Lewis structures fluorine

Representing a two electron covalent bond by a dash (—) the Lewis structures for hydrogen fluoride fluorine methane and carbon tetrafluoride become... 

The fluoride SF4 forms when a mixture of fluorine and nitrogen gases is passed over a film of sulfur at 275°C in the absence of oxygen and moisture. Write the Lewis structure of sulfur retrafluoride and give the number of electrons in the expanded valence shell. 

These three frameworks and the framework for glycine in Figure 9 illustrate an important point about Lewis structures. Although Lewis structures show how atoms are connected to one another, a Lewis structure is not intended to show the actual shape of a molecule. Silicon tetrachloride is not flat and square, SO2 is not linear, and the fluorine atoms in CIF3 are not all equivalent. We describe how to use Lewis structures to determine the shapes of molecules later in this chapter. 

The difference between equatorial and axial positions determines the arrangement of bonding pairs and lone pairs around an atom with a steric number of 5. An example is provided by sulfur tetrafluoride, a colorless gas that has industrial uses as a potent fluorinating agent. The Lewis structure of SFq shows four S—F bonds and one lone pair of electrons on the sulfur atom. These five pairs of electrons are distributed in a trigonal bipyramid around the sulfur atom. 

Stable noble gas compounds are restricted to those of xenon. Most of these compounds involve bonds between xenon and the most electronegative elements, fluorine and oxygen. More exotic compounds containing Xe—S, Xe—H, and Xe—C bonds can be formed under carefully controlled conditions, for example in solid matrices at liquid nitrogen temperature. The three Lewis structures below are examples of these compounds in which the xenon atom has a steric munber of 5 and trigonal bipyramidal electron group geometry. 

C09-0083. Fluorine forms compounds whose chemical formula is XF4 with elements from groups 14, 16, and 18. Determine the Lewis structure, describe the shape, and draw a ball-and-stick model of Gep4, SeF4, and Xep4. 

C09-0132. Sulfur and fluorine form seven different molecules SF2, SSF2, FSSF, F3 SSF, SF4, F5 SSF5, and SFg. Draw the Lewis structure of each molecule, and identify the geometry around each inner sulfur atom. 

The Lewis structure of hydrogen fluoride shows three lone pairs on the fluorine atom. These nonbonding electrons are localized in atomic orbitals that belong solely to fluorine. Remembering that one of the fluorine 2 p orbitals is used to form the H—F bond, we conclude that the three lone pairs must occupy the remaining pair 2 p orbitals and the 2 s orbital of the fluorine atom. 

A Following the strategy outlined in the textbook, we begin by drawing a plausible Lewis structure for the cation in question. In this case, the Lewis structure must contain 20 valence electrons. The skeletal structure for the cation has a chlorine atom, the least electronegative element present, in the central position. Next we join the terminal chlorine and fluorine atoms to the central chlorine atom via single covalent bonds and then complete the octets for all three atoms. 

NBOs 3-5 are the three fluorine lone pairs ( LP ). As shown by the occupancies and hybrid composition, these lone pairs are inequivalent. LP(1) is the s-rich sigma-type sp0 26 lone pair (nF(cr) 79% s character), directed along the bond axis. LP(2) andLP(3) are the p-rich pi-type lone pairs (nF(7t) and nF(rf) 99.97% p-character), perpendicular to the bond axis. The lone pairs have occupancies slightly less than 2.000 00 (due to weak delocalization into Rydberg orbitals of the adjacent H), but overall, the correspondence with the elementary Lewis-structure description is excellent. 

Now we will begin drawing the Lewis structure of XeF4. Xenon will be the central atom, and we will arrange fluorine atoms around it. In this way, we avoid attaching identical atoms to each other. We will need a bond between the central xenon and each of the fluorine atoms. This arrangement means there will be at least four bonds. 

Even though we have an exception, we can still complete the Lewis structure. We need to draw a bond from each of the fluorine atoms to the central xenon. This gives us 4 bonds and uses 8 electrons. Each fluorine atom needs to complete its octet. The bond accounts for 2 electrons, so we need 6 more electrons (3 pairs) for each. Therefore, we add 3 separate pairs to each of the fluorine atoms. Six electrons per fluorine times 4 fluorine atoms accounts for 24 electrons. Our Lewis structure now contains 8 + 24 = 32 electrons. The number of available electrons (A) is 36, so we still need to add 36 - 32 = 4 electrons. These 4 electrons will give us 2 pairs. The xenon atom will get these pairs and become an exception to the octet rule. The actual placement of the pairs is not important as long as it is obvious that they are with the central atoms and not one of the fluorine atoms. The final Lewis structure is ... 

Individual atoms of hydrogen and fluorine are highly reactive, and readily bond together to form molecules of hydrogen fluoride. Draw a Lewis structure for hydrogen fluoride. Label the bonding and lone pairs, and explain why this molecule is stable. 

Lewis structure shows the connectivity between atoms in a molecule by a number of dots equal to the number of electrons in the outer shell of an atom of that molecule. A pair of electrons is represented by two dots, or a dash. When drawing Lewis structures, it is essential to keep track of the number of electrons available to form bonds and the location of the electrons. The number of valence electrons of an atom can be obtained from the periodic table because it is equal to the group number of the atom. Eor example, hydrogen (H) in Group lA has one valence electron, carbon (C) in Group 4A has four valence electrons, and fluorine (E) in Group 7A has seven valence electrons. 

STRATEGY A fluorine atom forms only single bonds, so we anticipate that the Lewis structure consists of a shared pair between the central S atom and each of the four surrounding F atoms. However, each F atom has three lone pairs and supplies one bonding electron, and the S atom already has six elec-5 trons in its valence shell. So, there are two extra electrons. Because sulfur is in Period 3 and has empty 3d-orbitals available, it can expand its octet. 

A boron trifluoride molecule, BF3, has the Lewis structure shown in (5). There are three bonding pairs attached to the central atom and no lone pairs. According to the VSEPR model, as illustrated in Fig. 3.4, the three bonding pairs, and the fluorine atoms they link, lie at the corners of an equilateral triangle. Such a structure is trigonal planar, and all three F—B—F angles are 120° (6). 

These Lewis structures show the formation of bonds between one atom of calcium and two atoms of fluorine. 

Note that in each of these resonance structures (and the many other similar ones that can be drawn) the sulfur atom always has an octet of electrons around it. The true structure is a composite of these equivalent resonance structures. The overall bonding is described as a combination of covalent and ionic contributions. The covalent contribution to the bonding involves four electron pairs spread out over the six sulfur-fluorine bonds. The ionic contribution to the bonding arises as follows. Note that for each Lewis structure two of the fluorines have -1 formal charges and the sulfur has a +2 formal charge,... 

C1F3 and BrF3 are both used to fluorinate uranium to produce UF6 in the processing and reprocessing of nuclear fuel. Draw Lewis structures for C1F3 and BrF3. 

The relatively large change in size in going from the first to the second member of a group also has important consequences for the Group 7A elements. For example, fluorine has a smaller electron affinity than chlorine. This violation of the expected trend can be attributed to the fact that the small size of the fluorine 2p orbitals causes unusually large electron-electron repulsions. The relative weakness of the bond in the F2 molecule can be explained in terms of the repulsions among the lone pairs, shown in the Lewis structure ... 

To draw a Lewis structure for a diatomic molecule like HF, recall that hydrogen has one valence electron and fluorine has seven. H and F each donate one electron to form a two-electron bond. The resulting molecule gives both H and F a filled valence shell. 

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