Drawing Lewis Structures: A Step-by-step Approach Wan-Yaacob Ahmad and Siraj Ornar Universiti Kebangsaan Malaysia, Sabah Campus, Locked Bag No. 62, 88996 Kota Kinabalu, Sabah, Malaysia The Lewis theory of chemical bond formation through the sharine of electron ~ a i r ( shas ) remained an i m ~ o r t a n t topic in inkoductory cLemistry cumiculum. ~ l t h o i the ~h Lewis structure seems a s a rather simple presentation, undoubtedly it is very useful as a means of predicting molecular shape and geometry when combined with the Valence Shell Electron Pair Repulsion Theory (VSEPR).The drawing of Lewis structures is a straightforward task to experienced chemists who understand and recognize structural patterns in molecules. This is not so for many students who have not mastered the skill, where writing even a simple Lewis structure could be as troublesome as predicting its geometry. In addition, the topic dealing with Lewis theory is not clearly presented in many general chemistry textbooks, and too much emphasis is often given to the descriptive part of the octet rule. No textbook has yet presented a simole mide for derivine Lewis structures. As a result many stu&nts resort to the'trial and error' to solve Lewis structures. Some of them even have difficultv choosing the central atom from which to start. A number of articles dealing with Lewis structures have appeared in literature (14). This paper proposes a simple step-by-step approach for deriving Lewis structures for students studying Introductory Chemistry. Students are required to know only the location of elements in the Periodic Table and their trend in electronegativity. This approach of drawing Lewis structures is based on the knowledge of octet rule and the minimization of the formal charge distribution. Formal Charge Formal charge is defmed as the residual charge on an i (or a partial structure)Gbtained atom in a ~ e w istructure bv subtract in^ the number of valence electrons of that atom in its isolated neutral state with the total number of valence electrons assigned to that atom in the structure. The latter can he determined as follows: .all -bared
(nonbonding)valence electrons surrounding a particular atom are assigned to that atom; onlv one-half of the total number of bonding - electrons engaged by a panleular atom, whleh 1s smular to the number ufelectron-pambonds, are assl~medtu that atom
For re~resentativeelements in the Denodic table. the number i f valence electrons belonging tb a neutral atom is eaual to its eroun number. The formal charee a t an atom in'the stru&re Ls calculated as follows:
-
(Formal charge) = (Group No.) - (No. of unshared eledmns + No. of electron-pair bonds) (1)
ALewis structure where no formal charge (zero residual charge) is given to any atom is preferable. If this is not possible, then a formal charge of +I andlor -1 may he assigned to the fewest number of atoms possible. Note that the net formal charge should be zero for neutral molecules or the same as the charee on ionic soecies. The molecular structure can be consildered stabie if negative formal charees are assiened to atoms with hieher electroneeativity, &d positive ?omal charges are as&ned to atomswith lower eleetronegativity.
A Checklist for Drawing Lewis Structures The Lewis structure is constructed in the following steps. (1) Chwse a central atom from the molecular formula. As a general rule, this would be the atom that is the least in electronegativity and its number. The monovalent atom, H, and the most electronegative atom, F,always occupy the peripheral position in Lewis structures. For small c e
valent inorganic species, a close cyclic structure may be omitted due to high angle strain. ( 2 ) Construct a sigma skeletal structure by drawing a single bond connectine the central atom to each terminal atom. (3) Three unshared (nonhonding) electron pairs are then added to complete the actet an each terminal atom (except HI. All valence electrons should be accounted for. and if not, the remaining pairk) are placed on the central atom to give a skeletal Lewis structure. (4) Ry to complete the octet on the central atom (if it has not been achieved) by moving to it one or more unshared electron uairs from the multivalent terminal atomfs) . . to farm pi b&d(s). The other peripheral atoms are normally comprised of the secand-mostelectronegative element, 0,and N atom. ( 5 ) Calculate the formal charge on each bonded atom in the sketched Lewis structure according to eq 1 discussed above. As far as possible zero (orminimum)formal charge should be achieved by converting one or more unshared electron p a i n into bandingpairs. The conversion of a lone pair of electrons on the terminal atom into a pi bond to the central atom results in the change of the formal charges one unit more positive for the former and one unit more negative for the latter as the respective electron ownership decreases and increases by one electron. More than eight electrons (bonding and nonbonding)may appear on the central atom made up of third period elements (such as P and S)or higher. This state of transactet or expanded octet is possible due to the presence of empty nd orbitals in those atoms. Deriving Lewis Structures The following discussion on Lewis structure is based on several useful examples chosen to represent a wide spectrum of covalent molecules a s well-as polyatomic ions. Bonding and nonbonding electron pairs are shown as short dashed lines. Let us consider the nitrous oxide molecule, NzO. First, we determine the central atom according to step (1).Because N is less electronegative than 0 , one of the N atoms can serve as the central atom. Using a total of 16 valence shell electrons (eieht airs) in N90. we follow t h r o u ~ h steps (2) and (3) toobtain the skelitai structure 1.At tgis stage, the central atom has only four surrounding electrons. Then, we proceed to step (4) in order to complete the octet on the central atom. This can be done in three different ways: 1. two unshared electron nairs from the terminal N are moved to form two bondkg pairs with the central atom to give the structure 2; 2. one unshared electron pair each from the terminal N and 0 are moved to form two double bonds as in structure 3; 3. two nonbonding electron pairs from the terminal 0 form pi bonds with the central atom and yield structure 4. Volume 69 Number 10 October 1992
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To check for the preferred Lewis structure, we calculate the formal charges on each atom in all three proposed structures according to eq 1. Obviously, structure 2 is the preferred structure, because it has the minimum and the best distributed formal charges where a residual +1 is placed on the less electronegative central N while a -1 charge is placed on the more electronegative terminal 0. In structure 3, a formal charge of -1 is placed on the terminal N instead of on the 0. This structure can be considered as a resonance form to N20 besides 2. The third resonance structure 4 can be discarded. The formal charge distribution is high and will make the molecule unstable. The existence of a structural hybrid between 2 and 3 is supported by spectroscopic data. The nitrogen-nitrogen bond length is between a double and a triple bond, while the nitrogenoxygen is between a single and a double bond. An ozone molecule ( 0 3 ) with its resonance structures 5 and 6 i s among several other species where a n octet state is achieved by all atoms in the Lewis structures.
The method presented can be applied to numerous examples, and these include polyatomic species or complex ions with expanded, or reduced octets, and odd numbers of valence electrons in which the central atoms containing the representative, inert, or transition metal atoms. The Lewis structures for specific examples of each of these categories are eiven in structures 7-15, or 16, 17, and 18/19. 1;should b e k t e d that 18 and 19 are resonance structures for the odd molecule KO,. Thc central atoms in structures 7,8, and 11are surrounaed by 10 electrons; 9, 10, 12, 13, and 15 by 12 electrons; 14 by 14 electrons; whereas, in structures 16, 17 and 18/19, there are four, six and seven valence electrons, respectively, around the central atoms. Conclusions The derivation of the Lewis structure based on formal charge distributions can be applied to all simple inorganic species. No complicated mathematical formulation needs to be memorized. Students are required only to know the position of each element involved in the Periodic Table, and its relative electronegativity and then decide on the central atom accordingto step (1).Once this is achieved, its sigma skeleton can be drawn to tie all terminal atoms followed by the completion of their octets (except for H) according to steps (2) and (3). To complete the octet on the central atom (ifit has not been achieved),one only needs to carry out step (4). Finally, the process of formal charge minimization is carried out according to step (5). In most cases, steps (4) and (5)are carried out simultaneously. The octet rule is normally obeyed by elements of the second pe792
Journal of Chemical Education
riod, with a few exceptions with Be and B, while elements of the higher periods in the Periodic Table may attain the state of expanded octet as shown in many of the examples given above. Finally, from the numher of valence shell electron pairs on the central atom (bonding and nonbonding) and knowledge of the Valence Shell Electron Pair Repulsion Theory (VSEPR),the geometrical shape of the molecules and polyatomic ions can be oredicted (9.10).For examole. both theory and cxperimenial data ha& demonstrilGd (hat ozone (5.6,. sulfur dioxide 17, and nitroeen dioxide I 18. 19, have %-shape3.Other examples, like &trons oxide (2; 3,4) and beryllium chloride (16)are linear molecules; chromium trioxide (15) and boron trifluoride (17) have a trigonal shape; sulfur tetrafluoride (8)is an irremlar tetrahedron due to lone pair electrons on S; permanganate ion (14) and phos~ h o r u soxvchloride (11)are tetrahedron: xenon tetrafluoride (12) is a square planar; aluminium. hexafluoride ion (10) and sulfur hexafluoride (9) are octahedron, and finally, xenon oxotetrafluoride (13) has a square pyramidal shape with the 0 atom at the apex.
a
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2. Clark T J. J Cham. Educ 1984,61,100. 3. Zandler. M. E.:Talatv E. R.J. Cham.Educ. 19%. 61.124-127. 4. Carrol1,J.A.J.~ h a m E d u e1986.63, . 2W0. 5. DeKock,R. L. J. C h m Educ 1987.64.934-941. 6 . Snsdden, R. B.Edue. Cham. 1997,24,81-83. 7 . Maiench, C. J J ChamEduc 1887.64.403. 8. Parda. J. 0. J Chrm Educ. 1989.66.456458. 9. Gillespie, R. J. J Cham. Educ. 1970,47. 18-23. 10. Gillespie, R. J. Molecular Gromrfr).;Van Nostrand Reinhold Co.: Landon. 1972.
