UNIT IV chemical bonding (1) (2)

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UNIT IV: CHEMICAL BONDING
Lewis Structures (Diagrams) of covalent compounds and polyatomic ions show the bonding
arrangement of those compounds or ions.
(Chang 13th edition, section 9.6 and 9.9 read pages 381-384 and 389-395)
Practice Problems from the text: #9.43-9.47 pages 402-403
Lewis structures are based on the idea of electron sharing between atoms in order to
complete an electron octet. They show a two-dimensional bonding picture of a molecule or an
ion.
The valence electrons are represented as follows:
 non-bonded electron pairs: 
(free or lone pairs)
 bonded electron pair(s): – (the dash represents two bonded electrons)
(shared between two atoms)
LEWIS STRUCTURES OF COVALENT COMPOUNDS AND POLYATOMIC IONS
1. A probable arrangement of the atoms can be deduced as follows:
(note: this is for neutral molecules)
 The structure assumed must be consistent with the known valences of the elements:
e.g. H forms one bond,
F forms one bond,
O forms two bonds,
N forms three bonds,
C forms four bonds, etc. (
 In a molecule with the formula AXn the atom A is usually the central atom to which X
are attached:
e.g. CH4, BeH2, PCl5, SF6, etc.
 Atom A is also normally the least electronegative atom or the atom with the highest
valence.
ex. H2O, NH3, BF3, etc.

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1. Draw a diagram of the molecule or polyatomic ion, showing the atoms connected by single
bonds to the central atom
2.
 Add the number of valence electrons of each atom in the molecule to find the total
number of valence electrons.
 Subtract the number of electrons needed to form the single bonds from the total
number of electrons.
 Use the remainder to complete octets around each atom except hydrogen.
 Exceptions to the octet rule: compounds of Be and B have less than 8 electrons in their
valence shell. Be ends up with 4 electrons and B has 6 electrons.
The rule is that if there are insufficient electrons to complete all the octets, complete
those of the more electronegative atoms first.
1) Incomplete octet: BeCl2, BF3, SnF2, and HgCl2
2) Molecules/ions with an odd number of electrons: NO, NO2
3) Expanded octet: PCl5, SF6, XeF4 etc… (Noble gases only have expanded octets)
Examples:
CH4
H2O
NH3

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Exceptions to the octet rule (Incomplete octet):
BeCl2
BeH2
BF3
BH3

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3. If the molecule is charged, and/or a polyatomic ion, add one electron for each negative charge
or subtract one electron for each positive charge.
Examples:
H3O+
NH4+
BCl4–
NH2–

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4. Atoms from the 3rd or later period:
 If the central atom is from period 3 or a later period, the octet rule may not apply. These
atoms can have more than 8 electrons in their valence shell: (Expanded octet)
Examples:
PCl5 SF6
 If there are more than enough electrons to complete all octets, add the remaining
electrons in pairs to the central atom (which must be from the 3rd or a later period).
Examples:
SCl4 ClF3
XeF2 BrF4+
SbCl52–