Chemical properties of halogens

Fluorine

Fluorine gas is especionally chemically active; Fluorine molecule is the strongest chemical oxidant that has:

• Low energy of F₂  2F dissociation (159 kJ/mol). This energy is low due to the strong repulsion of nonbonding electrons in the molecule:

For instance, O₂ dissociation energy is 493.2 kJ/mol and Cl₂ dissociation requires 242.4 kJ.

• High affinity to the electron F + e^-  F^‑ (332.7 kJ/mol). This value is lower than in case of chlorine but dissociation energy compensates it.

These two factors (as well as small size of fluorine atom) determine a greater strength of covalent bonds of fluorine with other elements. For example, H—F bond energy is equal to 564.3 kJ/mol against 459.8 kJ/mol for H—O bond and 430,5 kJ/mol for H—Cl.

Coal, paper, and wood burn in fluorine atmosphere. The stable substances: glass (in the form of wool) as well as water self-ignite in fluorine either:

SiO₂ + 2 F₂  SiF₄ + O₂ (O₃ also forms); H = -657 kJ/mol

2H₂O + 2F₂  4HF + O₂ (O₃ is co-product); H = - 598 kJ/mol

Therefore, dissolution of gaseous F₂ in water is impossible. In addition, F₂ substitutes hydrogen in hydrogen compounds of other elements.

Interaction of fluorine with many elementary substances proceeds vigorously. It interacts with H₂ even at -220 ^oC (this temperature characterises transition to solid hydrogen). Fluorine reacts with sulfur and phosphorus at -190 ^oC (liquid air temperature).

S + 3F₂  SF₆ H = - 1207 kJ/mol

2P + 5F₂  2PF₅

Most metals (uncluding Au, Pt) burn in the atmosphere of fluorine on cold or at heating. At the same time, fluorine has no interaction with metallic Fe, Cu, Ni, Al, and Zn in ambient conditions. This unexpected behaviour can be explained by very tight protective coatings of fluorides formed on that metals surface.

When heating, fluorine oxidizes heavy noble gases (Kr, Xe, Rn), for example:

Xe + 3F₂  XeF₆

Elementary fluorine gas is inert towards n2 and o2, as well as He, Ar, Ne.

Fluorine substitutes other halogens in their compounds

NaI + F₂  NaF + I₂

All binary compounds of fluorine belong to one class of fluorides, in which fluorine is found in the oxidation state -1. Binary fluorides are divided obviously enough into ionic (NaF, MgF₂) and covalent (SiF₄, ClF₃). There are also a few compounds with the intermediate (ionic-covalent) character of bond (ZnF₂, MnF₂, CoF₂, NiF₂). Ionic fluorides are basic, and the covalent fluorides have acidic nature:

SiF₄ (acidic fluoride) + 2NaF (basic) = Na₂SiF₆

For comparison purposes, a reaction of interaction of basic and acid oxides is shown below:

Li₂O + CO₂ (acid oxide) = Li₂CO₃

Basic fluorides form alkaline medium due to hydrolysis:

NaF + H₂O HF + NaOH

and acidic fluorides, in turn, provide acid environment:

SiF₄ + 3H₂O H₂SiO₃ + 4HF

Compounds of fluorine

Hydrogen fluoride. Fluorine gas interacts with H₂ even at low temperatures explosively:

H₂ + F₂ = 2HF;  H = - 535 kJ g/mol

This direct synthesis has no practical significance for HF preparation. However, it can be the source of rocket fuel. The flame temperature at burning of F₂ + H₂ mixture exceeds 4000 ^oC.

Obtaining (in industry). This is the fluorite reaction with concentrated H₂SO₄

CaF₂ + H₂SO₄ = CaSO₄ + 2HF (at 120-300 ^oC)

Properties. Hydrogen fluoride, HF, is a colourless, very poisonous gas (m.p. = -83 ^оС, b.p = 19 ^oC). Moreover, this dangerous substance has rather weak smell. The liquid HF skin contact leads to the heaviest injuries dissolving albumens and penetrating deeply into the tissues, causing very painful and severe burns form there which heal slowly (especially under nails).

HF molecule is strongly polar [ = 1.91 D = 0.64^.10^-29 C^.m]. This value exceeds the value of electric dipole moment of water, SO₂ and NH₃. A molecule of HF has strong tendency to association as a result of formation of the closely-coupled chain hydrogen bonds.

It should be noted that formation of hydrogen bonds is determined by: 1) extraordinarily small size of the positively charged atom of hydrogen; 2) ability of hydrogen to penetrate deeply into the electronic shell of neighbouring atom (that doesn’t form covalent bond with it). Hereupon, electrostatic attraction together with donor-acceptor interaction is the reason of hydrogen bond formation.

Hydrogen bond energy in HF is a little bit stronger (33 kJ g/mol) than in H₂O molecules. That is why HF consists of mixture of polymers (HF)₂, (HF)₃, (HF)₄, (HF)₅, (HF)₆ even in the gaseous state. B.p. measurements show that hydrogen fluoride molecules have average composition (HF)₄.

Inorganic compounds mostly are well soluble in liquid HF. As a rule, such solutions are good conductors of electric current. Liquid HF is a strong ionising solvent. The dissolved substance accepts protons of HF molecules, increases HF₂^- ions concentration so it behaves as a base, for example:

KNO₃ + 2HF K+ + HNO₃ + HF₂^-

Chemical activity of HF depends substantially on the absence or presence even traces of water. Dry HF has no reaction with many metals and with metal oxides. On the other hand, when a reaction with an oxide will start even to negligibly small degree, it proceeds with autoacceleration farther as a result of the reaction:

MO + 2HF = MF₂ + H₂O

since the content of water grows.

HF dissolves in water without limits and a lot of heat (58,5 kJ/mol) liberates. The formation of azeotropic mixture (that boils without separation of components) takes place that contains 38.3% of hydrogen fluoride and boils at 112 ^oC.

Ionization of HF molecules occurs at dissolution in water with H₃O⁺ and F^- ions formation. The latter interact with HF molecules, as a result, hydrogen fluoride-ions are formed:

HF + H₂O H₃O+ + F- K = 7.2.10^-4  CH₃COOH

HF + F ^- HF₂^- K = 5.1 - due to hydrogen bonds

or total equation,

2HF + H₂O H₃O+ + H F₂.^-

Water solution of HF (hydrofluoric acid) is acid of intermediate strength. Its interesting feature is ability to interact with glass and quartz:

SiO₂ + HF = SiF₄ + 2H₂O

Therefore it is forbidden to store it in glass bottles. For this purposes bottles should be coated with polyethylene, lead or rubber. HF is used for chemical etching of glass, making on it various pictures.

Mixture of concentrated HF and HNO₃ (1:3) is stronger than “aqua regia” and able to dissolve Si, W, Ta, Mo, and Nb with formation of complex acid of these metals:

3Ta + 2IHF + 5HNO₃ = 3H₂[TaF₇] + 5NO + 10H₂O

FLUORIDES of OXYGEN. At 100-1000 К fluorine does not react directly with O₂. Unlike other halogens fluorine has no oxygencontaining acids (single exception being HFO synthesized recently).

Compounds of fluorine with oxygen are not stable and exist only at low temperatures. Fluorine is the most electronegative in these compounds therefore oxygen carries a positive charge (+2). For this reason they are named «fluorides of oxygen». Fluorides of oxygen are endothermic compounds of fluorine:

Compound	O F2	O2F2	O3F2
Hof, 298, kJ g/mol	+ 16.7	+ 21.2	+ 26.1

Among them OF₂ exists only room temperature– light yellow gas with a characteristic smell of fluorine, it dissolves very sparingly in water, is unstable, less active than fluorine, however it is a very strong oxidant. The covalent molecule OF₂ has triangular structure: d (F—O) = 0.142,  FOF = 103^o. OF₂ is formed when gaseous fluorine passes through cold dilute solution of alkali:

2 + 2 = + 2 + H₂O (1-2% solution of NaOH)

1e^.2 4e

In this case we put coefficients before an oxidant and product of oxidation (it can be put before reductant and the product of reduction)

Fluoride of oxygen OF₂ can be considered as an anhydride of hypofluorous acid HFO.

It reacts with non-metals, metals, inorganic and organic compounds. With water and alkalis it enters into the reaction of selfreduction-selfoxidation:

+ 2 = 2HF +

+ 2 = 2NaF + + H₂O

Other fluorides of oxygen O_nF₂ (n = 2-6) are obtained due to the effect of high-voltage electric discharge on the mixture of oxygen and fluorine under the diminished pressure and low temperatures. All of them decompose completely at low temperatures.

O₂F₂ is an orange toxic gas (d_F-O = 0.1575 nm, d_O-O = 0.1217 nm; OOF 109,5^о). m.p. = -154^оC, b.p. = -57 ^oC, the half-decay period O₂F₂ at -50^оС is 3 hours.

All fluorides of oxygen are very strong oxidants and fluorinating agents. So, O₂F₂ reacts with Xe, with many metals (among them Ag, Au, Pt). It forms salts of dioxygenyls with fluorides, for example:

PtF₄ + O₂F₂ = O₂PtF₆ (O₂⁺ -dioxygenyl)

The length of bond O—O (0.1123 nm) in it is less than in O₂ (0.1207 nm). All dioxygenyls contain unpaired electron, all of them are therefore paramagnetics. The best studied among them is O₂PtF₆ forming orange-red crystals, m.p. = 219^оC (with decomposition).