Group 17 trends and chlorine

From the pool to the toothpaste

The smell at the side of a swimming pool comes from chlorine compounds that kill germs. A tube of toothpaste has a fluoride in it that makes teeth harder. Both come from group 17, the halogens. The word means salt former. A halogen joins with a metal and makes a salt: sodium chloride, calcium fluoride, potassium iodide.

The group has fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). Astatine is radioactive and very rare, so we do not study it. Each atom has the outer arrangement $ns^2np^5$. It needs just one more electron to reach the eight of a noble gas. That is why a halogen takes an electron so easily.

The trends in the group

Let us read the numbers of the four common halogens, one by one.

• F: a pale yellow gas. Covalent radius 64 pm. First ionisation enthalpy 1680 kJ/mol. Electronegativity 4.0.
• Cl: a greenish yellow gas. Radius 99 pm. Ionisation enthalpy 1256 kJ/mol. Electronegativity 3.2.
• Br: a red-brown liquid. Radius 114 pm. Ionisation enthalpy 1142 kJ/mol. Electronegativity 3.0.
• I: a violet solid. Radius 133 pm. Ionisation enthalpy 1008 kJ/mol. Electronegativity 2.7.

The radius grows and the ionisation enthalpy and the electronegativity fall, as in every group. The state changes from gas to liquid to solid, because the dispersion forces between the larger molecules are stronger. Some books print 3.0, 2.8 and 2.5 for chlorine, bromine and iodine (older scales). The values above follow the Pauling scale as most texts now print it.

Electron gain enthalpy is the energy change when an atom of a gas takes one electron. For a halogen it is negative, which means that energy is released. The values in kJ/mol are $-328$ for F, $-349$ for Cl, $-325$ for Br and $-295$ for I. NCERT prints $-333$ for F and $-296$ for I, and the accepted data are a few kJ/mol different. The surprise is that chlorine is more negative than fluorine. The fluorine atom is so small that the electron it takes feels strong repulsion from the seven electrons in the small $2p$ shell. In the larger chlorine atom, the repulsion is less.

The bond dissociation enthalpy of $\text{X}_2$ follows the same pattern. It is 159 kJ/mol for F2, 243 kJ/mol for Cl2, 193 kJ/mol for Br2 and 151 kJ/mol for I2. The F-F bond is short, so the lone pairs of the two atoms are very close and repel each other. That weakens the bond.

Oxidising power and displacement

Every halogen is an oxidising agent: it takes electrons and becomes a halide ion. The standard electrode potential $E^\circ$ measures this. A bigger value means a stronger oxidising agent. The values for $\text{X}_2 + 2e^- \rightarrow 2\text{X}^-$ are 2.87 V for F2, 1.36 V for Cl2, 1.09 V for Br2 and 0.54 V for I2. The power falls down the group.

Why is fluorine the strongest, when its electron gain enthalpy is less negative than that of chlorine? Three things add up. The F-F bond is weak, so it is easy to break. The electron gain is large. Most important, the small fluoride ion is held very strongly by water. Its hydration enthalpy is $-515$ kJ/mol, compared with $-381$ kJ/mol for chloride. This large release of energy gives the fluorine its top place.

A halogen with a higher potential takes electrons from the halide ion of a halogen with a lower potential. This is a displacement reaction:

$\text{Cl}_{2}\ +\ 2\text{KBr}\ \rightarrow \ 2\text{KCl}\ +\ \text{Br}_{2}$

$\text{Cl}_{2}\ +\ 2\text{KI}\ \rightarrow \ 2\text{KCl}\ +\ \text{I}_{2}$

The bromine appears as an orange-brown colour in the solution. The iodine turns the solution brown, and it dissolves in a little hexane to give a violet layer. Fluorine cannot be used in this test in water, because it attacks water itself.

Choose chlorine and bromide. The cell potential is $1.36 - 1.09 = +0.27$ V, so the reaction goes. Now swap them: bromine with chloride gives $-0.27$ V, so nothing happens. The sign of the cell potential decides.

Why fluorine is different

Fluorine is the odd one in the family.

• It shows only the oxidation state $-1$ in its compounds, apart from 0 in $\text{F}_2$. It has no d orbitals. The others also show $+1$, $+3$, $+5$ and $+7$ in compounds with oxygen and with fluorine.
• Hydrogen fluoride forms hydrogen bonds, so HF is a liquid that boils near room temperature (293 K), while HCl is a gas.
• HF is a weak acid in water. The others are strong acids.
• Fluorine reacts with water: $2\text{F}_{2}\ +\ 2\text{H}_{2}\text{O}\ \rightarrow \ 4\text{HF}\ +\ \text{O}_{2}$. Some ozone also forms: $3\text{F}_{2}\ +\ 3\text{H}_{2}\text{O}\ \rightarrow \ 6\text{HF}\ +\ \text{O}_{3}$.
• HF attacks glass: $\text{SiO}_{2}\ +\ 4\text{HF}\ \rightarrow \ \text{SiF}_{4}\ +\ 2\text{H}_{2}\text{O}$.

Chlorine, the useful halogen

Making chlorine. In the laboratory, hydrochloric acid is oxidised by manganese dioxide:

$\text{MnO}_{2}\ +\ 4\text{HCl}\ \rightarrow \ \text{MnCl}_{2}\ +\ \text{Cl}_{2}\ +\ 2\text{H}_{2}\text{O}$

A mixture of common salt, manganese dioxide and concentrated sulphuric acid works too: $4\text{NaCl}\ +\ \text{MnO}_{2}\ +\ 4\text{H}_{2}\text{SO}_{4}\ \rightarrow \ \text{MnCl}_{2}\ +\ 4\text{NaHSO}_{4}\ +\ 2\text{H}_{2}\text{O}\ +\ \text{Cl}_{2}$. Potassium permanganate needs no heat: $2\text{KMnO}_{4}\ +\ 16\text{HCl}\ \rightarrow \ 2\text{KCl}\ +\ 2\text{MnCl}_{2}\ +\ 8\text{H}_{2}\text{O}\ +\ 5\text{Cl}_{2}$.

In industry, the Deacon process oxidises hydrogen chloride with air over a copper(II) chloride catalyst at about 723 K: $4\text{HCl}\ +\ \text{O}_{2}\ \xrightarrow{\text{CuCl}_{2},\ \ 723\ \text{K}}\ 2\text{Cl}_{2}\ +\ 2\text{H}_{2}\text{O}$. Most chlorine today comes from the electrolysis of brine, which gives chlorine at the anode, hydrogen at the cathode and sodium hydroxide: $2\text{NaCl}\ +\ 2\text{H}_{2}\text{O}\ \rightarrow \ 2\text{NaOH}\ +\ \text{Cl}_{2}\ +\ \text{H}_{2}$.

With water and bleaching. Chlorine dissolves in water and reacts a little: $\text{Cl}_{2}\ +\ \text{H}_{2}\text{O}\ \rightarrow \ \text{HCl}\ +\ \text{HOCl}$. Hypochlorous acid breaks down and gives nascent oxygen, which means oxygen atoms that are ready to react. They oxidise the colouring matter to colourless compounds. This bleaching is permanent. Sulphur dioxide bleaches by reduction, and the colour comes back in air, so its action is temporary.

With alkali. The product depends on the temperature and on the strength of the alkali. This is an example of disproportionation, where one substance is oxidised and reduced at the same time.

• Cold dilute sodium hydroxide: $2\text{NaOH}\ +\ \text{Cl}_{2}\ \rightarrow \ \text{NaCl}\ +\ \text{NaOCl}\ +\ \text{H}_{2}\text{O}$. The chlorine goes to $-1$ and $+1$.
• Hot concentrated sodium hydroxide: $6\text{NaOH}\ +\ 3\text{Cl}_{2}\ \rightarrow \ 5\text{NaCl}\ +\ \text{NaClO}_{3}\ +\ 3\text{H}_{2}\text{O}$. The chlorine goes to $-1$ and $+5$.
• Slaked lime: $2\text{Ca}(\text{OH})_{2}\ +\ 2\text{Cl}_{2}\ \rightarrow \ \text{Ca}(\text{OCl})_{2}\ +\ \text{CaCl}_{2}\ +\ 2\text{H}_{2}\text{O}$. The solid is bleaching powder.

With other substances.

• Excess ammonia: $8\text{NH}_{3}\ +\ 3\text{Cl}_{2}\ \rightarrow \ 6\text{NH}_{4}\text{Cl}\ +\ \text{N}_{2}$. Excess chlorine gives the explosive nitrogen trichloride: $\text{NH}_{3}\ +\ 3\text{Cl}_{2}\ \rightarrow \ \text{NCl}_{3}\ +\ 3\text{HCl}$.
• Hydrogen sulphide: $\text{H}_{2}\text{S}\ +\ \text{Cl}_{2}\ \rightarrow \ 2\text{HCl}\ +\ \text{S}$.
• Iron(II) sulphate: $2\text{FeSO}_{4}\ +\ \text{H}_{2}\text{SO}_{4}\ +\ \text{Cl}_{2}\ \rightarrow \ \text{Fe}_{2}(\text{SO}_{4})_{3}\ +\ 2\text{HCl}$.
• Iron: $2\text{Fe}\ +\ 3\text{Cl}_{2}\ \rightarrow \ 2\text{FeCl}_{3}$.
• Methane in light: $\text{CH}_{4}\ +\ \text{Cl}_{2}\ \rightarrow \ \text{CH}_{3}\text{Cl}\ +\ \text{HCl}$, a substitution. With ethene it adds to give 1,2-dichloroethane.

Chlorine is poisonous. It is used to make drinking water safe, to bleach paper and cotton, and to make plastics such as PVC and many solvents.

Practise with new numbers

Try it in Python

The program goes through every pair of halogens and uses the electrode potentials to decide whether the first can oxidise the halide of the second.

Common mistakes

• Saying that fluorine has the most negative electron gain enthalpy. It is chlorine. Fluorine is the best oxidising agent for other reasons.
• Using the wrong alkali conditions. Cold dilute sodium hydroxide gives hypochlorite. Hot concentrated sodium hydroxide gives chlorate.
• Calling the bleaching by chlorine temporary. It is permanent, because it is oxidation.
• Giving fluorine positive oxidation states. In compounds it is always $-1$.

Remember these

• The halogens are $ns^2np^5$. F2 and Cl2 are gases, Br2 is a liquid and I2 is a solid.
• Chlorine has the most negative electron gain enthalpy. Fluorine is the strongest oxidising agent: $E^\circ = +2.87$ V.
• F2 has a weaker bond than Cl2, because of lone pair repulsion.
• Chlorine: $\text{MnO}_{2}\ +\ 4\text{HCl}$, Deacon process over CuCl2 at about 723 K, electrolysis of brine.
• Cold dilute alkali gives hypochlorite. Hot concentrated alkali gives chlorate. Slaked lime gives bleaching powder.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Why is the electron gain enthalpy of fluorine less negative than that of chlorine?

2. Chlorine water is added to potassium iodide solution. What do you see and what is the equation?

3. Why does chlorine bleach permanently while sulphur dioxide bleaches only for a time?

4. What does chlorine give with hot concentrated sodium hydroxide, and what is the oxidation state of chlorine in each product?

References

• NCERT, Chemistry Part I for Class XII, the chapter The p-Block Elements, on group 17 and on chlorine.
• OpenStax, Chemistry 2e, Chapter 18, on the halogens; the electrode potentials are in Appendix L.