Anomalous behaviour of Fluorine: Oxidation State, Reactivity, Examples

Anomalous Behaviour of Fluorine: Fluorine with atomic number \(9\) is the first member of the halogen family. Where electronegativity, ionisation enthalpy, and electrode potentials are high for Fluorine compared to other halogens. Its ionic and covalent radii, m.p. and b.p., bond dissociation enthalpy, and electron gain enthalpy are quite low than expected. Fluorine has an extremely small atomic size. But why does Fluorine behave so abnormally? In this article, we will learn about fluorine and its anomalous behaviour.

Anomalous Behaviour of Fluorine

Fluorine shows exceptional behaviour as compared to other members of the family because of its-

1. Extremely small atomic size.
2. Very high values of electronegativity and ionisation energy.
3. Absence of vacant d-orbitals in its valence shell.
4. Low F-F bond dissociation energy \((158.8 \mathrm{KJ} /\)mole).
5. Strong tendency to form Hydrogen bonds.

Anomalous Behaviour of Fluorine: Properties

• Oxidation State- Due to its high electronegativity and absence of vacant d orbitals, Fluorine exhibits only a \(-1\) oxidation state in its compounds. However, other halogens exhibit \(-1, +1, +3, +5\) and \(+7\) oxidation states due to low electronegativities and the presence of vacant d-orbitals in their outermost shells.
• Covalency- Fluorine exhibits a covalency of only \(+1\) (in F-F molecule), while other members of the halogen family can show covalency even up to \(+7\) due to vacant d-orbitals in their valence shell.
• Hydrogen Bonding- The high electronegativity of Fluorine accounts for its strong tendency to form H-bonds in its compounds with hydrogen. While the other members of the halogen group form oxyacids such as \({\rm{HClO}},{\rm{HBrO}},{\rm{HIO}},{\rm{HCl}}{{\rm{O}}_3},{\rm{HBr}}{{\rm{O}}_3},{\rm{HI}}{{\rm{O}}_3},{\rm{HCl}}{{\rm{O}}_4}\) and \({\rm{HI}}{{\rm{O}}_4}\) Hydrogen fluoride is an associated molecule due to hydrogen bonding.
• Formation of hexafluorides- Due to its high electronegativity, Fluorine easily transforms S into \(a+6\) oxidation state, forming a stable hexafluoride \(\mathrm{SF}_{6}\). No other halogens form the hexahalide.
• Oxidising Nature- Due to its high electronegativity and electron-affinity, Fluorine has an increased tendency to accept an electron. Hence, it acts as a stronger oxidising agent than other members of the halogen family. Though the electron gain enthalpy of Fluorine is less negative as compared to chlorine, it acts as a stronger oxidising agent than chlorine. This is because the standard reduction potential of Fluorine (i.e. \(+2.87 \mathrm{~V})\) is more positive than chlorine (i.e.\(=1.36 \mathrm{~V})\)
  Oxidising power: \({{\rm{F}}_{\rm{2}}}{\rm{ > C}}{{\rm{l}}_{\rm{2}}}{\rm{ > B}}{{\rm{r}}_{\rm{2}}}{\rm{ > }}{{\rm{I}}_{\rm{2}}}\)

• Reactivity- Fluorine is the most reactive of all halogens. This is due to the low F-F dissociation enthalpy of the \(\mathrm{F}_{2}\) molecule. It can displace other halogens from their salt solutions.

1. Reaction with hydrogen: Fluorine reacts with hydrogen in the dark at a low temperature, whereas other halogens do not react with hydrogen in the dark.
2. Action with metals: Fluorine reacts with gold and platinum, but other halogens do not react with noble metals.
3. Action with non-metals: Fluorine combines directly with the non-metals like carbon, silicon, nitrogen etc., to give their fluorides. However, the other halogens do not combine directly with these elements.

\({\rm{C + 2}}{{\rm{F}}_2} \to {\rm{C}}{{\rm{F}}_{\rm{4}}}\)

\({{\rm{N}}_{\rm{2}}}{\rm{ + 3}}{{\rm{F}}_{\rm{2}}} \to {\rm{2N}}{{\rm{F}}_{\rm{3}}}\)

Action with water: Fluorine reacts with water forming \({\rm{HF,}}{{\rm{O}}_{\rm{2}}}\) and \({{\rm{O}}_3}\) However, the other halogens do not give ozone to water.

\({\rm{2}}{{\rm{F}}_{\rm{2}}}{\rm{ + 2}}{{\rm{H}}_{\rm{2}}}{\rm{O}} \to {\rm{4HF + }}{{\rm{O}}_{\rm{2}}} \uparrow \)

\({\rm{3}}{{\rm{F}}_{\rm{2}}}{\rm{ + 3}}{{\rm{H}}_{\rm{2}}}{\rm{O}} \to 6{\rm{HF}} + {{\rm{O}}_3} \uparrow \)

Action with alkalies: Fluorine reacts with caustic alkalies such as \(\mathrm{NaOH}\) to form oxygen difluoride.

\(2\;{{\rm{F}}_2} + 2{\rm{NaOH}} \to 2{\rm{NaF}} + {{\rm{H}}_2}{\rm{O}} + {\rm{O}}{{\rm{F}}_2}\)

The other members of the halogen family react with cold and dilute alkalies to form hypohalites.

\({\rm{C}}{{\rm{l}}_2} + 2{\rm{NaOH}} \to {\rm{NaCl}} + {\rm{NaClO}} + {{\rm{H}}_2}{\rm{O}}\)

However, with hot and concentrated alkalies, they form higher oxy salts, halates.

\(3{\rm{C}}{{\rm{l}}_2} + 6{\rm{NaOH}} \to 5{\rm{NaCl}} + {\rm{NaCl}}{{\rm{O}}_3} + 3{{\rm{H}}_2}{\rm{O}}\)

Formation of oxyacids: Fluorine is the strongest oxidising agent hence does not form any oxyacid, while the remaining members of the halogen family form four types of oxyacids. These are-

HXO-Hypohalous acid 

\({\rm{HX}}{{\rm{O}}_2} – \) Halous acid 

\({\rm{HX}}{{\rm{O}}_3} – \) Halic acid 

\({\rm{HX}}{{\rm{O}}_4} – \) Per-halic acid

Behaviour of hydracids: 

a. The hydracid of Fluorine HF is a liquid, whereas the hydracids of other halogens HCl, HBr, HI exist as gases at ordinary temperatures.

b. HF is a weak acid, while the other hydracids are strong acids.

c. The hydracid of Fluorine HF is the most stable of all the hydracids of remaining halogens. 

d. HF forms acid salts such as \(\mathrm{NaHF}_{2}\) and complex acids such as \({\rm{HB}}{{\rm{F}}_{{4^\prime }}}{{\rm{H}}_2}{\rm{Si}}{{\rm{F}}_6}\) while the other halogens acids do not form such salts and acids.

e . HF is the only acid to react with silica, silicates and hence attacks glass. 

f . Order of bond dissociation energy: \({\rm{HI > HBr > HCl > HF}}\)

g. Order of acidic strength: \({\rm{HF < HCl > HBr > HI}}\)

1. Silver fluoride, AgF is soluble in water, whereas other silver halides (AgCl, AgBr, Agl) are insoluble. 
2. The fluorides of calcium, strontium and barium are insoluble in water, whereas corresponding salts of other halogens are soluble. 
   Fluorides have the highest ionic character. Other halides have covalent characters.
3. Electron affinity: \(\mathrm{Cl}>\mathrm{F}>\mathrm{Br}>1\)

Summary

Fluorine is positioned in group VII of the periodic table. The group VII elements of the periodic table are known as halogens. Hence, Fluorine with atomic number \(9\) and electronic configuration \(2,7\) is the first member of the halogen family. However, its behaviour and properties do not match with the other members of the halogen family. It behaves abnormally due to its extremely small atomic size, high electronegativity, low F-F bond dissociation enthalpy and absence of vacant d-orbital. It is the strongest reducing agent and exhibits an oxidation state of \(-1\). In this article, we learnt how Fluorine differs from other members of its family. 

Frequently Asked Questions

Q.1. What do you mean by anomalous behaviour?
Ans: Anomalous behaviour is a behaviour that deviates from its original order. It exhibits different properties compared to other members of its groups. 

Q.2. Why is the behaviour of fluorine different from other halogens?
Ans: Fluorine differs from the rest of the members of the halogen family due to:
(i) its small atomic size, (ii) the highest electronegativity, (iii) low bond dissociation enthalpy and (iv) absence of d-orbitals in its valence shell. 

Q.3. Why does fluorine exhibit a \(-1\) oxidation state only?
Ans: As per the general electronic configuration, there are \(7\) electrons in the outermost shell of the fluorine atom. It gains one electron to attain stability. Hence, it exhibits a \(-1\) oxidation state. As there are no d-orbitals available in its valence shell, fluorine atoms do not exhibit any other oxidation state.

Q.4. Does fluorine show a positive oxidation state?
Ans: Fluorine does not have a vacant d orbital, so it cannot exhibit a positive oxidation state.

Q.5. Is fluorine or chlorine more reactive?
Ans: Chlorine has more electrons as compared to fluorine. Hence, it suffers electronic repulsions making it less likely to react.

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