Acid-Base Chemistry
Assignment #7
Acids and Bases - 3 Definitions

Arrhenius

Bronsted-Lowry

Lewis
Arrhenius Definition

Acid = proton donor
HA = H+ + A-

Base = hydroxide donor
BOH = B+ + OH-

Dilemma: NH3
Bronsted-Lowry Definition

Acid = proton donor

Base = proton acceptor
NH3 + H+ = NH4+

Dilemna: ferrocene (organometallics)
Lewis Definition

Acid = electron pair acceptor (electrophile)

Base = electron pair donor (nucleophile)
Acid and Base Strength
Strong acids and Bases completely
dissociate (ionize) in aqueous solution
 EX:
HClaq -> H+aq + Cl-aq

Weak acids and Bases incompletely
dissociate in aqueous solution
 EX:
HC2H3O2aq = H+aq + C2H3O2-aq

Not All Mineral Acids are
Strong!


HCN
HCN = H+ + CN-
Ka = 2.1 x 10-9
HF
HF = H+ + F-
Ka = 6 x 10-4
Polyprotic Acids
Protons are always lost one at a time!
 acids produced by proton loss from
polyprotic acids are weak acids,
characterized by a unique Ka value


H2SO4 = HSO4- + H+ Ka1 not measurable
HSO4- = SO42- + H+
Ka2 = 1.1 x 10-2
Strong Acids
Most mineral acids
 HCl
 H2SO4
 HNO3
 HClO4


Not: HF, H3PO4
Weak Acids

Organic acids (need C and usually have
COOH)
EXAMPLES:
formic acid
acetic acid
propionic acid
Strong Bases

Inorganic hydroxides containing metals
from families IA or IIA

Note: names of these families: Alkali
metals, alkaline earth metals
Weak Bases

NH3, organic amines, and hydroxides other
than group Ia or IIa hydroxides

organic amines contain amino group:
R-NH2 + H+ = R-NH3+
Scheme for Identification of
Acids and Bases
Proton donor
or proton acceptor?
If proton donor,
then acid
Inorganic
If proton acceptor,
Then base
Organic
Strong except
HF, HCN, H2S, H3PO4
Weak acid
Inorganic
Weak acid
Organic
If IA or IIA,
Weak base
then strong
If NH4OH or not IA and IIA
Then weak
Practice







Identify the acid/base nature of the following
compounds:
Hydroxyl amine
Calcium hydroxide
Carbon dioxide
Pthalic acid
Hydrogen sulfide
pyridine
Conjugates

Acids and Bases exist in a conjugate
relationship:
HA = H+ + Aacid
base
BOH = B+ + OHbase acid
Example:

NH4OH = NH4+ + OHbase
acid

HC2H3O2 = H+ + C2H3O2acid
base
Problem
Identify the conjugate acid-base pairs for
each of the following compounds:
 Ammonium hydroxide
 Diethylamine
 Iodic acid
 Formic acid
 HPO4
Amphoterism

Some compounds can function as both acids
or bases depending on the situation
e.g., H2O

HC2H3O2 + H2O = H 3O+ + C2H3O2acid
base acid
base

NH3 + H2O = NH4+ + OHbase acid acid base
Dissociation Constants for Weak
Acids and Bases
Recall for HA:
 Ka = [H+][A-]
[HA]


The bigger Ka, the _____ the [H+] and the
_____ the [HA]
Question
What is the comparatively strongest weak
acid on Table A?
 ANS: iodic acid, Ka = 0.18
 What is the comparatively strongest weak
base on Table B?
 ANS: diethylamine and piperidine are
equally strong, Kb = 0.0013

Conjugate Acid-Base Strength






For HA = H+ + ARecall,
Ka = [H+][A-]
[HA]
HA is the conjugate _____ and A- is its conjugate
____
HA is a ____ ____
A- is a ____ ____
If HA is a relatively strong weak acid, then A- is a
comparatively ____ ____ base
Problem

HSO4- is
a) the conjugate acid of SO4-2
b) a strong acid
c) the conjugate base of H2SO4
d) a strong base
e) the conjugate acid of H2SO4
Problem:

Which of the following is/are amphoteric:
a) H2PO4b) C2H3O2c) CH3CH2NH2
d) HCO3e) CH3CH(NH2)COOH
Conjugate Acid-Base Strength

The stronger the
conjugate acid is, the
weaker the conjugate
base is and vice versa
HA
A-
Salts

There are four kinds of salts:
 salts
of weak acids
example: sodium acetate
 salts of weak bases
example: ammonium chloride
 salts of strong acids and strong bases
example: sodium chloride
 salts of weak acids and weak bases
example: ammonium acetate
Identification of Salts
Salts hydrolyze in water:
Salt + water  acid + base

EX:
NaCl + HOH  Na+ + OH- + H+ + Cl-
Identification of Salts
Salts are obtained by reacting acids and
bases:
Acid + Base = Salt + water

Note: this is just the reverse of the
hydolysis reaction:
Salt + water = Acid + Base

Identification of Salts


So, salts are classified based on their parent acids
Their parents are the acids and bases used to form
them
EX: NaCl
NaOH + HCl  Na+ + Cl- + H2O
So, parents are strong acid and strong base and
NaCl is the salt of a strong acid and strong base

Examples:

NaC2H3O2 – salt of ________________

NH4Cl – salt of ___________________

NH4COOH – salt of _______________

LiF – salt of _____________________
Homework Problem #1:

Classify each of the following as a strong or
weak acid, base, or salt (identify parents):
CsOH
AgOH
sodium salicylate H2PO4HClO4
H2CO3
ferric hydroxide
oxalic acid
NH4C2H3O2
trimethylamine
The Autoionization of Water
HOH(l) + HOH(l) = H3O+ + OHhydronium ion
 This is an equilibrium process and is
characterized by an equilibrium constant,
Kw :

Kw = [H3O+][OH-] = 10-14 at 250C
Kw DOES vary with Temperature
0
Temperature, C Kw
0
1.13 x 10
-15
lower
-14
25
1.008 x 10
60
9.614 x 10-14
higher
The Relationship between [H+]
and [OH-]

Kw = [H+][OH-] = 10-14
Useful Equivalent forms:
 [H+] = 10-14/[OH-]

[OH-] = 10-14/[H+]
The pH Scale

pH = -log [H+]

[H+] = antilog[-pH]

pH of pure water = 7
; no units
A Brief Review of LOG Math

Taking a logarithm corresponds to
answering the question:
To what power do I raise 10 in order to
represent the number of interest?
log 100 => 10? = 100

NOTE: LOG is not same thing as LN
natural log is based on e?= number
A Brief Review of LOG Math

log (ab) = log a + log b

log (a/b) = log a - log b

log ab = b log a
The pOH scale

pOH = -log [OH-]

[OH-] = antilog[-pOH]
; no units
What is the pOH in pure water?
 ANS: pOH = -log (10-7) = 7

How is pH related to pOH?

recall:
Kw = [H3O+][OH-] = 10-14 at 250C
Derivation
 call pKw = -logKw = 14 at 250C


then:
pKw = pH + pOH = 14
The pH of Some Common
Substances
Common Substance
lemons
seawater
blood
urine
saliva
gastric juices
pH
2.3
8.5
7.4
5.5 - 7.0
6.5 - 7.5
1.0 - 3.0
Aspirin and Ibuprofen
COOH
OH

(H3C)2HCH2C
CH2COOH
salicylic acid
ibuprofen
 active
ingredient 0.5%
salicylic acid
O
CH3
O
Clean and Clear
Sensitive Skin Deep
Cleaning Astringent

Advil
O
COOH
acetyl salicylic acid (aspirin)
 active
ingredient
ibuprofen
Orange Juice

Tropicana Pure
Premium with
Calcium
 active
ingredients:
calcium hydroxide,
malic acid and citric
acid
Toothpaste

Aquafresh Whitening
Toothpaste
 active
ingredient:
sodium fluoride

Colgate Maximum
Cavity Protection
Fluoride Toothpaste
 active
ingredient:
0.76% sodium
monofluorophosphate
The Relationship between pH
and pOH
[H+], M pH pOH [OH-], M
10-1
1
??
10-13
-10
-4
??
10
10-7
7
7
??
??
11
3
10-3
10-14
14
??
1
10
10
acidic
neutral
basic
Acidic Solution
high [H+] concentration
 low pH value
value below 7

low [OH-] concentration
 high pOH value
value greater than 7

Basic Solution
high [OH-] concentration
 low pOH value
value less than 7

low [H+] concentration
 high pH value
value greater than 7

pKa and pKb

pKa ≡ - log Ka

pKb ≡ - log Kb

pKa * pKb = pKw
Conjugate Acid-Base Strength



The stronger the conjugate
acid is, the weaker the
conjugate base is and vice
versa
Ka * Kb = 10-14
pKa + pKb = 14
HA
A-
pKa , pKb, and Weak Acid/ Base
Strength



The lower the pKa the
______ the weak acid
The higher the pKa, the
______ the weak acid
The lower the pKa of a
weak acid, the ______ the
pKb of its conjugate weak
base and the _____ its
conjugate base
HA
A-
pKa * pKb = 14
Question

Which is the comparatively stronger weak
acid ammonium or pyridinium?
Strong Acids and Strong Bases



HA → H+ + ABOH → B+ + OH-
Strong acids and strong
bases completely
dissociate in water so
their concentration
gives us the [H+] in
solution directly
sea slugs secrete H2SO4
sea squirts squirt HNO3
Photographs from Atkins, P.W. Molecules; W.H. Freeman: New York, 1987.
Problem:

Determine the pH of the following
solutions. Are the solutions acidic or basic?
a) 0.001 M HCl solution
b) a solution whose [OH-] is 10-3 M
c) 0.0001 M NaOH solution

ANS:
a) pH 3, acidic; b) pH 11, basic; c) pH 10,
basic
Homework Problem #1:

Calculate the hydrogen ion concentration
and pH of a solution prepared by placing
11.5 g of HClO4 (perchloric acid; FW 100)
in a 500 mL volumetric flask subsequently
filled to the mark with water.
(Hint: What kind of acid is HClO4???)
Problem:

What are the pH and pOH of a solution
prepared by mixing 25 mL of 0.20 M NaOH
with 60 mL of 0.10 M HCl?
(Suggestion: Draw a picture)

ANS: pH = 1.93; pOH = 12.07
Weak Acids and Bases
H+
A-
HA =
+
 BOH = B+ + OH

Fire ants venom contains formic acid
Weak acids and weak bases incompletely
dissociate so their concentration does not
provide meaningful insight into the aqueous
pH of their solutions
Photograph from Atkins, P.W. Molecules; W.H. Freeman: New York, 1987.
Weak Acids (Table A)
Weak Acid
Ka
HF
7.2 x 10-4
-4
HCOOH (formic) 1.8 x 10
-5
HC2H3O2
1.8 x 10
HCN
6.2 x 10-10
-10
H3BO3 (boric)
5.8 x 10
Weak Acids

The stronger a weak acid the greater the
[H+] in solution

The weaker a weak acid, the greater the
[HA] in solution

Recall: pKa = -log Ka
Acid-Base Strength
A-
HA
Reminder: Weak Acids
Weak Acid
Ka
pKa
Acetic
1.75 x 10-5
4.76
Hydrocyanic
2.1 x 10
-9
8.68
The stronger the weak acid, the larger the
Ka and the lower the pKa
Calculating the pH of Weak
Acids

the dissociation of weak acids determines
the equilibrium concentration of H+ and
therefore the pH

the dissociation of weak acids is
characterized by the equilibrium constant Ka

Derivation
Problem:

What is the pH and pOH of a solution of
0.05 M butyric acid (CH3CH2CH2COOH)
given the pKa for butyric acid 4.81?

Q: Where found in nature?
Weak Bases

The stronger a weak base the greater the
[OH-] in solution

The weaker a weak base, the greater the
[BOH] in solution
Define: pKb = -log Kb
 Derivation

Problem

What is the pH of an 0.026 M solution of
hexamethylenetetramine (Kb 10-9)?
Salts

There are four kinds of salts:
 salts
of weak acids
example: sodium acetate
 salts of weak bases
example: ammonium chloride
 salts of strong acids and strong bases
example: sodium chloride
 salts of weak acids and weak bases
example: ammonium acetate
Salts of Strong Acids and Bases

Example: NaCl

NaCl + H2O -> Na+aq + Cl-aq + H2O

pH  7

These salts do not directly perturb the water
equilibrium
Salts of Weak Acids

Example: NaC2H3O2

NaC2H3O2 + H2O = Na+ + HC2H3O2 + OH-

pH = 0.5(14 + pKa + log[salt])

solutions are weakly basic
Salts of Weak Acids
Derivation
 pH = 0.5(14 + pKa + log[salt])


the weaker the weak acid, the more basic
the solution (the more tightly the weak acid
holds onto the H+ and the higher the [OH-]
in solution
Question:

Solutions of which salt would be more basic
- sodium acetate or sodium cyanide?
Salts of Weak Bases

Example: NH4Cl

NH4Cl + H2O = NH4OH + H+ + Cl-


Derivation
pH = 0.5(14 - pKb - log[salt])

solutions are weakly acidic
Salts of Weak Bases
pH = 0.5(14 - pKb - log[salt])
 The weaker the weak base, the lower the
solution pH (the more tightly the weak base
holds onto the OH- and the higher the [H+]
in solution)

Salts of Weak Acids and Weak
Bases

Example: NH4C2H3O2

NH4C2H3O2 + H2O = NH4OH + HC2H3O2

Derivation
pH = 0.5(14 + pKa - pKb)


solution pH depends on the relative strength of
the acid vs. the base
Homework Problem #2
Calculate the pH of an 0.1 M solution of
each of the following compounds:
 A) sodium acetate
 B) sodium nitrate
 C) sodium hydroxide
 D) hydrogen fluoride
 E) pyridine

Problem:

Identify the following compounds in terms
of their acid/base properties and predict
whether aqueous solutions of these
compounds will be acidic, basic, or neutral:
a) sodium cyanide
b) ammonium nitrate
c) potassium nitrate
Buffers

DEFINITION:
a solution containing both a weak acid/base
and its salt which resists change in pH due
to:
 temperature
 dilution
 and
addition of SMALL amounts of strong acid or
base
Examples:
HC2H3O2 and NaC2H3O2
 HCOOH and HCOONa
 NaH2PO4 and Na2HPO4
 pyridine and pyridinium chloride
 ammonia and ammonium chloride

Problem:

Solutions are made by combining equal
volumes of the following. Which is/are a
buffer(s)?
a) 0.1 M NH4Cl + 0.1 M NH4+
b) 0.1 M HF + 0.05 M NaOH
c) 0.05 M HF + 0.1 M NaOH
d) 0.1 M NaF + 0.05 M HCl
e) 0.1 M NaF + 0.05 M Na+
Henderson Hasselbalch Equation

pH = pKa + log [conj. base/conj. acid]
or
pH = pKa + log [salt/acid]
for a weak acid and its salt
Recall:
pKa = -log Ka
 derivation

Problem:

Calculate the pH of a solution that is
0.25 M sodium acetate and 0.30 M
acetic acid given Ka = 1.8 x 10-5 for
acetic acid.
(hint: what is the pKa?)

ANS:
pH = pKa + log [salt/acid]
pH = 4.74 + log(0.25/0.30) = 4.74 - 0.08
pH = 4.66
Question:

If we have a base buffer containing
ammonia and ammonium chloride, what is
the correct form of the HendersonHasselbach equation based on these
species?

pH = pKa + log [ammonia]/[ammonium
chloride]
Problem:
Calculate the pH of the solution that results
when 200 mL of 0.300 M ammonium
hydroxide are mixed with 250 mL of 0.150
M ammonium chloride.
 ANS:
pH = pKa + log (conj. base/conj. acid)
= 9.25 + log [ammonia/ammonium]
= 9.25 + log (0.06/0.0375) =
= 9.25 + 0.20 = 9.45

Buffers - Effect of Dilution

Consider the pH of the solution that results
from mixing 100 mL 0.1 M NaH2PO4 and
100 mL 0.1 M Na2HPO4? What is the pH if
the solution is diluted by a factor of 2? 5?
Generalize your findings.
Buffers - Effect of Temperature
Name
pKa @ 200C
pKa/0C
MES
6.15
-0.011
HEPES
7.55
-0.014
Tris
8.30
-0.031
Phosphoric acid
(K2)
7.21
-0.0028
Buffers Calbiochem Corp., Doc. No. CB0052-591; Perrin & Dempsey
Buffers for pH and Metal Ion Control Chapman & Hall: London, 1979.
Comparison - Effect of Addition of
SMALL Amount of Strong Acid or Base
Buffer itself (100 mL 0.1 M NaH2PO4 and
100 mL 0.1 M Na2HPO4) pH 7.20
vs. 200 mL water pH 7.00
 Upon addition 0.005 moles strong acid,
buffer pH 6.72
water pH 1.60
 Upon addition 0.005 moles strong base,
buffer pH 7.68
water pH 10.40

Preparing Buffers - Useful
References

Perrin & Dempsey “Buffers for pH and
Metal Ion Control” New York: Wiley,
1974.

Chemical Company booklets.
Example: Calbiochem Doc. No. CB0052591
Practical Preparation

Practical:
 Identify
 Prepare
reagent based on pKa
appropriate molarity and add
NaOH/HCl to adjust pH
Special Types of Buffers

GOOD Buffers - temperature resistant

Volatile Buffers - can be removed by freeze
drying

Universal Buffers - wide effective pH range
GOOD Buffers


zwitterionic - have
both amino and
sulfonyl groups
EXAMPLES:
MES pK 6.15
HEPES pK 7.55
HEPES:
MES:
HOCH2CH2N
O
NCH2CH2SO3H
NCH2CH2SO3H
GOOD Buffers
pK typically 6-8 (physiological pH)
 No complexation with metal ions (no
inhibition of enzymes)
 High aqueous solubility
 Minimal salt effects
 No UV-vis absorption (240-280 nm)

Desirable Characteristics of
GOOD Buffers
pH independent of temperature
 Compare with TRIS:

Temperature, C
pH
4 (cold room)
8.8
20 (room temp)
7.0
37 (incubation)
5.95
Volatile Buffers

Can be removed by simple evaporation or
lyophilization

good for electrophoresis or preparative ion
exchange chromatography
Volatile Buffers
 EXAMPLES:
ammonium
acetate pH 4-6
pyridinium
formate pH 3-6
ammonium
carbonate pH 8-10
Universal Buffers
Mixture of two or more buffers
 Effects of buffers are additive

 greater
buffering capacity
 wider effective pH range

EXAMPLE:
citric acid 3.13, 4.76, 6.40
phosphoric acid 2.15, 7.20, 12.15
boric acid 9.24, 12.74, 13.80
Amino Acids

Given:
 pKa
(COOH) 2.3
 pKa (NH3+) 9.6

What form do amino
acids assume at pH 7?
Titrimetry

Purpose:
 Determine
concentration of an acid or base of unknown
concentration (MAVA = MBVB)
 Identification of unknown acid or base based on pKa
(pH = pKa at ½ volume at equivalence point)

Method: volumetrically using biuret

At endpoint: moles acid = moles base
Ma Va = Mb Vb
Terminology
Titrant = standardized strong acid or base
delivered from a biuret
 Standardized = concentration made known
both in terms of accuracy and precision
 Endpoint=pH at which visual indicator
changes color
 Equivalence point=pH at which moles of
acid equal moles of base

Typical Experimental
Methodology - Weak Acid

Standardize titrant (NaOH)
titrate NaOH with KHP of known
concentration

Titrate unknown (weak acid)
titrate unknown with standardized titrant
Indicators
Organic weak acids that have different
colors in their acid and conjugate base
forms
 EX:
phenolphthalein
HA = H+ + Acolorless
pink

Indicators

pH = pKa + log[A-/HA]

Your eye can detect color for 10-fold excess of A/HA

At equivalence point pH changes rapidly

Bottom line: endpoint may not be equivalence
point if indicator pKa not near equivalence point
Titration of a Weak Acid

Let’s calculate the pH of the solution
produced by adding 0, 10, 20, 25, 50, and
70 mL of 0.1 M sodium hydroxide to 50 mL
of 0.1 M formic acid
Titration of Weak Acid
1 – weak acid
pH = 0.5 (pKa – log[acid]

2 – buffer
pH = pKa + log[base/acid]

3 – equivalence point; salt of
weak acid
pH = 0.5 (14 + pKa + log[salt])

4 – strong base
pH = 14+log[OH-]
4
pH

X
1
3
2
X
Vol. Titrant, mL
Identification of Weak Acid

At equivalence point:
Macid Vacid = Mbase Vbase

At ½ volume corresponding to equivalence
point:
pH = pKa
Titration of a Weak Base

Let’s calculate the pH of the solution
produced by adding 0, 10, 20, 25, 50, and
70 mL of 0.1 M hydrochloric acid to 50 mL
of 0.1 M ammonium hydroxide
Titration of Weak Base
1 – weak base
pH = 14 - 0.5 (pKb – log[base]

2 – buffer
pH = pKa + log[base/acid]

3 – equivalence point; salt of
weak base
pH = 0.5 (14 - pKb - log[salt])

4 – strong acid
pH = -log[H+]
1
X
2
pH

X
3
4
Vol. Titrant, mL