John E. McMurry
http://www.cengage.com/chemistry/mcmurry
Chapter 2
Polar Covalent Bonds; Acids
and Bases
Javier E. Horta, M.D., Ph.D. • University of Massachusetts Lowell
Polar Covalent Bonds: Electronegativity
• Polar covalent bonds
• Bond polarity is due to difference in electronegativity (EN)
•
A covalent bond in which the electron distribution between
atoms is unsymmetrical
∆EN < 0.5
0.5 ≤ ∆EN < 2.0
∆EN ≥ 2.0
Polar Covalent Bonds: Electronegativity
Electronegativity (EN)
• The ability of an atom to attract shared electrons in a
covalent bond
• Generally increases across the periodic table from left to
right and from bottom to top
Polar Covalent Bonds: Electronegativity
Electrostatic potential maps
2.2 Polar Covalent Bonds: Dipole Moments
Molecules as a whole are often polar
Molecular polarity results from the vector summation of all
individual bond polarities and lone-pair contributions in the
molecule
• Strongly polar substances are soluble in polar solvents like water
•
Dipole moment ()
A measure of the net polarity of a molecule
• Arises when the centers of mass of positive and negative charges
within a molecule do not coincide
• Measured in debyes (D)
•
Polar Covalent Bonds: Dipole Moments
Polar Covalent Bonds: Dipole Moments
Factors Affecting Dipole
Moments
Lone-pair electrons on
oxygen and nitrogen project
out into space away from
positively charged nuclei
giving rise to a considerable
charge separation and
contributing to the dipole
moment
• Symmetrical structures of
molecules cause the
individual bond polarities and
lone-pair contributions to
exactly cancel
•
Worked Example 2.1
Predicting the Direction of a Dipole Moment
Make a three-dimensional drawing of methylamine,
CH3NH2, and show the direction of its dipole
moment ( = 1.31)
Worked Example 2.1
Predicting the Direction of a Dipole Moment
Strategy
•
•
•
Look for any lone-pair electrons
Identify any atom with an electronegativity
substantially different from that of carbon (usually O,
N, F, Cl, or Br)
Electron density will be displaced in the general
direction of the electronegative atoms and the lone
pairs
Worked Example 2.1
Predicting the Direction of a Dipole Moment
Solution
•
Methylamine has an
electronegative nitrogen
atom and a lone pair of
electrons. The dipole
moment thus points
generally from –CH3 toward
the nitrogen
2.3 Formal Charges
Formal charge
•
•
•
•
The difference in the number of electrons owned by an atom in a molecule and
by the same atom in its elemental state
Formal charges do not imply the presence of actual ionic charges
Device for electron “bookkeeping”
Assigned to specific atoms within a molecule
• Dimethyl sulfoxide CH3SOCH3
•
•
Sulfur atom has three bonds rather than the usual two and has a formal
positive charge
Oxygen atom has one bond rather than the usual two and has a formal
negative charge
Formal Charges
Formal Charge Determination
Formal Charges
Formal Charges
2.4 Resonance
Two different ways to draw the acetate ion
• Double bond placement
• Neither structure correct by itself
• True structure is intermediate between the two
• Two structures are known as resonance forms
Resonance
Resonance forms
• Individual line-bond
structures of a molecule or
ion that differ only in the
placement of  and
nonbonding valence
electrons
• Indicated by “
”
• Resonance forms
contribute to a single,
unchanging structure
that is the resonance
hybrid of the individual
forms and exhibits the
characteristics of all
contributors
Resonance
Benzene has two equivalent resonance forms
• The true structure of benzene is a hybrid of the two
individual forms, and all six carbon-carbon bonds are
equivalent
• Symmetrical distribution of electrons is evident in an
electrostatic potential map of benzene
2.5 Rules for Resonance Forms
Rule 1 – Individual resonance forms are imaginary, not
real
•
Real structure is a composite
Rule 2 – Resonance forms differ only in the placement of
their or nonbonding electrons
Rules for Resonance Forms
•  electrons in double bonds of benzene move
• Electron movement is indicated by curved arrow formalism
• Curved arrows indicate electron flow, not the movement of
atoms
• A curved arrow indicates that a pair of electrons moves
from the atom or bond at the tail of the arrow to the atom or
bond at the head of the arrow
Rules for Resonance Forms
Rule 3 – Different resonance forms of a substrate do not
have to be equivalent
Rules for Resonance Forms
Rule 4 – Resonance forms obey normal rules of valency
(follow the octet rule)
Rule 5 – The resonance hybrid is more stable than any
individual resonance form
• Resonance leads to stability
2.6 Drawing Resonance Forms
In general any three-atom grouping with a p orbital on each atom
has two resonance forms:
•
•
The atoms X,Y, and Z in the general structure might be C,N,O,P,or S
The asterisk (*) on atom Z for the resonance form on the left might
mean that the p orbital is:
•
•
•
Vacant
Contains a single electron
Contains a lone pair of electrons
Drawing Resonance Forms
Reaction of pentane-2,4-dione with a base
•
•
H+ is removed
An anion is formed
Resonance of the anion product:
Worked Example 2.2
Drawing Resonance Forms of an Anion
Draw three resonance forms for the carbonate ion, CO32-.
Worked Example 2.2
Drawing Resonance Forms of an Anion
Strategy
• Look for one or more three-atom groupings that
contain a multiple bond next to an atom with a p
orbital. Then exchange the positions of the multiple
bond and the electrons in the p orbital
•
In the carbonate ion, each of the singly bonded oxygen
atoms with its lone pairs and negative charge is next to
the C=O double bond, giving the grouping O=C-O:-
Worked Example 2.2
Drawing Resonance Forms of an Anion
Solution
• Exchanging the position of the double bond and an electron
lone pair in each grouping generates three resonance
structures:
2.7 Acids and Bases: The Brønsted-Lowry
Definition
Two frequently used definitions of acidity in organic chemistry
The Brønsted-Lowry definition
• Lewis definition
Brønsted-Lowry acid
• A substance that donates a hydrogen ion (proton; H+) to a base
Brønsted-Lowry base
• A substance that accepts a hydrogen ion (proton; H+) from an acid
•
Acids and Bases:
The Brønsted-Lowry Definition
Water can act either as an acid or as a base
Acid and Base Strength
Acidity constant Ka
• A measure of acid strength in water
• The concentration of water, [H2O], remains nearly constant
at 55.5 M at 25 °C
• For any weak acid HA, the acidity constant is given by the
expression Ka

 A - + H3O+
HA + H2O 

H3O+   A - 
K a = K eq H2O =
HA 
Acid and Base Strength

 A - + H3O+
HA + H2O 

• Equilibria for stronger acids favor the products (to the right)
and thus have larger acidity constants
• Equilibria for weaker acids favor the reactants (to the left)
and thus have smaller acidity constants
Acid strengths are normally expressed using pKa values
pKa
• The negative common logarithm of the Ka
pKa = -log Ka
• Stronger acids (larger Ka) have smaller pKa
• Weaker acids (smaller Ka) have larger pKa
Acid and Base Strength
Acid and Base Strength
• Strong acid (BrØnsted-Lowry)
•
•
•
One that loses H+ easily
Conjugate base holds on to the H+ weakly (weak
base)
Strong acid has weak conjugate base
• Weak acid (BrØnsted-Lowry)
•
•
•
One that loses H+ with difficulty
Conjugate base holds on to the H+ strongly (strong
base)
Weak acid has strong conjugate base
Predicting Acid-Base Reactions from pKa
Values
• Acid-base reactions always proceed in the direction of the
weakest acids or bases
Worked Example 2.4
Predicting Acid Strengths from pKa Values
Water has pKa = 15.74, and acetylene has pKa = 25. Which
is the stronger acid? Does hydroxide ion react with
acetylene?
Worked Example 2.4
Predicting Acid Strengths from pKa Values
Strategy
• In comparing two acids, the one with the higher
pKa is weaker – that side of the reaction is favored
• The reaction will NOT proceed as written
2.10 Organic Acids and Organic Bases
Most biological reactions involve organic acids and organic bases
Organic acid
•
Has positively polarized hydrogen atom(s)
•
Two main kinds of organic acids
1.
2.
Contains a hydrogen atom bonded to an oxygen atom (O-H)
Contains a hydrogen atom bonded to a carbon atom next to a C=O
double bond (O=C-C-H)
Organic Acids and Organic Bases
Conjugate base
• Anion stabilized by having its negative charge on a highly
electronegative atom, OR
• Anion stabilized by resonance
Methanol
Acetic Acid
Acetone
Organic Acids and Organic Bases
• Conjugate bases from methanol, acetic acid, and acetone
• The electronegative oxygen atoms stabilize the negative
charge in all three
Organic Acids and Organic Bases
Carboxylic acids
• Contain the –CO2H grouping
• Occur abundantly in all living organisms
• Involved in almost all metabolic pathways
• At cellular pH of 7.3 carboxylic acids are usually dissociated
and exist as their carboxylate anions, –CO 2
Organic Acids and Organic Bases
Organic bases
• Characterized by the presence of an atom with a lone pair of
electrons that can bond to H+
• Nitrogen-containing compounds are common organic bases
and are involved in almost all metabolic pathways
• Oxygen-containing compounds can act both as acids and as
bases
2.11 Acids and Bases: The Lewis Definition
The Lewis definition is broader than the Brønsted-Lowry definition
Lewis acid – an electrophile
A substance with a vacant orbital that can accept an electron pair
from a base
• All electrophiles are Lewis acids
•
Lewis base – a nucleophile
A substance that donates an electron lone pair to an acid
• All nucleophiles are Lewis bases
•
Acids and Bases: The Lewis Definition
Lewis Acids and the Curved Arrow Formalism
• To accept an electron pair a Lewis acid must have either:
• A vacant, low-energy orbital
• A polar bond to hydrogen so that it can donate H+
• Various metal cations, such as Mg2+, are Lewis acids
because they accept a pair of electrons when they form a
bond to a base
Acids and Bases: The Lewis Definition
• Compounds of group 3A elements, such as BF3 and AlCl3
are Lewis acids
• Have unfilled valence orbitals and can accept electron
pairs from Lewis bases
• Many transition metals, such as TiCl4, FeCl3, ZnCl2, and
SnCl4 are Lewis acids
Acids and Bases: The Lewis Definition
Curved arrow formalism
Acids and Bases: The Lewis Definition
Further examples of Lewis acids
Acids and Bases: The Lewis Definition
Lewis bases
• A compound with a pair of nonbonding electrons that it can
use in bonding to a Lewis acid
• Definition of Lewis base similar to Brønsted-Lowry definition
• H2O acts as a Lewis base
• Has two nonbonding electrons on oxygen
Acids and Bases: The Lewis Definition
• Most oxygen- and nitrogen- containing organic compounds
are Lewis bases
•
They have lone pair electrons
Acids and Bases: The Lewis Definition
• Some compounds can act as both acids and bases
• Some compounds have more than one atom with a lone pair
of electrons
•
•
•
Reaction normally occurs only once in such instances
The more stable of the two possible protonation products is
formed
Occurs with carboxylic acids, esters, and amides
Worked Example 2.6
Using Curves Arrows to Show Electron Flow
Using curved arrows, show how acetaldehyde,
CH3CHO, can act as a Lewis base.
Worked Example 2.6
Using Curves Arrows to Show Electron Flow
Strategy
• A Lewis base donates an electron pair to a Lewis
acid
•
•
Locate the electron lone pairs on acetaldehyde
Use a curved arrow to show the movement of a pair
toward the H atom of the acid
Worked Example 2.6
Using Curves Arrows to Show Electron Flow
Solution
2.12 Noncovalent Interactions between
Molecules – PLEASE READ!!
Noncovalent Interactions
• Also called intermolecular forces or van der Waals forces
• Types of noncovalent interactions include:
•
Dipole-Dipole forces
•
Occur between polar molecules as a result of electrostatic
interactions among dipoles
Forces are either attractive or repulsive
•
Attractive geometry is lower in energy and therefore predominates
•
Attractive
Repulsive
Noncovalent Interactions between Molecules
•
Dispersion forces
•
•
•
•
Attractive dispersion forces in nonpolar molecules are caused by
temporary dipoles
One side of the molecule may have a slight excess of electrons
relative to the opposite side, giving the molecule a temporary dipole
Temporary dipole in one molecule causes a nearby molecule to
adopt a temporarily opposite dipole resulting in a small attraction
between the two molecules
Arise because the electron distribution within molecules is constantly
changing
Noncovalent Interactions between Molecules
•
Hydrogen Bonds
• A weak attraction between a hydrogen atom bonded to an
electronegative O or N and an electron lone pair on another O
or N atom
• Strong dipole-dipole interaction involving polarized O-H and
N-H bonds
• Important noncovalent interaction in biological molecules
Noncovalent Interactions between Molecules
•
Effects of Hydrogen Bonding
•
•
•
Causes water to be a liquid rather than a gas at room
temperature
Holds enzymes in the shapes necessary for catalyzing biological
reactions
Causes strands of deoxyribonucleic acid (DNA) to pair up and
coil into a double helix
Noncovalent Interactions between Molecules
Hydrophilic (water-loving)
• Dissolves in water
• Table sugar
•
Has ionic charges, polar –OH groups, in its structure
Hydrophobic (water-fearing)
• Does not dissolve in water
• Vegetable oil
•
Does not have groups that form hydrogen bonds