Fluorine, the lightest of the halogens, has several unique features to its chemistry which merit individual discussion:

The bond strength of the F₂ molecule is unexpectedly low when compared to the other dihalogens. This is due to the repulsion between the lone pairs on the fluorine atoms (non-bonding pair repulsion) which becomes important when they are so close together. In the larger halogens, the atoms are not so close together and this extra repulsion is not so important. Also, there may be extra bonding contributions from pπ-dπ interactions in the larger halogens.

It forms very strong bonds to other elements when this non-bonding pair repulsion is minimal, and when π-bonding from fluorine to empty orbitals on the bonded element may occur.

The boiling points of fluorine containing compounds tend to be lower than similar hydrogen containing compounds, and are always lower than the chlorine containing compound (eg. BP(CF₄) = -127.9^oC, BP(CH₄) = -161.5^oC, BP(CCl₄) = 76.7^oC).

The volatility of a compound reflects the degree of attractive intermolecular forces. For these, closed-shell, species the principal interactions are van der Waals forces, or instantaneous dipole-dipole interactions. The small F atoms mean that the electrons are tightly held, and hence undergo instantaneous distortion poorly and so the intermolecular forces are small. This low polarizability (the ability to distort) also leads to the general inertness of fluorine containing compounds.

The fluorine atom acts as a good electron withdrawing agent, due to its high electronegativity. The substitution of a CH₃ group by CF₃ can enhance the Bronsted acidity of the parent compound, eg. HSO₃CF₃ has pK_a = 3.0 compared to HSO₃CH₃ which has pK_a = 6.0. These F atoms may also enhance the Lewis acidity: the lone pairs on the F atom can act as electron pair donors, eg. SbF₅ is a much stronger Lewis acid than SbCl₅.

The fluoride ion is very good at stabilizing high oxidation states, as seen in compounds such as IF₇, PtF₆, BiF₅ and AgF₄^- (Ag^III is very uncommon). On the other hand, fluorine tends to disfavour low oxidation states (the fluoride ion is an oxidizing agent), as shown by the existence of CuCl, CuBr and CuI, but the non-existence of CuF. The high lattice enthalpy of CuF₂, containing the small, doubly charged cation Cu²⁺ drives CuF to disproportionate into Cu and CuF₂.

The ability of fluorine to stabilize high oxidation states is also linked to its small size: this sterically allows the large numbers of F atoms to bond to the central atom to stabilize the high oxidation state.