Acids and Bases Introduction and pH
 You probably know that you encounter acids and
bases on a daily basis. Acids flavor both food and
drink (soda, for example) and line your stomach to
help digest that food. Bases are commonly found in
most household cleaners, as well as the soap you
use to wash everyday.
 Properties of Acids and Bases differ quite a bit.
o Common 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
releases Hydrogen gas, H2
 Acids react with bases to produce salts
and water
o Common Properties of Bases
 Aqueous solutions of bases have a bitter
taste
 Bases change the color of acid-base
indicators
 Dilute aqueous solutions of bases feel
slippery
 Bases react with acids to produce salts
and water
 Bases conduct electric current
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Acids and Bases Introduction and pH
Acid Nomenclature
A binary acid is an acid that contains only two different
elements—hydrogen and one of the more
electronegative nonmetal elements. Examples of
binary acids include all of the hydrogen halides: HF,
HCl, HBr, and HI. The hydrogen halide acids are all
gases when they are pure, so we name them as we do
all other compounds. When they are in aqueous
solutions, we name them differently, and this is the way
that I want you to name them:
 The name of a binary acid begins with the prefix
hydro-.
 The root of the name is the root of the second
element, and this follows the prefix.
 The name of the acid ends with the suffix –ic.
An oxyacid is an acid that is a compound of hydrogen,
oxygen, and a third element (usually a nonmetal).
Because they have more then 2 elements in them, they
are called ternary acids (the root ter- means “three.”).
When writing the formula for a ternary acids, they
should be visualized as the attachment of a hydrogen
atom to a polyatomic ion class called oxyanions. You
should study the chart below, and compare it to the
compound naming flowchart that you got the first half
of the year (yes, we are going back to that). We talked
briefly in our nomenclature unit about naming
oxyanions, so let’s start with their naming:
You need to understand the naming of oxyanions—
polyatomic ions that contain oxygen. In several cases,
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Acids and Bases Introduction and pH
two or more different oxyanions can form from the
same two elements.
Let us use the example of nitrogen and oxygen with the
anions nitrate, NO , and nitrite, NO .

2

3
 When naming oxyanions, the most common ion that
forms is given the suffix –ate.
 The one containing one fewer oxygen atom is given
the suffix –ite.
 If a further reduction in oxygen atoms is possible,
the anion with one fewer oxygen atom than the –ite
anion keeps the –ite suffix and adds the prefix
hypo-.
 An oxyanion with one more oxygen than the –ate
anion is also given the prefix per-. An example of
this is given in the text with the oxyanions formed
from oxygen and chlorine. The table below
summarizes the naming of common binary and
ternary acids:
HF
hydrofluoric acid
HNO2
nitrous acid
HClO
hypochlorous acid
HCl
hydrochloric acid
HNO3
nitric acid
HClO2
chlorous acid
HBr
hydrobromic acid
H2SO3
sulfurous acid
HClO3
chloric acid
HI
hydroiodic acid
H2SO4
sulfuric acid
HClO4
perchloric acid
H3PO4
phosphoric acid
CH3COOH
acetic acid
H2CO3
carbonic acid
 You will notice that when phosphoric acid is
stripped of all of its hydrogen ions, a phosphate ion
remains. The same is true for sulfuric acid (sulfate
ion forms), and chloric acid (chlorate ion forms).
Therefore the most common oxyanion will have a
suffix of “-ic” when it forms an acid. You will notice
that when nitrous acid is stripped of its hydrogen
ion, it leaves a nitrite ion. Therefore the “—ite”
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Acids and Bases Introduction and pH
oxyanions take the suffix “—ous” when they
become acids.
 What is an acidic and basic solution?
o Acidic Solution: a solution containing more
hydrogen ions [H+] than hydroxide ions [OH-].
o Basic Solution: a solution containing more
hydroxide ions [OH-] than hydrogen ions [H+].
 There are two “models” or difinitions that provide a
better understanding of acids and bases.
o Arrhenius Model (Svante Arrhenius-1884)
 An acid is a substance that has hydrogen
in its formula and ionizes to produce
hydrogen ions in aqueous solutions (i.e.
H3O+)
 HCl (g)  H+ (aq) + Cl- (aq)
 This increases the concentration of H+
 A base is a substance that contains a
hydroxide group (OH-) and ionizes to
produce a hydroxide ion in aqueous
solutions
 NaOH (s)  Na+ (aq) + OH- (aq)
 This increases the concentration of
OH-
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Acids and Bases Introduction and pH
o Bronsted-Lowry Model (Johannes Bronsted
and Thomas Lowry-1923; proposed
independently)
 Acids are simply a hydrogen ion (H+) donor
HCl(aq) + H2O(l)  Cl-(aq) + H3O+(aq)
 Bases are simply a hydrogen ion (H+)
acceptor
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
 This donating and accepting of hydrogen
ions produces products that are also acids
and bases. These acids and bases are
called “conjugates”. This means
substances are “joined together” by their
acid-base relationship.
 Conjugate acids: the species formed when
a hydrogen ion is added to a base
 Conjugate bases: the species formed
when a hydrogen ion is removed from an
acid
HF(aq)
Acid
+ H2O(l)
Base
↔
F-(aq)
+
H3O-(aq)
Conj. Base
Conj. Acid
 These two sets of substances are called
“conjugate acid-base pairs”.
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Acids and Bases Introduction and pH
Sometimes a substance can be considered both an
acid and a base (water is a great example). These
substances are called amphoteric.
 Strength vs. Concentration: What is the
difference?
o The ease with which a proton is released (for
an acid) or accepted (for a base) from the
compound is the indication of the acid’s or
base’s strength.
o The number of protons (H+) in solution is the
indication of the solution’s acidity.
SA
SB
All
o Remember concentration is the
HCl hydroxides
amount of a substance dissolved in a
of Group 1
specified amount of solvent.
HNO3 Mg(OH)2
H2SO4 Ca(OH)2
o Strength is the degree of ionization or
HBr Sr(OH)2
dissociation of the acid or base in
aqueous solutions.
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Acids and Bases Introduction and pH
What is pH? (note the small p and large H!)
 A measure of just how concentrated H+ or H3O+
is in solution is needed. Chemists’ just aren’t
satisfied with “acidic” or “basic” to describe
solutions. They need numbers!!!!
 Remember that water “autoionizes” itself as
follows:
H2O(l) + H2O(l) ↔ OH-(aq) + H3O+(aq)
Or
H2O(l) ↔ OH-(aq) + H+(aq)
 Electrical conductivity measurements of pure
water show that concentrations of H3O and
OH are 1.0  107 M at 25C.

(aq )

(aq )
 We designate the aqueous hydrogen ion molar
concentration as [H+] or [H3O+]. Thus, [H3O+] = 1.0 
107 M, and [OH] = 1.0  107 M.
 Scientists have found that the product of [H3O+] 
[OH] = 1.0  1014 M 2 in water and all aqueous
solutions at constant temperature. We designate
this constant value as the ionization constant or ion
product constant of water, Kw:
K w  [OH  ]  [ H 3O  ] or [OH  ]  [ H  ]
Where,
K w  [OH  ]  [ H  ]  (1.0  107 )(1.0 107 )  1.0 1014
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Acids and Bases Introduction and pH
Because the hydronium and hydroxide ion
concentrations are the same in pure water, it is
neutral. The presence of acid solutes however, will
increase [H3O+], while the presence of basic solutes
will increase [OH]. When [H3O+]>[OH], the solution will
be acidic. When the reverse is true, the solution is
basic. We can use our knowledge of acid and base
strength to determine the relative concentrations of
each ion.As the concentration of one increases, the
other must decrease so that the Kw = 1.0  1014. We
can use this to calculate concentrations of [H+] and
[OH] in solution.
Practice Exercise:
Indicate whether each of the following solutions is
neutral, acidic, or basic:
Remember that the concentration of [H+] = [OH] = 1.0 
107M.
(a) [H+] = 2  105 M > 1.0  107, so this solution is
acidic
(b) [OH] = 3  109 < 1.0  107, so this solution is
acidic
(c) [OH] = 1.0  107 = 1.0  107, so this solution is
neutral
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Acids and Bases Introduction and pH
So, when differing (additional or removed) amounts of
hydroxide or hydrogen ions are present in solution, we
can use Kw to determine if solutions are acidic, basic or
neutral.
EX1
At 298 K, the H+ concentration of an aqueous solution
is 1.0 x 10-13 M. What is the OH- concentration and is
the solution acidic, basic or neutral?
Step 1: Write knowns, unknowns
[H+] = 1.0 x 10-13 M
Kw = 1.0 x 10-14
[OH-] = ???
Step 2: Write equation.
K w  [OH  ]  [ H  ]
Step 3: Isolate unknown variable.
K
[OH  ]  w
[H ]
Step 4: Solve.
1.0  1014
1
[OH ] 

1
.
0

10
M
13
1.0  10

[OH-] > [H+]; therefore solution is basic
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Acids and Bases Introduction and pH
Practice Exercise:
Calculate the concentration of OH in a solution in
which
(a) [H+] = 2  106 M: Kw = [H+][OH] = 1.0  1014

(aq )
1.0  10 14 1.0  10 14
 [OH ] 

 5  10 9 M

6
[H ]
2  10

(b) [H+] = 1  107 M: Kw = [H+][OH] = 1.0  1014
1.0  10 14 1.0  10 14
 [OH ] 

 1  10 7 M

7
[H ]
1  10

(c) [H+] = 100  [OH]: Kw = [H+][OH] = 1.0  1014
 [OH

]
1.0  10 14
1.0  10 14
 2

[
OH
]

 1  10 16  [OH  ]  1  10 8 M
100
100  [OH  ]
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Acids and Bases Introduction and pH
pH scale:
 As you can see, concentrations of hydrogen and
hydroxide can be very small numbers. The
negative exponents can be very confusing and
difficult to manipulate, so chemists developed a
scale to simplify the values of these
concentrations.
 The concentrations have both negatives and
exponents in them: a double whammy! So they
invented a scale which removed them both. The
basis of this scale is:
pH = - log[H+]
This changes a concentration of hydrogen from a
value like 4.5 x 10-4 M to pH = - log(4.5 x 10-4) = 3.3.
 SO, acids have a pH under 7, bases have a pH of
over 7, and neutral solutions have a pH of 7. You
should be able to see that these numbers come
right from the negative exponents in the hydrogen
ion concentrations.
 Also, hydroxide concentrations can be dealt with in
the same way. So, pOH = - log[OH-]
 This “boils down” to:
pH + pOH = 14
and
[H+] x [OH-] = 1.0 x 10-14
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Acids and Bases Introduction and pH
Calculating pH of SA or SB solutions:
 The pH of strong acids and bases are very easy to
calculate.
 Since there is 100% dissociation of the acid or
base, the reaction can be written with a single
forward arrow. We saw this before with HCl:
HCl(aq) + H2O(l)  Cl-(aq) + H3O+(aq)
Because this is a one way process (and not at
equilibrium), only the equilibrium of water is
important here.
H2O(l) ↔ OH-(aq) + H+(aq)
So, any concentration of acid donated by the HCl,
totally shifts the equilibrium of water, and is the
primary source of hydrogen ion.
EX2
Calculate the pH of a 0.28 M solution of HCl.
Step 1: Knowns, unknowns
[H+] = [HCl] = 0.28 M
pH = ???
Step 2: Choose equation
Write reaction (above)
pH = -log[H+]
Step 3: Isolate variable, if necessary.
All okay here already.
Step 4: Plug and chug.
pH = -log(0.28) = 0.55
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Acids and Bases Introduction and pH
EX3
Calculate the pH of a 0.28 M solution of Mg(OH)2.
Write reaction:
Mg(OH)2(s)  Mg+2(aq) + 2OH-(aq)
***Note, the concentration of OH- is 2x that of Mg(OH)2
Step 1: Knowns, unknowns
[OH-] = 2[Mg(OH)2] = 2 x 0.28 M = 0.56 M
pH = ???
Step 2: Choose equation(s)
pOH = -log[OH-]
pH + pOH = 14
Step 3: Isolate variable, if necessary.
First equation okay.
Second equation, pH = 14 - pOH
Step 4: Plug and chug.
First equation, pOH = -log(0.56) = 0.25
Second, 14 - 0.25 = 13.75 (negate rules for SF’s here—
power functions have
their own rules anyway!)
Calculating pH of WA or WB solutions:
 Weak acids and weak bases are solutions whose
ionization process is NOT one way, but two.
 This means that there is an equilibrium which
exists in each of their dissociations. For example:
HNO2(g) + H2O(l)  H3O+(aq) + NO2-(aq)
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Acids and Bases Introduction and pH
 Today, we will just reduce this to the following:
o A weak acid or base is only partially ionized.
o This can be stated in a Percent Dissociation
value.
o All we need to do to calculate the [H+] is to
apply the percent dissociation to the solution
concentration.
o Then, calculate pH as before.
EX4
HClO2 is a weak acid which dissociates by 40% in
solution. Calculate the pH of a solution which is 0.0400
M HClO2.
Write reaction:
HClO2(g) + H2O  ClO2-(aq) + H3O+(aq)
Step 1: Knowns, unknowns
[H+] = % x [HClO2] = 40% x 0.0400 M = 0.0160 M
pH = ???
Step 2: Choose equation(s)
pH = -log[H+]
Step 3: Isolate variable, if necessary.
Equation okay.
Step 4: Plug and chug.
pH = -log(0.0160) = 1.80
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