Fundamentals of General, Organic,
and Biological Chemistry
5th Edition
Chapter Ten
Acids and Bases
James E. Mayhugh
Oklahoma City University
2007 Prentice Hall, Inc.
Outline
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10.1 Acids and Bases in Aqueous Solution
10.2 Some Common Acids and Bases
10.3 The Brønsted–Lowry Definition of Acids and Bases
10.4 Water as Both an Acid and a Base
10.5 Some Common Acid–Base Reactions
10.6 Acid and Base Strength
10.7 Acid Dissociation Constants
10.8 Dissociation of Water
10.9 Measuring Acidity in Aqueous Solution: pH
10.10 Working with pH
10.11 Laboratory Determination of Acidity
10.12 Buffer Solutions
10.13 Buffers in the Body
10.14 Acid and Base Equivalents
10.15 Titration
10.16 Acidity and Basicity of Salt Solutions
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10.1 Acids and Bases in Aqueous
Solution
► An acid is a substance that produces hydrogen ions,
H+, when dissolved in water. (Arrhenius definition)
► Hydronium ion: The H3O+ ion formed when an acid
reacts with water.
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► A base is a substance that produces hydroxide ions,
OH-, when dissolved in water. (Arrhenius definition)
► Bases can be metal hydroxides that release OH- ions
when they dissolve in water or compounds that
undergo reactions with water that produce OH- ions.
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10.2 Some Common Acids and
Bases
► Sulfuric acid, H2SO4, is manufactured in greater
quantity than any other industrial chemical. It is the
acid found in automobile batteries.
► Hydrochloric acid, HCl, is “stomach acid” in the
digestive systems of most mammals.
► Phosphoric acid, H3PO4, is used to manufacture
phosphate fertilizers. The tart taste of many soft
drinks is due to the presence of phosphoric acid.
► Nitric acid, HNO3, is a strong oxidizing agent that is
used for many purposes.
► Acetic acid, CH3CO2H, is the primary organic
constituent of vinegar.
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► Sodium hydroxide, NaOH, or lye, is used in the
production of aluminum, glass, and soap. Drain
cleaners often contain NaOH because it reacts with
the fats and proteins found in grease and hair.
► Calcium hydroxide, Ca(OH)2 , or slaked lime, is
made industrially by treating lime (CaO) with water.
It is used in mortars and cements. An aqueous
solution is often called limewater.
► Magnesium hydroxide, Mg(OH)2, or milk of
magnesia, is an additive in foods, toothpaste, and
many over-the-counter medications. Many antacids
contain magnesium hydroxide.
► Ammonia, NH3, is used primarily as a fertilizer. A
dilute solution of ammonia is frequently used around
the house as a glass cleaner.
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10.3 The Brønsted-Lowry Definition
of Acids and Bases
► A Brønsted–Lowry acid can donate H+ ions.
► Monoprotic acids can donate 1 H+ ion, diprotic
acids can donate 2 H+ ions, and triprotic acids can
donate 3 H+ ions.
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► A Brønsted–Lowry base accepts H+ ions.
► Putting the acid and base definitions together, an
acid–base reaction is one in which a proton is
transferred. The reaction need not occur in water.
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► Conjugate acid–base pair: Two substances whose
formulas differ by only a hydrogen ion, H+.
► Conjugate base: The substance formed by loss of
H+ from an acid.
► Conjugate acid: The substance formed by addition
of H+ to a base.
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10.4 Water as Both an Acid and a
Base
► Water is neither an acid nor a base in the Arrhenius
sense because it does not produce appreciable
concentrations of either H+ or OH-. In the Brønsted–
Lowry sense water is both an acid and a base.
► In its reaction with ammonia, water donates H+ to
ammonia to form the ammonium ion.
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► When water reacts as a Brønsted–Lowry base, it
accepts H+ from an acid like HCl.
► Substances like water, which can react as either an
acid or a base depending on the circumstances, are
said to be amphoteric.
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10.5 Some Common Acid-Base
Reactions
► Acids react with metal hydroxides to yield water and
a salt in a neutralization reaction. The H+ ions and
OH- ions are used up in the formation of water.
HA(aq) + MOH(aq)H2O(l) +MA(aq)
► Carbonate and bicarbonate ions react with acid by
accepting H+ ions to yield carbonic acid, which is
unstable, and rapidly decomposes to yield carbon
dioxide gas and water.
H2CO3(aq)H2O(l) + CO2(g)
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► Acids react with ammonia to yield ammonium salts,
most of which are water-soluble.
► Living organisms contain a group of compounds
called amines, which contain ammonia-like nitrogen
atoms bonded to carbon. Amines react with acids just
as ammonia does. Methylamine, an organic
compound found in rotting fish, reacts with HCl:
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10.6 Acid and Base Strength
► Strong acid: An acid that gives up H+ easily and is
essentially 100% dissociated in water.
► Dissociation: The splitting apart of an acid in water
to give H+ and an anion.
► Weak acid: An acid that gives up H+ with difficulty
and is less than 100% dissociated in water.
► Weak base: A base that has only a slight affinity for
H+ and holds it weakly.
► Strong base: A base that has a high affinity for H+
and holds it tightly.
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► The stronger the acid, the weaker its conjugate base;
the weaker the acid, the stronger its conjugate base.
► An acid–base proton transfer equilibrium always
favors reaction of the stronger acid with the stronger
base, and formation of the weaker acid and base.
► The proton always leaves the stronger acid and
always ends up in the weaker acid, whose stronger
conjugate base holds the proton tightly.
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10.7 Acid Dissociated Constants
► The reaction of a weak acid with water, like any
chemical equilibrium, can be described by an
equilibrium equation.
Ka = [H3O+][A-]/[HA]
► Strong acids have Ka >> 1 because dissociation is
favored and weak acids have Ka << 1 because
dissociation is not favored.
► Donation of each successive H+ from a polyprotic
acid is more difficult than the one before it, so Ka
values become successively lower.
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Most organic acids, which contain the group
-COOH, have Ka values near 10-5.
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10.8 Dissociation of Water
► Like all weak acids, water is slightly dissociated into
H+ and OH- ions. At 25oC, the concentration of each
ion is 1.00 x 10-7 M in pure water.
► The ion product constant for water, kw, is:
kw = [H3O+][OH-]= 1.00 x 10-14 at 25oC.
► Acidic solution: [H3O+] > 10-7 M and [OH-] < 10-7 M
► Neutral solution: [H3O+] = 10-7 M and [OH-] = 10-7 M
► Basic solution: [H3O+] < 10-7 M and [OH-] > 10-7 M
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10.9 Measuring Acidity in Aqueous
Solution: pH
► The concentrations of H3O+ or OH- in solution can
vary over a wide range. A logarithmic scale can be
easier to use.
► pH = -log [H3O+ ]
pOH = -log [OH- ]
► [H3O+ ] = 10-pH
[OH- ] = 10-pOH
► Acidic solution: pH < 7 pOH > 7
► Neutral solution: pH = 7 pOH = 7
► Basic solution: pH > 7 pOH < 7
► pH + pOH = 14.00 at 25C
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10.10 Working with pH
► An antilogarithm has the same number of digits that the
original number has to the right of the decimal point.
► A logarithm contains the same number of digits to the
right of the decimal point that the original number has.
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10.11 Laboratory Determination of Acidity
(a) The color of universal indicator in solutions of
known pH from 1 to 12. (b) Testing pH with a paper
strip. Comparing the color of the strip with the code
on the package gives the approximate pH.
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►A much more accurate way to determine pH uses an
electronic pH meter like the one shown below.
►Electrodes are dipped into the solution, and the pH is
read from the meter.
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10.12 Buffer Solutions
► Buffer: A combination of substances that act
together to prevent a drastic change in pH; usually a
weak acid and its conjugate base.
► Rearranging the Ka equation shows that the value of
[H3O+] depends on the ratio [HA]/[A-].
[H3O+] = Ka [HA]/[A-]
► Most H3O+ added is removed by reaction with A- ,so
[HA] increases and [A-] decreases. As long as these
changes are small, the ratio [HA]/[A-] changes only
slightly, and there is little change in the pH.
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When 0.010 mol of acid and 0.010 mol of base are
added to 1.0 L of pure water and to 1.0 L of a 0.10 M
acetic acid–0.10 M acetate ion buffer, the pH of the
water varies between 12 and 2, while the pH of the
buffer varies only between 4.85 and 4.68.
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10. 13 Buffers in the Body
► The pH of body fluids is maintained by three major
buffer systems. The carbonic acid–bicarbonate
system, the dihydrogen phosphate–hydrogen
phosphate system, and a third system that depends
on the ability of proteins to act as either proton
acceptors or proton donors at different pH values.
► The carbonic acid–bicarbonate system is the
principal buffer in blood serum and other
extracellular fluids. The hydrogen phosphate system
is the major buffer within cells.
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► A change in the breathing rate provides a quick
adjustment in the bicarbonate buffer system.
► When the CO2 concentration in the blood starts to
rise, the breathing rate increases to remove CO2,
thereby decreasing the acid concentration.
► When the CO2 concentration in the blood starts to fall,
the breathing rate decreases and acid concentration
increases.
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Each day, acid produced in the body is excreted in the
urine. The kidney returns HCO3- to the extracellular
fluids, where it becomes part of the bicarbonate reserve.
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10.14 Acid and Base Equivalents
► We think in terms of ion equivalents when we are
primarily interested in the ion itself rather than the
compound that produced the ion. For similar reasons,
we use acid or base equivalents.
► One equivalent (Eq) of an acid is equal to the molar
mass of the acid divided by the number of H+ ions
produced per formula unit. Similarly, one equivalent
of a base is the weight in grams that can produce one
mole of OH- ions.
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► One Eq of a monoprotic acid is the molar mass of the
acid. One equivalent of a diprotic acid is the molar
mass of the acid divided by 2 since each mole can
produce two moles of H+.
► One equivalent of any acid neutralizes one
equivalent of any base; this is convenient when only
the acidity or basicity of a solution is of interest
rather than the identity of the acid or base.
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► Because acid–base equivalents are so useful, clinical
chemists sometimes express acid and base
concentrations in normality rather than molarity.
► The normality (N) of an acid or base solution is
defined as the number of equivalents of acid or base
per liter of solution.
► For any acid or base, normality is always equal to
molarity times the number of H+ or OH- ions
produced per formula unit.
► N of acid = (M of acid)  (# of H+ ions produced)
► N of base = (M of base)  (# of OH- ions produced)
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10.15 Titration
►The pH of a solution gives the H+ concentration but
not necessarily the total acid concentration.
►The H+ concentration gives only the amount of acid
that has dissociated into ions, whereas total acid
concentration gives the sum of dissociated plus
undissociated acid.
► In a 0.10 M solution of acetic acid the total acid
concentration is 0.10 M, yet the H+ concentration is
only 0.0013 M because acetic acid is a weak acid that
is only about 1% dissociated.
►The total acid or base concentration of a solution can
be found by carrying out a titration procedure.
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► (a) A measured volume of the acid solution is placed
in the flask along with an indicator.
► (b) The base of known concentration is then added
from a buret until the color change of the indicator
shows that neutralization is complete.
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► Reading the buret gives the volume of base. The
molarity and volume of base yields the moles of base.
► The coefficients in the balanced equation allow us to
find the moles of acid neutralized. Dividing the moles of
acid by the volume of the acid gives the acid molarity.
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10.16 Acidity and Basicity of Salt
Solutions
► Salt solutions can be neutral, acidic, or basic, depending
on the ions present, because some ions react with water
to produce H+ and some ions react with water to
produce OH-.
► To predict the acidity of a salt solution, it is convenient
to classify salts according to the acid and base from
which they are formed in a neutralization reaction.
► The general rule for predicting the acidity or basicity of
a salt solution is that the stronger partner from which
the salt is formed dominates.
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► A salt formed from a strong acid and a weak base yields
an acidic solution because the strong acid dominates.
► A salt formed from a weak acid and a strong base yields
a basic solution because the base dominates.
► A salt formed from a strong acid and a strong base yields
a neutral solution because neither dominates.
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Chapter Summary
►According to the Brønsted–Lowry definition, an acid
is a substance that donates a hydrogen ion and a base
is a substance that accepts a hydrogen ion.
►Thus, the generalized reaction of an acid with a base
involves the reversible transfer of a proton.
►A strong acid gives up a proton easily and is 100%
dissociated in aqueous solution.
►A weak acid gives up a proton with difficulty and is
only slightly dissociated in water.
►A strong base accepts and holds a proton readily,
whereas a weak base has a low affinity for a proton.
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Chapter Summary Cont.
► The two substances that are related by the gain or loss
of a proton are called a conjugate acid–base pair.
► A proton-transfer reaction always takes place in the
direction that favors formation of the weaker acid.
► Water is amphoteric; it can act as an acid or a base.
► The acidity or basicity of an aqueous solution is given
by its pH, defined as the negative logarithm of the
hydronium ion concentration.
► A pH below 7 means an acidic solution; a pH equal to 7
means a neutral solution; and a pH above 7 means a
basic solution.
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Chapter Summary Cont.
► The pH of a solution can be controlled through the
use of a buffer that acts to remove either added H+
ions or added OH- ions. Most buffer solutions
consist of roughly equal amounts of a weak acid and
its conjugate base.
► Acid (or base) concentrations are determined in the
laboratory by titration of a solution of unknown
concentration with a base (or acid) solution of
known strength until an indicator signals that
neutralization is complete.
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End of Chapter 10
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