Chapter 16 Notes- Acid-Base Equilibria
16.1 Acids and Bases: A Brief Review[p.668]
1. Acids properties
a. Sour taste
b. Causes certain dyes to change color
c. Arrhenius acid donates H+ ions in an aqueous solution
2. Bases
a. Bitter
b. Slippery
c. Arrhenius base donates OH- ions in an aqueous solution
3. Combining acids and bases characteristic properties begin to disappear
4. Arrhenius acid definitions is restricted to aqueous solutions
16.2 Bronsted-Lowry Acids and Bases [p.668]
1. The H+ ions in water
a. H+ ion is aqueous solution interact with the polar water molecule forming a
Hydronium Ion, H3O+
b. H3O+ can form additional hydrogen bonds to form large clusters of hydrated ions
c. H+ and H3O+ used interchangeably but H3O+ better represents what actually happens
2. Proton-Transfer Reactions
a. Bronsted-Lowry definitions
i. Bronsted-Lowry Acid
a. is a substance (molecule or ion) that donates a proton to another
substance
b. Must have a hydrogen ion ( proton) to donate
ii. Bronsted-Lowry Base
a. is a substance that accepts a proton
b. Must have a nonbonding pair of electrons that can bond to the proton
b. Acid and base reactions do not have to occur in an aqueous solution
c. Acid and bases always work together in pairs
d. Some substances can act as an acid or base (amphiprotic)
2. Conjugate Acid-Base Pairs
a. Found in acid-base equilibrium
b. Conjugate acid-base pair – acid and base differ only in the presence or absence of a
proton
i. Conjugate Base formed by removing a proton from the acid
ii. Conjugate Acid formed by adding a proton to a base
3. Relative Strength of Acids and Bases
a. Strength of acid is dependent on the ability of an acid or base to dissociate or ionize
b. The stronger an acid, the weaker is its conjugate base. The weaker the acid, the
stronger its conjugate acid
c. Three broad categories of acids and bases
i. Strong acids – completely transfer protons (conjugate base weak)
ii. Weak acids- partially dissociate (conjugate base is also weak)
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Chapter 16 Notes- Acid-Base Equilibria
iii. Negligible acidity- does not demonstrate acid behaviors (conjugate base is
strong)
d. In every acid-base reaction the position of the equilibrium favors transfer of the
proton from the stronger acid to the stronger base to form the weaker acid and the
weaker base
NOTE memorize strength of acid and base chart on p.672
16.3 The Autoioniaztion of Water [p. 673]
1. Water is unique because it can be a Bronsted-Lowry acid or base depending on circumstances
a. Autoionization – process whereby water spontaneously forms low concentrations of
H+ and OH- by transfer of a proton
b. Percentage of ions formed is very low making water a poor conductor
2. The Ion Product of Water
a. Equilibrium-constant Kc for water is symbolized by Kw (ion-product constant)
b. Formula excludes liquid water ---------Kw=[H3O+][OH-] or Kw=[H+][OH-]
c. Kw at 25˚C = 1.0 X 10-14
d. Kw equation is applicable to pure water and any aqueous solution
e. When [H+] = [OH-], the solution is neutral
f. In acidic solutions, the [H+] exceeds the [OH-]
g. In basic solutions, the [OH-] exceeds the [H+]
16.4The pH Scale [p.675]
1. pH scale
a. Molecular concentration of H+ in an aqueous solution is usually very small and
typically expressed in terms of the pH scale
b. pH is the negative logarithm in base 10 of [H+]
pH= -log[H+]
c. If [H+] = 1.0 x 10-7, then the pH is 7.00 (sig figs are represented by the two decimal
places.
d. Note: pH decreases as [H+] increases
i. pH = 7.00 means neutral
ii. pH < 7.00 indicate acid
iii. pH > 7.00 indicate base
e. Changes in the [H+] concentration affect the kinetic rate law. Therefore small changes
in the hydrogen ion concentration can have dramatic effects
c. pOH and Other “p” scales
a. pOH is the negative logarithm in base 10 of [OH-]
pH= -log[OH-]
b. pH meters are used to measure the actual pH of a substance
c. indicators can be used to find the pH range
16.5 Strong Acids and Bases [p.679]
1. Strong Acids
a. Strong electrolytes
b. Completely dissociate (not equilibrium)
c. Strong monoprotic acids (one H+)
i. HCl
ii. HBr
iii. HI
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Chapter 16 Notes- Acid-Base Equilibria
iv. HNO3
v. HClO3
vi. HClO4
d. Strong diprotic acid H2SO4
e. In aqueous solutions containing strong acids, acid is the only source of [H+]
2. Strong Bases
a. Most common are ionic hydroxides containing a group 1 metal or heavier group 2 metal
(NaOH, KOH, Ca(OH)2)
b. Completely dissociate
c. LiOH , RbOH, and CsOH not common
d. Calcium hydroxide, strontium hydroxide, and barium hydroxide have limited
solubilities
e. Strong bases are formed from ionic oxides
16.6 Weak Acids [p. 681]
1. General information regarding weak acids
a. Most acids are weak therefore they only partially ionize in an aqueous solution
b. Since an equilibrium be established, the equilibrium-constant known as the aciddissociation constant (Ka) is calculated the same way as the (Kc). Water is dropped as it is
a pure liquid used as the solvent.
[ H 3O  ][ A  ] [ H  ][ A  ]
Kc  Ka 

[ HA]
[ HA]
c. Many weak acids are organic compounds composed of carbon, hydrogen, and oxygen
d. Acidic behavior of an organic compound is never the result of hydrogen attached to a
carbon. It is due to the hydrogen attached to oxygen.
e. Magnitude of Ka indicates the tendency of an acid to ionize. Larger the value, stronger the
acid.
2. Calculating Ka from pH
a. Use same techniques used in calculating chemical equilibriums
b. Important to realize the proton-transfer is rapid
3. Percent Ionization
a. Another measurement for the strength of an acid. Stronger the acid, greater the percent
ionization.
concentration.ionized
percent _ ionization 
x100
original .concentration
b. For acid, concentration ionized = [H+]
]equilibrium
percent.ionization  [ H[HA
]initial
4. Using Ka to Calculate pH
a. Process
i. Step 1: write ionization equilibrium
ii. Step 2: write equilibrium-constant expression
iii. Step 3: express concentrations that are involved in the equilibrium reaction
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Chapter 16 Notes- Acid-Base Equilibria
iv. Step 4: substitute the equilibrium concentrations into the equilibrium-constant
expression
b. General rule- if the quantity of x is greater than 5% of the initial value it is better to use the
quadratic formula
c. pH of a weak acid with the same concentration should be higher than the pH of strong acid
d. properties of acid relate directly to [H+] are less evident with weak acids
e. The percent ionization decreases as the concentration increases therefore % ionized v acid
concentration is a inverse relationship.
4. Polyprotic Acidsa. Acids containing more than one ionizable H atom

H 2 SO3 
 H   HSO3 ............K a1  1.7 x10 2

HSO3 
 H  SO3
2
...............K a 2  6.4 x10 8
b. It is always easier to remove the first proton from a polyprotic acid than to remove the
second. (rule hold true for the removal of 3 H atoms)
c. Remember the larger the Ka, the stronger the acid. H2SO4 completely ionizes the 1st H
atom but hydrogen sulfate ion is a weak acid.
16.7 Weak Bases [p. 690]
1. Many substances behave as weak bases in water (substance abstracts H atom form H2O)
2. Equilibrium-constant expression for a base that reacts with water to form corresponding
conjugate base acid and OH- is called the base-dissociation constant
[ BH ][OH ]
Kb 
[ B]
3. Bases contain one or more lone pairs of electron that can bond with H+
4. Types of weak bases
a. 1st category: neutral substances that have a nonbonded pair of electrons
i. Usually contain a nitrogen
ii. Examples include ammonia and amines
b. 2nd category: anions of weak acids
16.8 Relationship Between Ka and Kb [p. 693]
1. The quantitative relationship of the conjugate acid-base pair is seen when two equations one for
the acid and one for the base are added together. After removal of the intermediates the overall
equation becomes
H 2O 
 H   OH 
2. Equilibrium for two equations added together = Ka x Kb
[ NH 3 ][ H  ] [ NH 4 ][OH  ]
K a xKb  (
)(
)  [ H  ][OH  ]  K w

[ NH 3 ]
[ NH 4 ]
3. The product of the acid-base dissociation constant for an acid and the base-dissociation
constant for the conjugate base equals the ion-product constant for water
4. Important to remember that this is only true of conjugate acid-base pairs
16.9 Acid-Base Properties of Salt Solutions [p. 695]
1. General background
a. Ions from dissociated ionic compounds can be acidic or basic
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Chapter 16 Notes- Acid-Base Equilibria
b. Ions from ionic compounds can reach with H+ (aq) and OH- (aq). This process is
known as hydrolysis.
2. An Anion’s Ability to React with Water
a. In general an anion X- in solution, can be consider the conjugate base to an acid.
b. If the X- forms a strong acid when a proton is added, then the anion has virtually no
effect on the pH.
c. If the acid formed by the additional of a proton to X- is weak, then the anion is a weak
base and will have some effect on the pH
d. Amphiprotic anions like HSO4- can act as an acid or base depending on the relative
magnitude of the Ka and Kb
i. Ka > Kb solution is acidic
ii. Ka < Kb solution is basic
3. A Cation’s Ability to React with Water
a. Cations just like anions can affect pH
b. NH3’s conjugate acid is NH4+. The ammonium ion can lower the pH.
c. Most metal ions can react with water to decrease pH. Exceptions would be alkali metals
and heavier alkaline earth metals.
4. Combined Effect of Cation and Anion in Solution- when a solution contains both the conjugate
base of a weak acid and the conjugate acid of a weak base, the ion with the larger equilibrium
constant will have the greater influence on the pH.
16.10 Acid-Base Behavior and Chemical Structure [p.699]
1. Factors that affect Acid Strength
a. Polar molecules will transfer a proton
i. Ionic hydrides will not because hydride ion is negative (behaves as a proton
acceptor.
ii. Nonpolar H-X bonds have no effect
b. Strength of bond- stronger H-X bond, lower dissociation.
c. HX ionization is also affected by the stability of the conjugate base. Greater the stability
of the conjugate base, the stronger the acid.
2. Binary Acids
a. Strength of H-X bond decrease as X increases in size causing acidity to increase and
basicity to decrease down a group.
b. Acidity increases with an increase in the Electronegativity of X across a period left to
right. The opposite is true for a base.
3. Oxyacids
a. Oxyacids that have the same # of OH groups as O atoms, acid strength increases with
increasing electronegativity.
b. Oxyacids that have the same central atom show an increase in acidity as the number of
O atoms attached to the central atom increases. [Acidity increases as the oxidation
number of the central atom increases.
4. Carboxylic
a. Acids that contain a carboxyl group –COOH
b. –COOH attachment increases the polarity of the molecule causing the formation of a
stable conjugate base
c. Conjugate base exhibits resonance which also increases stability
16.11 Lewis Acids and Bases [p. 704]
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Chapter 16 Notes- Acid-Base Equilibria
1. General information
a. Lewis acid is an electron-pair acceptor
b. Lewis base is an electron-pair donor. Electron-pair does not have to be donated to a
hydrogen
c. The broadening of the definition allow for acids and bases to be considered that do not
contain H+ and OH2. Hydrolysis of Metal Ions
a. Metals are acids based on the Lewis acid definition because cations accept electrons.
b. During hydration, the metal ion acts as the acid and the water molecule act as the base.
c. Ka for hydrolysis reactions are generally increase with an increasing charge and a
decreasing ionic radius
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