Spencer L. Seager
Michael R. Slabaugh
www.cengage.com/chemistry/seager
Chapter 9
Acids, Bases, and Salts
Jennifer P. Harris
DEFINITIONS OF 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
• HC2H3O2(aq) + H2O(l)
H3O+ (aq) + C2H3O2-(aq)
• In this reaction, HC2H3O2 behaves as a Brønsted acid by
donating a proton to the H2O. The H2O behaves as a
Brønsted base by accepting the proton.
BRØNSTED ACIDS & BASES
• NH3(aq) + H2O(l)
NH4+ (aq) + OH-(aq)
• In this reaction, NH3 behaves as a Brønsted Base by
accepting a proton from the H2O. The H2O behaves as a
Brønsted acid by donating a proton.
EXAMPLES OF ACIDS
Memorize the following acids by name and formula
HCl
Hydrochloric Acid
H2SO4
Sulfuric Acid
HNO3
Nitric Acid
H3PO4
Phosphoric Acid
HC2H3O2
Acetic Acid (an organic acid)
H2CO3
Carbonic Acid (carbonated water)
[salts of bases – later]
EXAMPLES OF ACIDS
A strong acid
or base
dissociates100%,
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. [similar to an equilibrium constant covered in
Chapter 8]
• 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 B 


Ka
HB
-
EXAMPLES OF BASES
Memorize the following bases by name and formula
Metal hydroxides, e.g., NaOH (lye), Ca(OH)2, Mg(OH)2 (milk of
magnesia)
NH3
HCO3-
Ammonia
Bicarbonate compounds (NaHCO3 – is common
baking soda)
Strong Bases – NaOH, KOH, Ca(OH)2
Weak Bases – Mg(OH)2, NH3, HCO3- [note: your author describes
Mg(OH)2 as a strong base]
[salts of acids – later]
PROPERTIES OF ACIDS and BASES
•
All acids
• Taste sour
• Turn blue litmus red
• Produce H3O+ ions when dissolved in water (or donate Hydrogen ions)
• React with oxides and hydroxides to form water and a salt (the reaction with
hydroxides is the classic neutralization reaction)
• React with carbonates, and bicarbonates to form a salt, water and carbon
dioxide
• Non-metallic oxides plus water form oxyacids (e.g., SO2 acid rain, CO2)
• React with active metals to form hydrogen gas and a salt.
PROPERTIES OF ACIDS and BASES
•
All bases
• Taste bitter
• Turn red litmus blue
• Feel slippery (due to their caustic nature)
• Produce OH- when added to water (or accept hydrogen ions)
• Can be formed by a reaction between a metallic oxide and water (e.g., K2O,
CaO)
• React with acids to form water and a salt (this is the classic neutralization
reaction).
• Bases also react with fats and oils and convert them into smaller, soluble
molecules (soap – related to 3rd bullet above).
PROPERTIES OF ACIDS (continued)
• Acids can react with and dissolve active 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+.
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.
• Define amphoteric
THE ION PRODUCT OF WATER (continued)
This gives rise to an expression called the ion product of
water and is termed Kw.

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
-
Hydrolysis of salts
• When an simple acid (or base) is added to water, the solution
becomes acidic (or basic).
• But when a salt is added to water the solution could be
neutral, acidic or basic depending on the nature of the salt.
Why?
• This is related to
• Conjugate acids and bases
• Hydrolysis of salts
• Buffers
[the next three topics]
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.
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+).
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.
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.
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) or in general the conjugates of a weak acid/base pair.
• 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.
Common buffers
• Acetic acid/acetate
• Bicarbonate/carbonate
• Carbonic acid(aq. CO2)/bicarbonate (a common buffer in
blood)
• Dihydrogen phosphate/monohydrogen phosphate
Explain/show how a buffer can resist pH change using
acid/base theory and Le Chatelier’s principle.
• The formal Chapter 9 lecture will stop about here.
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
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