CHEMICAL BONDING
Chemical Bonds
Lewis Symbols
Octet Rule
whenever possible, valence electrons in covalent compounds distribute so that
each main-group element is surrounded by 8 electrons (except hydrogen which
wants 2 electrons around it).
Ionic Bonding
Electron Configurations of Ions
Sizes of Ions
Ex 1
Which of the following atoms and ions is largest:
2–
2–
(a)
S , S, O
(b)
Rb , Sr , Y
+
2+
3+
Covalent Bonding
Elements which are neither highly electropositive nor highly electronegative are less likely to form
ionic compounds. Consider CH4 , methane; carbon and hydrogen have similar electronegativities. It
would be inappropriate to describe the bonds between carbon and hydrogen in methane as ionic.
Instead, the bonding between these two atoms involves a sharing of electrons (not an electron
transfer) and is termed a covalent bond.
Lewis Structures
Multiple Bonds
Ex 2
(a)
H2O
(b)
PH3
(c)
CH4
(d)
CO 2
(e)
N2
Bond Polarity
Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself.
Electronegativity and Bond Polarity
Ex 3
Which bond is more polar: S—Cl, S—Br, Se—Cl, or Se—Br?
Dipole Moments
Drawing Lewis Structures
Lewis structures help to explain or predict trends in chemical reactivity.
Procedure for drawing a Lewis Structure (from your text):
1. Count up total number of valence electrons from all atoms. If the species is an ion, either add
the negative charge to A (for anion) or subtract the positive charge from A (for cation).
2. Draw a skeleton structure.
3. Complete the octets of the atoms bonded to the central atom.
4. Place any leftover electrons on the central atom, even if that means giving the central atom more
than an octet of electrons.
5. If not enough electrons to give central atom an octet, try using multiple bonds.
Alternative procedure for drawing a Lewis Structure:
1. Count up total number of valence electrons available (A). If the species is an ion, either add the
negative charge to A (for anion) or subtract the positive charge from A (for cation).
2. Calculate total number of electrons needed (N) to give each atom an octet
(8 × # non-hydrogen atoms + 2 × # hydrogen atoms)
3. Calculate the number of bonding electrons (S):
S =N -A
4. Draw a molecular skeleton for the substance and then assign two bonding electrons to each
connection between atoms. (Note that when constructing a molecular skeleton, H and F are
always terminal atoms.)
5. If any bonding electrons are left over, use them for double or triple bonds. Double bonds form
only between carbon, nitrogen, oxygen, and sulfur. Triple bonds are usually restricted to carbon
and nitrogen.
6. Distribute remaining electrons around atoms to complete their octets. (All atoms want an octet
except H.)
7. Determine the formal charge on each atom, and write it next to the atom. Check to see that the
formal charges add up to the overall charge on the molecule or ion. (See below)
Ex 4
Draw Lewis structures for the following:
(a) POCl 3
(b) SOCl 2
(c) ClO
−
(d) CH2O
Formal charge
Unequal sharing of electrons between atoms leads to partial charges on some of the atoms in a
molecule. Determination of the formal charge on atoms in a molecule can serve as a check of the
reasonableness of the Lewis structure we have drawn.
–
–
Formal charge = group number − # e in lone pairs − 1/2(# of shared e )
Ex. 5
Compute the formal charges on the atoms in the following Lewis structure for the
-
azide ion N3 .
[: N˙˙ = N = N˙˙ :]
–
Resonance Structures
Some molecules may have two or more equivalent Lewis structures. These are known as resonance
structures. This does not mean that the molecule jumps back and forth between these structures. In
reality, the true structure of the molecule will lie somewhere in between these resonance structures.
Ex 6 Draw the Lewis structures (including any resonance structures) for the following:
(a) SO 2
(b) SO 3
−
(c) NO 3
Exceptions to the Octet Rule
Odd Number of Electrons
Electron Deficient Molecules
Expanded Valence
Shapes of Molecules
Lewis structures give us a picture of how atoms are connected to each other, but they don’t tell us the
actual shapes of the molecules. Molecules are three-dimensional. We need an extension to Lewis
theory to help us predict the arrangement of the atoms in space.
VSEPR Theory (Valence Shell Electron Pair Repulsion)
•
Lone pairs occupy more space than bonding pairs
• Repulsive forces among valences pairs diminish in the following order:
lone pair vs. lone pair > lone pair vs. bond pair > bond pair vs. bond pair
•
Lone pairs prefer equatorial positions to axial positions in trigonal bipyramid.
Ex. 7
Predict the structures of the following:
(a)
water
(b)
SO 2
(c)
ClO 3
+
(d)
IF5
Shapes of Molecules with Bond Pairs ONLY, No Electron Pairs
109.5˚
120˚
180˚
AX 2
Linear
Planar
(CO2 )
AX 3
Trigonal
Planar
(BF3 )
AX 4
Tetrahedral
(CH4 )
90˚
90˚
120˚
AX 5
Trigonal
Bipyramidal
AX 6
Octahedral
(PCl5 )
(SF 6 )
Shapes of Molecules with . . .
Three Electron Pairs
Cl
Cl
B
Ge
•
•
No Lone Pairs
Trigonal Planar
One Lone Pair
Bent
(BCl3 )
(GeCl2 )
Four Electron Pairs
• •
C
• •
N
H
O
•
•
No Lone Pairs
Tetrahedral
One Lone Pair
Pyramidal
Two Lone Pairs
Bent
(CCl4 )
(NH 3 )
(H2 O)
Polarity of Polyatomic Molecules
Ex. 8
Predict whether the following have a dipole moment:
(a)
NH3
(b)
SF6
Some general rules of thumb regarding dipole moments:
•
a central atom with no lone pairs on it, surrounded by other atoms of all the
same type will NOT have a dipole moment (e.g. CH 4 , CO 2 , S F 6 )
•
a central atom with no lone pairs surrounded by other atoms of not all the same
type WILL usually have a dipole moment (e.g. CH 3 Cl, NO 2 Cl, OCN)
•
a central atom with lone pairs on it, surrounded by other atoms of any type
WILL almost always have a dipole moment (e.g. H 2 O, NH 3 , ICl 3 )
•
a diatomic molecule containing the same atoms will NOT have a dipole moment
(e.g. O 2 ); a diatomic molecule containing different atoms WILL have a dipole
moment (e.g. CO)