Chlorine Lewis structure

Sei f-Test 2.4A Write the Lewis structure for the interhalogen compound chlorine... 

The Lewis structure of the PCI, molecule is shown in (22), and we see that it obeys the octet rule. However, when phosphorus trichloride reacts with more chlorine (Fig. 2.10), phosphorus pentachloride, a pale yellow crystalline solid, is produced ... 

Write the Lewis structure of each of the following reactive species, all of which are found to contribute to the destruction of the ozone layer, and indicate which are radicals (a) chlorine monoxide, CIO (b) dichloroperoxide, Cl—O—O—Cl ... 

In the gas phase, BeCl2 forms a dimer by forming chlorine-atom bridges like those in the AICI dimer. Draw the Lewis structure for the BeCl2 dimer and assign formal charges. 

The Lewis structure of chlorine trifluoride is treated in Example. Determine the molecular geometry and draw a three-dimensional picture of the molecule. 

Use the Lewis structure of CIF3 to determine the steric number of the chlorine atom. Obtain the molecular shape from the orbital geometry after placing lone pairs in appropriate positions. 

C09-0084. Iodine forms three compounds with chlorine ICl, ICI3, and ICI5. Determine the Lewis structure, describe the shape, and draw a ball-and-stick model of each compound. 

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. 

Once the Lewis diagram is complete, we can then use the VSEPR method to establish the geometry for the electron pairs on the central atom. The Lewis structure has two bonding electron pairs and two lone pairs of electrons around the central chlorine atom. These four pairs of electrons assume a tetrahedral geometry to minimize electron-electron repulsions. 

Structure (3.226c), for example, depicts a central heptavalent Cl atom (Fa = 7), exceeding the normal valence octet by six electrons (These excess electrons are assumed to be accommodated in chlorine 3d orbitals, whereas d-orbital participation is prevented in first-row compounds.) Hypervalent structures such as (3.226a)-(3.226c) are claimed to be justified by the electroneutrality principle, which stipulates that second-row central atoms have zero formal charge (whereas first-row oxyanion Lewis structures commonly violate this principle).148... 

A) The Lewis structure of a chlorine molecule (B) The formal counting of electrons in a chlorine molecule for oxidation number purposes... 

You have seen examples of how Lewis structures can be used to assign oxidation numbers for polar molecules such as water, non-polar molecules such as chlorine, and polar polyatomic ions such as the cyanide ion. 

During the long Antarctic night, appreciable amounts of molecular chlorine, Cl, and hypochlorous acid, HOCl, accumulate within the polar vortex. When the sun returns during the spring (in September in Antarctica), ultraviolet radiation decomposes the accumulated molecular chlorine and hypochlorous acid to produce atomic chlorine. Cl. Atomic chlorine is a free radical. Free radicals are atoms or molecules that contain an unpaired or free electron. The Lewis structures of free radicals contain an odd number of electrons. The unpaired electron in free radicals makes them very reactive. The free radical Cl produced from the decomposition of CI2 and HOCl catalyzes the destruction of ozone as represented by the reaction ... 

A highly toxic gas that has been used in chemical warfare gives the following elemental analysis figures 12.1% carbon, 16.2% oxygen, and 71.7% chlorine by mass. Its molar mass is 98.9 g-moU1. Write the Lewis structure of this compound. 

Sulfur and chlorine combine in various proportions forming S2CI2, SCI2, and SCI4. Draw the Lewis structures of these molecules and, using VSEPR, predict their shapes. 

Only CIO2 is known. Chlorine dioxide is an odd-electron molecule. Theoretical calculations suggest that the odd electron is delocalized throughout the molecule, and this probably accounts for the fact that there is no evidence of dimerization in solution, or even in the liquid or solid phase. Its important Lewis structures are shown below ... 

Yellowish, toxic chlorine gas (Cl2) is called a diatomic molecule. Looking at its Lewis structure, you can see why it needs to be diatomic in order to have octet stability. This is an example of a nonpolar covalent bond because the electronegativity difference is zero. 

Show Lewis structures for the simplest neutral compounds formed from these elements a) Carbon and chlorine b) Hydrogen and bromine... 

Phosphorus forms two compounds with chlorine, PC13 and PC15. The former follows the octet rule, but the latter does not. Show Lewis structures for each of these compounds. For the corresponding nitrogen compounds, explain why NC13 exists but NCI5 does not. 

The splitting of a Cl2 molecule is an initiation step that produces two highly reactive chlorine atoms. A chlorine atom is an example of a reactive intermediate, a short-lived species that is never present in high concentration because it reacts as quickly as it is formed. Each Cl- atom has an odd number of valence electrons (seven), one of which is unpaired. The unpaired electron is called the odd electron or the radical electron. Species with unpaired electrons are called radicals or free radicals. Radicals are electron-deficient because they lack an octet. The odd electron readily combines with an electron in another atom to complete an octet and form a bond. Figure 4-1 shows the Lewis structures of some free radicals. Radicals are often represented by a structure with a single dot representing the unpaired odd electron. 

When you draw Lewis structures to show the formation of a bond, you can use different colours or symbols to represent the electrons from different atoms. For example, use an "x" for an electron from sodium, and an "o" for an electron from chlorine. Or, use open and closed cirlces as is shown here. This will make it easier to see how the electrons have been transferred, v /... 

Therefore, instead of transferring electrons, the two atoms each share one electron with each other. In other words, each atom contributes one electron to a covalent bond. A covalent bond consists of a pair of shared electrons. Thus, each chlorine atom achieves a filled outer electron energy level, satisfying the octet rule. Examine Figure 3.15 to see how to represent a covalent bond with a Lewis structure. 

These Lewis structures show the formation of a bond between two atoms of chlorine. 

Describe, in detail, what happens when an ionic bond forms between calcium and chlorine. Use Lewis structures to illustrate your description. 

This diagram showing how the valence electrons interact is called a Lewis structure. In this case both hydrogen atoms have satisfied their need to have a full outermost principal energy level. Because both hydrogen atoms have the same electronegativity, the atoms will share the electrons equally. This will be the case with any diatomic molecule, such as chlorine gas (see Figure 5.6). 

The traditional Lewis structure for PC15 is shown in the margin. Five pairs of electrons around the phosphorus atom require a trigonal bipyramidal arrangement (see Table 13.8). When the chlorine atoms are included, a trigonal bipyramidal molecule results. 

The Lewis structure for PC15 shows that each chlorine atom is surrounded by four electron pairs. This requires a tetrahedral arrangement, which in turn requires a set of four sp3 orbitals on each chlorine atom. 

FC102 and F3C10 can both gain a fluoride ion to form stable anions. F3C10 and F3C102 can also lose a fluoride ion to form stable cations. Draw Lewis structures and describe the hybrid orbitals used by chlorine in these four ions. 

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 ... 

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