Chapter 9:
Acids, Bases, and Salts
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ARRHENIUS ACID
• An Arrhenius acid is any substance that provides hydrogen
ions, H+, when dissolved in water.
ARRHENIUS BASE
• An Arrhenius base is any substance that provides
hydroxide ions, OH-, when dissolved in water.
EXAMPLES OF AN ARRHENIUS ACID AND BASE
• HNO3 is an acid: HNO3(aq)
• KOH is a base: KOH(aq)
H+ (aq) + NO3- (aq)
K+ (aq) + OH- (aq)
2
BRØNSTED ACID
• A Brønsted acid is any hydrogen-containing substance that
is capable of donating a proton (H+) to another substance.
BRØNSTED BASE
• A Brønsted base is any substance capable of accepting a
proton from another substance.
EXAMPLE OF A BRØNSTED ACID AND BASE
• HNO2(aq) + H2O(l)
H3O+ (aq) + NO2-(aq)
• In this reaction, HNO2 behaves as a Brønsted acid by
donating a proton to the H2O. The H2O behaves as a
Brønsted base by accepting the proton.
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• The reaction was written using a double arrow that points
to both the right and the left, signifying that the reaction
can go in either direction and establish an equilibrium.
• HNO2(aq) + H2O(l)
H3O+ (aq) + NO2-(aq)
• When the reaction is read from right to left, the H3O+
behaves as a Brønsted acid by donating a proton to the
NO2- ion, which behaves as a Brønsted base by accepting
the proton.
• The behavior noted above is characteristic of Brønsted
acids and bases. When a substance such as HNO2
behaves as an acid and donates a proton, the species that
remains (the NO2- ion in this case) is a Brønsted base.
Similarly, the H2O that behaved as a base and accepted
the proton from the HNO2 was converted into the acid
H3O+.
4
CONJUGATE ACIDS AND BASES
• The base formed (NO2-) when a substance (HNO2) acts as
a Brønsted acid is called the conjugate base of the acid.
Similarly, the acid formed (H3O+) when a substance (H2O)
acts as a Brønsted base is called the conjugate acid of
the base.
CONJUGATE ACID-BASE PAIRS
• A Brønsted acid (such as HNO2) and its conjugate base
(NO2-) form what is called a conjugate acid-base pair.
• The same name is given to a Brønsted base (such as
H2O) and its conjugate acid (H3O+).
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BINARY ACIDS
• Solutions of binary acids such as HCl (aq) are formed by
dissolving binary compounds such as HCl gas in water.
• The HCl gas before being dissolved in water is said to be
anhydrous (meaning without water). The anhydrous gas is
named hydrogen chloride.
• The water solution of the gas is called hydrochloric acid.
• Some hydrogen-containing compounds such as HCl, HI,
HBr, and H2S form acidic solutions when they are dissolve
in water. The names of the acid solutions, such as the
hydrochloric acid given above, can be obtained by
following four rules.
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RULES FOR NAMING BINARY ACIDS
• Rule 1: Drop the word hydrogen from the anhydrous
compound name. For example, HI, called hydrogen iodide
becomes "iodide".
• Rule 2: Add the prefix hydro- to the result of step 1.
"Iodide" becomes "hydroiodide".
• Rule 3: Drop the suffix -ide from the result of step 2 and
replace it with the suffix -ic. "Hydroiodide" becomes
"hydroiodic".
• Rule 4: Add the word acid to the end of the name as a
separate word. The final name is "hydroiodic acid".
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RULES FOR NAMING ACIDS CONTAINING
POLYATOMIC IONS
• Rule 1: All hydrogen atoms that are written as the first part of
the formula of the acid are removed. The hydrogens are
removed in the form of H+ ions.
• Rule 2: The polyatomic ion that remains after the H+ ions are
removed is named by referring to sources such as Table 4.7.
• Rule 3: When the remaining polyatomic ion has a name
ending in the suffix -ate, the suffix is replaced by the suffix
–ic, and the word acid is added.
• Rule 4: When the remaining polyatomic ion has a name
ending in the suffix -ite, the suffix is replaced by the suffix
-ous, and the word acid is added.
• Rule 5: If the polyatomic ion contains sulfur or phosphorus,
the stems -sulf or -phosph that remain when the suffixes -ate
or -ite are replaced, are expanded to -sulfur and -phosphor
before the –ic or –ous suffixes are added.
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THE SELF-IONIZATION OF WATER
• A sample of absolutely pure water does not contain only
H2O molecules. In addition, small but equal amounts of
H3O+ and OH- ions are also present.
• The reason for this is that in one liter of pure water 1.0 x
10-7 moles of water molecules behave as Brønsted acids
and donate protons to another 1.0 x 10-7 moles of water
molecules, which act as Brønsted bases.. The reaction is:
H2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
• As a result, absolutely pure water contains 1.0 x 10-7 mol/L
of both H3O+ and OH-.
• The term neutral is used to describe any water solution in
which the concentrations of H3O+ and OH- are equal.
• Thus, pure water is neutral because each of the ions is
present at a concentration of 1.0 x 10-7 M.
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THE ION PRODUCT OF WATER
• The reaction given earlier for the formation of H3O+ and
OH- in pure water is called the self-ionization of water. The
reversible nature of the reaction (indicated by the double
arrow) means that an equilibrium is established and an
equilibrium expression can be written for the reaction. The
equilibrium expression is:
• This expression contains the square of the molar
concentration of water in the denominator. However, only a
tiny amount of water reacts to establish the equilibrium, so
the concentration of water remains essentially constant.
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• The equilibrium expression can be rearranged to give:
• Because the concentration of water is essentially constant,
the product of K multiplied by the square of the water
concentration is equal to another constant designated as
Kw, and called the ion product of water. The equation
then becomes:
• Because the molar concentration of both H3O+ and OH- in
pure water is 1.0 x 10-7 , the numerical value for Kw can be
calculated:
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• Even though this equilibrium equation was derived on the
basis of pure water, it is true for any solution in which water
is the solvent.
ACIDIC SOLUTION
• An acidic solution is a solution in which the concentration
of H3O+ is greater than the concentration of OH-. It is also
a solution in which the pH is less than 7.
BASIC OR ALKALINE SOLUTION
• A basic or alkaline solution is a solution in which the
concentration of OH- is greater than the concentration of
H3O+. It is also a solution in which the pH is greater than
7.
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EXAMPLE OF ACID-BASE CALCULATION
• Calculate the molar concentration of OH- in a solution that
has an H3O+ concentration of 1.0 x 10-5. Classify the
solution as acidic or basic.
• Solution: The molar concentration of H3O+ will be
substituted into the equilibrium expression for water, the
resulting equation will be solved for [OH-]:
• The molar concentration of OH- is seen to be smaller than
the molar concentration of H3O+, so the solution is
classified as being acidic.
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THE pH CONCEPT
• It is often the practice to express the concentration of H3O+
in an abbreviated form called the pH rather than to use
scientific notation.
• It is also a common practice to represent the H3O+ ion by
the simpler H+ ion.
• The pH notation is defined below, using H+ in place of
H3O+:
pH = -log[H+], or in alternate form [H+]= 1x10-pH
• Thus, the pH is seen to be the negative of the exponent
used to express the molar concentration of H+ using
scientific notation.
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EXAMPLES OF pH CALCULATIONS
• Example 1: Calculate the pH of a solution in which
[H+]= 1.0x10-9.
• Solution: Because the pH is the negative of the exponent
on 10 used to express [H+] using scientific notation,
pH= -(-9) = 9.
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EXAMPLES OF pH CALCULATIONS
Example 2: Calculate the [OH-] for a solution with a pH = 4.
• Solution: Because pH is the negative of the exponent on
10 used to express [H+] in scientific notation, the exponent
must by -4. Then, [H+]= 1.0 x 10-4. This value is
substituted into the equilibrium expression for water, and
the equation is solved for [OH-]:
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EXAMPLES OF pH CALCULATIONS
• Example 3: Calculate the pH of a solution in which
[H+]= 3.6x10-4.
• Solution: Use the pH equation, pH = -log [H+], to find
pH= -log [3.6 x 10-4], then evaluate with a calculator.
• The pH of the solution is 3.44.
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EXAMPLES OF pH CALCULATIONS
• Example 4: Calculate the [H+] of a solution in which
pH = 5.92.
• Solution: Use the alternate pH equation, [H+] = 1 x 10-pH,
to find [H+] = 1.0 x 10-5.92, then evaluate with a calculator.
• The [H+] of the solution is 1.2 x 10-6.
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PROPERTIES OF ACIDS
• All acids have certain properties in common such as tasting
sour and producing H3O+ ions when dissolved in water.
• In addition, all acids undergo characteristic doublereplacement reactions with solid oxides, hydroxides,
carbonates, and bicarbonates. These reactions are given
below, using hydrochloric acid, HCl(aq), as a representative
acid.
• Reaction with metal oxide:
2HCl(aq) + MgO(s)
MgCl2(aq) + H2O(l)
• Reaction with metal hydroxide:
2HCl(aq) + Mg(OH)2(s)
MgCl2 (aq) + 2 H2O(l)
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• Reaction with metal carbonate:
2HCl(aq) + MgCO3(s)
MgCl2 (aq) + CO2 (g) + H2O(l)
• Reaction with metal bicarbonate:
2HCl(aq) + Mg(HCO3)2(s)
MgCl2 (aq) + 2CO2 (g)
+ 2H2O(l)
• Notice each of these reactions can be rewritten in net ionic
form with the chloride spectator ions removed, which
shows that all acids share this reactivity.
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Marble, a naturally occurring form of
CaCO3, reacts with hydrochloric acid,
HCl.
Eggshells are also made of CaCO3.
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• Acids can react with and dissolve certain metals to yield
hydrogen gas in a redox reaction.
• The activity series is a tabular representation of the
tendencies of metals to react with H+.
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• The reaction of zinc metal with hydrochloric acid can be
written as follows:
• The chloride ion (Cl-) is a spectator ion.
• The hydrogen ion gains an electron to be reduced, and
therefore, the HCl is the oxidizing agent.
• The zinc metal loses electrons to be oxidized, and
therefore, the zinc metal is the reducing agent.
• This reaction occurs because zinc is above the reactivity
line that divides lead (reactive) from copper (unreactive) in
the activity series.
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Metals vary in their ability to reduce hydrogen ions (H+) to hydrogen
gas (H2). The difference is apparent when iron, zinc, and magnesium
(left to right) are put into hydrochloric acid (HCl) of the same molarity.
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PROPERTIES OF BASES
• Basic solutions feel soapy or slippery to the touch and
contain the OH- ion.
• Basic solutions also change the color of litmus from red to
blue.
• Their most characteristic chemical property is their ability
to react readily with acids in what is called a
neutralization reaction.
• Bases also react with fats and oils and convert them into
smaller, soluble molecules.
• Most household cleaning products contain basic
substances. (e.g. lye (NaOH) in drain cleaner and
ammonia (NH3) in liquid household cleaners)
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CLASSIFICATION OF HOUSEHOLD PRODUCTS
Weak Acids
Weak Bases
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NEUTRALIZATION REACTIONS
• In neutralization reactions, an acid reacts with a base to
produce a salt and water. The following are typical
neutralization reactions involving the base sodium
hydroxide, NaOH, which is also known commercially as
lye.
• Reaction with hydrochloric acid:
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
• The salt produced in this reaction is sodium chloride,
commonly called table salt.
• Reaction with nitric acid:
NaOH(aq) + HNO3(aq)
NaNO3(aq) + H2O(l)
• The salt produced in this reaction is sodium nitrate.
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• Reaction with sulfuric acid:
2NaOH(aq) + H2SO4(aq)
Na2SO4(aq) + 2H2O(l)
• The salt produced in this reaction is
sodium sulfate.
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SALTS
• At room temperature, salts are solid crystalline ionic
compounds that contain the cation (positive ion) of a base
and the anion (negative ion) of an acid in their formulas.
• Sodium chloride, NaCl, contains one Na+ cation from the
base NaOH, and one Cl- anion from the acid HCl in its
formula.
• Sodium nitrate, NaNO3, contains one Na+ cation from the
base NaOH, and one NO3- anion from the acid HNO3 in its
formula.
• Sodium sulfate, Na2SO4, contains two Na+ cations from the
base NaOH, and one SO42- from the acid H2SO4 in its
formula.
• The cation of a salt can be any positive ion, except H+, and
it will usually be a simple metal ion or NH4+.
• The anion of a salt can be any negative ion, except OH-.
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• Salts can be formed in a number of reactions:
SALT HYDRATES
• When salts are obtained from water solutions by
evaporating the water away, specific numbers of water
molecules are retained in some cases as a part of the
recovered solid salt. These solids, called hydrates, have
formulas that indicate the number of water molecules
retained by the solids called water of hydration.
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• A number of hydrates are very useful as indicated by the
entries in the following table:
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EQUIVALENT OF A SALT
• An equivalent of a salt is the amount of salt that will
produce 1 mole of positive (or negative) electrical charges
when dissolved and dissociated into ions.
• The number of moles of salt in an equivalent depends
upon the charges of the ions that make up the salt.
• For a salt like NaCl that dissociates into one Na+ ion and
one Cl- ion, 1 mole of salt produces 1 mole of positive
charges.
Thus, 1 mole of NaCl = 1 equivalent of NaCl.
• For a salt like MgCl2 that dissociates into one Mg2+ ion and
two Cl- ions, 1 mole of salt produces 2 moles of positive
charges.
Thus, 1 mole of salt = 2 equivalents of salt,
or 1/2 mole of salt = 1 equivalent of salt.
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THE STRENGTH OF ACIDS AND BASES
• The strength of an acid or base is determined by the extent
to which dissolved acid or base dissociates to form ions. A
strong acid or base dissociates 100%, while a weak or
moderately weak one dissociates less than 100%.
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ACID DISSOCIATION CONSTANTS
• An acid dissociation constant is the equilibrium constant for
the dissociation of a weak acid. It is represented by the
symbol Ka.
• The dissociation of a weak acid in solution is represented by
the following equation in which HB represents the weak
acid, and B- is the conjugate base of the acid.
HB(aq) + H2O(l)
H3O+(aq) + B-(aq)
• The equilibrium expression for this reaction is:
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• The molar concentration of water in the solution is
essentially constant and can be multiplied times K to form a
new constant Ka.
• When [H+] is substituted for [H3O+] a simplified form of the
equation results:
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MONOPROTIC, DIPROTIC AND TRIPROTIC ACIDS
• Monoprotic acids give
up only one proton per
molecule when
dissolved in water.
• Diprotic acids give up a
maximum of two protons
per molecule when
dissolved in water.
• Triprotic acids give up
a maximum of three
protons per molecule
when dissolved in water.
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COMMON BASES
• Ammonia (NH3) is the weak base most often encountered
in addition to the anions of strong acids.
• The most common strong bases are the hydroxides of
group IA(1) metals (NaOH, KOH, etc.) and the hydroxides
of group IIA(2) metals (Mg(OH)2, Ca(OH)2, etc.).
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Indicators
• An indicator changes color with changes in pH (the
numbers on the tubes).
Methyl red goes from red
at low pH to orange.
Phenolphthalein goes from
colorless to pink.
Bromthymol blue from low
pH to high pH.
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ANALYZING ACIDS AND BASES
• The analysis of acid solutions to determine the amount of
acid they contain is an important procedure done in many
laboratories.
• An acid-base titration is one commonly-used method of
analysis.
• When a titration is done, an accurately-measured volume
of acid is put into a flask using a pipet.
• A few drops of indicator solution is added, then a base
solution of known concentration is carefully added from a
buret until all the acid has been reacted (equivalence
point).
• The point at which all the acid has reacted is shown by a
color change (endpoint) in the indicator.
• The concentration of the base and the volume required in
the titration allow the concentration of acid to be
determined.
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TITRATION TECHNIQUE
40
pH METER
• A pH meter can also be used to detect the equivalence
point of a titration.
At the beginning, the
pH meter gives the
pH of the acid
solution being
titrated.
Partway through the titration, the
pH meter reading is of a solution
of unreacted acid and
the salt produced by the
reaction.
At the end of the titration, the pH
meter gives the pH of the salt
solution formed by the complete
reaction of acid with base.
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TITRATION CALCULATIONS
• Titration calculations are dependent upon knowledge of two
things: the stoichiometry of the reaction that occurs between
the acid and base, and the equation defining molarity.
• An example of a reaction equation is:
H2SO4(aq) + 2NaOH(aq)
Na2SO4(aq) + 2H2O(l)
• Such an equation provides the relationship between the
number of moles of acid and base that react. In this reaction it
is seen that 1 mole of H2SO4 acid reacts with 2 moles of NaOH
base.
• The molarity equation may be rearranged to allow the
calculation of the number of moles of solute contained in a
specific volume of solution or the volume of solution that
contains a specific number of moles of solute.
42
M x liters of solution = moles of solute
43
HYDROLYSIS REACTIONS OF SALTS
• Salts consist of the cation of a base and the anion of an
acid. The cation of a base is the conjugate acid of the
base from which it came. Similarly, the anion of an acid is
the conjugate base of the acid from which it came.
• The strength of a conjugate acid or base depends upon
the strength of the base or acid from which they came.
The stronger an acid is, the weaker is its conjugate base.
Similarly, the stronger a base is, the weaker is its
conjugate acid.
• The pH of a water solution of a salt depends on the
strength of the salt cation as an acid and the strength of
the salt anion as a base.
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• Example 1: A solution containing the dissolved salt NaCl
has a pH the same as the water used as a solvent for the
solution.
• This is because the Na+ ion is the conjugate acid of the
strong base NaOH and is a very weak acid.
• Similarly, the Cl- ion is the conjugate base of the strong
acid HCl and is a very weak base.
• Neither the Na+ cation nor the Cl- anion will react
appreciably with water to produce OH- or H+.
45
• Example 2: A solution containing the dissolved salt sodium
carbonate, Na2CO3, has a pH significantly higher than that
of the water used as a solvent for the solution.
• The Na+ ion is a weak acid as was discussed on the
previous slide.
• The CO32- ion is the conjugate base of the weak acid
HCO3- and as a result is a significant base that will react
with water as follows:
CO32-(aq) + H2O(l)
HCO3- + OH-
• This reaction, called a salt hydrolysis reaction, is seen to
produce OH- ions which causes the pH to be higher than
water and the solution is basic.
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PURE WATER vs. SODIUM ACETATE
• Samples of pure water (left) and sodium acetate dissolved
in water (right) behave differently when phenolphthalein
indicator is added. The acetate ion hydrolyzes in water to
form a basic solution that turns phenolphthalein to a pink
color.
47
BUFFERS
• Buffers are solutions with the ability to resist changing pH
when acids (H+) or bases (OH-) are added to them.
• Many useful buffers consist of a solution containing a
mixture of a weak acid and a salt of the acid (e.g. acetic
acid and sodium acetate).
• Any added acid (H+ ions) react with the anion from the salt,
which also happens to be the conjugate base of the weak
acid.
• Any added base (OH- ions) react with the nonionized weak
acid.
• The buffer capacity is the amount of acid (H+) or base
(OH-) that can be absorbed by a buffer without causing a
significant change in pH.
48
UNBUFFERED vs. BUFFERED SOLUTIONS
The solution on the left is not buffered;
the one on the right is; universal indicator
has been added to each solution.
Sodium hydroxide has been added to each solution
Hydrochloric acid has been added to two
fresh samples that originally looked like
the first pair of samples.
49
pH and BUFFERS
• The pH of buffers made this way can be calculated using
the Henderson-Hasselbalch equation:
• In this equation, pH = -log[H+], pKa= -log Ka, and [B-] is the
molar concentration of the salt of the weak acid HB, and [HB]
is the molar concentration of the weak acid.
• If [B-] is equal to [HB], then the pH is equal to the pKa.
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