Functional Groups, Orbitals, and
Geometry
Resonance Structures
Bond Polarity - Part I
 A bond is polar when the charge is not
equally shared between the two atoms.
 The more electronegative atom will
have a partial negative charge (δ-).
The arrow
shows the
dipole
moment.
Here we
show
partial
charges.
Acids and Bases-Definitions
 Arrhenius acid: A substance
which dissolves in water to
produce H+.
 Brønsted-Lowry acid: a proton
donor
 H+ is a proton.
 Lewis acid: an electron pair
acceptor
Arrhenius Acids and Bases
 Arrhenius acid:
which dissolves
produce H+.
 Arrhenius base:
which dissolves
produce OH-.
A substance
in water to
A substance
in water to
 Limited to aqueous solutions.
 Does not explain a reaction such as
NH3(g) + HCl(g)  NH4Cl(s)
Brønsted-Lowry Acids and
Bases
 B-L acids are proton donors.
 B-L bases are proton acceptors.
 The emphasis is on the transfer of
the H+. This links acids and
bases.
 A B-L acid HB has a conjugate base:
HB  H+ + B:This is the equation for HB acting as
an acid.
Brønsted-Lowry Acids and
Bases
 HB  H+ + B: This is the equation for HB acting as an
acid. B:- is the conjugate base.
 B:- +H2O  HB + OH This is the equation for B- acting as a
base in water.
 B:- + HA  HB + A This is the equation for B- acting as a
base with an acid other than water.
 Be able to write these types of equations
for any B-L acid or base.
Brønsted-Lowry Acids and
Bases
 Ammonia acting as an acid:
 NH3  NH2- + H+
 Ammonia acting as a base:
 NH3(aq) + H2O  NH4+(aq) + OH(aq)
 What is the conjugate acid and what is
the conjugate base of ammonia?
 Is ammonia a conjugate acid or base?
Acid Strength and pKa
HB
H+ + B:-
Ka = acid dissociation constant
Ka = [H+][B-]
[HB]
pKa = -log Ka
 The more completely an acid dissociates in
water, the stronger it is.
 The stronger the acid, the larger its Ka and the
smaller its pKa.
Comparing Acid Strengths
 Which is the stronger acid,
ammonia or water?
 There are two ways to find an
answer:
 The quantitative way: compare pKa
values.
 The qualitative way: compare the
stabilities of the conjugate bases.
Comparing Acid Strengths
 The quantitative way: compare
pKa values.
 NH3  NH2- + H+ pKa = 36
 H2O(l)  H+(aq) + OH-(aq)
pKa = 15.7
 Water is the stronger acid.
Comparing Acid Strengths
 The qualitative way: compare
stabilities of the conjugate bases.
 NH3  NH2- + H+
 H2O(l)  H+(aq) + OH-(aq)
 The more stable the conjugate base is
in water, the stronger the acid.
 The amide ion is such a strong base it
cannot exist in water, therefore ammonia is
the weaker acid.
Comparing Acid Strengths
 You will find it very helpful in
studying organic chemistry to have
a good idea of the relative strengths
of some of the more common
compounds acting as acids.
 Please become VERY familiar with
Table 1-5.
Comparing Acid Strengths by
Comparing Structures
 How does the structure of a compound
affect its acid/base properties?
 Look at the stability of the conjugate
base. The more stable the conjugate
base, the stronger its acid.





Electronegativity
Size/polarizability
Resonance Stabilization
Induction
Hybrid orbital containing electrons
Comparing Acid Strengths by Comparing
the Stabilities of the Conjugate Bases
 Electronegativity (e.n.)
 A more electronegative atom holds
negative charge more easily. Many
bases are anions. The more stable
the anion, the weaker the base:
 e.n.(C) < e.n.(N)<e.n.(O)<e.n.(F)
 Base strength: CH3->NH2->OH->F Acid strength: CH4<NH3<H2O<HF
Comparing Acid Strengths by Comparing
the Stabilities of the Conjugate Bases
 Size
 A larger anion is more stable:
 Size/stability: F- < Cl- < Br- < I Acid strength: HF < HCl < HBr < HI
 Base strength: F- > Cl- > Br- > I-
Comparing Acid Strengths by Comparing
the Stabilities of the Conjugate Bases
 Resonance Stabilization
 An anion stabilized by resonance
has a stronger conjugate acid.
Comparing Acid Strengths by
Comparing Structures
 Induction
 Look at nearby atoms.
 Electronegative atoms “pull” electron
density away (induction). This can stabilize a
negative charge. (Note: they must be very
close to the negative charge to be effective.)
Trichloroacetic acid is
stronger than acetic
acid.
more stable
Comparing Acid Strengths by
Comparing Structures
 Hybrid orbital containing electrons
 Acetylene (H-C≡C-H), believe it or
not, can act as an acid with certain
really strong bases.
 H-C≡C-H + B:-  H-C≡C:- + HB
 The sp orbital is short (50% s
character) and stabilizes the anion
by holding the electrons closer to
the nucleus.
Lewis Bases and Acids
 Lewis looked at acid/base behavior from
the viewpoint of the bonds that are
formed instead of the transfer of a proton.
Lewis Bases and Acids
 Lewis bases have nonbonding electrons
that can be donated to form new bonds.
 Lewis bases are nucleophiles (lovers of nuclei
+++).
 Lewis acids accept these electrons.
 Lewis acids are electrophiles (lovers of
electrons ---).
Two Bases Worth Knowing
 NaH and NaNH2
sodium
hydride
sodium
amide
sodium
methoxide
sodium
ethoxide
Given the reactants, be able to write the products of any
acid/base reaction!
Identifying Bases






NaH and NaNH2
Amines
Hydroxide ion, OHAlkoxide ions, e.g. CH3OAlcohols
Water
Identifying Acids






Inorganic (the seven strong acids)
Carboxylic acids
Phenols
Alcohols
Water
These are pretty much in order
from strongest to weakest.