Lecture 8. Acid and Base
theory. Protolitic processes
Prepared by PhD Halina Falfushynska
Acid-Base Reactions
• Termed neutralization reactions.
• Involve an acid and a base.
• Double replacement (or metathesis)
reaction
• A molecular compound (water) is a
common product along with a salt (ionic
compound).
Common Acids and Bases
All the other acids and bases are weak
electrolytes (important for net ionic
equations).
• Definitions of acids and bases
–Arrhenius acid - produces H+ in
solution
–Arrhenius base - produces OH in
solution
–More inclusive definitions:
• Brønsted acid - proton donor
• Brønsted base - proton acceptor
– Examples of a weak base and weak acid
• Ammonia with water:
• Hydrofluoric acid with water:
• Types of acids
– Monoprotic: one ionizable hydrogen
HCl + H2O  H3O+ + Cl
– Diprotic: two ionizable hydrogens
H2SO4 + H2O  H3O+ + HSO4
HSO4 + H2O  H3O+ + SO42
– Triprotic: three ionizable hydrogens
H3PO4 + H2O  H3O+ + H2PO4
H2PO4 + H2O  H3O+ + HPO42
HPO42 + H2O  H3O+ + PO43
– Polyprotic: generic term meaning
more than one ionizable hydrogen
• Types of bases
– Monobasic: One OH group
KOH  K+ + OH
– Dibasic: Two OH groups
Ba(OH)2  Ba2+ + 2OH
Acid-Base Neutralization
• Neutralization: Reaction between an acid
and a base
Acid + Base  Salt + Water
Molecular equation:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Ionic equation:
H+(aq)+ Cl(aq) + Na+(aq) + OH(aq)
 Na+(aq) + Cl(aq) + H2O(l)
Net ionic equation:
H+(aq) + OH(aq) H2O(l)
Solutions of acetic acid and lithium
hydroxide are mixed. Write the net ionic
reaction.
HC2H3O2 (aq) + OH (aq)
C2H3O2 (aq) + H2O(l)
Acids & Bases
Acids:
acids are sour tasting
Arrhenius acid: Any substance that, when dissolved
in water, increases the concentration of hydronium
ion (H3O+)
Bronsted-Lowry acid: A proton donor
Lewis acid: An electron acceptor
Bases:
bases are bitter tasting and slippery
Arrhenius base: Any substance that, when dissolved
in water, increases the concentration of hydroxide
ion (OH-)
Bronsted-Lowery base: A proton acceptor
Lewis acid: An electron donor
Lone Hydrogen ions do not exist by themselves
in solution. H+ is always bound to a water
molecule to form a hydronium ion
Brønsted-Lowry Theory of Acids & Bases
Conjugate Acid-Base Pairs
General Equation
Brønsted-Lowry Theory of Acids & Bases
Brønsted-Lowry Theory of Acids & Bases
Brønsted-Lowry Theory of Acids & Bases
Notice that water is both an
acid & a base = amphoteric
Reversible reaction
Conjugate Acid-Base Pairs
Conjugate Acid-Base Pairs
Acids & Bases
For the following Identify the conjugate acid and the
conjugate base. The conjugate refers to the acid or
base produced in an acid/base reaction. The acid
reactant produces its conjugate base (CB).
CB
CA
a. Al(OH)3 + 3 HCl  AlCl3 + 3 H2O
b. Ba(OH)2 + 2 HC2H3O2 
CB
CA
Ba(C2H3O2)2 + 2
CB
CA
c. 2 KOH + H2SO4  K2SO4 + 2 H2O
CA
d. NH3 + H2O  NH4
CB
+ +
OH-
H2O
Water Equilibrium
“Dissociation” of water
22
The Equilibrium Constant for Water
Dissociation, Kw
H2O  H+ + OH
K eq 
+

H
OH
  
 H 2O 

+

H
OH
  
1.0
 1014.0  K w
[H+]x[OH- ] = 1.8 x 10-16 x 55.5 = 10-14
pH = log10(H+)
Water Equilibrium
 Water or water solutions in which [H+] = [OH-] = 10-7 M
are neutral solutions.
 A solution in which [H+] > [OH-] is acidic
 A solution in which [H+] < [OH-] is basic
pH
A measure of the hydronium ion
• The scale for measuring the hydronium ion concentration [H3O+]
in any solution must be able to cover a large range. A logarithmic
scale covers factors of 10. The “p” in pH stands for log.
• A solution with a pH of 1 has [H3O+] of 0.1 mol/L or 10-1
• A solution with a pH of 3 has [H3O+] of 0.001 mol/L or 10-3
• A solution with a pH of 7 has [H3O+] of 0.0000001 mol/L or 10-7
pH = - log [H3O+]
Manipulating pH
Algebraic manipulation of:
pH = - log [H3O+]
allows for:
[H3O+] = 10-pH
If pH is a measure of the hydronium ion
concentration then the same equations could be
used to describe the hydroxide (base)
concentration.
[OH-] = 10-pOH
pOH = - log [OH-]
thus:
pH + pOH = 14 ; the entire pH range!
pH
[H+]= 10-1M
pH = – log [H+]
Log scale
means 10X
change per
unit!
[H+]= 10-9 M
27
Henderson Hasselbalch equation
H2O + CO2  H2CO3  H+ + HCO3Looking at 2nd half of equation:
K (equilibrium constant) = [H+] x [HCO3-]
[H2CO3]
Rearranging and taking negative logs:
pH = pK + log10 [HCO3-]
[H2CO3]
Assuming the 1st half of the equation is at equilibrium:
[H2CO3] = 0.03 x Pa.CO2
Then, the Henderson Hasselbalch equation is:
pH = pK + log10
[HCO3-]
0.03 x Pa.CO2
28
• Calculate the pH of the following
• 0.25 M NH3 and 0.40 M NH4Cl
(Kb = 1.8 x 10-5)
• Ka = 1 x 10-14
1.8 x 10-5
Ka = 5.6 x 10-10
• remember its the ratio base over acid
 0.25 M 
pH = -log(5.6 x 10 ) + log 

 0.40 M 
-10
lpH = 9.05
Indicators
• Weak acids that change color when they
become bases.
• weak acid written HIn
• Weak base
• HIn
H+ + Inclear
red
• Equilibrium is controlled by pH. End point when the indicator changes color.
• Try to match the equivalence point
Indicators
• Since it is an equilibrium the color change is
gradual.
• It is noticeable when the ratio of
[In-]/[HI] or [HI]/[In-] is 1/10
• Since the Indicator is a weak acid, it has a Ka.
• pH the indicator changes at is.
• pH=pKa +log([In-]/[HI]) = pKa +log(1/10)
• pH=pKa - 1 on the way up
Indicators
• pH=pKa + log([HI]/[In-]) = pKa + log(10)
• pH=pKa+1 on the way down
• Choose the indicator with a pKa 1 more
than the pH at equivalence point if you
are titrating with base.
• Choose the indicator with a pKa 1 less
than the pH at equivalence point if you
are titrating with acid.