filled shell, so they react readily with metals and almost everything else; in nature they occur as halide ions (F^-, Cl^-, Br^-, I^-) in salts such as NaCl, CaF₂, and seawater bromides.
Rank F₂, Cl₂, Br₂, and I₂ by oxidizing strength, and use that to explain how bromine is prepared.
Show answer
F₂ > Cl₂ > Br₂ > I₂. A stronger halogen displaces a weaker one, so Cl₂ oxidizes Br^- to Br₂ (Cl₂ + 2 Br^- → 2 Cl^- + Br₂) — the industrial route to bromine.
Which hydrogen halide is the weak acid, and why?
Show answer
HF. The H–F bond is very strong (the most polar and strongest hydrogen–halogen bond), so HF barely ionizes in water; HCl, HBr, and HI ionize almost completely.
How many grams of I₂ form when 71.0 g of Cl₂ reacts with excess I^- (I₂ = 253.80 g/mol)?
Show answer
Moles Cl₂ = 71.0 ÷ 70.90 = 1.00 mol; 1:1 ratio with I₂; mass = 1.00 × 253.80 = 253.8 g I₂.
What are the physical states of the four common halogens at room temperature?
Show answer
F₂: pale-yellow gas; Cl₂: greenish-yellow gas; Br₂: red-brown liquid; I₂: gray-black solid that sublimes.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Halogen: Group 17 element, one electron short of a filled shell
Halide: X^- ion or salt of a halogen (e.g., Cl^-, NaCl)
Displacement reaction: A more reactive element replaces a less reactive one in a compound
Chlor-alkali process: Electrolysis of brine producing NaOH, H₂, and Cl₂
Hypochlorite: ClO^-, chlorine in +1 with oxygen
Sublimation: Direct solid → gas transition
Hydrofluoric acid: HF(aq), the only weak hydrogen halide acid

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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