Chapter 11
Acids, Bases, & Salts
Acid/Base Equilibrium
Properties of Acids
• sour or tart taste
• strong acids burn; weak acids feel similar
to H2O
• acid solutions are electrolytes
• acids react with most metals to release H2
• acids cause indicators to change color
– Acids turn litmus red
Properties of Bases
•
•
•
•
•
•
bases taste bitter
basic solutions do not burn
basic solutions feel smooth and slippery
basic solutions are also electrolytes
bases usually do not react with metals
bases also cause indicators to change
color
– Bases turn litmus blue
Three definitions
• Arrhenius
– Deals with H+ and OH-
• Bronsted Lowry
– Deals with protons
• Lewis
– Deals with pairs of electrons
Arrhenius Definition
• Arrhenius defined an acid as a substance
that ionizes in water to produce H+ ions
– Ex. HCl
– Any of our acids with H+ in the beginning
• Arrhenius defined a base as a substance
that ionizes in water to produce OH- ions
– Ex. NaOH
– Any of the metal hydroxides
Bronsted Lowry Definition
• NH3 is known to turn litmus blue, but it
does not have hydroxide ion
– Needed a new definition
– Substances do not need to be in water
• A B/L acid is any substance that can
donate a proton (H+)
• A B/L base is any substance that can
accept a proton (H+)
Conjugate pairs
HC2H3O2 + H2O → C2H3O2- + H3O+
NH3 + H2O → NH4+ + OH-
Amphoteric Substances
• Notice in the examples, that H2O can act
as a B/L acid and as a B/L base.
• It is know as an amphoteric substance.
• Additional examples include
– HSO3– HCO3– H2PO4– HPO42-
These are the ions that result
when polyprotic acids release
H+ ions one at a time. Notice
that they can accept an H+ or
they can release an H+.
Lewis definition
•
•
•
•
This is the most broad of the definitions
A Lewis acid is an electron pair acceptor.
A Lewis base is an electron pair donor.
Example Lewis acid/base reaction….
Ion product constant for H2O
• Remember [H+][OH-] = 1.0 x 10-14 = Kw
• So if you know [H+] you can find [OH-] and
vice versa
• If [H+] = 9.3 x 10-4 M, what is [OH-]
pH or pOH
• pH = -log[H+] or pOH = -log[OH-]
• pH + pOH = 14
• What is the pH of a solution that contains
4.9 x 10-9 M OH-?
What is the [H+] of a solution that is pH4.82?
What is the [OH-] of a solution that is pH8.56?
You must understand that…
• When [H+] = [OH-], the solution is neutral, and the pH = 7
• When [H+] > [OH-], the solution is acidic, and the pH <7
• When [H+] < [OH-], the solution is basic, and the pH > 7.
• Increasing pH means decreasing [H+]…fewer H+
ions floating in solution….less acidic.
• Decreasing pH means increasing [H+]…more H+
ions floating in solution….more acidic.
Weak Acids
• When a weak acid is placed in water, a small
fraction of its molecules will dissociate into
hydrogen ions and conjugate base ions.
– Most of the acid molecules will remain in
solution as undissociated aqueous particles.
• The dissociation constant, Ka, is a measure of
the strength of weak acids.
• For the reaction…HA ↔ H+ + AKa =
Weak Bases
• When a weak base is placed in water, a small
fraction of its molecules will dissociate into
hydroxide ions and conjugate acid ions.
– Most of the base molecules will remain in
solution as undissociated aqueous particles.
• The dissociation constant, Kb, is a measure of
the strength of weak bases.
• For the reaction…B ↔ HB+ + OHKb =
Acid Equilibrium problems
• Things to know
– All of our acids will be monoprotic, that is, give
off only one hydrogen ion.
– x will represent the amount of acid that
dissociates.
• Therefore x also represents the [H+]
– We will use ICE to determine concentrations
for the Ka expression
– We can use HA to represent the acid and A- to
represent the conjugate base.
Ka example
What is the pH of a 0.20M HC2H3O2, with Ka = 1.8 x 10-5?
(also find % ionization)
Polyprotic acids
• Polyprotic acids are weak acids (except for
H2SO4 which is a strong acid).
– They release their H+ ions one at a time.
– Each has its own Ka
– Ex. H3PO4
H3PO4 → H+ + H2PO4H2PO4- → H+ + HPO42HPO42- → H+ + PO43-
Ka = 7.5 x 10-3
Ka = 6.2 x 10-8
Ka = 4.8 x 10-13
Kw revisited
• Water comes to equilibrium with its ions
according to the following reaction…
H2O(l) ↔ H+(aq) + OH-(aq)
Kw = [H+][OH-] = 1.0 x 10-14
Kw = KaKb
pKa + pKb = 14
Strong Acids/Bases
Strong Acids
HCl
HBr
HI
HNO3
HClO4
H2SO4
Strong Bases
LiOH
NaOH
KOH
RbOH
CsOH
Ba(OH)2
Sr(OH)2
Strong Acids/Base description
• Strong acids and bases completely
dissociate in water, therefore no Ka or Kb
– The dissociations do not reverse.
• Oxoacids are acids that contain oxygen.
– The greater the number of oxygen atoms
attached to the central atom in an oxoacid,
the stronger the acid.
• That’s because increasing the number of oxygen
atoms that are attached to the central atom
weakens the attraction that the central atom has
for the H+ ion.
Strong Acid/Base calculations
• Since these acids and bases completely
dissociate in water, the final concentration
of H+ ions is the same as the original
concentration.
• So you can always find the pH of a strong
acid solution directly from its
concentration.
– What is the pH of .20 M HCl?
pH = -log(.20) =
Titration
• When an acid and a base are mixed, a
neutralization reaction occurs.
– Acid + base → salt + water
• Neutralizations reactions are generally
performed by titration, where a base of known
concentration is slowly added to an acid (or vice
versa)
• The progress of a neutralization reaction can be
shown in a titration curve.
• The equivalence point is the point in the titration
when exactly enough base has been added to
neutralize all the acid that was initially present.
• An indicator will be used to mark the equiv. pt.
• Strong acid/strong base titration.
– Equivalence point at pH7
• Notice the shape of the curve if a strong
acid is added to a strong base
– The equivalence point is also pH7
• Weak acid vs strong base
– The equivalence point is > pH7
• Weak base vs strong acid
– Notice the equivalence point is < pH7
Which indicator to use?
• The indicator needs to change color close
to the equivalence point.
pH of dissolved salts
• If a salt is composed of the conjugates of a
strong base and a strong acid, its solution will be
neutral. (NaCl)
• If a salt is composed of the conjugates of a weak
base and a strong acid, its solutions will be
acidic. (NH4Cl)
• If a salt is composed of the conjugates of a
strong base and a weak acid, its solution will be
basic. (NaC2H3O2)
• If a salt is composed of the conjugates of a weak
base and a weak acid, the pH of its solution will
depend on the relative strengths of the
conjugate acid and base of the specific ions in
the salt. (NH4C2H3O2)
Anhydrides
• An acid anhydride is a substance that
combines with water to form an acid.
– Generally nonmetal oxides.
• Ex. CO2 + H2O → H2CO3
• A basic anhydride is a substance that
combines with water to form a base.
– Generally metal oxides.
• Ex. Na2O + H2O → 2NaOH
What to know for the test?
• Find pH from either [H+] or [OH-]
• Find [H+] from pH or pOH
• Find the # of grams of a base to make a
solution of a certain pH
• Use titration info to find molar mass acid
• Ka problems
– Find pH
– Find concentrations
– Find % ionization
- find Ka from pH