ACIDS & BASES
All acids and bases are solutions. Therefore, a solid understanding of Solution Chemistry and
Equilibrium will guarantee you success in this unit.
I. Defining Acids & Bases:
A. General Characteristics of Acids & Bases:
Properties of Acids
Properties of Bases
Ionic compounds which ______________ in H2O
to give a solution that:
Ionic compounds which dissociate in H2O to
give a solution that:
1. tastes _____________
2. conducts _________________
3. causes __________________ to change
colour
Loses the above properties when neutralized
with a base. But the ending salt solution will still
conduct ___________________.
1. Tastes _________________
2. Conducts ____________________.
3. Causes certain indicators to change
colour.
4. Feels __________________
5. Loses the above properties when
neutralized with an acid. But resulting salt
solution will still conduct electricity.
B. Definitions of Acids & Bases:
Arrhenius Definition:
Bronsted – Lowry Definition
Acid: an ionic compound that dissociates to
give __________ ions in solution.
Acid: species which _____________ a _________ [H+ (aq)].
Ex: HNO3 (aq)+ H2O (l)
Ex: HCl (aq)
Base: an ionic compound that dissociates to
give ____________ions in solution.
Exc: CH3COOH (aq) + H2O (l)
Ex: NaOH (aq)
Base: species which _______________ a proton.
Ex: NH3 (aq)+ H2O (l)
Hydronium ion (H3O+): is a hydrated H+ ion,
meaning it is an H+ ion
dissolved in water.
H3O+
H+ (aq)
Proton
Exc: CN-1(aq) + H2O (l)
HCN (aq) + OH-1(aq)
NOTE: H2O can act as both a Bronsted-Lowry acid and
a base. Therefore, water is considered
___________.
Lewis Acids & Bases
Lewis Acid: any compound that is an ____ __________ acceptor.
Lewis Base: any compound that is an electron pair
________________.
This definition lends itself to describing acids and bases that do NOT have OH – or H+ in their chemical
makeup.
Example: BF3 (g) + NH3 (g)
Example: SO3 (g) + H2O (l)
BF3NH3 (l)
H2SO4 (aq)
II. Relationships between Acids & Bases:
A. Conjugate Acid – Base pairs: molecules and ions that differ only by one proton (or H+).
Conjugate means __________________.
Example: HCN (aq) + H2O (l)
H3O+ (aq) + CN- (aq)
Conjugate Acid of CN- :
Conjugate Base of HCN:
Conjugate Acid of H2O:
Conjugate Base of H3O+:
Example: HCO3- (aq) + H2PO4-(aq)
H2CO3 (aq) + HPO42-(aq)
B. Polyprotic Acid: an acid capable of donating _____________________________________.
Example: H3PO4 (aq) + H2O (l)
H2PO4-(aq) + H2O (l)
HPO42- (aq) +
H2PO4- (aq) + H3O+ (aq)
HPO42- (aq)
+
H2O (l) → PO43- (aq) + H3O+ (aq)
H3O + (aq)
Amphiprotic Species: a species can be amphiprotic only if it has the ability to BOTH ______________ a
proton and __________________ a proton.
Example:
III.
Preparation & Properties of some Common Acids and Bases:
A. Acids:
Categories of Acids:
Binary Acids: Those that have 2 elements. Can be made by directly combining element with
______ gas.
Example: H2 (g) + Br2 (g)
2 HBr (g)
Covalent Oxide (Acidic Anhydride): an oxygen – containing _______________compound that
reacts with water to produce an ____________________ solution.
Example: P4O10 (s) + H2O
4 H3PO4 (aq)
(l)
Example: SO3(g) + H2O(l)
H2SO4 (aq)
(1) Sulfuric Acid: produced by the _____________________ Process.
(a) S (l) + O2 (g)
(b) 2 SO2 (g) + O2
(g)
SO2 (g)
catalyst
2 SO3 (g)
(c)
(d)
SO3 (g) + H2SO4 (l)
H2S2O7 (l) + H2O
H2S2O7 (l)
2 H2SO4 (l)
V2O5
Properties: ______________ acid – Good neutralizer and great electrolyte.
Good _____________________ agent.
Great affinity with ___________ and can be used as a __________________ agent.
(2) Hydrochloric acid: soluble in all _______________ ions.
2 NaCl (s) + H2SO4
H2O (l) + HCl (g)
(l)
Na2SO4 (aq) + 2 HCl
(g)
H3O+ (aq) + Cl- (aq)
(3) Nitric Acid: made by the _______________ Process:
4 NH3 (g) + 5 O2 (g)
2 NO (g)
+ 2 H2O (l)
3 NO2 (g) + H2O (l)
4 NO (g) + 6 H2O (g)
2 NO2 (g) + 2 H2 (g)
2 HNO3 (aq) + NO (g)
Properties: Has a suffocating odour.
Leaves a ______________ stain on skin due to NO2 being exposed to ____________.
B. Bases:
Sodium Hydroxide (NaOH): is in most household cleaning products. Reacts with fats and
proteins to make soap, a substance that is water soluble.
Reaction proceeds fastest in warm temperature.
NaOH has ability to absorb moisture from the air – Thus it is a good
___________________________ agent.

NaOH is made by electrolysis of NaCl dissolved in water. H2 collects at the cathode and Cl2
collects at the anode.
NaCl (aq) + 2 H2O (l)
2 NaOH (aq) + H2 (g) + Cl2 (g)
Basic Anhydrides (Metallic Oxides): an oxygen containing _______________compound that
reacts with water to give a basic solution.

Example: Na2O(s) + H2O(l)
NaOH (aq)
BASIC ANHYDRIDE + ACIDIC ANHYDRIDE
Example: N2O5(g) + Na2O(s)
SALT
2 NaNO3(s)
IV. Strengths of Acids & Bases:
A. Strong & Weak Acids:

Strength of an acid depends on the ______________ of _______ions produced per mole of acid.
1) Strong Acid: ionizes _______________ in water.
Reaction goes to completion. NO ____________ exists.
Ex:
Initial:
Final:
HNO3 (aq)
H+ (aq) + NO3- (aq)
1.2 M
0M
0M
1.2M
0M
1.2M
Reasons why Strong Acids do NOT exhibit equilibrium:
Common Strong Acids:
HNO3
HBr
HCl
HI
HClO4
H2SO4
The Leveling Effect:

Table shows all strong acids dissociating 100% to form H3O+(aq). Water is said to have “leveled”
all the strong acids to the same strength, they are all essentially solns of H3O+ - This means that H3O+ is
the strongest acid that can exist. ALL the strong acids have identical strengths. Their chronological
order is purely arbitrary.

Same can be said for the two strong bases at bottom of table.
O-2 and NH2- cannot exist on their own in water, they immediately react with H2O to form NH3
and OH- in soln. - Again, H2O has “leveled” the strong bases to the same strength, all solns of OH.
2) Weak Acids: do not _____________________ completely in water.
Equilibrium exists between the ___________ and the undissociated molecule.
Example: CH3COOH (aq) + H2O
CH3COO-
(l)
(aq)
+ H3O+ (aq)
Reasons why Weak Acids go to Equilibrium:
B. Strong & Weak Bases:
1) Strong Base: ionizes completely when dissolved in water. Goes to ___________________.
Presence of the ________ion gives solution its basic characteristic.
Common Strong Bases: Bases formed from _____________metals and heavier ________ metals. Not
all Group 2 bases are completely soluble in water.
NaOH
KOH
Example: KOH (aq)_
LiOH
Sr(OH)2
NH2-1
O-2
K+ (aq) + OH- (aq)
2) Weak Bases: do not completely ionize.
Equilm exists. A large amount of the compound has not ____________________.
Ex: CN- (aq) + H2O (l)
HCN (aq) + OH- (aq)
C. Comparing Strengths of Acids and Bases:
1) Left side of “Relative Strength of Acid and Bases” table: species acting as _______________
2) Right side of table:
species acting as _______________
When comparing strengths of species:
a) The stronger an acid is, the ______________ its conjugate base is & vice versa.
Ex: HIO3 is strong acid, therefore its conjugate base, IO3- is weak. 
HIO3 is a strong acid, so HIO3 must be a weak base!

Ex: HCO3- and HPO42-: both species are amphiprotic.
This means they appear on both LEFT and RIGHT sides of the table.
i) Relative ACID strength: HCO3- appears higher than HPO42- on _________ side – HCO3stronger acid.
ii) Relative BASE strength: HPO42- appears lower on _____________ side HPO42- stronger base.
D. Weak Acids/Bases and Le Chatelier’s Principle:
All acid-base rxns. can be seen as _________________ for ____________. The stronger base takes the
H+ from the reacting partner. The species that does not behave as the base is automatically
considered as the ____________:
Ex: Write the balanced equilibrium that results when HSO3-1 and HCO3-1 react. Determine
which direction is favoured at equilibrium.
Le Chatelier’s Principle.:A strong acid will more likely want to donate H+ than a weaker acid.
Therefore, at equilm, rxn. will ________________ the direction of the
_________________ of the ___________________ acid.
V. Quantitative Analysis of Acids & Bases:
A. Ionization of Water and K w :


H2O is considered a ____________________, meaning it won’t conduct ________________.
However, a VERY! small amount of H2O molecules do dissociate to form ions.
This means that aside from the acid or base dissociating, our solvent is also slightly dissociating
into ions:
H2O(l)
↔
H+ (aq) +
OH-1(aq)
-7
1x 10 M
1x 10-7M
(@ 250C)
NOTE: The above reaction is at equilibrium. This means that it too has an Equilibrium Constant.
Keq = [H+] [OH-]
Kw = [H+] [OH-]
Temperature Dependence of H2O:
H2O (l) + Energy
H+(aq)
+ OH-
(aq)
Increase in Temperature: Shifts equilibrium ____________, and increases concentrations of
______________.  Kw ______________.
Decrease in Temperature: shifts reaction left and Kw ____________________.
The addition of either an acid or base to H2O will disturb this 1 H+:1 OH-1 ratio, therefore, resulting in a
shift in equilibrium:
Example: Add 1.0M HCl to pure water:
Ex: What is the [H+] and [OH-] for a soln of 0.0010M HCl?
[H+] = 0.0010M
B. Measuring [H3O+] or [OH-1]:


Conc. of acids/base can range from extremely high to extremely low it is easier to express
these
[ ] as ______________________.
We call these logged values of [acid] or [base] _______ or ____________.
Logarithms: turns an exponential value of base 10 into an integer.
Ex: 1.00 x 10-5M
1.0 x 102 M
-log (1.00x 10-5) = 5.000
log (1.0 x 102) = 2.00
Kw = [H+] [OH-]
-log Kw = -log [H+] + -log [OH-]
pKw = pH + pOH
pH = -log [H+]
pOH = -log [OH-]
Sig. Figs for Logs:
Only the digits _______________ the decimal place of a log value is significant.
Ex: - log (5.28 x 10-5) = -log (5.28) + -log (10-5)
-0.723 + – (-5) = 4.277 (Has 3 sig. figs)
The pH Scale:
Please note that the scale is based on pH, meaning the reference point is that for _____________:
|
pH 7.00
pH 1.00
pH 14.00
[H+]:
Most acidic
Neutral
Most basic (@
0
25 C)
NOTE: A change of 1 pH value is equivalent to __________ the concentration change of an
acid/base.
Ex: What is the pH of a 0.010M nitric acid soln?
[H+] =
pH =
Ex: A soln of phosphoric acid has a pH of 3.00.
(a) What is [H+] of the soln?
pH = - log [H+]
(b) What is the pOH of the soln and its [OH-]?
pH + pOH = 14.00
pOH = - log [OH-]
C. Weak Acid Equilibrium & Ka:


For weak acids, ionization is ________ complete, and it is NOT possible to use the initial conc. of
the acid in the solution.
For weak acids, the incomplete ionization results in _______________________.
Example: HF(aq) + H2O (l)
H3O+
(aq)
+ F-(aq)
The equilibrium expression is as follows:
Keq = [H3O+] [F-]
[HF]
Ka = [H3O+] [F-]
[HF]
, where Ka - Acid Ionization Constant for Weak
Acids.
NOTE: The higher the Ka value, the _______________ the ionization for the particular acid,  the stronger
the acid.
As __________________ changes, the ability to ionize changes.
Example: What pH results when 0.25 mol of acetic acid is dissolved in enough water to make a 2.0L
solution?
1. Write the balanced dissociation equation.
2. Look up the ionization constant , Ka, for acetic acid:
3.
ICE Table:
[I]
[C]
[E]
[solution] = 0.25M
Because acetic acid is a weak acid, only a small amount of the solution dissociates.
Therefore, the change in [CH3COOH]is negligible. The concentration of acetic acid at
equilibrium can be considered the same as its initial equilibrium.
5% Rule: The above assumption can be made for any weak acid/weak base as long as the Keq value
for that particular acid/base is 1000X smaller than its initial concentration.
Ka = [H3O+] [CH3COO-]
= 1.8 x 10-5
0.25M
D. Ka & Kb for Conjugate Pairs:
Ka x Kb = Kw
Example: Determine Kb for the weak base F-.
Notice that Kb for F- is unavailable. But its conjugate acid HF, is available, with a Ka value of 6.7 x 10-4.
E. Weak Base Equilibrium & Kb:


Weak Base: does NOT ionize completely in aqueous solutions. Different bases have varying
degrees of ionizability.
Every weak base also has a corresponding Base Ionization Constant, Kb.
Example: NH3 (aq) + H2O (aq)
NH4+
(aq)
+ OH- (aq)
Kb = [NH4+] [OH-]
[NH3]
Example: Determine the pH for a solution of 0.100M NH3 solution.
1. Balanced equation:
2. Construct an ICE table. (Use 5% Rule if possible)
[I]
[C]
[E]
VI. Typical Reactions of Acids & Bases:
A. Formation of Salts (Neutralization Reactions):

Salts: the product of ___________________ reactions.
B. Hydrolysis:
Hydrolysis: when the ions of a soluble salt reacts further with __________.
The resultant equilibrium may be ______________, or __________________.
1. KCl:
KCl (aq)
K+ (aq) + Cl- (aq)
CATION RULE: Any cation from Family 1 and 2, are ions of strong bases, and will NOT hydrolyze.
ANION RULE: Conjugate bases of strong acids do NOT hydrolyze. And this is the case of Cl-. It is the
conjugate of the strong acid HCl.
2. NH4NO3 : NH4NO3 (aq)
NH4+(aq) + NO3-(aq)
NH4+ : ________________________________.
NO3-: ________________________________.
3. NaHCO3: NaHCO3 (aq)
Na+
(aq)
+ HCO3- (aq)
Na+: ___________________________________.
HCO3-1: _________________________________.
Example: Predict whether the solution of the salt NH4NO2 will be acidic, basic, or neutral.
NOTE: Some salts contain BOTH a cation and an anion that will hydrolyze. In such cases, the resulting
solution may be acidic, basic, or neutral, depending on the relative extent of the hydrolysis.

NH4+ : will hydrolyze -
NO2-1: will hydrolyze To determine the overall acidity or alkalinity of the solution, compare Ka to Kb values.

Example: Determine the pH of a 0.50M NH4HC2O4 solution.
VII. Buffer Solutions:

Buffer solutions: equilibrium system that maintains a relatively constant _______ when excess
amounts of acid or base is added.
A. Buffer Preparation & Action:

Buffers: can be prepared by combining appreciable amounts of a weak acid/base with its
_________.
Example: CH3COOH + H2O
High conc.
H3O+ + CH3COO- (aq)
Low Conc. Low conc.
(Before addn of salt)
Salt to be added:
pH of Buffer Solution before addition of excess acid or base:
(i) Addition of excess acid: 0.010 mol of HCl added to 1.0L of buffer solution.
(ii) Addition of excess Base: 0.010 mol of NaOH added to 1.0L of buffer solution.
Example: NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
VIII. Indicators:

Indicators are weak acids or bases. Usually however, they are weak _____________.
HInd (aq) + H2O(aq)
HInd+ (aq) + H2O (l)
H3O+ (aq) + Ind- (aq) (Weak Acid Indicator)
H3O+ (aq) + Ind (aq) (Weak Base Indicator)
Depending on the form of the indicator at the moment, it can take on different _________________.
Ionization Constant for Indicators: Ka = [H3O+] [Ind-]
[HInd]
…..at the transition point, the colour just begins to change because [HInd] = [Ind-].
 Ka = [H3O+]
Example: Determine the Ka for the chemical indicator phenolphthalein.
Phenolphthalein pH range: 8.2 – 10.6
Assume the transition point lies ___________________ these two pH values.
IX. Titration:
HCl (aq) + NaOH (aq)


NaCl (aq) + H2O (aq)
Equivalence Point: the point at which the __________ ratio of reactants in the beaker is
equivalent to the ratio described by the _________________ equation.
In a rxn. of a strong acid & strong base, there is a dramatic change in ________ at the equivalence
point – This fact aides in the finding of the equivalence point - Thus titration is born.

Standard Solutions: a solution of _______________ concentration.
Example: It is found that 42.5 mL of 1.02M NaOH have been added to 50.0 mL of vinegar when the
phenolphthalein in the solution just turns pink. What is the [vinegar]?
(i) Primary Standards: when a highly _____________ substance of known composition is used to
react with a substance to be _____________________.
Example: Na2CO3 (s) + 2 HCl (aq)
2 NaCl (aq) + H2O (l) + CO2 (g)
A solution of HCl is standardized using pure Na2CO3 (s) as a primary standard. What is the [HCl] solution
if 60.00 mL of this solution are needed to titrate a 1.000g sample of sodium carbonate?
(ii) Equivalence Point vs. Transition Point:



Equivalence Point (Stoichiometric Point): refers to the ____________ point of the titration.
Transition Point: refers to the point where the indicator just begins to change ______________.
Ideally, the end and equivalence point should be reached at the same time.
(a) Strong Acid/Strong Base Titration Curves:
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l) (Indicator = Phenolphalein)
pH
Volume of NaOH titrated


For a strong acid/strong base titration, the equivalence point is always __________. This means that
an ______________ whose transition point is 7 should _______________be used.
The ___________ from the neutralization reaction does not _______________ because it is the product
of strong acid and base.
(b) Weak Acid/Strong Base Titration Curves:
CH3COOH(aq) + NaOH(aq)_
NaCH3COO(aq) + H2O(l)
(Indicator: Phenolphathlein)
pH
Volume of NaOH titrated
NOTE: A weak acid is capable of demonstrating ____________________ action when a __________ is
added. This buffer action is what causes for the ___________________ rise in pH in this region of
the curve.
At the Equivalence Point: The solution is ___________ neutral!
CH3COOH (aq)_ + NaOH (aq)
NaCH3COO(aq) + H2O (l)_
The salt sodium acetate has been formed at the equivalence point and ____________________ to give
Na+ and CH3COO - . The CH3COO- ion will __________________ in the aqueous environment to give the
following ions:
CH3COO- (aq) + H2O

(aq)
CH3COOH (aq) + OH- (aq)
It is because of the OH- ions produced that causes the solution to be slightly ____________.
Example: Determine the pH at the following points when 50.0mL of 0.10M acetic acid is titrated with
0.10M NaOH.
(a) The pH of the solution of acetic acid when no NaOH is yet added.
(b) When 10.0mL of 0.10M NaOH has been added.
(c) At the mid-point of the titration.
(d) At the equivalence point when 50.0mL of NaOH has been added.
NOTE: At the equivalence point is where _______ of the original CH3COOH has been consumed.
The equilibrium then is as follows:
(c) Strong Acid/Weak Base Titration:
NH3 (aq) + H+(aq) ↔ NH4+
Indicator: methyl red
(aq)
+ H2O (l)
pH
Volume of HCl added (mL)
Example: Determine the pH at the following points of the titration between 100.0mL of 0.050M NH3
with 0.10M HCl.
(a) Before any HCl is added.
(b) At mid-point of the titration:
(c) At equivalence point where 50.0mL of HCl has been added.
(d) When 60.0mL of HCl has been added.
X. Acid Rain and its Impact:

Normal rainwater has a pH close to ______ due to the natural reaction of H2O with CO2 gas
from the environment, forming _____________ carbonic acid:

Acid Rain: has a pH below _______ due mainly to reaction of the rainwater with gaseous
__________ of industrial processes:
Industrial Process
Reaction with Rain Water
Resulting Acidic Soln:
1) Combustion of ___________ fuels:
C(s) + O2 (g) → CO2 (g)
2 C8H18 (l) + 25 O2 (g) → 16 CO2 +18 H2O
C2H5OH + 3 O2 → 2 CO2 + 3 H2O
2) Roasting of Sulfur Ores and Coal
____________________
Fired Power Stations:
2 ZnS (s) + 3 O2 (g) → 2 ZnO (s) + 2 SO2 (g)
_____________________
NOTE: SO2 is also produced industrially
by the Contact process used
to make _______________ acid:
S (s) + O2(g) → SO2 (g)
3) Incomplete combustion Inside
Engines at High Temperatures:
N2(g) + O2(g) → 2 NO (g)
2 NO (g) + O2(g) → 2 NO2 (g)
Negative Outcomes of Acid Rain:
1) Buildings or ______________ made of solid __________ will gradually dissolve:
2) Leach ________ from soils into ground water and lakes and rivers, poisoning aquatic life.
3) Can disrupt _______________________ in plants. Low acid concentrations will ___________ down the
process of photosynthesis and concentrated levels of acid will __________ the plant.