Step 1: Understanding the Concept:
Fluorine is the first member of the halogen group (Group 17). It exhibits several anomalous properties compared to the rest of the group due to its extremely small atomic size, highest electronegativity, low F-F bond dissociation enthalpy, and the absence of vacant d-orbitals in its valence shell.
Step 2: Key Formula or Approach:
Approach: Evaluate each given statement against the established chemical and physical properties of the element fluorine.
Step 3: Detailed Explanation:
Let's analyze each statement:
- (A) It is highly electronegative element: True. Fluorine is the most electronegative element in the entire periodic table.
- (B) It exhibits only -1 oxidation state: True. Because it is the most electronegative element and lacks d-orbitals to expand its octet, it can only attract electrons and thus exclusively shows a $-1$ oxidation state in all its compounds (and $0$ in its elemental form $\text{F}_2$).
- (C) It has high bond dissociation enthalpy among all halogens: False. The F-F bond dissociation enthalpy is anomalously low, actually being lower than that of Cl-Cl and Br-Br. This occurs because the fluorine atom is very small, leading to strong interelectronic repulsions between the non-bonding electron pairs on the two closely situated fluorine atoms in the $\text{F}_2$ molecule. The correct trend is $\text{Cl}_2>\text{Br}_2>\text{F}_2>\text{I}_2$.
- (D) It form only one oxoacid: True. Due to its small size and very high electronegativity, it forms only a single oxoacid, which is fluoric(I) acid or hypofluorous acid ($\text{HOF}$).
Step 4: Final Answer:
The false statement is that fluorine has a high bond dissociation enthalpy among all halogens.