The Halogens: Occurrence and Physical Properties – Trends in Chemical Properties of Group 17 Elements and Their Compounds (Senior 4 Chemistry) | Ganzaa

Trends in Chemical Properties of Group 17 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 17 Elements and Their Compounds

5 lessons

Lesson 1 of 5

The Halogens: Occurrence and Physical Properties

25 min

A family you can see

Most groups of the periodic table are hard to tell apart by eye, but the halogens announce themselves: fluorine is a pale yellow gas, chlorine a yellow-green gas, bromine a dark red-brown liquid that gives off orange fumes, and iodine a shiny grey-black solid that turns into a violet vapour when warmed. You meet their compounds every day: chloride in table salt, fluoride in toothpaste, iodine added to salt to protect the thyroid gland, chlorine in the water that comes out of a tap. "Halogen" means "salt-former", and all four combine with metals to make salts. This lesson describes where the halogens come from and how their physical properties change down the group.

Definition

Halogens (Group 17): fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At) and the artificial element tennessine (Ts). Each atom has seven outer electrons, $ns^2\,np^5$ , one short of a noble gas. The elements exist as diatomic molecules, $\mathrm{X_2}$ , and gain an electron to form halide ions, $\mathrm{X^-}$ . Example: $\mathrm{Cl}$ is $[\mathrm{Ne}]\,3s^2\,3p^5$ ; $\mathrm{Cl^-}$ is $[\mathrm{Ar}]$ .

First electron affinity: the energy change when one mole of gaseous atoms each gain one electron to form one mole of gaseous 1− ions. Example: $\mathrm{Cl(g) + e^- \rightarrow Cl^-(g)}$ , $\Delta H = -349\ \mathrm{kJ\,mol^{-1}}$ .

Where they occur

The halogens are too reactive to occur free; they are found as halide ions, or in the case of iodine also as iodate.

Element	Main sources
Fluorine	fluorite (fluorspar), $\mathrm{CaF_2}$ ; fluorapatite; cryolite, $\mathrm{Na_3AlF_6}$
Chlorine	sodium chloride in sea water, salt lakes and underground rock salt
Bromine	bromide ions in sea water and salt-lake brines
Iodine	brines; seaweed; sodium iodate in Chilean nitrate deposits
Astatine	only tiny, short-lived traces from radioactive decay

Physical properties down the group

Table 12.1: The halogens

Property	F	Cl	Br	I
Atomic number	9	17	35	53
Outer configuration	$2s^2\,2p^5$	$3s^2\,3p^5$	$3d^{10}\,4s^2\,4p^5$	$4d^{10}\,5s^2\,5p^5$
Appearance at 25 °C	pale yellow gas	yellow-green gas	red-brown liquid, orange vapour	grey-black solid, violet vapour
Covalent radius / pm	64	99	114	133
First ionisation energy / kJ mol⁻¹	1681.0	1251.2	1139.9	1008.4
Electronegativity (Pauling)	3.98	3.16	2.96	2.66
First electron affinity / kJ mol⁻¹	−328.2	−348.6	−324.5	−295.2
X–X bond energy / kJ mol⁻¹	159	243	193	151
Melting point / °C	−219.6	−101.5	−7.2	113.7
Boiling point / °C	−188.1	−34.0	58.8	184.4

Melting and boiling points rise down the group (Figure 12.1). All four exist as $\mathrm{X_2}$ molecules held by London forces; bigger molecules with more electrons have stronger London forces, so the state changes from gas to liquid to solid.
Colour deepens down the group, from pale yellow to violet-black, because the energy gap for absorbing visible light gets smaller.
Electronegativity falls as the atoms get bigger: the nucleus attracts bonding electrons less strongly through more shielding shells. Fluorine is the most electronegative element of all.
Electron affinity becomes less exothermic from chlorine to iodine, because an incoming electron is further from the nucleus. Fluorine is the exception: its electron affinity is smaller than chlorine's, because its tiny 2p shell is crowded, and an extra electron is repelled by the seven already there.

Figure 12.1: Melting and boiling points of the halogens.

The weak F–F bond

The X–X bond energy falls from chlorine to iodine as the atoms get larger and the shared electrons are further from both nuclei (Figure 12.2). On that trend, the F–F bond should be the strongest of all. Instead it is weaker than Cl–Cl. The F–F bond is so short that the lone pairs on the two fluorine atoms are very close together and repel each other strongly, weakening the bond. This weak bond is one reason fluorine is so violently reactive: very little energy is needed to split $\mathrm{F_2}$ into atoms.

Figure 12.2: X–X bond energies of the halogens.

Worked example 1: light that splits chlorine

A mixture of hydrogen and chlorine explodes in bright sunlight, because light splits $\mathrm{Cl_2}$ into atoms. Using the Cl–Cl bond energy of 243 kJ mol⁻¹, calculate the longest wavelength of light that can break one Cl–Cl bond. $h = 6.626 \times 10^{-34}\ \mathrm{J\,s}$ , $c = 3.00 \times 10^{8}\ \mathrm{m\,s^{-1}}$ , $N_A = 6.022 \times 10^{23}\ \mathrm{mol^{-1}}$ .

Solution. Energy per bond: $E = \frac{243 \times 10^{3}\ \mathrm{J\,mol^{-1}}}{6.022 \times 10^{23}\ \mathrm{mol^{-1}}} = 4.04 \times 10^{-19}\ \mathrm{J}$ $\lambda = \frac{hc}{E} = \frac{6.626 \times 10^{-34} \times 3.00 \times 10^{8}}{4.04 \times 10^{-19}} = 4.92 \times 10^{-7}\ \mathrm{m} = 492\ \mathrm{nm}$ Any light with a wavelength shorter than about 492 nm (blue, violet and ultraviolet) can split chlorine molecules, so sunlight starts the chain reaction.

Worked example 2: predicting astatine

Use the trends in Table 12.1 to predict, with reasons, the state at room temperature, the colour, and the electronegativity of astatine compared with iodine.

Solution.

State: a solid, with a melting point above iodine's 113.7 °C, because $\mathrm{At_2}$ would be larger, with more electrons and stronger London forces.
Colour: darker than iodine, probably black, continuing the deepening of colour down the group.
Electronegativity: lower than iodine's 2.66, because its outer electrons are further from the nucleus and better shielded. (Astatine is so radioactive that few of these properties have been measured.)

Warning

Common mistakes. (1) Explaining the rise in boiling point by "stronger covalent bonds": boiling separates $\mathrm{X_2}$ molecules; the X–X bonds do not break, and they actually get weaker. (2) Saying fluorine has the highest electron affinity: chlorine's is larger. (3) Writing halogens as single atoms ("Cl") in equations for the element: they are diatomic, $\mathrm{Cl_2}$ . (4) Confusing halogen (the element) with halide (the ion).

Safety. Fluorine is too dangerous for school laboratories. Chlorine and bromine are toxic and corrosive: chlorine attacks the lungs, and liquid bromine causes severe burns and its vapour is poisonous; both are used only in a fume cupboard, with gloves and goggles, and in small amounts. Iodine stains skin and its vapour irritates the eyes. If someone breathes in chlorine, move them to fresh air and get medical help.

Tip

Halogen trends in one line: down the group, bigger molecules (higher m.p./b.p., darker colour), less pull on electrons (lower IE, EN, electron affinity). Fluorine breaks two rules: a weak F–F bond and a smaller electron affinity than chlorine, both caused by its tiny, crowded atom.

Professor Dave Explains: Periodic Table Part 8, the Halogens (7 min). Watch for the colours and states of the halogens, why they are so reactive, and where their compounds are used.

✦

Practice

Write the electron configurations of a bromine atom and of a bromide ion.
Describe the appearance of each halogen at room temperature.
Explain why the boiling points of the halogens increase down the group.
Explain why the electronegativity of the halogens decreases down the group.
Explain why the F–F bond is weaker than the Cl–Cl bond.
Why is the electron affinity of fluorine less exothermic than that of chlorine?
Calculate the energy needed to break all the bonds in 1.00 g of chlorine molecules. $M(\mathrm{Cl_2}) = 70.90$ ; bond energy 243 kJ mol⁻¹.
Using the I–I bond energy (151 kJ mol⁻¹), calculate the longest wavelength of light that can split an iodine molecule, and say whether visible light can do it.

Show answersHide answers

Br: $[\mathrm{Ar}]\,3d^{10}\,4s^2\,4p^5$ ; $\mathrm{Br^-}$ : $[\mathrm{Ar}]\,3d^{10}\,4s^2\,4p^6$ (= [Kr]).
F₂ pale yellow gas; Cl₂ yellow-green gas; Br₂ red-brown liquid (orange vapour); I₂ shiny grey-black solid (violet vapour).
The molecules contain more electrons, so London forces between molecules become stronger and more energy is needed to separate them.
The atoms get bigger with more inner shells; the bonding pair is further from the nucleus and better shielded, so it is attracted less.
The F–F bond is very short, so the lone pairs on the two atoms are close together and repel strongly, weakening the bond.
Fluorine's 2p shell is small and crowded; the incoming electron is repelled by the electrons already there, so less energy is released than for the larger chlorine atom.
$n = \frac{1.00}{70.90} = 0.01410$ mol; energy $= 0.01410 \times 243 = 3.43\ \mathrm{kJ}$ .
$E = \frac{151 \times 10^3}{6.022 \times 10^{23}} = 2.51 \times 10^{-19}\ \mathrm{J}$ ; $\lambda = \frac{6.626 \times 10^{-34} \times 3.00 \times 10^8}{2.51 \times 10^{-19}} = 7.92 \times 10^{-7}\ \mathrm{m} = 792\ \mathrm{nm}$ . All visible light (400–700 nm) has enough energy.

Quick check

Why do the boiling points of the halogens increase from fluorine to iodine?

Key points

Halogens are $ns^2\,np^5$ , exist as $\mathrm{X_2}$ molecules and form $\mathrm{X^-}$ ions; they occur as halides (fluorite, sea salt, brines) and, for iodine, iodates.
Down the group, melting and boiling points rise (stronger London forces) and colour deepens: F₂ pale yellow gas, Cl₂ yellow-green gas, Br₂ red-brown liquid, I₂ grey-black solid.
Ionisation energy, electronegativity and electron affinity decrease down the group (F is the most electronegative element).
Fluorine is anomalous: its F–F bond (159 kJ mol⁻¹) is weaker than Cl–Cl (243), and its electron affinity is smaller than chlorine's, because the atom is so small.

Sources and credits

Image: Chlorine-sample.jpg by Benjah-bmm27, via Wikimedia Commons, public domain.
Image: Bromine-ampoule.jpg by Jurii, via Wikimedia Commons, licensed under CC BY 3.0.
Image: Sample of iodine.jpg by LHcheM, via Wikimedia Commons, licensed under CC BY-SA 3.0.
Data: ionisation energies from Wikipedia's ionisation energies data page; electron affinities from the electron affinity data page; covalent radii from the atomic radii data page; bond energies (from Greenwood and Earnshaw) from Halogen; all CC BY-SA 4.0, adapted by Ganzaa. Melting and boiling points and electronegativities from the PubChem Periodic Table.
Video: Periodic Table Part 8: Halogens (F, Cl, Br, I, At, Tn), Professor Dave Explains, embedded from YouTube.
The text, charts, tables, worked examples and Practice items were written and drawn by Ganzaa.
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1A family you can see

2Where they occur

3Physical properties down the group

4The weak F–F bond

A family you can see

Where they occur

Physical properties down the group

The weak F–F bond