Download Chapter Ten

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 use in the petroleum and pharmaceutical
industry’s, and 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. Chapter Ten
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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-thecounter medications. Many antacids contain
magnesium hydroxide, can also be used in mortars and
cements.
► Ammonia, NH3, is used primarily as a fertilizer. A dilute
solution of ammonia is frequently used around the
house
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Chapter Ten
Soap is manufactured by the reaction of vegetable oils and
animal fats with the bases NaOH and KOH. Sodium
hydroxide is also called caustic soda or lye and is the most
commonly used of all bases.
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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Acetic acid is an organic acid, which donates a hydrogen
ion in solution. Just one of the hydrogen's on acetic acid
is acidic. The hydrogen attached to the electronegative
oxygen atom is the acidic hydrogen on acetic acid.
► 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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Which of the following would you expect to be a
BrØnsted-Lowry acid or BrØnsted-Lowry base?
1.
2.
3.
4.
Fe3+
H2CrO4
NH3
NO3–
Which of the following would you expect to be a
BrØnsted-Lowry acid?
1.
2.
3.
4.
Fe3+
H2CrO4
NH3
NO3–
(referred to as a Lewis Acid)
acid: 2H+ + CrO42base: NH4+
base: HNO3
Which of the following would you expect to be a
BrØnsted-Lowry base?
1.
2.
3.
4.
HNO2
NH4+
Ni2+
PO43–
Which of the following would you expect to be a
BrØnsted-Lowry base?
1.
2.
3.
4.
HNO2
NH4+
Ni2+
PO43–
► 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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►Indentify the conjugate acid/base pairs
HCO3-(aq) + H2O(l) ⇄ CO32-(aq) + H3O+(aq)
HF(aq) + HPO42-(aq) ⇄ F-(aq) + H2PO4-(aq)
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►Indentify the conjugate acid/base pairs
acid
base
HCO3-(aq) + H2O(l) ⇄ CO32-(aq) + H3O+(aq)
base
acid
acid
base
HF(aq) + HPO42-(aq) ⇄ F-(aq) + H2PO4-(aq)
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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Shown below is the zwitterion of the amino acid
alanine. The zwitterion can serve as
1.
2.
3.
4.
both an acid and a base.
only an acid.
only a base.
neither an acid not a base.
Shown below is the zwitterion of the amino acid
alanine. The zwitterion can serve as
1.
2.
3.
4.
both an acid and a base.
only an acid.
only a base.
neither an acid not a base.
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
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Most organic acids, which contain the group
-COOH, have Ka values near 10-5.
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The weaker the acid, the stronger its conjugate
base. The reverse reaction is favored as indicated
by the arrow size.
HC2H3O2 (acetic acid) Ka = 1.8 × 10-5
The equilibrium constant for the reaction shown below is
called the base dissociation constant, Kb, and is a measure of
the base strength of the acetate ion. For acetate ion
Kb = 5.6  10–10. In this reaction the strongest acid/strongest
base are
1.
2.
3.
4.
acetic acid/acetate ion.
acetic acid/hydroxide ion.
water/acetate ion.
water/hydroxide ion.
The equilibrium constant for the reaction shown below is
called the base dissociation constant, Kb, and is a measure of
the base strength of the acetate ion. For acetate ion
Kb = 5.6  10–10. In this reaction the strongest acid/strongest
base are
1.
2.
3.
4.
acetic acid/acetate ion.
acetic acid/hydroxide ion.
water/acetate ion.
water/hydroxide ion.
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-] but [H2O] = 1, so much of it, so
[H2O]
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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What is the [H3O+] if a NaOH(aq) is 2.1 × 10-5 M at 25 °C?
(Kw = 1.0 × 10-14)
Kw = [H3O+] [OH-]
1.0 × 10-14 = [H3O+] 2.1 × 10-5
[H3O+] = 1.0 × 10-14
2.1 × 10-5
[H3O+] = 4.8 × 10-10
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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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A aqueous solution has a pH of 12.26 at 25 °C, what is
the pOH? Is this acidic or basic?
pH + pOH = 14
12.26 + pOH = 14
pOH = 14 - 12.26 = 1.74
A baking soda solution has [H3O+] of 1.3 × 10-8, what
are the pH and pOH? Is this acidic or basic?
pH = -log [H3O+ ]
pH + pOH = 14
pH = -log 1.3 × 10-8
7.9 + pOH = 14
pH = -log 1.3 × 10-8 = 7.9
pOH = 14 - 7.9 = 6.1
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A 0.10 M solution of ammonium ion, NH4+, has a pH of
5.6. Ammonium ion is a
1.
2.
3.
4.
strong acid.
weak acid.
strong base.
weak base.
A 0.10 M solution of ammonium ion, NH4+, has a pH of
5.6. Ammonium ion is a
NH4+ + H2O ⇄ NH3 + OH1.
2.
3.
4.
strong acid.
weak acid.
strong base.
weak base.
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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A aqueous solution has a pH of 12.26 at 25 °C. What
is the [OH-]. Is the solution acidic or basic?
pH + pOH = 14
12.26 + pOH = 14
pOH = 14 – 12.26 = 1.74
pH = -log [H3O+ ] or the inverse [H3O+ ] = 10-pH
pOH = -log [OH- ] or the inverse [OH- ] = 10-pOH
[OH- ] = 10-pOH
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[OH- ] = 10-1.74 = 1.82 × 10-2
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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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10.78 & 10.79
►10.78. What is the pH of buffer system that contains
0.200 M hydrocyanic (HCN) and 0.150 M sodium
cyanide (NaCN)? The Ka of HCN is 4.9 × 10-10
HCN + H2O ⇄ CN- + H3O+
pH = -log [6.5 × 10-10] = 9.18
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10.78 & 10.79
►10.79. What is the pH of 10.78 after .020 mol of HCl
is added.
HCN + H2O ⇄ CN- + H3O+
The solution started as a 1 liter solution so that
0.200 M HCN is .200 mol and 0.150 M cyanide is
.150 mol CN-.
The HCl will react with CN- to make more HCN.
HCN will go up from .200 to .220 mol or .220 M
- will go down from .150 to .130 mol or .130 M
CN
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10.78 & 10.79
►10.79. What is the pH of 10.78 after .020 mol of HCl
is added.
HCN will go up from .200 to .220 mol or .220 M
CN- will go down from .150 to .130 mol or .130 M
HCN + H2O ⇄ CN- + H3O+
pH = -log [8.3 × 10-10] = 9.08
ΔpH = 9.18 - 9.08 = .1
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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.
CO2 + H2O ⇄ H2CO3 ⇄ H+ + HCO3-
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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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10.85
►How many equivalencies of NaOH are needed to
react with 0.035 Eq of the triprotic acid H3PO4?
H3PO4 + 3 NaOH ⇄ Na3PO4 + 3 H2O
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10.91
►What are the molarity and the normality of a
solution made by dissolving 25.0 g of citric acid
(C6H5O7H3, a triprotic acid) in enough water to make
800 mL of solution?
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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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Learning Check
Calculate the mL of 2.00 M H2SO4 required to
neutralize 50.0 mL of 1.00 M KOH.
H2SO4 + 2KOH
K2SO4 + 2H2O
69
Solution
Calculate the mL of 2.00 M H2SO4 required to
neutralize 50.0 mL of 1.00 M KOH.
H2SO4 + 2KOH
K2SO4 + 2H2O
1) 12.5 mL
0.0500 L x 1.00 mole KOH x 1 mole H2SO4 x
1L
2 mole KOH
1L
x 1000 mL = 12.5 mL
2 mole H2SO4
1L
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Learning Check
A 25 mL sample of phosphoric acid is neutralized by
40. mL of 1.5 M NaOH. What is the molarity of the
phosphoric acid solution?
3NaOH + H3PO4
71
Na3PO4 + 3H2O
Solution
2) 0.80 M
0.040 L x 1.5 mole NaOH x 1 mole H3PO4
1L
3 mole NaOH
= 0.020 mole H3PO4
0.020 mole H3PO4 = 0.80 mole/L = 0.80 M
0.025 L
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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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Acid-Base Properties of Salts
MX + H2O ----> acidic or basic solution?
Consider NH4Cl
NH4Cl(aq) ----> NH4+(aq) + Cl-(aq)
Reaction of Cl- with H2O
Cl- +
base
X
H2O <----> HCl
acid
acid
+
OHbase
Cl- ion is a VERY weak base because its conjugate
acid is strong.
Therefore, Cl- ----> neutral solution
Acid-Base Properties of Salts
Determine if the following solutions are acidic,
basic, or neutral.
KBr
K+ + HOH <--> KOH + H+
neutral
Br + HOH <--> HBr + OH
X
X
CrCl3
Cr3+ + HOH <--> CrOH2+ + H+
Cl- + HOH <--> HCl + OH-
X
acidic
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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