Group 17 Elements: On the periodic table, the halogens are to the left of the noble gases. Fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (A) are the five poisonous non-metallic elements that make up Group 17 of the periodic table (At). Despite the fact that astatine is radioactive and has only short-lived isotopes, it behaves similarly to iodine and is frequently classified as a halogen. Because halogen elements have seven valence electrons, forming a complete octet requires only one extra electron. Because of this, they are more reactive than other non-metal groups.

In the periodic table, the Group 17 elements exist as the second column from the right side. The group 17 elements consist of Fluorine, Chlorine, Bromine, Iodine, Astatine, and Tennessine. The group 17 elements are called halogens. We will study the importance of group 17 elements, their properties, and their uses in this article. Some of the classic uses of halogens are as follows: Iodine is used as an antiseptic, chlorine is a disinfectant, fluorine is used in toothpaste in the form of sodium fluoride as it prevents tooth decays.

Group 17 Elements Introduction

Fluorine, Chlorine, Bromine, Iodine, Astatine, and Tennessine constitute group \({\rm{VII A}}\) (\(17\) group elements) of the long form of the periodic table.

The elements of the group are known as halogens (halos \( = \) salt, gen \( = \) product). These are all p block elements with the differentiating electron entering the \({\rm{P}}\)-orbital. The outer electronic configuration is \({\rm{n}}{{\rm{s}}^2}{\rm{n}}{{\rm{P}}^5}.\) The first element, fluorine, is very reactive, and it differs significantly in properties from other elements. Fluorine is regarded as super halogen because of its high reactivity.

The group 17 elements electronic configuration, names, symbols, and atomic numbers are presented below.

Elements	Atomic Number	Electronic Configuration	With Inert gas core
Fluorine \(({\rm{F}})\)	\(9\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{5}}}\)	\({\rm{[He]2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{5}}}\)
Chlorine \(({\rm{Cl}})\)	\(17\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{5}}}\)	\({\rm{[Ne]3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{5}}}\)
Bromine \(({\rm{Br}})\)	\(35\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{p}}^{\rm{5}}}\)	\([{\rm{Ar}}]3\;{{\rm{d}}^{10}}4{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{5}}}\)
lodine \(({\rm{I}})\)	\(53\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{d}}^{{\rm{10}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^{\rm{5}}}\)	\({\rm{[Kr]4\;}}{{\rm{d}}^{{\rm{10}}}}{\rm{5\;}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^5}\)
Astatine \(({\rm{As}})\)	\(85\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{f}}^{{\rm{14}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^{\rm{6}}}{\rm{5}}{{\rm{d}}^{{\rm{10}}}}{\rm{6}}{{\rm{s}}^{\rm{2}}}{\rm{6}}{{\rm{p}}^{\rm{5}}}\)	\([{\rm{Xe}}]4{{\rm{f}}^{14}}5\;{{\rm{d}}^{10}}6{{\rm{s}}^2}6{{\rm{p}}^5}\)

These elements have one electron short of the inert gas configuration. Therefore, out of five electrons in the \({\rm{P}}\)-orbital, one electron is unpaired. Consequently, it takes part in chemical bonding.

General Characteristics of Halogens

The general physical characteristics of group 17 elements are given below:

Let us first understand the group \(17\) elements properties based on their physical characteristics.

Position of Halogens in the Periodic Table

Halogens occupy group \({\rm{VII A}}\) in the periodic table. Therefore, these are classified as representative elements. The differentiating electron enters the \({\rm{P}}\)- orbitals in these elements, and hence these are called \({\rm{p}}\)-block elements.

Physical State

Halogens in the uncombined state exist as diatomic covalent molecules. The forces existing between the individual molecules are weak Van der Waals forces that explain the volatile nature of these elements (gases).

Melting Points, Boiling Points and Densities- The melting point and the boiling points increase while volatility decreases with an increase in atomic number. As a result, the density increases from Fluorine to Iodine \(\left( {1.5,\,1.66,\,3.19,\,4.94\left( {\;{\rm{g}}\;{\rm{c}}{{\rm{m}}^{ – 3}}} \right)} \right.\) respectively for \({\rm{F, Cl, Br, I}}\)).

Size of the Atoms and Ions

The size of the atoms and the ions increases gradually from fluorine to Iodine, as shown below:

Element	Co-valent radius \(\left( {\mathop {\rm{A}}\limits^{\rm{o}} } \right)\)	Ionic radius \({{\rm{X}}^ – }\left( {\mathop {\rm{A}}\limits^{\rm{o}} } \right)\)
\({\rm{F}}\)	\(0.64\)	\(1.33\)
\({\rm{Cl}}\)	\(0.99\)	\(1.84\)
\({\rm{Br}}\)	\(1.14\)	\(1.96\)
\({\rm{l}}\)	\(1.33\)	\(2.20\)

Ionisation Energy

The ionisation energies of the halogens decrease as the atoms increase in size. However, the values are very high, and there is little tendency for the atoms to lose electrons and form positive ions.

Element	First ionisation energy \(({\rm{kJ}}/{\rm{mol}})\)
\({\rm{F}}\)	\(1680\)
\({\rm{Cl}}\)	\(1256\)
\({\rm{Br}}\)	\(1142\)
\({\rm{l}}\)	\(1008\)

Because of its small size, the ionisation energy of Fluorine is abnormally higher as compared to other halogens. Therefore, the electrons are firmly held in the small atom. Although the iodine atom is the largest, its ionisation energy is the least.

Bond Energy

Halogens form diatomic molecules. The bond energies are presented in the table:

Element	Bond dissociation enthalpy \(({\rm{kJ}}/{\rm{mol}})\)	Bond length of \({{\rm{X}}_2}({\rm{pm}})\)
\({{\rm{F}}_2}\)	\({\rm{158}}{\rm{.8}}\)	\(143\)
\({\rm{C}}{{\rm{l}}_2}\)	\({\rm{242}}{\rm{.6}}\)	\(199\)
\({\rm{B}}{{\rm{r}}_2}\)	\({\rm{192}}{\rm{.8}}\)	\(228\)
\({{\rm{I}}_2}\)	\({\rm{151}}{\rm{.1}}\)	\(266\)

It would be expected that the bond energy in the \({{\rm{X}}_2}\) molecules would decrease as the atom becomes larger since the increased size results in less effective overlap of orbitals. Chlorine, bromine and iodine show the expected trend, but the bond energy for fluorine does not fit this expected trend. Instead, the bond energy in fluorine is abnormally low. It is largely responsible for its high reactivity. Two explanations are given for this unexpectedly low energy.

1. Mulliken proposed that some multiple bonding occurs in chlorine, bromine, and Iodine, making the bonds stronger due to \({\rm{d, P}}\)-orbitals hybridisation in these molecules, compared to fluorine. This hybridisation is not possible in fluorine because there are no d orbitals in the valence shell of \({\rm{F}}\).

2. Coulson suggested that since fluorine atoms are small, the \({\rm{F  –  F}}\) distance is also small. Hence internuclear repulsion is high. The large electron-electron repulsion between the lone pairs of electrons on the two fluorine atoms weakens the bond. Coulson explanation is now widely accepted. This repulsion is less in other atoms because they are larger in size.

Electronegativity and Electron Affinity

The electronegativity of halogens is very high. For example, fluorine has the highest electronegativity value of \(4.0\) on the Pauling scale.

Element	Electronegativity	Electron affinity \(\left( {{\rm{kJ / mol}}} \right)\)
\({\rm{F}}\)	\(4.0\)	\(-333\)
\({\rm{Cl}}\)	\(3.2\)	\(-349\)
\({\rm{Br}}\)	\(3\)	\(-325\)
\({\rm{l}}\)	\(2.7\)	\(-296\)
\({\rm{At}}\)	\(2.2\)	\(-270\)

The electronegativity decreases gradually from fluorine to Iodine since the size of atoms increases. The electron affinity values for all halogens are negative. It shows that the energy is evolved when a halogen atom gains an electron. Thus, all the halogens form halide ions easily.

Chlorine has maximum electron affinity in this group. Fluorine is more electronegative than chlorine, but fluorine has a low electron affinity compared to chlorine. This lower value is attributed to the small size and compact \({\rm{2p}}\) subshell of fluorine. Because of this small size, there is repulsion between the electron pairs already present and added electrons. Therefore, fluorine exhibits a lesser tendency to accept electrons and to form fluoride ions. Thus, its electron infinity value is less than chlorine.

Types of Bonds and Oxidation States

Most compounds formed between the halogens and the metals are ionic. However, covalent halides are formed in few cases where the metals are very small and have a high charge. For example, \({\rm{BeC}}{{\rm{l}}_2}\) and \({\rm{AlC}}{{\rm{l}}_3}\) are examples of covalent compounds. It is explained by Fajans rule. The halogens have high electronegative values.

Hence when they react with metals, there will be a large electronegativity difference between them, leading to the formation of ions. Halide ions are easily produced because of their high affinity for electrons. When two halogen atoms form a molecule, they form a covalent bond. Most compounds between the halogens and non-metals are also covalent. Fluorine is always univalent since it is the most electronegative element. Its oxidation number is always \(-1\) with chlorine, bromine, and Iodine. A covalence of \(1\) is the most common.

Oxidation States

Halogens exhibits \( – 1,{\rm{ }} + 1,{\rm{ }} + 3,{\rm{ }} + 5\) and \( + 7\) oxidation states in their compounds.

As the halogen atoms contain seven electrons in their valency shells, they require only one electron to attain the nearest noble gas configuration. The shortage of one electron can be met either by gaining one electron or by sharing an electron. Due to highly electronegativity, halogens can easily take an electron from a lesser electronegative atom and exhibit a \(-1\) oxidation State.

However, even if they share one electron with the lesser electronegative atom, the oxidation state shown by them is \(-1\) because the shared electrons are counted with the more electronegative atom for determining the oxidation state. Thus, the \(-1\) oxidation state is the most common oxidation state of halogens. On the other hand, when a halogen atom shares an electron with the more electronegative element, the oxidation state exhibited by it is \(+1\).

Fluorine, being most electronegative, can only gain an electron. Moreover, it does not possess any vacant orbital in its valency shell and cannot excite its ns or np electron pairs. Therefore, fluorine is unable to show higher positive oxidation states. Hence fluorine always exists in a \(-1\) oxidation state in all of its compounds.

Except for fluorine, all other halogen atoms possess vacant \({\rm{d}}\)-orbitals in their valence shells and can therefore excite their ns and np electron pairs. It leads to the unpairing of electrons. The unpaired electrons can form covalent bonds with the lesser electronegative atoms. It enables chlorine and other heavier halogen to exist in the higher positive oxidation States.

The energy required for electrons unpairing and excitation to nd-orbitals is compensated from the energy released in the bond formation by the unpaired electrons. Therefore, in addition to the above-mentioned oxidation states, chlorine and bromine also exhibit \(+4\) and \(+6\) oxidation states in their oxides and oxoacids.

Colour

Halogens are coloured. The colour darkens on moving down the group.

Trends in Chemical Reactivity

The halogens are very reactive elements and react with metals, non-metals, and many other compounds. The reactivity decreases \({{\rm{F}}_2}\) is the most reactive amongst all the halogens, and this reactivity decreases from \({{\rm{F}}_2}\) to \({{\rm{I}}_2}.\) \({{\rm{F}}_2}\) and \({\rm{C}}{{\rm{l}}_2}\) often oxidise elements further than do \({\rm{B}}{{\rm{r}}_2}\) and \({{\rm{I}}_2}.\) Some typical reactions of halogens with hydrogen, oxygen, metals, water and alkalis are given below.

Halogens Reaction with Hydrogen

Halogens combine with hydrogen to form hydrogen halides of the type \({\rm{HX}}\). The reactivity decreases from \({{\rm{F}}_2}\) to \({{\rm{I}}_2}\).

\({{\rm{H}}_2} + {{\rm{F}}_2} \to 2{\rm{HF}}\)
\({{\rm{H}}_2} + {\rm{C}}{{\rm{l}}_2} \to 2{\rm{HCl}}\)
\({{\rm{H}}_2} + {\rm{B}}{{\rm{r}}_2} \to 2{\rm{HBr}}\)
\({{\rm{H}}_2} + {{\rm{I}}_2} \to 2{\rm{HI}}\)

The stability order of hydrogen halide is
\({\rm{HF}} > {\rm{HCl}} > {\rm{HBr}} > {\rm{HI}}\)

Halogens Reaction with Oxygen

All the halogens react with oxygen to form oxides. The oxides of fluorine are called oxygen fluorides instead of fluorine oxides. It is because fluorine is more electronegative than oxygen. The compounds of other halogens with oxygen are called halogen oxides. Halogen exhibit \(+ 1\) to \(+ 7\) oxidation states in these oxides.

In any series of oxides, the acidic character increases with an increase in the oxidation state of the halogen.

Halogens Reaction with Metals

Most metals form halides when reacted with halogen. \({{\rm{F}}_2}\) that reacts even with gold and Platinum. The general reaction is written as

\({\rm{n}}{{\rm{X}}_2} + 2{\rm{M}} \to 2{\rm{M}}{{\rm{X}}_{\rm{n}}}\)

Halogens Reaction with Water

All halogens react with water, but the reactivity decreases gradually from fluorine to Iodine.

\({{\rm{F}}_2}\) reacts with \({{\rm{H}}_2}{\rm{O}}\) and decomposers it to \({{\rm{O}}_2}\) and \({{\rm{O}}_3}.\)

\(2\;{{\rm{F}}_2} + {{\rm{H}}_2}{\rm{O}} \to 4{\rm{HF}} + {{\rm{O}}_2}\)
\(3\;{{\rm{F}}_2} + 3{{\rm{H}}_2}{\rm{O}} \to 6{\rm{HF}} + {{\rm{O}}_3}\)

\({\rm{C}}{{\rm{l}}_2}\) gives hydrochloric acid and hypochlorous acid

\({\rm{C}}{{\rm{l}}_2} + {{\rm{H}}_2}{\rm{O}} \to {\rm{HCl}} + {\rm{HOCl}}\)

But \({\rm{HOCI}}\) is unstable and decomposes to give \({\rm{HCI}}\) and nascent oxygen.

\({\rm{HOCl}} \to {\rm{HCl}} + ({\rm{O}})\)

Halogens Reaction with Alkalis

\({{\rm{F}}_2}\) reacts with cold and dilute alkalis to form fluoride and oxygen fluoride.

\(2\;{{\rm{F}}_2} + 2{\rm{KOH}} \to 2{\rm{KF}} + {\rm{O}}{{\rm{F}}_2} + {{\rm{H}}_2}{\rm{O}}\)

The other halogens, Chlorine, Bromine, and Iodine, react with a cold dilute solution of alkalis and give the corresponding halides and the hypohalites.

\({\rm{C}}{{\rm{l}}_2} + 2{\rm{KOH}} \to {\rm{KCl}} + {\rm{KClO}} + {{\rm{H}}_2}{\rm{O}}\)
\({\rm{B}}{{\rm{r}}_2} + 2{\rm{KOH}} \to {\rm{KBr}} + {\rm{KBrO}} + {{\rm{H}}_2}{\rm{O}}\)
\({{\rm{I}}_2} + 2{\rm{NaOH}} \to {\rm{NaI}} + {\rm{NaIO}} + {{\rm{H}}_2}{\rm{O}}\)

Are Group 17 Elements Metal or Non-Metal?

Group \(17\) elements are strong non-metals. They have \(7\) electrons in their outermost shell, so they need to gain one electron to get an electron configuration like the noble gas elements.

Uses of Halogens

1. Fluorine is used to prepare compounds such as Teflon, freon, etc.

2. Chlorine is a good bleaching agent. It bleaches wood pulp, rayon and cotton. Hence it is used in the paper and the textile industries as a bleaching agent.

3. Bromine is used in dyestuffs, pharmaceuticals, etc.

4. Iodine is used in printing inks, disinfectants, and dyes.

Summary

Group 17 elements and their compounds show gradation in their physical and chemical properties. Hence, their inclusion in group 17 of the periodic table is justified. In this article, we studied the seventh group elements, properties, and trends. We also learned the uses of various halogens.

FAQs – Group 17 Elements

Q.1. What family is Group \(17\) on the periodic table?
Ans: The non-metallic elements \({\rm{F, Cl, Br, I}}\) and \({\rm{At}}\) are grouped to form group \(17\) of the periodic table, and they belong to the halogen family.

Q.2. Which element of the \({\rm{VII A}}\) group has the highest electronegativity, and what is the value?
Ans: Fluorine shows the highest electronegativity, and its value is \(4.0\).

Q.3. Which general trend is demonstrated by the group \(17\) elements considered in order from top to bottom on the Periodic Table?
Ans: The halogen elements have seven electrons in their valence shell. As the atomic radius increases from top to bottom, the electronegativity decreases.

Q.5. What element is in group 17 period 3?
Ans: Chlorine is a chemical element that belongs to the periodic table’s period 3, group 17.