Goals
► 1. What are acids and bases? Be able to recognize acids
and bases and write equations for common acid–base
reactions.
► 2. What effect does the strength of acids and bases have
on their reactions? Be able to interpret acid strength using
acid dissociation constants Ka and predict the favored direction
of acid–base equilibria.
► 3. What is the ion-product constant for water? Be able to
write the equation for this constant and use it to find the
concentration of H3O+ or OH- .
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►4. What is the pH scale for measuring acidity? Be able to
explain the pH scale and find pH from the H3O+ concentration.
► 5. What is a buffer? Be able to explain how a buffer
maintains pH and how the bicarbonate buffer functions in the
body.
►
► 6. How is the acid or base concentration of a solution
determined? Be able to explain how a titration procedure
works and use the results of a titration to calculate acid or base
concentration in a solution.
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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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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 of 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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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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► 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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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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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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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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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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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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► 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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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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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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► 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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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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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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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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► 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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► 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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