Spencer L. Seager
Michael R. Slabaugh
www.cengage.com/chemistry/seager
Chapter 9
Acids, Bases, and Salts
Jennifer P. Harris
ARRHENIUS ACIDS & BASES
• 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)
H+ (aq) + NO3- (aq)
• KOH is a base: KOH(aq)
K+ (aq) + OH- (aq)
BRØNSTED ACIDS & BASES
• 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.
EXAMPLE OF A BRØNSTED ACID & BASE
(continued)
HNO2 (aq) + H2O (l) ⇆ H3O+ (aq) + NO2− (aq)
• 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.
• 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+.
CONJUGATE ACIDS & BASES
• 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+).
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 dissolved in water. The names of the acid
solutions, such as the hydrochloric acid given
above, can be obtained by following four rules.
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".
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.
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.
THE ION PRODUCT OF WATER
• The reaction given earlier for the formation of H3O+ and OHin 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:

H O OH 
K

-
3
H2O
2
• 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.
THE ION PRODUCT OF WATER (continued)
• The equilibrium expression can be rearranged to give:

KH2O  H3 O
2

OH 
-
• 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:

-

K W  H3 O
OH 
• 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:

K W  H3 O

OH   1.0  10 
-
7 2
 1.0  10
14
THE ION PRODUCT OF WATER (continued)
• 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.
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 M. 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-]:


K W  1.0  10 14  1.0  10 5 OH 



14
1
.
0

10
9
OH 
 1.0  10
5
1.0  10
• The molar concentration of OH- is seen to be smaller than
the molar concentration of H3O+, so the solution is classified
as being acidic.
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.
EXAMPLES OF pH CALCULATIONS
• Example 1: Calculate the pH of a solution in which
[H+]= 1.0x10-9 M.
• Solution: Because the pH is the negative of the exponent on
10 used to express [H+] using scientific notation,
pH = -log (1.0x10-9) = -(-9) = 9.00.
EXAMPLES OF pH CALCULATIONS
(continued)
• Example 2: Calculate the [OH-] for a solution with a pH = 4.0.
• Solution: Because pH is the negative of the exponent on 10
used to express [H+] in scientific notation, the exponent must be
-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-]:
1.0  10


14

 1.0  10
14
4
 OH 

1.0  10
10
OH 
 1.0  10 M
4
1.0  10
-
EXAMPLES OF pH CALCULATIONS
(continued)
• Example 3: Calculate the pH of a solution in which
[H+]= 3.6x10-4 M.
• 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.
EXAMPLES OF pH CALCULATIONS
(continued)
• 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 M.
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) + 2H2O(l)
PROPERTIES OF ACIDS (continued)
• 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.
2H (aq)  MgO(s)  Mg2  (aq)  H2O(l)
2H (aq)  Mg(OH)2 (s)  Mg2  (aq)  2H2O(l)
2H (aq)  MgCO3 (s)  Mg2  (aq)  CO2 (g)  H2O(l)
2H (aq)  Mg(HCO3 )2 (s)  Mg2  (aq)  2CO2 (g)  2H2O(l)
PROPERTIES OF ACIDS (continued)
Marble, a naturally occurring form of
CaCO3, reacts with hydrochloric acid,
HCl.
Eggshells are also made of CaCO3.
PROPERTIES OF ACIDS (continued)
• 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+.
PROPERTIES OF ACIDS (continued)
• The reaction of zinc metal with hydrochloric acid can be
written as follows:
Molecular equation : Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g)
+ 2H+ (aq)
Zn2+ (aq)
Total ionic equation : Zn (s)
→
+ H2 (g)
+ 2Cl− (aq)
2Cl− (aq)
Net ionic equation: Zn (s) + 2H+ (aq) → Zn2+ (aq) + H2 (g)
• 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.
PROPERTIES OF ACIDS (continued)
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.
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).
CLASSIFICATION OF HOUSEHOLD
PRODUCTS
Weak Acids
Weak Bases
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.
NEUTRALIZATION REACTIONS (continued)
• Reaction with sulfuric acid:
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
• The salt produced in this reaction is
sodium sulfate.
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- anion 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-.
SALTS (continued)
• Salts can be formed in a number of reactions:
Acid + metal
→ salt + H2
Acid + metal oxide
→ salt + H2O
Acid + metal hydroxide → salt + H2O
Acid + metal carbonate → salt + H2O + CO2
Acid + metal bicarbonate → salt + H2O + CO2
• 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.
SALTS (continued)
• A number of hydrates are very useful as indicated by the
entries in the following table:
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.
THE STRENGTH OF ACIDS & 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%.
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:

H O B 
K

3
-
HBH2O
ACID DISSOCIATION CONSTANTS (continued)
• The molar concentration of water in the solution is essentially
constant and can be multiplied times K to form a new
constant Ka.
KH2 O  K a

H O B 


3
-
HB
• When [H+] is substituted for [H3O+] a simplified form of the
equation results:

-
Ka

H B 

HB
MONOPROTIC, DIPROTIC &
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.
COMMON BASES
• Ammonia (NH3) is the weak base most often encountered in
addition to the anions of strong acids.
NH3 (aq) + H2O (l) ⇆ NH4+ (aq) + OH− (aq)
• 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.).
Indicators
• An indicator changes color with changes in pH (the numbers
on the tubes).
Methyl red goes from red
at low pH to orange.
Bromthymol blue from
low pH to high pH.
Phenolphthalein goes from colorless to pink.
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.
TITRATION TECHNIQUE
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 Partway through the titration,
titrated.
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.
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.
TITRATION CALCULATIONS
moles of solute
M
liters of solution
moles of solute
liters solution 
M
M x liters of solution = moles of solute
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 its conjugate base is.
Similarly, the stronger a base is, the weaker its conjugate
acid is.
• 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.
HYDROLYSIS REACTIONS OF SALTS
(continued)
• 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+.
HYDROLYSIS REACTIONS OF SALTS
(continued)
• 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− (aq) + OH− (aq)
• 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.
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.
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.
C2H3O2− (aq) + H+ (aq) ⇌ HC2H3O2 (aq)
• Any added base (OH- ions) react with the nonionized weak acid.
HC2H3O2 (aq) + OH− (aq) ⇌ C2H3O2− (aq) + H2O (l)
• 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.
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.
pH and BUFFERS
• The pH of buffers made this way can be calculated
using the Henderson-Hasselbalch equation:
BpH  pK a  log
HB
 
• In this equation, pH = -log[H+], pKa= -log Ka, [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.
Ka & pKa VALUES