NOTES – Acids, Bases and pH, Chapter 14
Terms:
Acid dissociation constant, Ka
Amphoteric substance
Basic oxide
the equilibrium constant for a reaction in which a proton is removed
from an acid by the H2O to form a conjugate base and H30+
a covalent oxide that dissolves in water to give an acidic solution
an organic base derived from ammonia in which one or more of the
hydrogen atoms are replaced by organic groups
a substance that can behave either as an acid or a base
an ionic oxide that dissolves in water to produce a basic solution
Carboxyl group
the –COOH group in an organic acid
Conjugate base
Conjugate acid
Conjugate acid-base pair
what remains of an acid molecule after a proton is lost
the species formed when a proton is added to a base
two species related to each other by the donating and accepting of a
single proton
an acid having two acidic protons (ex: sulfuric acid)
the H30+ ion; a hydrated proton
The equilibrium constant for the auto ionization of water; (Kw)=
[H+][OH-] at 25 ºC. (Kw)= 1.0x10-14
an electron-pair acceptor
an electron-pair donor
a water-softening method in which lime and soda ash are added to
water to remove calcium and magnesium ions by precipitation
an acid with one acidic proton
an acid in which the acidic proton is attached to an oxygen atom
an acid with a carbon-atom backbone; often contains the carboxyl
group
Acidic oxide
Amine:
Diprotic acid
Hydronium Ion
Ion- product constant, (Kw)
Lewis Acid
Lewis base
Lime-soda process
Monoprotic acid
Oxyacids
Organic acid
pH scale
Major species
Percent dissociation
Polyprotic Acid
Salt
Slaked lime
Strong acid
Strong bases
Triprotic Acid
Weak acid
Weak base
a log scale based on 10 and equal to –log[H+]; a convenient way to
represent solution acidity that ranges from 0 - 7(neutral/H2O) - 14
the components present in relatively large amounts in a solution
the ratio of the amount of a substance that is dissociated at equilibrium
to the initial concentration of the substance in the solution multiplied by
100.
an acid with more than one acidic proton. It dissociates in a stepwise
manner, one proton as a time
an ionic compound
calcium hydroxide
an acid that completely dissociates to produce an H+ ion and the
conjugate base
a metal hydroxide salt that completely dissociates into its ions in water
three protons that dissociate (ex: Phosphoric acid)
an acid that dissociates only slightly in aqueous solutions
a base that reacts with water to produce hydroxide ions to only a slight
extent in aqueous solution.
Misc Concepts
I.
Brønsted -Lowery
 Acid- H+ donor (labile O-H bond)
 Base- H+ acceptor (has lone e- pairs)
NOTES – Acids, Bases and pH, Chapter 14
Ex: Brønsted -Lowery acid: HNO2
Brønsted -Lowery base: SO32II.
Arrhenius
 Acid- forms H3O+ ions in solution
 Base- forms OH- ions in solution
III.
Lewis
 Acid- e- pair acceptor
 Base- e- pair donor
Ex: BF3 + NH3 → BF3NH3
L.A. L.B.
Acceptor Donor
IV.
Acids and Acid strength
 Strong acid- 100% dissociates in water
 HCl, HBr, HI, H2SO4, HNO3, HIO3, HClO4
 Weak acid < 100% dissociates in water
 Acid strength is based on strength of the bond that includes the “H” that becomes H+
 Stronger bond = weaker acid and weaker bond = stronger acid
 Dissociation equation for a strong acid: HA + H2O → H3O+ + A–
 Dissociation equation for a weak acid: HA + H2O
H3O+ + A–
+
[H O ][A ]
 Acid dissociation constant: Ka = 3
[HA]
V.
Bases and Base strength
 Strong bases- release OH- into solution
 100% dissociates in aqueous solution
 Group 1 and some of group 2 (Ca, Ba) hydroxide
Ex: NaOH, KOH (generically written as … MtOH)
 Dissociation equation for a strong base: MtOH (s) → Mt+ + OH–
 Dissociation equation for a weak base:
if no charge: B + H2O
HB+ + OH1or for an anion B1- + H2O
HB + OH1[HB+ ][OH - ]
[HB][OH- ]
 Base dissociation constant: Kb =
Kb =
[B]
[B- ]
 Weak bases- are typically molecules with CO21- in them or amines (have NH groups)
for example CH3CO21- and C2H5NH2
VI.
Water as an acid and a base
 Amphoteric- can be either an acid or a base
 Water is most common
Ex: 2H2O → H+ + OHKw = [H+][OH¯] if [H+] = [OH-] = 1.0 x 10-7 M
Kw = ( 1.0 x 10-7)(1.0 x 10-7)
Kw = 1.0 x 10-14
VII
The pH Scale
 The pH scale is a scale from 1-14 to represent solution acidity
 It is a log scale based on 10, where pH = -log [H+]
i.
pH changes by 1 for every power of 10 change in [H+]
NOTES – Acids, Bases and pH, Chapter 14

ii.
pH decreases as [H+] increases because pH = -log [H+]
Consider the log form of the expression: Kw = [H+][OH-]
Kw = [H+][OH-]
log Kw = log [H+] + log [OH-]
- log Kw = -log [H+] - log [OH-])
Therefore: pKw = p[H+] - p[OH-])
Since Kw = 1.0 X 10-14 ,
pKw = -log (1.0 x 10-14 ) = 14.00
Thus, for all aqueous solutions at 25°C, pH, and pOH add up to 14.00
pH + pOH = 14.00
ex: Calculate pH and pOH for each of the following solutions at 25°C
a. 1.0 X 10 –3 M OH–
[H+] =
Kw
1.0 x 10 -14
=
= 1.0 x 10 –11 M
[OH ] 1.0 x 10 -3
pH = -log [H+] = -log (1.0 x 10 –11 ) = 11.00
pOH = -log [OH-] = -log (1.0 x 10 –3 ) = 3.00
b. 1.0 M H+
[OH-] =
Kw
1.0 x 10 -14
=
= 1.0 x 10 –14 M

1.0
[H ]
pH = -log [H+] = -log (1.0) = 0.00
pOH = -log [OH-] = -log (1.0 x 10 –14 ) = 14.00
VIII
Calculating the pH of Acidic Solutions
a. Mainly deals with the solution components and their chemistry, so it is important to identify and
focus on the major species
i. 1.0 M HC is actually H+ and Cl- ions
b. Major species are those that are present in large amounts
i. In the solution 1.0 M HCl, the major species are H+, Cl-, and H2O
c. Strong acids are those that dissolve (nearly) completely in solution
EX: Calculate the pH of 0.10 M HNO3

d.
1. List major species: H+, NO 3 , and H2O
2. Consider the major source of acid: HNO3
3. [H+] = 0.10 M and pH = -log(0.01) = 1.00
Strong Acid Equilibrium Problems
i. List the major species in solution
ii. Choose the species that produce H+ ions, and write balanced equations for those rxns
iii. Using the values of the equilibrium constants (K) for the rxns you have written, decide
which is the major producer of H+ ions
iv. Write the equilibrium expression for the dominant equilibrium
v. Make an ICE table to find the change in equilibrium in terms of “x”
vi. Solve for x
vii. Check to see if approximation is valid (5%) rule
viii. Calculate [H+] and pH
EX: see pg. 673 for an in-depth example of this type of problem
e.
The pH of a Mixture of Weak Acids
i. Sometimes a solution may contain two weak acids of different strengths
NOTES – Acids, Bases and pH, Chapter 14
EX: Calculate the pH of a solution that contains 1.00 M HCN (Ka = 6.2 x 10-10) and 5.00 M
HNO2 (Ka = 4.0 x 10-4). (p 676)
1. Major species: HCN, HNO2 , and H2O
2. All rxns produce H+:
a. HCN  H+ + CNKa = 6.2 x 10-10
+
b. HNO2  H + NO2
Ka = 4.0 x 10-4
c. H2O  H+ + OHKa = 1.0 x 10-14
3. Because of the Ka values, it can be determined that HNO2 is the major producer
of H+ ions
4. Write the equilibrium equation for HNO2:
a.
5.
Ka = 4.0 x 10-4 =
[ H  ][ NO2 ]
[ HNO2 ]
Make and ICE table to find x = [H+]
a.
Initial:
Change:
Equilibrium:
b.

HNO2  H+ + NO2
5.00
0
0
-x
+x
+x
5.00 – x
x
x
Solve for x {let (5.00 - x = 5.00)} using the
Ka = 4.0 X 10-4 =
[ H  ][ NO2 ]
[ HNO2 ]
6.
f.
Therefore [H+] = x = 4.5 X 10-2 M and pH = 1.35
Percent Dissociation
i. Used to find the amount of weak acid that has dissociated in reaching equilibrium in
aqueous solution
ii. Percent dissociation =
Refer to previous example and find the percent dissociation of HNO2 where [H+] = x = 4.5 x 10-2
M and pH = 1.35
EX:
The percent dissociation is:
IX.
amount dissociated (mol/L)
 100%
initial concentration (mol/L)
Bases



a.
[H  ]
 100% = .9%
[HNO2 ]
According to the Bronsted and Lowry model, and the Arrhenius concept, a base both accepts
protons and yields OH- ions. This can be shown by the dissociation of NaOH:
NaOH  Na+(aq) + OH- (aq)
The dissociation of a base can also be represented by B + H20 BH + OHThe neutralization of an acid and base will always result in water and a salt.
Acid Base
Water Salt
HCl + NaOH  H2O + NaCl
Calculating pH of Strong Bases
Since strong, assume 100% dissociation. Calculate pOH, then substract from 14 to get the
pH
EX: Calculate the pH of a 0.050 M solution of KOH.
Because [OH-] =0.050 M:
pOH = - log (0.050) = 1.30
pH = 14.00 - 1.30 = 12.70
NOTES – Acids, Bases and pH, Chapter 14
b.

Calculating pH of Weak Bases
Weak bases are only partially ionized in solution while strong bases are completely ionized, so
need to determine the hydroxide concentration via equilibrium, then calculate the pOH and then
subtract from 14 to get pH
EX: Calculate the pH for a 15.0M solution of NH3 (Kb = 1.8 x 10-5)
This reaction includes NH3, H20, NH4 and OH-; H2O is negligible
Kb= [NH4+][OH-]
[NH3]
Use an ice table to find the values to plug into the equation:
Kb= [x][x]
[15.0-x]
1.8 x 10-5 =
x2
15.0
Therefore:
[OH-] = 1.6 x 10-2
[H+] =
Kw
1.0 x10 14
=
= 6.3 x 10-13

2
[OH ] 1.6 x10
pH= -log(6.3x10-13) = 12.20
X.
Polyprotic Acids
 Polyprotic acids have more than one proton that may be removed by reaction with a base. These acids
can be either be called diprotic or triprotic.
a. Diprotic Dissociation:H2CO3  H+ + HCO3b. Tripotic Dissociation: HCO3  H+ + CO32-
XI.
The Effect of Structure on Acid-Base Properties
 Such molecules as HF, which is a weak acid, does not dissociate in water because of its strong bond.
See tables 14.7 and 14.8 for information on bond strength and structures.
XII.
Acid-Base Properties of Oxides
 Acidic Oxides:
a. When a covalent oxide dissolves in water, an acidic solution forms
SO2 + H2O  H2SO3
 Basic Oxides
b. The opposite is true for ionic oxides; a basic solution is created.
O2- + H2O  2OH-
XIII.
The Lewis Acid-Base Model
 Suggested by G. N. Lewis in the early 1920’s
i. Lewis Acid: electron-pair acceptor
ii. Lewis Base: electron-pair donor
 Created because it covers many reactions that do not involve other types of acid/base relationships
 See Table 14.10 as a reference:
Model
Arrhenius
Brønsted-Lowry
Lewis
Definition of Acid
H+ producer
H+ donor
Electron-pair acceptor
Definition of Base
OH- producer
H+ acceptor
Electron-pair donor