1. Drawing Lewis Structures

2. Some issues about Lewis Structures to be discussed: (1)Drawing “valid” Lewis structures which follow the “octet” rule (holds almost without exception for first full row) (2)Drawing structures with single, double and triple bonds (3)Dealing with isomers (same composition, different constitution) (4)Dealing with resonance structures (same constitution, different bonding between atoms) (5)Dealing with “formal” charges on atoms in Lewis structures (6)Dealing with violations of the octet rule: Molecules which possess an odd number of electrons Molecules which are electron deficient Molecules which are capable of making more than four covalent bonds

3. The Lewis Model of Chemical Bonding In 1916 G. N. Lewis proposed that atoms combine in order to achieve a more stable electron configuration. Maximum stability results when an atom is isoelectronic with a noble gas. An electron pair that is shared between two atoms constitutes a covalent bond.

4. Covalent Bonding and Lewis Structures (1)Lewis “dot” (electron) structures of valence electrons for atoms (2)Use of Periodic Table to determine the number of “dots” (3)Use of Lewis structures to describe the electronic structures of atoms and molecules (4)Works best for covalent bonds and for elements in the first full row of the Periodic Table: H, He, Li, Be, B, C, N, O, F, Ne (5)Works with restrictions for second full row of the Periodic Table and beyond: Na, Mg, Al, Si, P, S, Cl, Ar

5. Valence electrons for Elements Represent the number of valence electrons as dots Valence number is the same as the Periodic Table Group Number

6. Lewis “dot-line” representations of atoms and molecules Electrons of an atom are of two types: core electrons and valence electrons The number of valence electrons is equal to the group number of the element for the representative elements. For atoms the first four dots are displayed around the four “sides” of the symbol for the atom. If there are more than four electrons, the dots are paired with those already present until an octet is achieved. Ionic compounds are produced by complete transfer of an electron from one atom to another. Covalent compounds are produced by sharing of one or more pairs of electrons by two atoms.

7. Covalent bonding and Lewis structures (1)Covalent bonds are formed from sharing of electrons by two atoms. (2)Molecules possess only covalent bonds. (3)The bedrock rule for writing Lewis structures for the first full row of the periodic table is the octet rule for C, N, O and F: C, N, O and F atoms are always surrounded by eight valence electrons. (4) For hydrogen atoms, the doublet rule is applied: H atoms are surrounded by two valence electrons.

8. Writing Lewis structures The skeletal structure of a polyatomic ion / molecule indicates the order in which the atoms are attached to one another It consists of one or more central atom(s) and at least 2 terminal atoms A central atom is bonded to two / more atoms in the structure A terminal atom is bonded to only one other atom In writing a skeletal structure the idea that every atom must be connected to the rest of the structure by at least one bond is applied.

9. Valence electrons and number of bonds Number of bonds elements prefers depending on the number of valence electrons. In general - X Family  # Covalent Bonds* Halogens F, Br, Cl, I Calcogens O, S Nitrogen N, P Carbon C, Si  O N C    1 bond often 2 bond often 3 bond often 4 bond always The above chart is a guide on the number of bonds formed by these atoms.

10. 10.7.00 6:16 PM 10 Lewis Structure Tutorial Lewis Structure, Octet Rule Guidelines When compounds are formed they tend to follow the Octet Rule. Octet Rule: Atoms will share electrons (e - ) until it is surrounded by eight valence electrons. Rules of the (VSEPR) game- i) O.R. works mostly for second period elements. Many exceptions especially with 3rd period elements (d-orbitals) ii) H prefers 2 e - (electron deficient) iii) :C:N::O::F: 4 unpaired 3unpaired2unpaired1 unpaired up = unpaired e- 4 bonds3 bonds2 bonds1 bond O=C=ON  NO = OF - F iv) H & F are terminal in the structural formula (Never central)........

11. Atomic Connectivity The atomic arrangement for a molecule is usually given. CH 2 ClF HNO 3 CH 3 COOHH 2 SeH 2 SO 4 O 3 HC F Cl H H N OO O H O S O H O O OO O H C C O H H H O HSe H In general when there is a single central atom in the molecule, CH 2 ClF, SeCl 2, O 3 (CO 2, NH 3, PO 4 3- ), the central atom is the first atom in the chemical formula. Except when the first atom in the chemical formula is Hydrogen (H) or fluorine (F). In which case the central atom is the second atom in the chemical formula. Find the central atom for the following: 1) H 2 Oa) Hb) O2) PCl 3 a) Pb) Cl 3) SO 3 a) Sb) O4) CO 3 2- a) Cb) O 5) BeH 2 a) Beb) H6) IO 3- a) Ib) O

12. Note the following Hydrogen atoms are nearly always terminal atoms, they form only one bond. (H has 2 electrons in valence shell) In polyatomic molecules and / ions, the central atom(s) usually have the lowest electronegativity except for hydrogen (that is always terminal even when bonded to a more electronegative atom) In oxo acids, hydrogen atoms are usually bonded to oxygen atoms. With the major exception of carbon compounds in which long chains of carbon atoms are common, polyatomic molecules and ions usually form compact structures

13. Steps for writing Lewis structures 1.Determine the total number of valence electrons. – The total number of valence electrons for a molecule is the sum of the valence electrons for each atom. N 2 O 4 ----- (2 x 5) + (4 x 6) = 34 valence electrons – For a polyatomic anion, which has one / more extra electrons, add one electron for each unit of negative charge NO 3 - ----- 5 + (3 x 6) + 1 = 24 valence electrons – For a polyatomic anion, which is missing one / more electrons, subtract one electron for each unit of the positive charge NH 4 + ----- 5 + (4 x 1) – 1 = 8 valence electrons

14. 2.From the chemical formula, determine the atom connectivity for the structure. Given a chemical formula, AB n, A is the central atom and B flanks the A atom. i.e., NH 3, NCl 3, NO 2. In these examples, N is central in the structure. H and F are never central atoms. 3. Write the skeletal structure and connect bonded atoms with an electron – pair bond (dash)

15. 4. Place electron pair around terminal atoms so that each atom (except Hydrogen) has an octet) 5. Assign any remaining electrons as lone pairs around the central atom(s). 6. If at this point a central atom has fewer than 8 electrons, a multiple bond(s) is likely. Move one or more lone pairs from a terminal atom(s) to a region between it and the central atom to form a double or triple bond.