Covalent compounds fluorine

Molecules are the fundamental units of the gaseous covalent compound fluorine, F2. Notice that in this model of a fluorine molecule, the spheres overlap, whereas the spheres shown earlier for ionic compounds do not. Now you know that this difference in representation is because of the difference in bond types. 

The element before carbon in Period 2, boron, has one electron less than carbon, and forms many covalent compounds of type BX3 where X is a monovalent atom or group. In these, the boron uses three sp hybrid orbitals to form three trigonal planar bonds, like carbon in ethene, but the unhybridised 2p orbital is vacant, i.e. it contains no electrons. In the nitrogen atom (one more electron than carbon) one orbital must contain two electrons—the lone pair hence sp hybridisation will give four tetrahedral orbitals, one containing this lone pair. Oxygen similarly hybridised will have two orbitals occupied by lone pairs, and fluorine, three. Hence the hydrides of the elements from carbon to fluorine have the structures... 

Many of the reactions of halogens can be considered as either oxidation or displacement reactions the redox potentials (Table 11.2) give a clear indication of their relative oxidising power in aqueous solution. Fluorine, chlorine and bromine have the ability to displace hydrogen from hydrocarbons, but in addition each halogen is able to displace other elements which are less electronegative than itself. Thus fluorine can displace all the other halogens from both ionic and covalent compounds, for example... 

Covalent compounds, arising from the attack of strong oxidizing systems, such as fluorine or Mn(VII), on graphite. The aromatic planarity of the graphite sheet is destroyed, and a buckled, sp -hybridized sheet is created. 

P. Venkateswarlu, Sodium biphenyl method for determination of covalently bound fluorine in organic compounds and biological methods. Anal. Chem. 54 (1982) 1132-1137. 

Lewis and many other chemists had recognized the shortcomings of the ionic bond. When diatomic molecules, such as or Cl, were considered, there was no reason why one atom should lose an electron and an identical atom should gain an electron. There had to be another explanation for how diatomic molecules formed. We have seen how the octet rule applies to the formation of ionic compounds by the transfer of electrons. This rule also helps explain the formation of covalent bonds when molecules (covalent compounds) form. Covalent bonds result when atoms share electrons. Using fluorine, F, as a representative halogen, we can see how the octet rule applies to the formation of the molecule. Each fluorine atom has seven valence electrons and needs one more electron to achieve the stable octet valence configuration. If two fluorines share a pair of electrons, then the stable octet configuration is achieved ... 

A substance composed of atoms held together by covalent bonds is a covalent compound. The fundamental unit of most covalent compounds is a molecule, which we can now formally define as any group of atoms held together by covalent bonds. Figure 6.13 uses the element fluorine to illustrate this principle. 

When writing electron-dot structures for covalent compounds, chemists often use a straight line to represent the two electrons involved in a covalent bond. In some representations, the nonbonding electron pairs are left out. This is done in instances where these electrons play no significant role in the process being illustrated. Here are two frequently used ways of showing the electron-dot structure for a fluorine molecule without using spheres to represent the atoms ... 

The highest electron affinities are found at the top right of the periodic table (Fig. 14.4 see also Fig. 1.47). The electron affinity of an element is a measure of the energy released when an ion is formed. Except for the noble gases, elements near fluorine have the highest electron affinities, so we can expect them to be present as anions in compounds with metallic elements. For the same reason, these elements commonly have negative oxidation states in the covalent compounds they form with other nonmetallic elements. 

This section has demonstrated the strikingly unique nature of fluorine, both as an atom and in ionic and covalent compounds. With this in mind, we move on to a discussion of electronegativity equalization and its impact upon molecular properties. 

Elementary fluorine or chlorine may act drastically on a number of covalent compounds, forming products in which the only bonds remaining are eiement-to-halogen bonds. Such reactions are called halogenolyses the chlorinolysis of carbon diselenide and the fluorinolysis of ethane are typical ... 

The covalent compound hydrogen fluoride, for example, would be represented by the symbol H joined to the symbol / by a single line, with three pairs (six more dots) surrounding the symbol F. The line represents the two electrons shared by both hydrogen and fluorine, whereas the six paired dots represent fluorine s remaining six valence electrons. 

What is the formula of the covalent compound formed between one nitrogen atom and as much fluorine as needed to form the compound ... 

In Group 7A(17), fluorine and chlorine have the condensed electron configuration [noble gas] ns np, as do the other halogens (Br, I, At). Little is known about rare, radioactive astatine (At), but all the others are reactive nonmetals that occur as diatomic molecules, X2 (where X represents the halogen). All form ionic compounds with metals (KX, MgX2), covalent compounds with hydrogen (HX) that yield acidic solutions in water, and covalent compounds with carbon (CX4). 

Because fluorine is so electronegative, it is always assumed to control any shared electrons. So fluorine is always assumed to have a complete octet of electrons and is assigned an oxidation state of -1. That is, for purposes of assigning oxidation states, fluorine is always imagined to be F in its covalent compounds. 

The halogens, particularly fluorine, have very high electronegativity values (see Table 20.17). They tend to form polar covalent bonds with other nonmetals and ionic bonds with metals in their lower oxidation states. When a metal ion is in a higher oxidation state, such as +3 or +4, the metal-halogen bonds are polar and covalent. For example, TiCl4 and SnCl4 are both covalent compounds that are liquids under normal conditions. 

Fluorine reacts with sulfur to form several different covalent compounds. Three of these compounds are SF2, SF4, and SF. Draw the Lewis structures for these compounds, and predict the molecular structures (including bond angles). Would you expect OF4 to be a stable compound ... 

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