John E.

McMurry

www.cengage.com/chemistry/mcmurry

Chapter 2
Polar Covalent Bonds;
Acids and Bases

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1. Polar Covalent Bonds:
Electronegativity
▪ Covalent bonds can have ionic character
▪ Polar covalent bonds: Bonding electrons are
attracted more strongly by one atom than by the
other
▪ Electron distribution between atoms is not
symmetrical

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Electronegativity

▪ Intrinsic ability of an atom to attract the shared

electrons in a covalent bond
▪ Differences in EN produce bond polarity
▪ F is most electronegative (EN = 4.0),
Cs is least (EN = 0.7)
▪ Metals on left side of periodic table attract
electrons weakly
▪ Halogens and other reactive nonmetals on right
side of periodic table attract electrons strongly
▪ EN of C = 2.5
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Figure 2.2 - Electronegativity Values
and Trends

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Bond Polarity and Inductive
Effect
▪ Difference in EN of atoms < 2 in polar covalent bonds
▪ C–H bonds are relatively non-polar (2.5-2.1=0.4) ΔEN < 0.5
▪ C–O, C–X bonds are polar (3.5-2.5=1.0) 0.5 < ΔEN < 2
▪ Difference in EN > 2 in ionic bonds
▪ NaCl (3.0-0.9=2.1)
▪ Bonding electrons toward electronegative atom
▪ C acquires partial positive charge, +
▪ Electronegative atom acquires partial negative
charge, -
▪ Inductive effect: Shifting of electrons in a σ bond in
response to EN of nearby atoms

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Electrostatic Potential Maps

▪ Show calculated
charge distributions
▪ Colors indicate
electron-rich (red) and
electron-poor (blue)
regions
▪ Arrows indicate
direction of bond
polarity

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Worked Example

▪ Which element in each of the following pairs is

more electronegative?
▪ (a) Li or H
▪ (b) Cl or I

▪ Solution:
▪ Using Figure 2.2
▪ (a) Li (1.0) vs. H (2.1)
▪ (b) Cl (3.0) vs. I (2.5)

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2.Polar Covalent Bonds:
Dipole Moments
▪ Molecules are often polar from vector summation of
individual bond polarities and lone-pair contributions
▪ Strongly polar substances are soluble in polar
solvents like water
▪ Nonpolar substances are insoluble in water
▪ Dipole moment (): Net molecular polarity, due to
difference in summed charges
▪ =Qxr
▪ Q: Magnitude of charge at end of molecular dipole
▪ r : distance between charges

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Polar Covalent Bonds: Dipole
Moments
▪  = Q  r, in debyes (D), 1 D = 3.336  10-30 coulomb meter
▪ Q : 1.60 × 10-19 C (전자 단위 전하)
▪ r = 100 pm (양전하 음전하 간의 거리)
D = 1.60 × 10-29 Cm, or 4.80 D

Dipole Moments of HX

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Polar Covalent Bonds: Dipole
Moments
▪ Large dipole moments
▪ EN of O and N > H
▪ Both O and N have lone-pair electrons oriented
away from all nuclei

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Dipole Moments in Water and
Ammonia

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Absence of Dipole Moments

▪ In symmetrical molecules, the dipole moments

of each bond have one in the opposite direction
▪ The effects of the local dipoles cancel each other

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Worked Example
▪ Draw three-dimensional drawing of H2C═CH2 molecules
▪ Predict whether it has dipole moment

▪ Solution:
▪ Drawing an arrow that points from the least
electronegative element to the most electronegative
element
▪ Has zero dipole moment

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3.Formal Charges

▪ At times it is necessary to have structures with

formal charges on individual atoms
▪ Bonding of the atom in the molecule is
compared to valence electron structure of atom

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Formal Charge for Dimethyl
Sulfoxide
▪ Atomic S has 6 valence
electrons
▪ Dimethyl sulfoxide S has
only 5
▪ It has lost an electron and
has positive charge (+1)

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Table 2.2 - A Summary of Common
Formal Charges

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Worked Example

▪ Calculate formal charges on the four O atoms in

the methyl phosphate di-anion

▪ Solution:

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Worked Example

▪ For oxygen 1: 4
FC = 6 - -4=0
2
4
▪ For oxygen 2: FC = 6 - - 4 = 0
2
2
▪ For oxygen 3: FC = 6 - - 6 = -1
2
2
▪ For oxygen 4: FC = 6 - - 6 = -1
2
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Worked Example

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4. Resonance

▪ Some molecules have structures that cannot be

shown with a single representation
▪ Represented by structures that contribute to the
final structure but differ in the position of the 
bond or lone pair
▪ Such structures are delocalized and are
represented by resonance forms

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Resonance Hybrids

▪ Structure with resonance forms, that does not

alternate between the forms
▪ Example - Benzene (C6H6) has two resonance
forms with alternating double and single bonds
▪ Is a hybrid of the two individual forms
▪ All six carbon–carbon bonds are equivalent

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5. Rules for Resonance Forms

1. Individual resonance forms are imaginary

▪ Real structure is a hybrid of different forms
2. Resonance forms differ only in the placement
of their  or nonbonding electrons
▪ Curved arrow indicates movement of electrons,
not of the atoms

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Rules for Resonance Forms

3. Different resonance forms of a substance do

not have to be equivalent
▪ When two resonance forms are nonequivalent,
the actual structure of the resonance hybrid
resembles the more stable form

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Rules for Resonance Forms

4. Resonance forms obey normal rules of valency

5. Resonance hybrid is more stable than any

individual resonance form
▪ Resonance leads to stability

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6. Drawing Resonance Forms

▪ Any three-atom grouping with a p orbital on

each atom has two resonance forms

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Drawing Resonance Forms

▪ Resonance forms differ by an exchange in

position of the multiple bonds and the asterisk
▪ From one end of the three-atom grouping to the
other
▪ Recognizing three-atom groupings within larger
structures help generate resonance forms,
symmetrically

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2,4-Pentanedione Anion

▪ Has a lone pair of electrons

▪ Has a formal negative charge on the central
carbon atom, next to a C═O bond on the left
and on the right
▪ Has three resonance structures

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Carbonate ion, CO3-2

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Worked Example

▪ Draw the indicated number of resonance forms

for:
▪ The allyl cation, H2C═CH-CH2+
▪ Solution:
▪ Locating three-atom groupings that contain a multiple
bond next to an atom with a p orbital
▪ Exchanging the positions of the bond and the
electrons in the p orbital, we have:

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Pentadienyl Radical

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Worked Example

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7. Acids and Bases: The
Brø nsted-Lowry Definition
▪ Idea that acids are solutions containing a lot of “H+”
and bases are solutions containing a lot of “OH-” is
not very useful in organic chemistry

▪ Brø nsted-Lowry theory defines acids and bases by

their role in reactions that transfer protons (H+)
between donors and acceptors

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Acids and Bases: The Brø nsted-
Lowry Definition
▪ Brø nsted-Lowry acid: Substance that donates a
hydrogen ion, H+
▪ Brø nsted-Lowry base: Substance that accepts a
hydrogen ion, H+
▪ Proton is a synonym for H+
▪ Loss of valence electron from H leaves only the

nucleus—a proton

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Acids and Bases: The Brø nsted-
Lowry Definition
▪ Conjugate base: Product that results from
deprotonation of a Brø nsted-Lowry acid
▪ Conjugate acid: Product that results from
protonation of a Brø nsted-Lowry base

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8. Acid Base Strength

▪ Acidity constant (Ka): Measure of acid strength

▪ For the reaction of an acid (HA) with water to
form hydronium ion
▪ Conjugate base (A-) is a measure related to the
strength of the acid
▪ Brackets [ ] indicate concentration in moles per liter

𝐻3 𝑂+ [𝐴− ]
𝐾𝑎 =
𝐻𝐴 [𝐻2 𝑂] 용매 농도 무시

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Acid and Base Strength

▪ Acid strengths are normally expressed using pKa

values
▪ pKa: Negative common logarithm of the Ka

pK a = - logK a
▪ Stronger acids have smaller pKa
▪ Weaker acids have larger pKa

(10-60 < Ka < 1015)

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Table 2.3 - Relative Strengths of Some
Common Acids and Their Conjugate Bases

Inverse relationship
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Worked Example

▪ The amino acid phenylalanine has pKa = 1.83,

and tryptophan has pKa = 2.83
▪ Which is the stronger acid?

▪ Solution:
▪ Stronger acid has a smaller pKa and a weaker
acid has a larger pKa
▪ Accordingly, phenylalanine (pKa = 1.83) is a
stronger acid than tryptophan (pKa = 2.83)

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9. Predicting Acid-Base
Reactions from pKa Values
▪ pKa values are related as logarithms to
equilibrium constants
▪ Useful for predicting whether a given acid-base
reaction will take place
▪ Difference in two pKa values is the log of the
ratio of equilibrium constants, and can be used
to calculate the extent of transfer

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Predicting Acid-Base
Reactions from pKa Values

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Worked Example
▪ Will the following reaction take place to a significant
extent as written, according to the data in Table 2.3?
?
▪ HCN + CH3CO2- Na+ → Na+ -CN + CH3CO2H

▪ Solution:
?
▪ HCN + CH3CO2 - Na+ → Na+ -CN + CH3CO2H
pKa= 9.3 pKa= 4.7
Weaker acid Stronger acid

▪ Since CH3CO2H is stronger than HCN the reaction

will not take place to a significant extent the direction
written

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10. Organic Acids

▪ Characterized by the presence of positively

polarized hydrogen atom

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Organic Acids
▪ Two main kinds, those that contain:
▪ Hydrogen atom bonded to a electronegative oxygen
atom(O–H)
▪ A hydrogen atom bonded to a carbon atom next to a
C═O bond(O═C─C─H)

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Carboxylic acids

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Organic Bases

▪ Have an atom with a lone pair of electrons that can

bond to H+
▪ Nitrogen-containing compounds derived from
ammonia are the most common organic bases
▪ Oxygen-containing compounds can react as bases
with a strong acid or as acids with strong bases

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11. Acids and Bases:
The Lewis Definition
▪ Lewis acid: Electron pair acceptors
▪ Lewis bases: Electron pair donors
▪ Brø nsted-Lowry definition 보다 포괄적인 개념

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Lewis Acids and the Curved
Arrow Formalism
▪ Lewis definition of acidity includes metal cations,
such as Mg2+
▪ They accept a pair of electrons when they form a
bond to a base

▪ Group 3A elements, such as BF3 and AlCl3, are

Lewis acids
▪ Have unfilled valence orbitals and can accept
electron pairs from Lewis bases
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Figure 2.5 - The Reaction of Boron
Trifluoride with Dimethyl Ether

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Lewis Acids and the Curved
Arrow Formalism
▪ Transition-metal compounds, such as TiCl4, FeCl3,
ZnCl2, and SnCl4, are Lewis acids
▪ Curved arrow means that a pair of electrons move
from the atom at the tail of the arrow
to the atom at the head of the arrow

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Some Lewis Acids

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Lewis Bases

▪ Compound with a pair of nonbonding electrons

that it can use to bond to a Lewis acid
▪ Can accept protons as well as Lewis acids
▪ Definition encompasses that for Brø nsted bases
▪ Oxygen-and nitrogen-containing organic
compounds are Lewis bases; they have lone
pairs of electrons
▪ Some compounds can act as both acids and
bases

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Lewis Bases

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Worked Example

▪ Using curved arrows, show how acetaldehyde,

CH3CHO, can act as a Lewis base
▪ Solution:
▪ A Lewis base donates an electron pair to a Lewis
acid
▪ Using a curved arrow to show the movement of a
pair toward the H atom of the acid

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12. Noncovalent Interactions
Between Molecules
▪ Noncovalent interactions: One of a variety of
nonbonding interactions between molecules
▪ Dipole–dipole forces
▪ Dispersion forces
▪ Hydrogen bonds

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Dipole–Dipole Forces

▪ Occur between polar molecules as a result of

electrostatic interactions among dipoles
▪ Depending on orientation of the molecules, the
forces can be either attractive or repulsive

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Dispersion Forces

▪ Occur between all neighboring molecules

▪ Arise due to constant change in electron
distribution within molecules
▪ Induced dipole

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Hydrogen Bond Forces

▪ Forces are the result of attractive interaction

between a hydrogen bonded to an
electronegative O or N atom and an unshared
electron pair on another O or N atom

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A Deoxyribonucleic Acid Segment

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Worked Example

▪ Of the two vitamins A and C, one is hydrophilic

and water-soluble while the other is hydrophobic
and fat-soluble
▪ Which is which?

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Worked Example
▪ Solution:

▪ Vitamin C has several polar ─OH groups that can

form hydrogen bonds with water
▪ It is water soluble(hydrophilic)

▪ Most of Vitamin A’s atoms can’t form hydrogen bonds

with water
▪ It is fat-soluble(hydrophobic)

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Summary

▪ Organic molecules often have polar covalent

bonds as a result of unsymmetrical electron
sharing caused by differences in the
electronegativity of atoms
▪ Polarity of a molecule is measured by its dipole
moment, 
▪ (+) and (−) indicate formal charges on atoms in
molecules to keep track of valence electrons
around an atom

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Summary

▪ Some substances must be shown as a

resonance hybrid of two or more resonance
forms that differ by the location of electrons
▪ A Brø nsted(–Lowry) acid donates a proton
▪ A Brø nsted(–Lowry) base accepts a proton
▪ Strength of Brø nsted acid is related to the
negative logarithm of the acidity constant, pKa
▪ Weaker acids have higher values of pKa

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Summary

▪ Lewis acid has an empty orbital that can accept

an electron pair
▪ Lewis base can donate an unshared electron
pair
▪ Non-covalent interactions have several types –
Dipole–dipole, dispersion, and hydrogen bond
forces

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