Chapter 10: Chemical Bonding I: Basic Concepts
• We will look at topics culminating in two skills:
– Drawing Lewis structures.
– Predicting molecular shapes using a Lewis structure & VSEPR.
The Lewis Theory of Chemical Bonding
• Valence electrons play a fundamental role:
– Main Group elements react to achieve a filled valence shell
octet of the form ns2 np6 (the ‘octet’ rule). NB: duet for H, He
– Elements in period 3 and higher can expand their ‘octets.’
– Electron transfer from one atom to another (metal  nonmetal)
forms positive and negative ions, leading to ionic bonds.
– Sharing of one or more pairs of electrons results in a covalent
bond.
Lewis Symbols of Main Group Atoms
• The Lewis symbol of an element shows the
valence electrons as dots arranged around
the element symbol.
– For Main Group elements, the number
of valence electrons equals the last digit
of the group number (1-2 and 13-18).
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• Lewis symbols of the s-block elements:
Element or Ion
# VE
Lewis Symbol (or Structure)
H
1
He
2
Li
1
Be
2
Na
1
Mg
2
paired b/c shell 1 filled
• Lewis symbols of the p-block elements:
Group =
13
14
15
16
17
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• The number of unpaired electrons in the Lewis Structure of an
element equals the preferred number of covalent bonds.
– Deviations for period 2 elements results in a charged species.
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Simple Lewis Structures
• For ionic compounds, separate the ions and draw a Lewis
structure for each.
• For covalent compounds we shall review several topics which
will feed into a general approach to drawing Lewis structures.
– Keep in mind that these covalently bonded species will almost
always be molecules or ions consisting only of nonmetals.
• Knowing the Octet Rule and observing the bonding patterns
(common ‘fragments’) for period 2 elements will help!
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• Electron pairs in Lewis Structures fall into one of two types:
– Bonding Pairs which are shared.
 Placed between element symbols.
 Almost always replaced with — (each — is 2 bonding e−‘s)
– Lone Pairs or Nonbonding Pairs which are not shared.
– Species with an odd number of valence electrons will have a
‘lone single’ instead of a ‘lone pair.’
 Treat like lone pair in VSEPR, but has less repulsion.
• Shared bonding pairs may be further classified:
– Coordinate Covalent Bonds are formed when a single element
uses one of its lone pairs to form a bond.
 The resulting bond is indistinguishable from a covalent
bond formed when both elements contribute electrons.
– Multiple Covalent Bonds are formed when more than one pair
of electrons is shared; this is done to satisfy the octet rule.
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Visualizing Electron Distribution and Electronegativity
• Electrostatic potential map gives information about the
distribution of electron charge in a molecule.
• Electronegativity is the tendency of an element to attract shared
electrons to itself; this can ‘polarize’ a covalent bond.
– Results in δ+ and δ− ends of the bond.
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Difference in Electronegativity and Bond Character
• Bonds can be classified by the absolute value of the difference
in electronegativity (∆EN) between the two elements:
– Non-polar covalent bonds have ∆EN = 0.
 Often relaxed, especially in organic chemistry (cf C−H).
– Polar covalent bonds can be classified as those with a ∆EN
between 0.1 and 1.7, with some exceptions.
 Bonding electrons are not shared equally.
 Some books use a range of 0.1 to 2.0.
– Ionic bonds (greater than 50% ionic character) have ∆EN
greater than 1.7.
• The ‘percent ionic character’ of a bond varies widely with ∆EN.
• A useful rule of thumb is that compounds of nonmetals only are
covalent, while those of metals and nonmetals are ionic.
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Formal Charge Determination
• The formal charge (FC) of an element in a compound is a
calculated bookkeeping number that tells us whether the
element is electron rich (FC < 0) or electron poor (FC > 0)
compared to the free element.
– FC is not necessarily equal to the electric charge!
FC= (# valence e− in free atom) − (# lone-pair e−) − (½ # bond-pair e−)
FC = e−‘s in free element − all lone pair e−‘s − one-half bonding pair e−‘s
FC = VE – LPE − ½BPE
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A Lewis Structure Heuristic
1. Determine if species is ionic or covalent.
If ionic, treat each ion separately.
2. Count the number of valence electrons.
If anion, add electrons.
If cation, subtract electrons.
[Hint: Try using pairs since electrons come as bonding or lone pairs.]
3. Assemble the single-bonding framework. Connect atoms w/single bonds.
Inner/central atom (more electropositive atom) bonded to >1 atoms.
Outer/terminal/ligands atoms (more electronegative) bonded to 1 atom.
Connect inner atom to outer atoms with an electron pair (single bond).
[Hint: Many species exhibit an ABn motif: A inner and n B outer atoms.]
[Hint: There may be multiple inner atoms as number of atoms grows.]
4. Place e− pairs around outer atoms to form octets, except in the case of H.
5. Assign any remaining electron pairs around the central atom(s) beginning
with the most electronegative; this is the ‘provisional structure.’
[Hint: If central atom is in period > 2, may have more than four pairs.]
6. Check for optimum electron configuration around the central atom(s).
[Hint: This is where multiple bonds may be formed to complete octets.]
If central atom lacks an octet, try completing octet by shifting lone pairs
on terminal atoms to make double or triple bonds to central atom.
Determine Formal Charge on central atoms to test structure.
Formal charge on central atoms should be 0 or positive.
Sum of FCs for all atoms equals the charge on species.
7. Identify all possible equivalent structures, that is, look for resonance
structures involving alternating double and single bonds.
[Hint: Only move e−‘s to form resonance structures.]
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Draw the Lewis structures of the following
NNO
H 3O +
XeF2
ClF3
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HCN
Resonance
• Resonance is not an equilibrium between structures!!
– Resonance tells us that one Lewis structure is not sufficient to
describe the bonding. Resonance forms demonstrate a
limitation of Lewis structures, a localized bonding model.
– Best resonance structures have the fewest and smallest FC’s.
 If all formal charges = 0, do not draw resonance structures!
NO2−
SO42−
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Exceptions to the Octet Rule
• Odd electron species have an odd number of VE to work with.
• Incomplete octets occur in electron-deficient compounds of
boron where the central atom is surrounded by only 6 VE.
• Expanded valence shell results from ‘octet expansion’ for
elements from period 3 and higher (equivalently, n > 2).
Shape of Molecule: VSEPR (Valence Shell Electron Pair Repulsion)
• The rationale behind VSEPR is that the electron pairs around a
central atom (bonding and nonbonding) position themselves so
as to minimize electron repulsion.
– This maximizes the distance between electron pairs.
– Single, double, and triple bonds are treated as ‘one pair.’
[I like to count number of BONDED atoms and number of LP.]
• What you need to do to determine the shape, or the geometric
arrangement of bonded atoms around the central atom.
1. Draw a correct Lewis structure. (w/out correct Lewis, VSEPR NG!)
2. Pick a central atom and count lone pairs (LP or E) on the
atom, and number of elements (CN or X) bonded to the atom.
3. Look up geometry in table! (Repeat until table memorized!)
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# e− groups
(bonding+LP)
# e− groups
(bonding+LP)
Electron Group
Geometry
Electron Group
Geometry
Drawing
Molecular
Geometry
Drawing
Molecular
Geometry
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SN=CN+LP VSEPR Angles
2
2
0
AX2
180°
3
3
0
AX3
120°
3
2
1
AX2E
120°
4
4
0
AX4
109.5°
4
3
1
AX3E
109.5°
4
2
2
AX2E2
109.5°
SN=CN+LP VSEPR Angles
5
5
0
AX5
90°
180°
5
4
1
AX4E
90°
180°
5
3
2
AX3E2
90°
5
2
3
AX2E3
180°
# e− groups
(bonding+LP)
Electron Group
Geometry
Drawing
Molecular
Geometry
SN=CN+LP VSEPR Angles
6
6
0
AX6
90°
6
5
1
AX5E
90°
6
4
2
AX4E2
90°
Predict the VSEPR shape (molecular geometry) of the following:
NNO
H 3O +
XeF2
ClF3
HCN
NO2−
SO42−
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Draw the Lewis structure and predict the shapes of the three
central atoms in ethanol (ethyl alcohol): CH3CH2OH
Molecular Shapes and Dipole Moments
• Dipole moment µ = δ x d
– Nonpolar molecules have µ = 0
– Polar molecules have µ > 0
• Polar bonds and polar molecules
– A polar molecule must contain polar bonds, BUT polar bonds
do not necessarily produce a polar molecule.
 Bond dipoles can cancel if the geometry is right!
• Examining molecular polarity: CO2, H2O, CCl4, and CHCl3
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Bond Order and Bond Lengths
• Bond order is equal to one-half the number of bonding electrons
• Bond length is the distance between atomic centers.
• Estimate bond length using tabulated data.
Estimate the length of a Cl—O bond
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Bond Energies
• Bond dissociation energy is the energy required to break a
chemical bond.
– Bond breaking is always endothermic, ∆H > 0.
– Bond forming is always exothermic, ∆H < 0.
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• Tabulated bond energies are averages
Bond Energies and Reaction Enthalpy (∆Hrxn)
• Signs of bond breakage and bond formation
∆Hrxn = ∆H (bonds broken) + ∆H (bonds formed)
∆Hrxn = ΣBE (reactants) − ΣBE (products)
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A brief diversion: the paramagnetism of liquid
oxygen and the limitations of Lewis structures
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