Did you know that chlorine and iodine both exist as diatomic molecules, X₂, yet chlorine is a gas and iodine is a solid at ordinary room temperature? Their intramolecular bonds are both covalent, and neither Cl₂ nor I₂ has a permanent molecular dipole. The difference is the strength of the London dispersion forces between separate molecules. The iodine molecule contains more electrons in a larger, more easily distorted electron cloud. Its temporary dipoles can produce stronger induced-dipole attractions between neighbouring molecules, helping iodine form a solid at room temperature while chlorine remains a gas.
For JC learners searching for an H2 Chemistry tutor in Hougang, why iodine is solid but chlorine is gas, Group 17 physical states, halogen boiling point trend, instantaneous dipole induced dipole, London dispersion forces or A-Level Chemistry 9476 intermolecular forces, this is a useful example of molecular structure explaining everyday physical observations.
I would ask students to draw two circles labelled Cl₂ and I₂, then point to two different kinds of interactions. First is the covalent bond holding the two atoms inside each molecule together. Second is the attraction between neighbouring molecules. Melting or boiling ordinarily changes the second category, not the first. If a student can keep those apart, the physical state trend becomes much easier to explain.
We will compare chlorine, bromine and iodine, use verified melting and boiling values, explain electron-cloud polarizability, distinguish boiling from covalent bond dissociation and connect the physical trend with—but not confuse it with—Group 17 oxidising strength and hydrogen halide chemistry.
The revised Singapore H2 Chemistry 9476 syllabus covers molecular structure, London dispersion forces and Group 17 chemical trends. The physical state comparison is a useful application of those concepts. At eduKate, our small groups of up to three students emphasise clear explanation and an independently solved changed example rather than a memorised table alone.
Hougang Chemistry learning library · Official 2027 H2 Chemistry syllabus · RSC chlorine physical data · RSC iodine physical data
First, what kind of particles make up the halogen elements?
The familiar Group 17 elements include:
Fluorine, F₂.
Chlorine, Cl₂.
Bromine, Br₂.
Iodine, I₂.
Under their ordinary elemental molecular forms, each consists of two atoms of the same halogen bonded together.
The bond within a diatomic halogen molecule is a covalent bond formed through sharing electron density.
That means elemental chlorine is not a giant ionic lattice of Cl⁺ and Cl⁻ ions. Elemental iodine is not a metal-like network of freely moving electrons.
Both are molecular substances, and the forces connecting one molecule to another affect their physical states.
A student who labels solid iodine as “ionic because it forms a crystal” has confused crystal formation with ionic bonding. A molecular crystal can contain neutral molecules held together by intermolecular forces.
Why both Cl₂ and I₂ are nonpolar molecules
A Cl₂ molecule contains two chlorine atoms of the same element.
Those atoms have the same electronegativity, so the bond does not have a permanent dipole caused by unequal electronegativity of two different bonded elements.
The same is true of I₂: two identical iodine atoms share their covalent bond without a permanent difference in electronegativity between the ends.
Therefore, in the familiar school model, Cl₂ and I₂ are nonpolar molecules.
That does not mean there is no attraction between molecules.
It means they lack a permanent molecular dipole of the usual kind found in a polar molecule such as HCl.
Nonpolar molecules still contain electrons, and the distribution of those electrons can fluctuate over time.
Those fluctuations create the London dispersion interactions that matter so much to the Group 17 physical trend.
What are London dispersion forces?
London dispersion forces are attractive interactions associated with temporary fluctuations of electron density and the induced dipoles produced in nearby particles.
A nonpolar molecule’s electrons are not permanently fixed in a perfectly even arrangement at every instant.
At a given moment, the electron distribution may fluctuate so that one region becomes relatively more electron-rich and another region relatively more electron-poor.
That gives a temporary instantaneous dipole.
The electric field from that temporary dipole can influence a neighbouring molecule, distorting its electron density and creating an induced dipole.
The correlated fluctuations can lead to an overall attraction between the molecules.
This interaction exists even in substances whose molecules have no permanent dipole, such as Cl₂ and I₂.
It is therefore incorrect to say:
“Iodine is nonpolar, so iodine molecules have no intermolecular attraction.”
The correct description is that London dispersion forces provide substantial attraction between neighbouring I₂ molecules.
Why iodine’s electron cloud is more polarizable
An iodine atom has 53 electrons when neutral, while a chlorine atom has 17 electrons.
Therefore an I₂ molecule contains 106 electrons, whereas a Cl₂ molecule contains 34 electrons.
The iodine atoms also have more occupied electron shells, giving an electron cloud that extends farther from the nuclei on average.
A larger, more diffuse electron cloud is generally more easily distorted by the instantaneous electric fields produced by nearby particles.
This property is called polarizability.
When polarizability increases, the dispersion interactions between similar molecules can become stronger.
So as we move down from Cl₂ to Br₂ to I₂:
Larger molecules and more diffuse electron clouds → greater polarizability → stronger London dispersion interactions.
The trend is not determined by electron count alone in every possible molecular comparison, but it is a strong qualitative explanation for the related Group 17 molecular series.
Why London forces affect melting and boiling points
A liquid boils when its vapour pressure reaches the external pressure under the specified conditions.
To leave the condensed phase, molecules need to overcome or rearrange the attractive interactions that hold them near one another.
Stronger intermolecular attractions generally make the liquid less volatile at a given temperature, increasing the normal boiling point.
Similarly, melting a molecular solid into a liquid involves changing how closely molecules are packed and the intermolecular interactions of the condensed phases.
For the halogens, stronger London dispersion interactions down the group contribute to higher melting and boiling points.
That is the key trend:
Fluorine and chlorine are gases at ordinary room temperature.
Bromine is a liquid.
Iodine is a solid.
All four are composed of molecules; their different states do not require a change from covalent bonding to ionic or metallic bonding as we move down the group.
Verified melting and boiling values
The Royal Society of Chemistry periodic table gives the following representative values for elemental chlorine and iodine, with the other familiar values included for comparison:
| Element | Formula | Approximate melting point | Approximate boiling point | State at about 20 °C and ordinary pressure |
|---|---|---|---|---|
| Fluorine | F₂ | −220 °C | −188 °C | Gas |
| Chlorine | Cl₂ | −101.5 °C | −34.0 °C | Gas |
| Bromine | Br₂ | −7.2 °C | 58.8 °C | Liquid |
| Iodine | I₂ | 113.7 °C | 184.4 °C | Solid |
Sources: RSC chlorine, RSC iodine and the Chemistry LibreTexts halogen physical properties.
The values are rounded for teaching. Exact values and the observed phase depend on pressure and the conditions stated in a problem.
Why chlorine is gaseous at 20 °C
Chlorine’s normal boiling point is approximately −34.0 °C.
At 20 °C, the ambient temperature is well above that boiling point under ordinary atmospheric pressure.
Therefore chlorine exists as a gas in the familiar room-temperature comparison.
The chlorine molecules still have an intact Cl–Cl covalent bond in the gas. The phase change does not require producing separate chlorine atoms.
What differs from a liquid or solid is the distribution and separation of the molecules and the influence of intermolecular forces at the specified temperature and pressure.
A student who concludes “chlorine gas must be made of individual Cl atoms because Cl₂ bonds break on boiling” has confused a molecular phase change with chemical dissociation of a covalent bond.
The gas contains Cl₂ molecules under the ordinary model.
Why iodine is solid at 20 °C
Iodine’s normal melting point is approximately 113.7 °C.
At 20 °C, the temperature is far below this melting point under ordinary pressure.
Elemental iodine is therefore a molecular solid at room temperature.
Its I₂ molecules are arranged in a crystalline structure and held together by intermolecular attractions, including strong London dispersion interactions.
The I–I bonds within the molecules remain covalent.
The grey-black shiny appearance of solid iodine does not mean iodine has turned into a metal. Its crystal appearance and molecular bonding must be interpreted separately.
The Royal Society of Chemistry lists iodine as a solid at 20 °C, with a melting point around 113.7 °C and boiling point around 184.4 °C.
That is the direct experimental support for the physical-state classification.
Bromine provides the middle case
Bromine is particularly useful because its room-temperature state sits between chlorine and iodine.
Its melting point is approximately −7.2 °C.
Its boiling point is approximately 58.8 °C.
Room temperature, say 20 °C, falls between these values:
−7.2 < 20 < 58.8.
Therefore bromine is a liquid at ordinary room temperature and pressure.
The sequence Cl₂ gas, Br₂ liquid, I₂ solid shows the trend clearly as molecular size and dispersion attraction increase down the group.
It is not a coincidence that bromine’s phase lies between chlorine and iodine in the familiar temperature range.
But it is still the actual measured melting and boiling points, not a simple alphabetical ordering of the halogens, that determine the state at a specified temperature.
How to identify physical state from a data table
If an exam question gives melting and boiling points, the correct classification method is:
T < melting point: solid.
Melting point < T < boiling point: liquid.
T > boiling point: gas.
Use a consistent pressure and the specific phase-transition data supplied.
For example, at 25 °C:
- Chlorine is above its normal boiling point: gas.
- Bromine lies between its normal melting and boiling points: liquid.
- Iodine is below its normal melting point: solid.
This is stronger than memorising “halogens get heavier so they become solid”, because it works with changed temperatures and hypothetical compounds.
It also distinguishes the actual classification from the qualitative explanation of why those temperatures change down the group.
Why the covalent X–X bond trend does not give the answer
Within each halogen molecule, there is a covalent bond connecting the two atoms.
Across chlorine, bromine and iodine, the X–X bond length generally increases, and the covalent bond dissociation energy typically becomes smaller down the familiar comparison.
Representative bond dissociation energies are roughly:
Cl–Cl: about 242 kJ mol⁻¹.
Br–Br: about 193 kJ mol⁻¹.
I–I: about 151 kJ mol⁻¹.
Notice the apparent paradox:
I–I has a weaker intramolecular covalent bond than Cl–Cl, yet iodine has much higher melting and boiling points than chlorine.
This is not a contradiction.
The covalent bond dissociation energy describes separating the two atoms within one diatomic molecule into gaseous atoms.
Boiling describes separating or redistributing intact X₂ molecules from liquid into vapour.
The observed physical state trend is controlled mainly by intermolecular dispersion attraction, not by the energy needed to split the X–X bond.
This distinction is one of the most important transferable ideas in molecular bonding.
Compare intra- and intermolecular energy
| Process | What changes | Relevant interaction |
|---|---|---|
| Cl₂(g) → 2Cl(g) | One Cl–Cl covalent bond breaks | Intramolecular bond dissociation |
| I₂(g) → 2I(g) | One I–I covalent bond breaks | Intramolecular bond dissociation |
| I₂(s) → I₂(g) | Molecules separate from a solid into vapour | Intermolecular forces and sublimation thermodynamics |
| Br₂(l) → Br₂(g) | Molecules transfer from liquid to gas | Intermolecular interactions and vaporisation |
A student who identifies the reaction in the first column before choosing an energy term will avoid many mistakes.
Why iodine can sublime
Iodine is well known for producing a purple vapour under suitable conditions.
The process of going directly from solid to gas is called sublimation.
Iodine can have an appreciable vapour pressure even below its melting point, so some I₂ molecules can leave the solid and enter the gas phase without the entire solid first melting.
That does not mean iodine can never melt. The normal melting point near 113.7 °C is a real physical property.
Under the appropriate pressure and heating conditions, solid iodine can also melt and liquid iodine can boil.
The classroom phrase “iodine sublimes” should therefore be interpreted as a useful description of its observable vapour formation, not the false claim that no liquid iodine can exist.
The purple vapour consists of iodine-containing molecular species in the familiar elemental vapour description, not simply dissolved purple pigment escaping from a solid.
This is another opportunity to distinguish physical state from molecular identity.
Why iodine vapour is purple but the solid is dark
Iodine’s appearance depends on its physical environment and how light interacts with the species present.
Solid iodine is often described as dark grey-black or shiny black, while iodine vapour is purple or violet.
The molecules in the condensed solid have interactions with neighbouring molecules that can affect electronic energy levels and light absorption.
In vapour, the optical properties differ because the molecules have a different surrounding environment.
A change in appearance between solid and vapour does not automatically mean the covalent iodine molecule has turned into a completely different element.
Colour is a property of how light is absorbed, transmitted or reflected under the specified conditions.
For an exam about physical states, the essential point is that iodine remains a molecular substance rather than a metallic or ionic crystal.
The detailed optical explanation is enrichment beyond the simple London-force trend.
What happens when iodine dissolves in a solvent?
I₂ is nonpolar in its familiar molecular form. In a suitable nonpolar solvent, iodine can dissolve and produce a violet or purple appearance.
In aqueous solutions containing iodide, iodine can also participate in:
I₂ + I⁻ ⇌ I₃⁻.
That forms the triiodide ion and changes the solution speciation and colour.
So “iodine solution is always purple” is not a reliable universal statement.
The observed colour depends on solvent, concentration and which iodine-related species are present.
This is relevant to H2 redox titrations, where iodine-containing solutions can appear brown and react with thiosulfate.
The chemical identity of a dissolved iodine species and the physical state of pure elemental iodine are related topics, but they should not be conflated.
A careful learner identifies the solvent and the species in the question.
Why these are London forces rather than hydrogen bonds
Conventional strong hydrogen bonding in the H2 model involves hydrogen attached to a strongly electronegative atom such as O, N or F interacting with a suitable acceptor site.
Elemental Cl₂, Br₂ and I₂ contain no hydrogen atoms at all.
Therefore they cannot form the familiar intermolecular hydrogen-bond network seen in water or alcohols.
They are also nonpolar diatomic molecules, so ordinary permanent dipole–dipole attraction is not the dominant explanation.
Their condensed-phase attraction is especially well described by London dispersion interactions.
A student who says “iodine has a higher boiling point because it forms more hydrogen bonds” has chosen the wrong intermolecular-force class.
Always inspect the molecular formula before selecting a force.
Why molecular mass is a clue, not the whole explanation
Cl₂ has approximate Mr 70.9.
Br₂ has approximate Mr 159.8.
I₂ has approximate Mr 253.8.
As we go down the group, molecular mass increases.
It is tempting to say the higher boiling point is caused by greater molecular mass alone.
That is an incomplete explanation because molar mass is not itself an intermolecular force.
In this related molecular series, higher molecular mass accompanies a larger and more polarizable electron cloud, which increases London dispersion attractions.
The stronger attractions provide the physical reason for the higher phase-transition temperatures.
The distinction matters when comparing molecules from different functional-group families. Water, for example, can boil higher than heavier hydrogen sulfide because its hydrogen bonding changes the intermolecular-force picture.
Molar mass should prompt us to consider polarizability, but the explanation should identify actual interactions.
Why iodine solid does not conduct like a metal
Iodine’s grey-black appearance may resemble some metallic solids, but elemental iodine is a molecular crystal.
Its molecules are held together by intermolecular interactions. It does not have the same sea of freely mobile conduction electrons as a conventional metal.
Solid iodine is therefore not an ordinary good metallic electrical conductor under familiar conditions.
That observation is consistent with its molecular bonding structure.
By contrast, graphite can conduct electricity because of mobile electrons associated with its extended delocalised carbon network, even though graphite is not a metal.
A student who says “every dark shiny solid is metallic” is relying on appearance rather than bonding evidence.
To classify a substance, combine structure, conductivity, melting behaviour and other relevant data rather than one visual feature.
Why Group 17 oxidising strength is a different trend
Down Group 17, the halogens generally become less powerful oxidising agents in the familiar aqueous comparison from fluorine to iodine.
For example, chlorine can oxidise bromide ions under suitable conditions:
Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂.
This redox trend depends on electron transfer and the energetic balance of the halogen/halide systems, including bond and hydration contributions.
It is not caused directly by chlorine being a gas while bromine is a liquid.
Physical state and oxidising ability both change down the group, but they involve different properties:
- State and boiling point: mainly intermolecular attractions among X₂ molecules.
- Oxidising power: reduction of X₂ to X⁻ and the relevant electrode potentials.
- Hydrogen halide thermal stability: the strength of H–X bonds.
A student should not use “iodine is solid” as the explanation for every Group 17 observation.
Each property requires its own chemical model.
Why the H–X bond trend also differs from boiling X₂
Hydrogen halides include HF, HCl, HBr and HI.
Their thermal stability generally decreases down Group 17 because the H–X bond becomes longer and weaker.
That is an intramolecular covalent-bond comparison.
The boiling points of pure elemental halogens Cl₂, Br₂ and I₂ increase down the group because dispersion forces between molecules become stronger.
So the two trends run in opposite directions for different reasons.
Students sometimes feel that this is inconsistent, but it is precisely what we should expect when one question concerns breaking covalent bonds inside molecules and another concerns separating intact molecules.
The same Periodic Table group can support more than one trend without one trend explaining every other.
A strong H2 answer states the measured property before selecting the model.
Worked example 1: classify a halogen at 25 °C
Suppose an exam supplies a substance X₂ with:
Melting point = −7 °C.
Boiling point = 59 °C.
At 25 °C, we have:
−7 < 25 < 59.
Therefore the substance is liquid under the specified ordinary pressure conditions.
Those rounded values are characteristic of bromine.
The explanation is not that every molecular substance with Mr around 160 must be liquid. It is that this particular substance’s measured melting and boiling points straddle the stated temperature.
This question tests data reading and physical state classification.
A second part might then ask why its boiling point is higher than chlorine’s. That requires London dispersion and polarizability, not merely the state-rule inequality.
Worked example 2: calculate a boiling-point difference
Use the approximate normal boiling points:
Cl₂: −34.0 °C.
I₂: +184.4 °C.
The difference is:
184.4 − (−34.0) = 218.4 °C.
A temperature difference of 218.4 °C is also a difference of 218.4 K.
But this does not prove that the London dispersion interaction energy per mole in iodine is exactly 218.4 units greater than that in chlorine.
Degrees Celsius measure temperature, while intermolecular interaction energies use joules per mole or related units.
The boiling-temperature difference is an observed property explained by stronger overall condensed-phase attraction, not a direct numerical measure of one pairwise molecular force.
This is an important unit-analysis lesson.
Worked example 3: compare bond energy and phase
Suppose an examiner gives:
D(Cl–Cl) = 242 kJ mol⁻¹.
D(I–I) = 151 kJ mol⁻¹.
Which covalent bond is stronger?
Cl–Cl, because its dissociation requires more energy under the stated bond-enthalpy comparison.
Which element has a higher normal boiling point?
Iodine, because its molecules experience stronger London dispersion attractions in the condensed phase.
Both answers are correct.
They concern different quantities:
Bond dissociation energy describes forming separate gaseous atoms from a gaseous molecule.
Boiling temperature describes liquid–vapour equilibrium of intact molecules at a specified pressure.
A student who predicts chlorine’s boiling point must be higher solely because Cl–Cl is stronger has confused those definitions.
An unfamiliar data table exercise
Imagine three unidentified diatomic elements P₂, Q₂ and R₂ under the same pressure.
Their hypothetical normal boiling points are:
P₂: −60 °C.
Q₂: +10 °C.
R₂: +120 °C.
Suppose the question says that they are chemically similar nonpolar diatomic molecules and that polarizability increases from P₂ to R₂.
At 20 °C, P₂ and Q₂ would be gases because the temperature is above their normal boiling points. R₂ would be below its normal boiling point, but the boiling point alone is insufficient to decide whether R₂ is a liquid or solid; we also need its melting point.
That missing-data point is worth remembering.
An exam may ask for the trend in boiling points, which we can explain by increasing dispersion. But if it asks for the exact state below the boiling temperature, we may need additional melting-point data.
Do not infer “solid” solely because the temperature is below the boiling point.
Why external pressure matters
The boiling point of a liquid changes with external pressure because boiling occurs when the liquid’s vapour pressure reaches that external pressure.
The quoted normal boiling points correspond to a conventional reference pressure, approximately ordinary atmospheric pressure.
At another pressure, the boiling temperatures may differ.
A pressure change can also influence sublimation and the solid–liquid–vapour equilibrium of iodine.
Therefore when classifying a phase from a data table, use the pressure convention and temperatures given.
The familiar chlorine-gas, bromine-liquid, iodine-solid sequence is grounded at ordinary room temperature and ordinary pressure.
It should not be treated as a promise about all possible industrial storage conditions, where pressure and temperature can be controlled to produce different phases.
A seven-question independent diagnostic
- Explain why Cl₂ and I₂ are classified as nonpolar diatomic molecules.
- What causes a temporary dipole in a nonpolar molecule?
- Why is an I₂ molecule more polarizable than a Cl₂ molecule?
- Explain why iodine has stronger London dispersion forces than chlorine.
- Classify chlorine, bromine and iodine at about 20 °C using their familiar melting and boiling points.
- Why does iodine have a higher boiling point even though the I–I covalent bond is weaker than Cl–Cl?
- Does a purple iodine vapour mean the I–I bond has been completely broken to make isolated I atoms?
Answers
1. Each molecule contains identical atoms, so it has no permanent bond electronegativity difference producing a permanent molecular dipole. 2. Fluctuating electron density creates instantaneous dipoles and induces correlated dipoles nearby. 3. I₂ has a larger, more diffuse electron cloud with more electrons. 4. Greater polarizability strengthens induced-dipole interactions. 5. Chlorine gas, bromine liquid, iodine solid. 6. Boiling depends on forces between intact molecules, not the covalent bond dissociation energy inside each molecule. 7. No. Sublimation produces iodine-containing molecular vapour under the familiar conditions.
The sixth question is the most revealing because it checks whether the student distinguishes intramolecular bond strength from intermolecular-force strength.
Common mistakes and what to repair
| Student statement | First error | Repair question |
|---|---|---|
| Iodine is solid because it has ionic bonds | Misidentified molecular crystal bonding | What are the particles in elemental I₂? |
| Nonpolar halogens have no intermolecular forces | Forgot dispersion | Can electrons fluctuate in Cl₂? |
| Heavier halogen molecules always have stronger X–X bonds | Confused covalent bond length with polarizability | What is the trend in Cl–Cl, Br–Br and I–I bond dissociation? |
| Boiling breaks every Cl–Cl bond | Confused phase transition with dissociation | Is chlorine gas composed of Cl₂ or isolated atoms? |
| Iodine can only sublime and never melt | Ignored measurable iodine melting point | What phase diagram conditions are relevant? |
| Molar mass itself is a force between molecules | Skipped the polarizability mechanism | How does electron-cloud size influence dispersion? |
| A normal boiling point alone tells whether a below-boiling substance is solid | Forgot melting-point threshold | Could the substance be liquid? |
| Iodine is a weaker oxidant because it is solid | Confused redox thermodynamics with physical state | What electrode potentials describe X₂/X⁻? |
A seven-session JC1–JC2 revision route
- Session 1: recognise the halogen elements as diatomic molecular substances and distinguish covalent X–X bonds from attractions between molecules.
- Session 2: draw a temporary instantaneous dipole and explain induction in a neighbouring nonpolar molecule.
- Session 3: compare electron-cloud sizes and polarizability for Cl₂, Br₂ and I₂.
- Session 4: use the London-force trend to explain increasing melting and boiling points.
- Session 5: classify room-temperature states from real or changed melting/boiling data.
- Session 6: distinguish the X–X covalent-bond-energy trend from the X₂ boiling-point trend.
- Session 7: answer a new Group 17 question in which one property concerns physical state, another concerns oxidising power and a third concerns hydrogen halide thermal stability.
The last task is valuable because it makes students choose the correct physical or chemical model for each observation instead of writing one generic explanation for every periodic trend.
What parents seeking H2 Chemistry tuition around Hougang can ask
Bring a marked Group 17 or intermolecular-forces question. Ask the tutor whether the student is confusing molar mass with London dispersion, covalent bond energies with boiling, physical states with oxidising strength or the meaning of sublimation.
An effective three-student tutorial should ask each learner to predict the state and explain the interaction trend for an unfamiliar molecule before comparing answers. Private lessons should end with the same independent transfer check.
Confirm current H2 Chemistry 9476 experience, actual lesson venue, timetable, small-group availability and travel arrangements. A Hougang-targeted educational guide is not evidence of a physical tuition classroom operating inside Hougang.
Why this matters beyond the A Levels
London dispersion interactions influence the physical behaviour of many gases, liquids, molecular solids and industrial materials.
Temperature and pressure also determine which phase a substance occupies, while molecular polarizability affects how particles attract one another.
The broader scientific habit is to identify the particles, distinguish forces between them from bonds within them, and use actual data to classify physical state.
That is preparation for more advanced Chemistry, materials science and chemical engineering without any promise about grades or university admission.
Frequently asked questions
Why is iodine solid but chlorine gas at room temperature?
I₂ molecules are larger and more polarizable, so they experience stronger London dispersion attractions, giving much higher melting and boiling points than Cl₂.
Why is bromine liquid?
At ordinary room temperature, its temperature lies between its melting and boiling points, roughly −7 °C and +59 °C.
Do halogen molecules form hydrogen bonds?
Not the familiar strong hydrogen-bond network, because pure Cl₂, Br₂ and I₂ contain no hydrogen atoms.
Are chlorine and iodine molecules polar?
They are nonpolar homonuclear diatomic molecules, but they still exhibit London dispersion interactions.
Does boiling chlorine break its Cl–Cl covalent bonds?
No. Normal vaporisation involves intact Cl₂ molecules moving into the gaseous phase.
Is the I–I covalent bond stronger than Cl–Cl?
No. The I–I bond is weaker in the usual bond-dissociation comparison. Iodine’s higher boiling point arises from intermolecular attractions, not stronger I–I covalent bonds.
Does iodine sublime?
Yes, iodine has a noticeable vapour pressure and can sublime under suitable conditions. It also has a measurable melting point near 113.7 °C and can exist as a liquid.
Continue learning and sources
Why water boils far higher than hydrogen sulfide · Why H–X bond strength affects hydrogen halide stability · Why fluorine is a stronger oxidant than chlorine · JC1 intermolecular forces and boiling points · Royal Society of Chemistry: chlorine properties · Royal Society of Chemistry: iodine properties · Chemistry LibreTexts: Group 17 halogen physical properties · Hougang Chemistry Tutor Library · eduKate Singapore JC Chemistry enquiries.
Safety and data note: Cl₂, Br₂ and I₂ can be hazardous, and chlorine gas in particular is highly toxic. This is a theoretical molecular-bonding and physical-properties guide, not a suggestion to handle or heat halogens. Transition temperatures are representative values for ordinary pressure and may change under different conditions.