Ch. 20: Acids and Bases
Ch. 20.1
 Ch. 20.2
 Ch. 20.3
 Ch. 20.4

KMHS Magnet Academy Staff - Tammy Brown
Describing Acids and Bases
Hydrogen Ions and Acidity
Acid-Base Theories
Strengths of Acids and Bases
Ch. 20.1 Describing Acids and Bases

Properties of Acids and Bases
– Acids
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Produce H+ ions when dissolved in water
Sour taste
Solutions are electrolytes (some strong, some weak)
React with metals to produce H2
React with a base to form water and salt
Turn litmus paper red
– Bases
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Produce OH- ions when dissolved in water
Bitter taste
Feel slippery
Solutions are electrolytes (strong and weak)
React with acids to form water and a salt
Turn litmus paper blue
Ch. 20.1 Describing Acids and Bases

Names and formulas of acids and bases
– Acids
• Acids have a hydrogen ion
• The general formula for an acid is HX, where the X is
a monatomic or polyatomic ion
– Bases
• Bases have an OH- ion
• Ionic compounds that are bases are named like any
other ionic compound
– See Table 20.1, pg. 578
Ch. 20.2 Hydrogen Ions and Acidity

Hydrogen ions from water
– Water molecules that gain a hydrogen
ion become a hydronium ion (H3O+)
– Water molecules that lose a hydrogen
ion become a hydroxide ion (OH-)
– In pure water, the concentration of H+
and OH- ions are each 1.0 x 10-7 M
• This means that the concentration of each
are equal in pure water
• Described as a neutral solution
Ch. 20.2 Hydrogen Ions and Acidity

Hydrogen ions from water
– In any aqueous solution, the [H+] and [OH-] are
interdependent
• When [H+] decreases, the [OH-] increases
• For aqueous solutions, [H+] x [OH-] = 1.0 x 10-14
• This is also known as the ion-product constant for
water (Kw)
• An acidic solution is one in which the [H+]
concentration is greater than the [OH-]
– Therefore, the [H+] is greater than 1 x 10-7
• A basic (alkaline) solution is one in which the [OH-]
is greater than than the [H+] concentration
– Therefore, the [H+] is less than 1 x 10-7
Ch. 20.2 Hydrogen Ions and Acidity

The pH concept
– Expressing concentration in molarity is
inefficient, so we use a pH scale
• The scale ranges from 1 to 14
 1 is very acidic, 7 is neutral, and 14 is very basic
• The pH of a solution is the negative logarithm of the
hydrogen-ion concentration
 pH = -log [H+]
• The pOH of a solution equals the negative logarithm
of the hydroxide-ion concentration
 pOH = -log [OH-]
• pH + pOH = 14
Ch. 20.2 Hydrogen Ions and Acidity

Calculating pH values
– Most pH values are not whole numbers
– You can calculate the hydrogen-ion
concentration of a solution if you know
the pH
• If the pH is 3, then [H+] = 1.0 x 10 –3
• If the pH is not a whole number, you will
need a calculator to find antilog
[H+] = -pH antilog
Ch. 20.2 Hydrogen Ions and Acidity

Measuring pH
– A pH meter is preferred for precise
measurements
• Must be calibrated first by dipping the electrodes in a
solution of known pH
• It is then rinsed and used to measure the pH of an
unknown solution
– Acid-base indicators
• An indicator is an acid or base that dissociates in a
known pH range
• See Fig. 20.8, pg. 590
• These have limitations
– Some have a specific temperature range
– Do not work well in colored/cloudy solutions
– Can be affected by dissolved salts
Ch. 20.3 Acid-Base Theories

Arrhenius acids and bases
– Acids dissociate in water to produce H+
ions
– Bases dissociate in water to produce OHions
– The Arrhenius definition is very broad
• Does not include certain chemicals that
should be classified as an acid or base
• NH3 and Na2CO3 are both bases but would
not be classified as such under the Arrhenius
definition
Ch. 20.3 Acid-Base Theories
– Types of acids
• Monoprotic acids – acids that contain one ionizable
hydrogen
• Diprotic acids – acids that produce two ionizable
hydrogens
• Triprotic acids – acids that contain three ionizable
hydrogens
– Not all compounds that contain H are acids
– Not all hydrogens in an acid may be released
– Acid and base strength is based on
solubility
• Greater dissociation means greater strength
• Group 1 metals are more soluble than Group 2
metals
Ch. 20.3 Acid-Base Theories

Bronsted-Lowry acids and bases
– Defines an acid as a hydrogen-ion donor
(proton donor)
– Defines a base as a hydrogen-ion acceptor
(proton acceptor)
– Conjugate acid-base pairs
• A conjugate acid is the particle formed when a base
gains a hydrogen ion
• A conjugate base is the particle that remains when an
acid has donated a hydrogen ion
• A conjugate acid-base pair is made up of two
substances related by the loss or gain of a single
hydrogen ion
– Water is amphoteric (amphoprotic) – it can accept or
donate a hydrogen ion
Ch. 20.3 Acid-Base Theories

Lewis acids and bases
– Focuses on the donation or acceptance of a
pair of electrons during a reaction
• More general than the Arrhenius or Bronsted-Lowry
definitions
• A Lewis acid is one that can accept a pair of
electrons to form a covalent bond
• A Lewis base is one that can donate a pair of
electrons to form a covalent bond
• See Table 20.6, pg. 598 for a summary of the three
definitions
Ch. 20.4 Strengths of Acids and
Bases

Strong and weak acids and bases
– Acids
• Strong acids are completely ionized in
aqueous solution
• Weak acids are only partially ionized in
aqueous solutions
– See Table 20.7, pg. 600 for a list of acids/bases
• Ka is the acid dissociation constant
 The ratio of the concentration of the dissociated
acid to the concentration of the undissociated
acid in a solution
[H  ] x [anion]
Ka 
Undissocia ted acid
Ch. 20.4 Strengths of Acids and
Bases

Ka
– Reflects the fraction of an acid formed
• If the Ka is small, then the then the degree of
dissociation is small (weak acid)
• If the Ka is large, then the degree of
dissociation is large (strong acid)
– Diprotic and triprotic acids lose their H+
ions one at a time
• Each reaction has its own Ka
Ch. 20.4 Strengths of Acids and
Bases

Strong and weak acids and bases
– Bases
• Strong bases dissociate completely into metal ions
and hydroxide ions in aqueous solutions
• Weak bases react with water to form the hydroxide
ion and the conjugate acid of the base
• Kb is the base dissociation constant
– The base dissociation constant (Kb) is the ratio of
the concentration of the hydroxide ion to the
concentration of the conjugate base
– The smaller the value of Kb, the weaker the base
[conjugate acid] x [OH  ]
Kb 
conjugate base
Ch. 20.4 Strengths of Acids and
Bases

Strong and weak acids and bases
– Concentrated and dilute refer to how
much of an acid or base is dissolved in
solution
• Moles of acid or base per liter
– Strong and weak refer to the extent of
ionization or dissociation of an acid or
base
• a solution of ammonia, whether
concentrated or dilute, will be a weak base
because the ionization NH3 will be small
Ch. 20.4 Strengths of Acids and
Bases

Calculating dissociation constants
(Ka and Kb)
– It is possible to calculate Ka and Kb
from experimental data
– First, measure the equilibrium
concentrations of all the substances
present at equilibrium
– Then put into the Ka or Kb formula
• See Sample Problem 20-8, pg. 604
The End