Acids and Bases
pH and Titrations
Overview
 Acid – Base Concepts
 Arrhenius
 Brønsted – Lowry
 Lewis
 Acid and Base Strengths
 Relative Strengths of Acids and Bases
 Molecular Structure and Acid Strength
 Self – Ionization of Water and pH
 Self – Ionization of Water
 Solutions of a Strong Acid or Base
 The pH of a Solution
II.
Theories of acid and bases
A.
Arrhenius theory
1.
Arrhenius Acid as known as a traditional acid
a.
2.
a chemical compound that contains
hydrogens and ionizes in aqueous solution
to form hydrogen ions
Arrhenius base
a.
any chemical that produces hydroxide ions
B.
Bronsted -Lowry theory 1923
(worked independently)
1.
Bronsted Acid
a.
2.
an ion or molecule that is a proton donor
Bronsted base
a.
anything which will accept a proton
3.
Conjugate acid - base pairs
HF +
Acid1
4.
HOH
Base2
===>
H3O+ + FAcid2
Base1
a.
Strong acids have weak conjugate bases
b.
Strong bases and weak conjugate acids
Autoionization - self-ionization
HOH <===> H+ + OH -
Conjugate Acid-Base Pairs
Fig. 17-1, p. 507
Conjugate Acid-Base Pairs
p. 507
Brønsted-Lowry Theory of Acids & Bases
Conjugate Acid-Base Pairs
General Equation
p. 507
Brønsted-Lowry Theory of Acids & Bases
p. 504
5.
C.
Amphiprotic ability to act as an
acid or base
Lewis theory
1.
Lewis acid
a.
2.
anything which can accept a pair of
electrons
Lewis Base
a.
anything which will donate a pair of
electrons
Lewis Theory of Acids & Bases
p. 506
Objectives
 Properties of acids and bases
 The pH scale
 Distinguish between strong and weak
acids and list the clinical uses of these
acids
 Distinguish between strong and weak
bases and list the clinical uses of these
acids
 Understand neutralisation and the
clinical applications of neutralisation
On page 355 Do questions 4,6,7,10,14,
17,18,22,23,25,26,28
Properties of Acids
Aqueous solutions of acids have a sour taste.
Acids change the color of acid-base indicators.
Some acids react with active metals to release
hydrogen gas.
Acids react with bases to produce salts and water.
Some acids conduct electric current.
PROPERTIES OF ACIDS & BASES
 Acids
Produce hydrogen ions (H+) in H2O
Taste sour
Act as electrolytes in solution
Neutralise solutions containing hydroxide
ions (OH -)
React with several metals
releasing H2(g)  corrosion
React with carbonates releasing CO2(g)
Destroy body tissue
How many foods can you
think of that are sour?
Chances are, almost all of the foods that you
associate with being sour, owe their sour taste
to an acid:
Lemons – citric acid
Grapefruit – citric acid
Apples – malic acid
Sour milk – lactic acid
Vinegar – acetic acid
Grapes – tartaric acid
Strength of acids
A strong acid is one that ionizes completely
in aqueous solutions.
The strength of an acid depends on the polarity of
the bond between hydrogen and the element to
which it is bonded and the ease with which that
bond can be broken.
Acid strength increases with increasing polarity
and decreasing bond energy.
Acids that are weak electrolytes are weak acids.
Strong acids are:
Strong electrolytes
~ 100% ionisation  good conductors
Severe burns to body tissue
*** Stomach lining protected against HCl
by mucus
STRENGTHS OF ACIDS
Strong Acids (very few)
Eg
HCl Hydrochloric Acid
~ Stomach acid
Marieb, Fig 26.11
HNO3 Nitric Acid
~ May be used to cauterise warts
~Drugs, explosives, fertilisers, dyes
H2SO4 Sulphuric Acid
~  conc. to treat stomach hypoacidity
~ Fertilisers, dyes, glues
Factors Affecting Acid Strength
Binary acids:
Bond strength is directly related to the acid
strength (bond size).
HI and HBr have larger bonds lengths and are
more acidic than HF and HCl, even though
fluorine is most electronegative.
For bonds of similar size the acid strength is
related to electronegativity difference.
Oxyacids are acidic substances that contain
oxygen and some other nonmetal, e.g. HNO3
An increase in the electronegativity of an
atom bound to oxygen increases in polarity
of the bond and makes it more acidic.
More oxygen = more polar.
Weak Acids (most acids in nature)
CH3COOH Acetic Acid
~ Vaginal jellies, antimicrobial solution  ears,
plastics, dyes, insecticides
H2CO3 Carbonic Acid
~Bicarbonate buffer system, carbonated drinks
H3PO4 Phosphoric Acid
~ Drugs, fertilisers, soaps, detergents, animal feed
Bases
Produce or cause an increase in
hydroxide ions (OH-) in H2O
Taste bitter
Turn red litmus  blue
Act as electrolytes in solution
Neutralise solutions containing
hydrogen ions (H +)
Have a slippery, ‘soapy’ feel
Destroy body tissue/ dissolve fatty (lipid) material
How many bases can you
think of?





Ammonia
Sodium hydroxide – lye – drain cleaner
Milk of magnesia – Mg(OH)2 – antacid
Aluminum hydroxide – antacid
Baking soda – sodium hydrogen carbonate
4. STRENGTHS OF BASES
Strong Bases
NaOH Sodium Hydroxide
~ Removes grease – drains, ovens
Mg(OH)2 Magnesium hydroxide
~ Antacid
~ Laxative
Al(OH)3 Aluminium hydroxide
~ Antacid
~ Absorbs toxins, gases,
~ Causes constipation
Strength of Bases
 The strength of a base depends on the
extent to which the base dissociates.
 Strong bases are strong electrolytes.
 Strong bases: calcium hydroxide, barium
hydroxide, sodium hydroxide, etc
 Weak bases: ammonia, aniline
 Table 15-4 page 461
Relative Strengths of Acids
and Bases and Extent of
Reaction
 The table of relative
strengths of acids
and conjugate bases
can be used to
predict if a reaction
will produce product.
E.g. Which will

NO3
produce product?
HNO3 + CN or HCN +
Strong bases are:
 Strong electrolytes
 ~ 100% dissociation in water  good
conductors
 Severe damage to skin & eyes
(Group 1A elements)
Weak Bases
Eg
NH3 Ammonia
~ Waste product of protein break down in body.
CO3 2-
In antacids
HCO3 –
In antacids, buffers
HPO4 2-
In buffers
Weak bases are:
Weak electrolytes
Do not contain OH – but react with H2O 
small numbers of OH –
Reaction with Water : Weak bases
NH3(g) +
H 2O
HCO3 – (aq) + H2O
(aq)
NH4 + (aq) + OH – (aq)
H2CO3
(aq)
+
OH-
Acids & Bases
STRONG
_ completely ionized
_ strong electrolyte
_ ionic/very polar bonds
bonds
Strong Acids:
HClO4
H2SO4
HI
HBr
HCl
HNO3
vs
WEAK
_ partially ionized
_ weak electrolyte
_ some covalent
Strong Bases:
LiOH
NaOH
KOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
Dissociation in Water : Strong bases
Metal hydroxides  ions
H 2O
Na+(aq) +
NaOH(s)
OH-(aq)
H 2O
Mg 2 + (aq) +
Mg(OH)2(s)
Al(OH)3(s)
H 2O
Al 3+(aq) +
OH- (aq)
OH- (aq)
5. ACID-BASE
NEUTRALISATION
 Neutralisation Reaction
Acid +
HCl +
H+ +
Base 
NaOH 
OH – 
Salt +
NaCl +
Neutralise each other
Must be equal concentrations
Water
H2O
H2O
 Antacids – clinical applications
(Check for side effects!!)
~ Neutralise excess stomach acid
~ Raise stomach pH > 4
Pepsin inactive
~ Assist with ulcer treatment
~  solubility in H2O but still produce high
% of ions
CaCO 3
2HCl +
CaCO 3 
CO 2 (g)
CaCl 2 + H2O +
~Also a Ca 2 + supplement
Long term overuse  Ca 2 + levels
 risk kidney stones (renal calculi)
Eg
Mg(OH) 2
2HCl +
Al(OH) 3
3HCl +
Milk of Magnesia & in
Mylanta
Mg(OH) 2 
MgCl 2 + 2H2O
In Mylanta
Al(OH) 3 
AlCl 3 + 3H2O
NaHCO 3
Baking Soda
Not recommended!!
HCl + NaHCO 3 
NaCl + H2O + CO2
(g)
~ Elderly tend to OD
 Stomach can ‘explode’
 Weak acids are:
 Weak electrolytes
 Small % ionisation  weak conductors
 Dissociation in Water : Weak acids
Polar covalent molecules
 Mainly stay as molecules
Dissociation in Water : Strong acids
Polar covalent molecules  ions
Eg.
HCl(l)
H2O
H+(aq) + Cl-(aq)
H2O
H+ (aq) + NO3- (aq)
HNO3(l)
H2O
H2SO4(l)
2H+ (aq) + SO42- (aq)
 Dissociation in water : Weak acids (cont)
H2O
H+ (aq) + CH3COO- (aq)
CH3COOH (l)
H2CO3 (l)
H2O
H+ (aq) + HCO3-(aq)
H2O
H3PO4 (l)
H+ (aq) + H2PO4- (aq)
2. THE pH SCALE
Ion Product of Water
Pure H2O at 25°C
Some molecules ionise
H2O  H+ + OH[H+ ] = 1 x 10-7 M = [OH- ]
On the pH scale, values below 7 are acidic, a value of 7 is neutral, and
values above 7 are basic.
The pH scale
The pH scale ranges from 1x 100 to 1 x 10-14 mol/L or
from 1 to 14.
pH = - log [H+]
1 2 3 4 5 6 7 8 9 10 11 12 13 14
acid
neutral
base
Stomach juice: pH = 1.0 – 3.0
Human blood: pH = 7.3 – 7.5
Lemon juice: pH = 2.2 – 2.4
Seawater: pH = 7.8 – 8.3
Vinegar: pH = 2.4 – 3.4
Ammonia: pH = 10.5 – 11.5
Carbonated drinks: pH = 2.0 – 4.0
0.1M Na2CO3: pH = 11.7
Orange juice: pH = 3.0 – 4.0
1.0M NaOH: pH = 14.0
 pH describes [H+ ] & [OH- ]
 Indicates if a fluid is :
0
Acidic
[H+ ] = 100
[OH- ] =10-14
7
Neutral
[H+ ] = 10-7
[OH- ] =10-7
14
Basic
[H+ ] = 10-14
[OH- ] = 100
Using the pH scale
Exponential values for [H+ ] & [OH- ]
inconvenient in a clinical workplace
Simplify  pH scale
 acid-base concentration
p  potential or Power
H  Hydrogen
Acidic solution
[H+ ] > [OH- ]
Neutral solution
[H+ ]
= [OH- ]
Basic solution
[H+ ]
< [OH- ]
Water Equilibrium
Kw = [H+] [OH-] = 1.0 x 10-14
Equilibrium constant for water
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
Autoionization of Water
 Water can act as both an acid and a
base equilibrium is:
H2O + H2O  H3O+ + OH.
Kw = [OH][H3O+] = 1.00 x 1014M2.
Since [OH] = [H3O+] [H3O+] = 1 x 107 M
(called a neutral solution)
 Ion Product of H2O:
[H+ ]
x [OH- ] = [1 x 10-7 ] x [1 x 10-7 ]
* Add exponents
= 1 x 10-14
Acidic [H3O+] > 1.00x107M
Neutral
[H3O+] = 1.00x107M
Basic [H3O+] < 1.00x107M
All acids/bases dissolved in water must obey
equation for the ionization of water.
They either add H3O+ or OH to water.
Most of the acids in this chapter will be stronger
than water and add significantly to the
hydronium ion concentration.
E.g. the hydronium ion concentration of an acidic
solution was 1.00x105 M. What was the [OH]?
E.g. what is the hydronium ion concentration
if the hydroxide concentration was 2.50x103 M?
The pH loop
Fig. 17-3, p. 516
The pH of+ a Solution
+
pH
 pH = log[H3O ] and [H3O ] = 10
 Acidic
pH < 7.00
 Neutral
pH = 7.00
 Basic pH > 7.00
E.g. determine the pH of a solution in which
[H3O+] = 5.40x106 M
E.g.2 determine the pH of a solution in which the
[OH] = 3.33x103 M
E.g.3 determine the pOH of a solution in which
the [OH] = 3.33x103 M
E.g.4 Determine the [H3O+] if the pH of the
solution is 7.35. The term pX is defined in exactly
the same way as pH
Eg.5 What is the pCa if [Ca2+] = 6.44x10-4
E.g. 6 what is pH and [OH] of 0.125 M Ba(OH)2.
[H3O+] of water is small compared to added [H3O+]
from the acid and ignored in the calculation.
pH scale
The scale for measuring the hydronium ion
concentration [H+] in any solution must be
able to cover a large range. A logarithmic
scale covers factors of 10. A solution with
a pH of 1 is 10 times stronger than a
solution with a pH of 2
A solution with a pH of 1 has [H+] of 0.1 mol/L
or 10-1
A solution with a pH of 3 has [H+] of 0.001
mol/L or 10-3
A solution with a pH of 7 has [H+] of 0.0000001
mol/L or 10-7
Manipulating pH
Algebraic manipulation of:
pH = - log [H+]
allows for:
[H+] = 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!
Table 17-3, p. 518
Methods of Measuring pH
pH paper is used that has compounds in it
which are change to different colors for different
pH ranges.
An colored indicator can be placed in the solution
and its color correlated with pH.
HIn(aq) + H2O(l)  H3O+(aq) + In(aq).
E.g. phenolphthalein is colorless in acid form but
pink in basic form.


The pH at which they change color depends on
their equilibrium constant.
More accurate and precise
measurements are made with a pH meter.
A combination of voltmeter and electrodes.
Indicators and pH meters
 Acid-Base indicators are compounds
whose color is sensitive to pH.
 The color of an indicator changes as the
pH of a solution changes.
 Indicators come in many colors.
 The pH range over which an indicator
changes color also varies. This pH range
is also called the transition interval.
Indicators and pH meters
con’t
 A universal indicator is made by mixing
several different indicators.
 If an exact value for the pH of a solution
is needed, a pH meter should be used.
 A pH meter determines the pH of a
solution be measuring the voltage
between the two electrodes that are
placed in the solution.
Sample Problems
Q1: Calculate the pH of a solution if
[H+] = 2.7 x 10-4 M
pH = -log[H+] pH = -log(2.7 x 10-4) = 3.57
Q2: Find the hydrogen ion concentration of a
solution if its pH is 11.62.
[H+] = 10-pH
[H+] = 10-11.62 = 2.4 x 10-12M
Q3: Find the pOH and the pH of a solution if its
hydroxide ion concentration is 7.9 x 10-5M
pOH = -log[OH-]
pOH = -log(7.9 x 10-5) = 4.10
pH + pOH = 14
pH = 14 - 4.10
pH = 9.9
A brief introduction to equilibrium constants and
Ionization constants
Equilibrium is defined as a state where concentrations of
products and reactants are unchanging
The equilibrium constant is a mathematical expression of the
ratio of the products compared to the reactants
K = { products}
The { } are read as concentration
{ reactants}
All species are considered to be at equilibrium
Equilibrium constant and Ionization constants
can be determine for many reactions
For acids
Ka = { H3O +} { A-}
{HA} {H2O}
Since the concentration of H2O is roughly 55 moles /liter
it can be deleted from the above equation therefore the
equation becomes
Ka = { H3O +} { A-}
{HA}
Titration
 A neutralization reaction occurs between
an acid and a base.
 Because acids and bases react, the
progressive addition of an acid to a base
or vice versa can be used to compare the
concentrations of the acid and the base.
Titrations con’t
 Titration is the controlled addition and
measurement of the amount of a solution
of known concentration required to react
completely with a measure amount of a
solution of unknown concentration
 Titration provides a sensitive means of
determining the chemically equivalent
volumes of acidic and basic solutions.
TITRATION
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION =
EQUIVALENCE POINT
At the end point for the titration of a strong acid with a
strong base, the moles of acid (H+) equals the moles
of base (OH-) to produce the neutral species water
(H2O). If the mole ratio in the balanced chemical
equation is NOT 1:1 then you must rely on the mole
relationship and handle the problem like any other
stoichiometry problem.
MOLES OF ACID = MOLES OF BASE
nacid = nbase
TITRATION
MAVA = MBVB
1. Suppose 75.00 mL of hydrochloric acid was required
to neutralize 22.50 mL of 0.52 M NaOH. What is the
molarity of the acid?
HCl + NaOH  H2O + NaCl
Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va
so Ma = (0.52 M) (22.50 mL) / (75.00 mL)
= 0.16 M
Now you try:
2. If 37.12 mL of 0.843 M HNO3 neutralized 40.50 mL of
KOH, what is the molarity of the base?
Mb = 0.773 mol/L
TITRATION
1. If 37.12 mL of 0.543 M LiOH neutralized 40.50
mL of H2SO4, what is the molarity of the acid?
2 LiOH + H2SO4  Li2SO4 + 2 H2O
First calculate the moles of base:
0.03712 L LiOH (0.543 mol/1 L) = 0.0202 mol LiOH
Next calculate the moles of acid:
0.0202 mol LiOH (1 mol H2SO4 / 2 mol LiOH)= 0.0101 mol
H2SO4
Last calculate the Molarity:
Ma = n/V = 0.010 mol H2SO4 / 0.4050 L = 0.248 M
2. If 20.42 mL of Ba(OH)2 solution was used to
titrate 29.26 mL of 0.430 M HCl, what is the
molarity of the barium hydroxide solution?
Mb = 0.308 mol/L
TITRATION
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION =
EQUIVALENCE POINT
At the end point for the titration of a strong acid with a
strong base, the moles of acid (H+) equals the moles of
base (OH-) to produce the neutral species water (H2O).
If the mole ratio in the balanced chemical equation is
1:1 then the following equation can be used.
MOLES OF ACID = MOLES OF BASE
nacid = nbase
Since M=n/V
MAVA = MBVB
Equivalence Point (E.P.)
 The point at which the two solutions used in
a titrations are present in chemically
equivalent amounts is the equivalence point
(e.p.).
 Indicators and pH meters cane be used to
determine the e.p.
 A pH meter will show a large voltage
change occurring at the e.p.
 An indicator changes color over a range
that includes the pH of the e.p.
End point
 The point in a titration at which an
indicator changes color is called the end
point.
Which indicator should
you use?
 Indicators that undergo transition at
about pH 7 are used to determine the
equivalence point of strong acid/ strong
base titrations because the neutralization
of strong acids with strong bases
produces a salt solution with a pH of
approximately 7.
Which indicator should
you use?
 Indicators that change color at pH lower
than 7 are useful in determining the
equivalence point of strong acid/ weak
base titrations. The equivalence point of
this titration is acidic because the salt
formed is a weak acid. Thus the salt
solution has a pH lower than 7.
Which indicator should
you use?
 Indicators that change color at pH higher
than 7 are useful in determining the
equivalence point of weak acid/ strong
base titrations. These reactions produce
salt solutions whose pH is greater than 7.
This occurs because the salt solution
formed is a weak base.
Which indicator should
you use?
 To determine the equivalence point for a weak
acid/ weak base, we use no indicator. The pH
of the equivalence point of weak acids and
weak bases may be almost any value,
depending on the relative strengths of the
reactants. The color transition of an indicator
helps very little in determining whether
reactions between weak acids and bases are
complete.
Molarity and titration
To calculate the molarity of a substance
using titration:
1. Start with the balanced equation and
determine the chemically equivalent
amounts of the acid and base.
2. Determine the moles of acid or base
from the known solution used in the
titration.
Molarity and titration
3. Determine the moles of solute of the
unknown solution used during the
titration.
4. Determine the molarity of the unknown
solution.
Predicting Acid-Base Reactions
Fig. 17-2, p. 511
Review questions (cont)
 Distinguish between strong & weak
bases; List clinical uses of these
bases & write equations for their
dissociation in water.
 Complete simple equations for the
neutralisation reaction of an acid & a
base; Discuss clinical applications of
acid-base neutralisation.
Review questions
 List the properties of acids & bases.
 Discuss the pH scale:
*Define the ion product of water & indicate
how this is determines the pH scale.
* Use the pH scale to determine if a given
solution is acidic, neutral or basic.
 Distinguish between strong acids & weak
acids: List clinical uses of these acids &
write equations for their dissociation in
water
Brønsted-Lowry Theory of Acids & Bases
Notice that water is both an acid & a
base = amphoteric
Reversible reaction
p. 504
Water Equilibrium
p. 514