Hydrogen fluoride.Fluorine gas interacts with H2 even at low temperatures explosively:
H2 + F2 = 2HF; D 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 îÇ, 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 [m = 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 theclosely-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 does not 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 limitsand 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 (hydrofluoricacid) 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 Äî 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:
|DHof, 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,
2 + 2 = + 2 + H2O (1-2% solution of NaOH)
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 hypofluorousacid 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 OnF2 (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.
O2F2is an orange toxic gas (dF-O = 0.1575 nm, dO-O = 0.1217 nm; ?OOF 109,5î). m.p. = -154îC, b.p. = -57 oC, the half-decay period O2F2 at -50îÇ 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. = 219îC (with decomposition).