Properties of Halogens

Halogen definition

The halogens are also known as group 7 or group 17.

Fig. 1 - The halogens, shown in the periodic table highlighted in green

Depending on who you ask, there are either five or six members of the halogen group. The first five are fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Some scientists also consider the artificial element tennessine (Ts) to be a halogen. Although tennessine follows many of the trends shown by the other halogens, it also acts strangely by showing some of the properties of metals. For example, it doesn’t form negative ions. Astatine also shows some of the properties of a metal. Because of their unique behaviour, we will largely ignore both tennessine and astatine for the rest of this article.

Like most of the groups in the periodic table, the halogens have certain shared characteristics. Let’s explore some of them now.

Physical properties of halogens

The halogens are all non-metals. They show many of the physical properties typical of non-metals.

• They are poor conductors of heat and electricity.

• When solid, they are dull and brittle.

• They have low melting and boiling points.

Physical appearance

The halogens have distinct colours. They are also the only group to span all three states of matter at room temperature. Take a look at the table below.

Here's a diagram to help you visualise these four halogens.

Fig. 2 - The physical appearance of the first four halogens at room temperature

Atomic radius

As you move down the group in the periodic table, the halogens increase in atomic radius. This is because they each have one more electron shell. For example, fluorine has the electron configuration 1s² 2s² 2p⁵, and chlorine has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁵. Fluorine has just two main electron shells, whilst chlorine has three.

Fig. 3 - Fluorine and chlorine with their electron configurations. Notice how chlorine is a larger atom than fluorine

Melting and boiling points

As you can tell from their states of matter shown in the table earlier on, melting and boiling points increase as you go down the halogen group. This is because the atoms get larger and have more electrons. Because of this, they experience stronger van der Waals forces between molecules. These require more energy to overcome and so increase the element's melting and boiling points.

Volatility

Volatility is very closely related to melting and boiling points - it is the ease with which a substance evaporates. From the data above, it is easy to see that the volatility of the halogens decreases as you move down the group. Once again, this is all thanks to van der Waals forces. As you move down the group, the atoms get larger and so have more electrons. Because of this, they experience stronger van der Waals forces, decreasing their volatility.

Chemical properties of halogens

Halogens also have some characteristic chemical properties. For example:

• They have high electronegativity values.
• They form negative anions.
• They take part in the same types of reaction, including reacting with metals to form salts, and reacting with hydrogen to form hydrogen halides.
• They are found as diatomic molecules.
• Chlorine, bromine, and iodine are all sparingly soluble in water. There is no point even considering the solubility of fluorine - it reacts violently the instant it touches water!

Let's look at some trends in chemical properties within the halogen group.

Electronegativity

Knowing what you know about atomic radius, can you predict the trend in Electronegativity as you go down the halogen group? Take a look at Polarity if you need a reminder.

As you move down the group in the periodic table, the halogens decrease in electronegativity. Remember that electronegativity is an atom’s ability to attract a shared pair of electrons. Let’s investigate why this is the case.

Take fluorine and chlorine. Fluorine has nine protons and nine electrons - two of these electrons are in an inner electron shell. They shield the charge of two of fluorine’s protons, so each electron in fluorine’s outer shell only feels a charge of +7. Chlorine has seventeen protons and seventeen electrons. Ten of these electrons are in inner shells, shielding the charge of ten protons. As in fluorine, each of the electrons in chlorine’s outer shell only feels a charge of +7. This is the case for all of the halogens. But as chlorine has a larger atomic radius than fluorine, the outer shell electrons feel the attraction towards the nucleus less strongly. This means that chlorine has a lower electronegativity than fluorine.

In general, as you go down the group, electronegativity decreases. In fact, fluorine is the most electronegative element on the periodic table.

Fig. 4 - Halogen electronegativity

Electron affinity

Factors affecting electron affinity include nuclear charge, atomic radius, and shielding from inner electron shells.

As we go down the group in the periodic table, the halogen’s nuclear charge increases. However, this increased nuclear charge is offset by extra shielding electrons. This means that in all of the halogens, the incoming electron only feels a charge of +7.

As you go down the group, atomic radius also increases. This means that the incoming electron is further away from the nucleus and so feels the nucleus’s charge less strongly. Less energy is released when the atom gains an electron. Therefore, electron affinity decreases in magnitude as you go down the group.

Fig. 5 - Halogen electron affinity

There is one exception - fluorine. It has a lower magnitude electron affinity than chlorine. Let’s look at it a little more closely.

Fluorine has the electron configuration 1s² 2s² 2p⁵. When it gains an electron, the electron goes into the 2p subshell. Fluorine is a small atom and this subshell isn’t very big. That means the electrons already in it are densely clustered together. In fact, their charge is so dense that they partially repel the incoming electron, offsetting the increased attraction from the decreased atomic radius.

Reactivity

To understand the reactivity of halogens, we need to look at two different aspects of their behaviour: their oxidising ability and their reducing ability.

Oxidising ability

Halogens tend to react by gaining an electron. This means that they act as oxidising agents and are reduced themselves.

As you move down the group, oxidising ability decreases. In fact, fluorine is one of the best oxidising agents out there. You can show this by reacting halogens with iron wool.

• Fluorine reacts vigorously with cold iron wool - well, to tell the truth, fluorine reacts instantly with almost anything!

• Chlorine reacts quickly with heated iron wool.

• Gently warmed bromine reacts more slowly with heated iron wool.

• Strongly heated iodine reacts very slowly with heated iron wool.

Reducing ability

Halogens can also react by losing electrons. In this case they act as reducing agents and are oxidised themselves.

The reducing ability of halogens increases as you go down the group. For example, iodine is a much stronger reducing agent than fluorine.

Overall reactivity

Because halogens mostly act as oxidising agents, their overall reactivity follows a similar trend - it decreases as you go down the group. Let's explore this a little further.

A halogen’s reactivity depends a lot on how well it attracts electrons. This is all to do with its electronegativity. As we’ve already discovered, fluorine is the most electronegative element. This makes fluorine extremely reactive.

We can also use bond enthalpies to show the trend in reactivity. Take the bond enthalpy of carbon, for example. Bond enthalpy is the energy required to break a covalent bond in gaseous state, and decreases as you move down the group. Fluorine forms much stronger bonds to carbon than chlorine does - it is more reactive. This is because the bonded pair of electrons is further from the nucleus, so the attraction between the positive nucleus and the negative bonded pair is weaker.

Bond strength

The final chemical property of halogens that we'll look at today is their bond strength. We'll consider both the strength of the halogen-halogen bond (X-X), and the hydrogen-halogen bond (H-X).

Halogen-halogen bond strength

Halogens forms diatomic X-X molecules. The strength of this halogen-halogen bond, also known as its bond enthalpy, generally decreases as you move down the group. However, fluorine is an exception - the F-F bond is much weaker than the Cl-Cl bond. Take a look at the graph below.

Fig. 6 - Halogen-halogen (X-X) bond enthalpy

Bond enthalpy depends on the electrostatic attraction between the positive nucleus and the bonding pair of electrons. This in turn depends on the atom's number of unshielded protons, and the distance from the nucleus to the bonding electron pair. All halogens have the same number of electrons in their outer subshell and so have the same number of unshielded protons. However, as you move down the group in the periodic table, atomic radius increases, and so the distance from the nucleus to the bonding electron pair increases. This decreases the bond strength.

Fluorine breaks this trend. Fluorine atoms have seven electrons in their outer shell. When they form diatomic F-F molecules, each atom features one bonding pair of electrons and three lone pairs of electrons. Fluorine atoms are so small that when two come together to form a F-F molecule, the lone pairs of electrons in one atom repel those in the other atom quite strongly - so much so that they decrease the F-F bond enthalpy.

Hydrogen-halogen bond strength

Halogens can also form diatomic H-X molecules. The strength of the hydrogen-halogen bond decreases as you move down the group, as you can see from the graph below.

Fig. 7 - Hydrogen-halogen (H-X) bond enthalpy

Once again, this is due to the increasing atomic radius of the halogen atom. As atomic radius increases, the distance between the nucleus and the bonding pair of electrons increases, and so bond strength decreases. But note that in this instance, fluorine follows the trend. Hydrogen atoms don't have any lone pairs of electrons, and so there isn't any additional repulsion between the hydrogen atom and the fluorine atom. Therefore, the H-F bond has the highest strength out of all of the hydrogen-halogen bonds.

Uses of halogens

To finish, we'll consider some of the uses of halogens. In fact, they have a number of applications.

• Chlorine and bromine are used as disinfectants in a range of situations, from sterilising swimming pools and wounds to cleaning dishes and surfaces. In some countries, chicken meat is washed in chlorine to rid it of any harmful pathogens, such as salmonella and E. coli.

• Halogens can be used in lights. They improve the lifespan of the bulb.

• We can add halogens to drugs to make them dissolve in lipids more easily. This helps them cross through the phospholipid bilayer into our cells.

• Fluoride ions are used in toothpaste, where they form a protective layer around tooth enamel and prevent it from acid attack.

• Sodium chloride is also known as common table salt and is essential to human life. Similarly, we also need iodine in our body - it helps maintain optimum thyroid function.