chemicd principle/ revi~i ted
Edited by Charles D. Mickey Texas A b M at Galveston Galveston, TX 77553
A Different Approach to Hybridization and Geometric Structure of Simple Molecules and Ions Diana Eberlln and Manus Monroe lndian Valley Colleges, lgnacio Boulevard, Novato, CA 94947 When studvine - " the c o n c e ~ tof orbital hvbridization in freshman and advanced secondary chemistry courses, many students become readily confused about orbital hybridization and geometric structures of simple molecules and ions. Simple molecules and ions (charged radicals) are those which contain from three to seven atoms, one of which is the central atom for the structure. In most textbooks, after several pages of discussion about structrlres and orbital hybridization, an extensive summary table is presented to help reduce student confusion. These summary tables often list the coordination numbers, the number of bonded pairs, and the number of lone pair electrons for molecules and ions without a specific and direct statement about hybridization and structure. Many students do not find these tables beneficial. We present here a successul step-by-step teaching tecbnioue. ,~~,with detailed comments for student learning. .,. which directly currelates hybridization with structure. b i r s t . we remind studt:nts that ill reference tu simole mulecules and ions that: ~~~~
(1) When evaluating a formula, the central atam is the single atom shown in the formula while the remainine atoms are lieands.
and SFBN. (2) The geometry is controlled by the number of single bonds and sets of lone pair electrons around the central atom. (3) Single bonds are formed from hybrid orbitals while doubleand triple bonds are formed from the overlap of p andlor d orbitals. Pauling, Linus, "General Chemistry," 3rd Edition, W. H. Freeman and Company, 1970, pp. 192-194. Pauling. Linus, and Pauling. Peter. "Chemistry," W. H. Freeman and Company, 1975, pp. 180-181. Durant. Philip and Durant, Beryl. "Introduction to Advanced Inorganic Chemistly." 2nd Edition, John Wiiey and Sons, 1970, pp. 530-531.
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valence electrons added to lhe sulfur atom toaccoum for the double negative charge
Multiple bonds affect bond lengths, not the type of geometric structure. (4) Only nonmetal atoms of the second row (8, C, N, 0, F) consistently follow the octet rule. These atoms have only s andp orbitals available for hybridization and can readily obtain an inert eas confieuration. All other central atoms. metal and nmmeral, ma" haw more r h n n eighr valrnr~clectnms a i they haw ,, p, and d orbitals nvailatde im hyhridimtion. Molerdm and ions u h ~ hove h a central atom that violates thr wrrr rub are called complex or coordinated Second, we review the concept of using formal charge to draw structures. We remind students that the formal charge of any atom in a molecule or ion is the difference between the number of valence electrons that atom has in its bonded state as compared with its nonbonded state, i.e., as if it were an isolated atom. Since we are followine the basic conceot of Linus Pauling's electroneutrality we call these structures "Pauling Structures." We remind students: (1) To first surround the central atom with its ligand atoms (e.g.,
sulfate ion, Figure 1A). (2) To place around each atom the number of valence electrons associated with its group number (Figure 1B). For noble gas atoms, assume eight valence electrons. For negatively charged ions, add one additional valence electron to the central atom for each negative charge (Figure 1C). For positively charged
Fogwe 2 A. Ammonl~mon wlth one valenceelectron removed bom the nobogen atom la account tor a s ngle pos.tlve cnarqa B Ammma wlth the one pair electrons affecting the location of bonding electrons.
Diana Eberlin is a stdem and Manus Monroe is an insbuctarat lndian Valley Colleges PIsend all wrilten inquiries to Dr. Manus m r o e . Chemistry Department, Indian Valley Colleges, lgnacio Baulevard, Novato, California 94947. This feature is aimed as a review of basic chemical .orinciDles . and as a reappram1 ol me state ol tne m Comments, s~ggesllonslor lopes m a conlroobt ons shobld oe sent to me lealure ed lor
Volume 59
Number 4
April 1982
285
Figure 3. Hydroxide ion.
Figure 4. Ammonium ion.
Figure 5. Paullng shrcture far boron trifluoride (resonance not shown).
ions, remove one valence electron from the central atom for each positive charge (Figure 2A). We now inform students that these rules are not applicable to the metal atoms of Group VIII. (3) To draw single bonds between the central atom and its ligands. Now, we tell students to count the number of electrons associated with each atom rememherine that in this teehniaue of an atomis the same as itsgroup number, they must assign that atom a formal charge of zero (e.g., ammonia molecule, Figure 28). If the total number of valence electrons associated with an atom is greater than its group number, then assign that atom a negative formal charge (e.g., hydroxide ion, Figure 3). If the total number of valence electrons associated with an atom is less than its group number, then assign that atom a formal positive charge (e.g., ammonium ion, Figure 4). (4) That, for molecules,all the formal charges must add up to zero. Generally, the formal charge on each atom in a molecule is zero. For ions, all the formal charges must add up to the actual charge on that ion. As a rule, atoms with higher electranegativity values most often have negative formal charges (Figure 3) while atoms with lower electroneeativitv values most often have positive formal charge+ F i ~ ~ l4). r e The nunmetal at
