See what "fluorine" is in other dictionaries. Reactivity of halogens Reaction of halogens with water

Reactions of adults Many parents do not always know how to react to the actions and some actions of their children. When we encounter problems, we react in three different ways.1. We pretend that nothing happened.2. We identify the enemy and attack.3. We are real

The hydrogen atom has the electronic formula of the outer (and only) electronic level 1 s 1 . On the one hand, by the presence of one electron in the outer electronic level, the hydrogen atom is similar to alkali metal atoms. However, just like halogens, it lacks only one electron to fill the external electronic level, since no more than 2 electrons can be located on the first electronic level. It turns out that hydrogen can be placed simultaneously in both the first and the penultimate (seventh) group of the periodic table, which is sometimes done in various versions of the periodic system:

From the point of view of the properties of hydrogen as a simple substance, it nevertheless has more in common with halogens. Hydrogen, as well as halogens, is a non-metal and forms diatomic molecules (H 2) similarly to them.

Under normal conditions, hydrogen is a gaseous, inactive substance. The low activity of hydrogen is explained by the high strength of the bond between hydrogen atoms in the molecule, which requires either strong heating or the use of catalysts, or both at the same time, to break it.

Interaction of hydrogen with simple substances

with metals

Of the metals, hydrogen reacts only with alkali and alkaline earth! Alkali metals include metals of the main subgroup of group I (Li, Na, K, Rb, Cs, Fr), and alkaline earth metals are metals of the main subgroup of group II, except for beryllium and magnesium (Ca, Sr, Ba, Ra)

When interacting with active metals, hydrogen exhibits oxidizing properties, i.e. lowers its oxidation state. In this case, hydrides of alkali and alkaline earth metals are formed, which have an ionic structure. The reaction proceeds when heated:

It should be noted that interaction with active metals is the only case when molecular hydrogen H2 is an oxidizing agent.

with non-metals

Of non-metals, hydrogen reacts only with carbon, nitrogen, oxygen, sulfur, selenium and halogens!

Carbon should be understood as graphite or amorphous carbon, since diamond is an extremely inert allotropic modification of carbon.

When interacting with non-metals, hydrogen can only perform the function of a reducing agent, that is, it can only increase its oxidation state:

Interaction of hydrogen with complex substances

with metal oxides

Hydrogen does not react with metal oxides that are in the activity series of metals up to aluminum (inclusive), however, it is able to reduce many metal oxides to the right of aluminum when heated:

with non-metal oxides

Of the non-metal oxides, hydrogen reacts when heated with oxides of nitrogen, halogens, and carbon. Of all the interactions of hydrogen with non-metal oxides, its reaction with carbon monoxide CO should be especially noted.

The mixture of CO and H 2 even has its own name - “synthesis gas”, since, depending on the conditions, such demanded industrial products as methanol, formaldehyde and even synthetic hydrocarbons can be obtained from it:

with acids

Hydrogen does not react with inorganic acids!

Of the organic acids, hydrogen reacts only with unsaturated acids, as well as with acids containing functional groups capable of being reduced by hydrogen, in particular aldehyde, keto or nitro groups.

with salts

In the case of aqueous solutions of salts, their interaction with hydrogen does not occur. However, when hydrogen is passed over solid salts of some metals of medium and low activity, their partial or complete reduction is possible, for example:

Chemical properties of halogens

Halogens are the chemical elements of group VIIA (F, Cl, Br, I, At), as well as the simple substances they form. Hereinafter, unless otherwise stated, halogens will be understood as simple substances.

All halogens have a molecular structure, which leads to low melting and boiling points of these substances. Halogen molecules are diatomic, i.e. their formula can be written in general form as Hal 2 .

It should be noted such a specific physical property of iodine as its ability to sublimation or, in other words, sublimation. sublimation, they call the phenomenon in which a substance in the solid state does not melt when heated, but, bypassing the liquid phase, immediately passes into the gaseous state.

The electronic structure of the external energy level of an atom of any halogen has the form ns 2 np 5, where n is the period number of the periodic table in which the halogen is located. As you can see, only one electron is missing from the eight-electron outer shell of the halogen atoms. From this it is logical to assume the predominantly oxidizing properties of free halogens, which is also confirmed in practice. As you know, the electronegativity of non-metals decreases when moving down the subgroup, and therefore the activity of halogens decreases in the series:

F 2 > Cl 2 > Br 2 > I 2

Interaction of halogens with simple substances

All halogens are highly reactive and react with most simple substances. However, it should be noted that fluorine, due to its extremely high reactivity, can react even with those simple substances with which other halogens cannot react. Such simple substances include oxygen, carbon (diamond), nitrogen, platinum, gold, and some noble gases (xenon and krypton). Those. actually, fluorine does not react only with some noble gases.

The remaining halogens, i.e. chlorine, bromine and iodine are also active substances, but less active than fluorine. They react with almost all simple substances except oxygen, nitrogen, carbon in the form of diamond, platinum, gold and noble gases.

Interaction of halogens with non-metals

hydrogen

All halogens react with hydrogen to form hydrogen halides with the general formula HHal. At the same time, the reaction of fluorine with hydrogen begins spontaneously even in the dark and proceeds with an explosion in accordance with the equation:

The reaction of chlorine with hydrogen can be initiated by intense ultraviolet irradiation or heating. Also leaks with an explosion:

Bromine and iodine react with hydrogen only when heated, and at the same time, the reaction with iodine is reversible:

phosphorus

The interaction of fluorine with phosphorus leads to the oxidation of phosphorus to the highest oxidation state (+5). In this case, the formation of phosphorus pentafluoride occurs:

When chlorine and bromine interact with phosphorus, it is possible to obtain phosphorus halides both in the + 3 oxidation state and in the + 5 oxidation state, which depends on the proportions of the reactants:

In the case of white phosphorus in an atmosphere of fluorine, chlorine or liquid bromine, the reaction begins spontaneously.

The interaction of phosphorus with iodine can lead to the formation of only phosphorus triiodide due to the significantly lower oxidizing ability than other halogens:

gray

Fluorine oxidizes sulfur to the highest oxidation state +6, forming sulfur hexafluoride:

Chlorine and bromine react with sulfur, forming compounds containing sulfur in oxidation states that are extremely unusual for it +1 and +2. These interactions are very specific, and to pass the exam in chemistry, the ability to write down the equations of these interactions is not necessary. Therefore, the following three equations are given rather for guidance:

Interaction of halogens with metals

As mentioned above, fluorine is able to react with all metals, even such inactive ones as platinum and gold:

The remaining halogens react with all metals except platinum and gold:

Reactions of halogens with complex substances

Substitution reactions with halogens

More active halogens, i.e. whose chemical elements are located higher in the periodic table, are able to displace less active halogens from the hydrohalic acids and metal halides they form:

Similarly, bromine and iodine displace sulfur from solutions of sulfides and or hydrogen sulfide:

Chlorine is a stronger oxidizing agent and oxidizes hydrogen sulfide in its aqueous solution not to sulfur, but to sulfuric acid:

Interaction of halogens with water

Water burns in fluorine with a blue flame in accordance with the reaction equation:

Bromine and chlorine react differently with water than fluorine. If fluorine acted as an oxidizing agent, then chlorine and bromine disproportionate in water, forming a mixture of acids. In this case, the reactions are reversible:

The interaction of iodine with water proceeds to such an insignificant degree that it can be neglected and considered that the reaction does not proceed at all.

Interaction of halogens with alkali solutions

Fluorine, when interacting with an aqueous solution of alkali, again acts as an oxidizing agent:

The ability to write this equation is not required to pass the exam. It is enough to know the fact about the possibility of such an interaction and the oxidizing role of fluorine in this reaction.

Unlike fluorine, the remaining halogens disproportionate in alkali solutions, that is, they simultaneously increase and decrease their oxidation state. At the same time, in the case of chlorine and bromine, depending on the temperature, flow in two different directions is possible. In particular, in the cold, the reactions proceed as follows:

and when heated:

Iodine reacts with alkalis exclusively according to the second option, i.e. with the formation of iodate, because hypoiodite is unstable not only when heated, but also at ordinary temperatures and even in the cold.

Halogens are the most reactive group of elements in the periodic table. They are made up of molecules with very low bond dissociation energies (see Table 16.1) and their atoms have seven electrons in their outer shell and are therefore very electronegative. Fluorine is the most electronegative and most reactive non-metallic element in the periodic table. The reactivity of halogens gradually decreases as you move towards the bottom of the group. The next section will consider the ability of halogens to oxidize metals and non-metals and show how this ability decreases in the direction from fluorine to iodine.

Halogens as oxidizing agents

When gaseous hydrogen sulfide is passed through chlorine water, sulfur is precipitated. The reaction proceeds according to the equation

In this reaction, chlorine oxidizes hydrogen sulfide, taking hydrogen from it. Chlorine also oxidizes to For example, if you mix chlorine with an aqueous solution of sulfate by shaking, sulfate is formed

The oxidative half-reaction that occurs in this case is described by the equation

As another example of the oxidizing action of chlorine, we present the synthesis of sodium chloride by burning sodium in chlorine:

In this reaction, sodium is oxidized as each sodium atom loses an electron to form a sodium ion:

Chlorine attaches these electrons, forming chloride ions:

Table 16.3. Standard electrode potentials of halogens

Table 16.4. Standard enthalpies of formation of sodium halides

All halogens are oxidizing agents, of which fluorine is the strongest oxidizing agent. In table. 16.3 shows the standard electrode potentials of halogens. From this table it can be seen that the oxidizing power of halogens gradually decreases towards the bottom of the group. This pattern can be demonstrated by adding a solution of potassium bromide to a vessel of chlorine gas. Chlorine oxidizes bromide ions, resulting in the formation of bromine; this causes a color to appear in a previously colorless solution:

Thus, it can be seen that chlorine is a stronger oxidizing agent than bromine. Similarly, if a solution of potassium iodide is mixed with bromine, a black precipitate of solid iodine is formed. This means that bromine oxidizes iodide ions:

Both reactions described are examples of displacement (substitution) reactions. In each case, the more reactive, that is, the stronger oxidizing agent, halogen displaces the less reactive halogen from solution.

Oxidation of metals. Halogens readily oxidize metals. Fluorine easily oxidizes all metals except gold and silver. We have already mentioned that chlorine oxidizes sodium, forming sodium chloride with it. To give another example, when a stream of chlorine gas is passed over the surface of heated iron filings, a brown solid chloride is formed:

Even iodine is capable, albeit slowly, of oxidizing metals below it in the electrochemical series. The ease of oxidation of metals by various halogens decreases when moving to the lower part of group VII. This can be verified by comparing the energies of formation of halides from the initial elements. In table. 16.4 shows the standard enthalpies of formation of sodium halides in order of movement to the bottom of the group.

Oxidation of non-metals. With the exception of nitrogen and most of the noble gases, fluorine oxidizes all other non-metals. Chlorine reacts with phosphorus and sulfur. Carbon, nitrogen and oxygen do not react directly with chlorine, bromine or iodine. The relative reactivity of halogens to non-metals can be judged by comparing their reactions with hydrogen (Table 16.5).

Oxidation of hydrocarbons. Under certain conditions, halogens oxidize hydrocarbons.

Table 16.5. Reactions of halogens with hydrogen

delivery. For example, chlorine completely removes hydrogen from the turpentine molecule:

The oxidation of acetylene can proceed with an explosion:

Reactions with water and alkalis

Fluorine reacts with cold water to form hydrogen fluoride and oxygen:

Chlorine slowly dissolves in water, forming chlorine water. Chlorine water has a slight acidity due to the fact that a disproportionation (see section 10.2) of chlorine occurs in it with the formation of hydrochloric acid and hypochlorous acid:

Bromine and iodine disproportionate in water in a similar way, but the degree of disproportionation in water decreases from chlorine to iodine.

Chlorine, bromine and iodine also disproportionate in alkalis. For example, in cold dilute alkali, bromine disproportionates into bromide ions and hypobromite ions (bromate ions):

When bromine interacts with hot concentrated alkalis, disproportionation proceeds further:

Iodate(I), or hypoiodite ion, is unstable even in cold dilute alkalis. It spontaneously disproportionates to form an iodide ion and an iodate(I) ion.

The reaction of fluorine with alkalis, like its reaction with water, is not similar to similar reactions of other halogens. In cold dilute alkali, the following reaction proceeds:

In hot concentrated alkali, the reaction with fluorine proceeds as follows:

Analysis for halogens and with the participation of halogens

Qualitative and quantitative analysis for halogens is usually performed using a silver nitrate solution. For example

For the qualitative and quantitative determination of iodine, a starch solution can be used. Since iodine is very slightly soluble in water, it is usually analyzed in the presence of potassium iodide. This is done because iodine forms a soluble triiodide ion with the iodide ion.

Solutions of iodine with iodides are used for the analytical determination of various reducing agents, for example, as well as some oxidizing agents, for example. Oxidizing agents shift the above equilibrium to the left, releasing iodine. Iodine is then titrated with thiosulfate (VI).

So let's do it again!

1. The atoms of all halogens have seven electrons in their outer shell.

2. To obtain halogens in the laboratory, the oxidation of the corresponding hydrohalic acids can be used.

3. Halogens oxidize metals, non-metals and hydrocarbons.

4. Halogens disproportionate in water and alkalis, forming halide ions, hypohalogenite and halogenate (-ions.

5. Patterns of changes in the physical and chemical properties of halogens when moving to the bottom of the group are shown in table. 16.6.

Table 16.6. Patterns of changes in the properties of halogens as the atomic number increases

6. Fluorine has anomalous properties among other halogens for the following reasons:

a) it has a low bond dissociation energy;

b) in fluorine compounds, it exists only in one oxidation state;

c) fluorine is the most electronegative and the most reactive among all non-metallic elements;

d) its reactions with water and alkalis differ from similar reactions of other halogens.

Fluorine

FLUORINE-A; m.[from Greek. phthoros - death, destruction] Chemical element (F), light yellow gas with a pungent odor. Add to drinking water f.

(lat. Fluorum), a chemical element of group VII of the periodic system, refers to halogens. Free fluorine consists of diatomic molecules (F 2); pale yellow gas with a pungent odor t pl –219.699°C, t bale –188.200°C, density 1.7 g/l. The most active non-metal: reacts with all elements except helium, neon and argon. The interaction of fluorine with many substances easily turns into combustion and explosion. Fluorine destroys many materials (hence the name: Greek phthóros - destruction). The main minerals are fluorite, cryolite, fluorapatite. Fluorine is used to obtain organofluorine compounds and fluorides; fluorine is part of the tissues of living organisms (bones, tooth enamel).

FLUORINE (lat. Fluorum), F (read "fluorine"), a chemical element with atomic number 9, atomic mass 18.998403. Natural fluorine consists of one stable nuclide (cm. NUCLIDE) 19 F. Outer electron layer configuration 2 s 2 p 5 . In compounds, it exhibits only the oxidation state –1 (valency I). Fluorine is located in the second period in group VIIA of the periodic system of elements of Mendeleev, refers to halogens (cm. HALOGENS).
The radius of the neutral fluorine atom is 0.064 nm, the radius of the F ion is 0.115 (2), 0.116 (3), 0.117 (4) and 0.119 (6) nm (the value of the coordination number is indicated in brackets). The successive ionization energies of a neutral fluorine atom are 17.422, 34.987, 62.66, 87.2, and 114.2 eV, respectively. Electron affinity 3.448 eV (the largest among atoms of all elements). According to the Pauling scale, the electronegativity of fluorine is 4 (the highest value among all elements). Fluorine is the most active non-metal.
In its free form, fluorine is a colorless gas with a pungent, suffocating odor.
Discovery history
The history of the discovery of fluorine is associated with the mineral fluorite (cm. FLUORITE), or fluorspar. The composition of this mineral, as now known, corresponds to the formula CaF 2 , and it is the first substance containing fluorine that began to be used by man. In ancient times, it was noted that if fluorite is added to ore during metal smelting, then the melting temperature of ore and slag decreases, which greatly facilitates the process (hence the name of the mineral - from Latin fluo - flow).
In 1771, by treating fluorite with sulfuric acid, the Swedish chemist K. Scheele (cm. SCHEELE Karl Wilhelm) prepared acid, which he called hydrofluoric acid. French scientist A. Lavoisier (cm. Lavoisier Antoine Laurent) suggested that this acid included a new chemical element, which he proposed to call "fluorine" (Lavoisier believed that hydrofluoric acid is a compound of fluorium with oxygen, because, according to Lavoisier, all acids must contain oxygen). However, he could not select a new element.
The new element was given the name "fluor", which is also reflected in its Latin name. But long-term attempts to isolate this element in a free form were not successful. Many scientists who tried to get it in a free form died during such experiments or became disabled. These are the English chemists brothers T. and G. Knox, and the French J.-L. Gay Lussac (cm. GAY LUSSAC Joseph Louis) and L. J. Tenard (cm. TENAR Louis Jacques), and many others. Sam G. Davy (cm. DEVI Humphrey), who was the first to receive sodium, potassium, calcium and other elements in a free form, as a result of experiments on the production of fluorine by electrolysis, he was poisoned and became seriously ill. Probably, under the impression of all these failures, in 1816, a name similar in sound, but completely different in meaning, was proposed for the new element - fluorine (from the Greek phtoros - destruction, death). This name of the element is accepted only in Russian, the French and Germans continue to call fluorine “fluor”, the British - “fluorine”.
Even such an outstanding scientist as M. Faraday could not obtain free fluorine (cm. FARADEUS Michael). Only in 1886 the French chemist A. Moissan (cm. Moissan Henri), using the electrolysis of liquid hydrogen fluoride HF, cooled to a temperature of -23 ° C (the liquid should contain a little potassium fluoride KF, which ensures its electrical conductivity), was able to obtain the first portion of a new, extremely reactive gas at the anode. In the first experiments, Moissan used a very expensive electrolyzer made of platinum and iridium to obtain fluorine. At the same time, each gram of the resulting fluorine "ate" up to 6 g of platinum. Later, Moissan began to use a much cheaper copper electrolyzer. Fluorine reacts with copper, but during the reaction a very thin film of fluoride is formed, which prevents further destruction of the metal.
Being in nature
The content of fluorine in the earth's crust is quite high and amounts to 0.095% by weight (significantly more than the closest analogue of fluorine in the group - chlorine (cm. CHLORINE)). Due to the high chemical activity of fluorine in the free form, of course, is not found. The most important fluorine minerals are fluorite (fluorspar), as well as fluorapatite 3Ca 3 (PO 4) 2 CaF 2 and cryolite (cm. CRYOLITE) Na 3 AlF 6 . Fluorine as an impurity is part of many minerals and is found in groundwater; in sea water 1.3 10 -4% fluorine.
Receipt
At the first stage of obtaining fluorine, hydrogen fluoride HF is isolated. Preparation of hydrogen fluoride and hydrofluoric acid (cm. HYDROFLUORIC ACID)(hydrofluoric) acid occurs, as a rule, along with the processing of fluorapatite into phosphate fertilizers. The gaseous hydrogen fluoride formed during the sulfuric acid treatment of fluorapatite is then collected, liquefied and used for electrolysis. Electrolysis can be subjected to both a liquid mixture of HF and KF (the process is carried out at a temperature of 15-20°C), and a KH 2 F 3 melt (at a temperature of 70-120°C) or a KHF 2 melt (at a temperature of 245-310°C) .
In the laboratory, to prepare small amounts of free fluorine, one can use either heating MnF 4, during which fluorine is eliminated, or heating a mixture of K 2 MnF 6 and SbF 5:
2K 2 MnF 6 + 4SbF 5 = 4KSbF 6 + 2MnF 3 + F 2 .
Physical and chemical properties
Under normal conditions, fluorine is a gas (density 1.693 kg / m 3) with a pungent odor. Boiling point -188.14°C, melting point -219.62°C. In the solid state, it forms two modifications: the a-form, which exists from the melting point to –227.60°C, and the b-form, which is stable at temperatures lower than –227.60°C.
Like other halogens, fluorine exists as diatomic molecules F 2 . The internuclear distance in the molecule is 0.14165 nm. The F 2 molecule is characterized by an anomalously low energy of dissociation into atoms (158 kJ/mol), which, in particular, determines the high reactivity of fluorine.
The chemical activity of fluorine is extremely high. Of all the elements with fluorine, only three light inert gases do not form fluorides - helium, neon and argon. In all compounds, fluorine exhibits only one oxidation state -1.
Fluorine reacts directly with many simple and complex substances. So, upon contact with water, fluorine reacts with it (it is often said that “water burns in fluorine”):
2F 2 + 2H 2 O \u003d 4HF + O 2.
Fluorine reacts explosively on simple contact with hydrogen:
H 2 + F 2 \u003d 2HF.
In this case, hydrogen fluoride gas HF is formed, which is unlimitedly soluble in water with the formation of a relatively weak hydrofluoric acid.
Fluorine interacts with most non-metals. So, in the reaction of fluorine with graphite, compounds of the general formula CF x are formed, in the reaction of fluorine with silicon, SiF 4 fluoride, and with boron, BF 3 trifluoride. When fluorine interacts with sulfur, compounds SF 6 and SF 4 are formed, etc. (see Fluorides (cm. FLUORIDE)).
A large number of fluorine compounds with other halogens are known, for example, BrF 3, IF 7, ClF, ClF 3 and others, moreover, bromine and iodine ignite in a fluorine atmosphere at ordinary temperature, and chlorine interacts with fluorine when heated to 200-250 ° C.
Do not react directly with fluorine, in addition to the indicated inert gases, also nitrogen, oxygen, diamond, carbon dioxide and carbon monoxide.
Nitrogen trifluoride NF 3 and oxygen fluorides О 2 F 2 and OF 2 were obtained indirectly, in which oxygen has unusual oxidation states +1 and +2.
When fluorine interacts with hydrocarbons, their destruction occurs, accompanied by the production of fluorocarbons of various compositions.
With slight heating (100-250°C) fluorine reacts with silver, vanadium, rhenium and osmium. With gold, titanium, niobium, chromium and some other metals, the reaction involving fluorine begins to proceed at temperatures above 300-350°C. With those metals whose fluorides are nonvolatile (aluminum, iron, copper, etc.), fluorine reacts with a noticeable rate at temperatures above 400-500°C.
Some higher metal fluorides, such as uranium hexafluoride UF 6 , are obtained by acting with fluorine or a fluorinating agent such as BrF 3 on lower halides, for example:
UF 4 + F 2 = UF 6
It should be noted that not only medium fluorides of the NaF or CaF 2 type, but also acidic fluorides - hydrofluorides of the NaHF 2 and KHF 2 types, correspond to the already mentioned hydrofluoric acid HF.
A large number of different organofluorine compounds have also been synthesized. (cm. organofluorine compounds), including the famous Teflon (cm. TEFLON)- material, which is a polymer of tetrafluoroethylene (cm. TETRAFLUOROETHYLENE) .
Application
Fluorine is widely used as a fluorinating agent in the production of various fluorides (SF 6 , BF 3 , WF 6 and others), including compounds of inert gases (cm. NOBLE GASES) xenon and krypton (see Fluorination (cm. FLUORINATION)). Uranium hexafluoride UF 6 is used to separate uranium isotopes. Fluorine is used in the production of Teflon and other fluoroplastics. (cm. Fluoroplastics), fluororubber (cm. fluororubbers), fluorine-containing organic substances and materials that are widely used in engineering, especially in cases where resistance to aggressive media, high temperatures, etc. is required.
Biological role
As a trace element (cm. MICROELEMENTS) Fluoride is found in all organisms. In animals and humans, fluorine is present in bone tissue (in humans, 0.2–1.2%) and, especially, in dentin and tooth enamel. The body of an average person (body weight 70 kg) contains 2.6 g of fluorine; the daily requirement is 2-3 mg and is met mainly with drinking water. A lack of fluoride leads to dental caries. Therefore, fluorine compounds are added to toothpastes, sometimes introduced into drinking water. Excess fluoride in water, however, is also harmful to health. It leads to fluorosis (cm. FLUOROSIS)- changes in the structure of enamel and bone tissue, bone deformation. MPC for the content of fluoride ions in water is 0.7 mg/l. Maximum concentration limit for gaseous fluorine in the air is 0.03 mg/m 3 . The role of fluorine in plants is unclear.

encyclopedic Dictionary. 2009 .

See what "fluorine" is in other dictionaries:

fluorine- fluorine, and ... Russian spelling dictionary

fluorine- fluorine/… Morphemic spelling dictionary

- (lat. Fluorum) F, a chemical element of group VII of the periodic system of Mendeleev, atomic number 9, atomic mass 18.998403, belongs to halogens. Pale yellow gas with a pungent odor, mp? 219.699 .C, tbp? 188.200 .C, density 1.70 g / cm & sup3. ... ... Big Encyclopedic Dictionary

F (from Greek phthoros death, destruction, lat. Fluorum * a. fluorine; n. Fluor; f. fluor; and. fluor), chem. element of group VII periodic. system of Mendeleev, refers to halogens, at. n. 9, at. m. 18.998403. In nature, 1 stable isotope 19F ... Geological Encyclopedia

- (Fluorum), F, chemical element of group VII of the periodic system, atomic number 9, atomic mass 18.9984; refers to halogens; gas, boiling point 188.2shC. Fluorine is used in the production of uranium, freons, medicines and others, as well as in ... ... Modern Encyclopedia