-5.1-CHAPTER 5. INTRODUCTION TO CHEMICAL BONDINGA. Types of bondingTwo simplified pictures are commonly employed to discuss chemical bonding - ionicbonding and covalent bonding.Ionic Bonding - bonding due to the electrostatic attraction of oppositely charged ions formed asa result of the transfer of one or more electrons from one atom to another. As an example of ionicbonding, consider the NaCl molecule: The electron configurations of Na and Cl atoms areCl 1s 2s 2p 3s 3p2 2 6 2 5Na 1s 2s 2p 3s .2 2 6 1The Cl atom needs one more electron to attain the stable outer octet structure of neon, while the Naatom has one additional electron beyond the inner octet formed by its L-shell. Thus, it is reasonableto expect that the transfer of the 3s electron of Na into the 3p orbital of Cl might result in a stableelectrostatic bond between the two resulting ions.Covalent Bonding - bonding that results from the sharing of electrons between two atoms. It isclear that one cannot evoke an ionic bonding model for homonuclear diatomic molecules like Cl2because it is unreasonable to expect that one Cl atom will have more of an attraction for an extraelectron than the other. In this case, as well as many others which involve heteronuclear atomshaving similar electronegativities, the bonding is considered to be associated with the sharing ofelectrons. It will be shown later that this sharing leads to a build-up of electron density between thetwo nuclei, thereby reducing the internuclear repulsion.In Cl , each atom may attain a stable outer-shell octet by sharing a pair of electrons. This is2illustrated by the so called Lewis dot structure shown below orin which all the valence electrons are arranged in pairs around the atoms so that each atom issurrounded by eight electrons. It is important to note that the simple concepts of ionic and-5.2-covalent bonding, although extremely useful for elementary discussions of bonding, are onlyqualitatively applicable to real bonding situations. In actual practice, most bonds must be treated asa combination of both types and quantum mechanical methods must be used to extract quantitativeinformation.B. Lewis Dot StructuresElectron dot structures were introduced by C. N. Lewis in 1916 as a simple way ofillustrating the following principle:Simple compounds form in such a way that the constituent atoms either lose, gain, or shareenough electrons to attain the valence configuration of the nearest noble gas.This principle is known as the octet rule. Although there are many exceptions to the octet rule,electron dot structures are still useful in understanding many aspects of molecular bonding.A procedure for writing electron dot structures is outlined as follows.1. Add up the number of valence electrons for each atom in the compound to get the totalnumber of electrons to be accommodated. Remember, the valence electrons are all theelectrons in the outermost shell. The element oxygen (ls 2s 2p ) has 6 valence electrons.2 2 42. Arrange the other atoms around the central atom.3. Place a shared pair of electrons between each atom and then add enough unshared pairs toachieve an octet around each atom. Unshared pairs of electrons are called lone pairs.4. Check to see that the total number of electrons corresponds to total number of valenceelectrons determined in step 1. If it doesn’t, try making one or more of the bonds into doubleor triple bonds. The bond order is equal to the total number of shared electron pairs. Thusa triple bond has a bond order of 3 and a double bond has a bond order of 2.Examples:CH C=4 4H = 4 x 1 = 4Total = 8-5.3-CO C = 432!O = 3 x 6 = 18Charge = 2!Total = 24Exercise: Determine the Lewis dot structure of CN!CN C = 4!N = 5Charge = 1!Total = 10A useful rule of thumb for determining how many electrons must be shared isshared = needed !!!! available,or S = N !!!! A.Applying this rule to CO ;32!S = 4 x 8 ! (3 x 6 + 4 + 2) = 32 ! 24 = 8.This number of electrons shared between four atoms requires one double bond.Exercise: Apply the above rule to the case of CN .!S = 2 x 8 ! (4 + 5 + 1) = 16 ! 10 = 6.This number of electrons shared between two atoms requires a triple bond.Exercises: Draw electron dot structures for each of the following:CH Br CH O N O (O not in middle)3 2 2FCN SF S O (O not in middle)3 2+CH CO HOCl C1 O3 2 2!BO SO F XeO3 3 33! !-5.4-C. Formal ChargeSometimes two or more electron dot structures are possible and it is necessary to decidewhich structure best represents the true bonding situation. The concept of formal charge isfrequently helpful in making such decisions. The formal charge of an atom in a compound iscalculated as follows:1. Count the number of valence electrons belonging to the atom of interest by assigning it oneelectron of each shared pair plus all of its lone pair electrons.2. The formal charge is the difference between the number of valence electrons in the free atomand the number of electrons determined in step 1.Example: Determine the formal charges of C and O in the CO molecule.Dot structure:C electrons = 2 + 1 + 1 + 1 = 5 so the formal charge of C is 4 ! 5 = !1.O electrons = 2 + 1 + 1 + 1 = 5 so the formal charge of O is 6 ! 5 = +1.In general, if there are several possible dot structures for a compound, the one which givesthe lowest formal charges is the best choice.Exercise: Determine the best dot structure for the CN molecule.In this case, it is not possible to satisfy the octet rule. Two possible structures are: orSince the one on the left has no formal charge, while the one on the right has formal chargesof -1 and +1 for C and N, respectively, the best choice is the structure on the left. Note: Thismolecule is an example of a radical - a molecule with an unpaired electron.-5.5-D. ResonanceThere are many molecules for which a single electron dot structure does not adequatelyrepresent the true bonding situation. Consider, for example, the case of SO . A dot structure that3satisfies the octet rule isThis structure predicts that SO should have one double bond and two single bonds. Measurements,3 however, show that all the S-O bonds have exactly the same strength and length, and that they areintermediate between double and single bonds. A better way of representing the true bondingstructure, therefore, is to draw the three so called resonance structures shown below;The true structure of SO may be considered to be a
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