1
Nature of Acids and Bases Notes
(D1) identify acids and bases through experimentation

list general properties of acids and bases
o Acids
1. _________________________________________
2. _________________________________________
3. _________________________________________
4. _________________________________________
5. _________________________________________
6. _________________________________________
7. _________________________________________
o Bases
1. _________________________________________
2. _________________________________________
3. _________________________________________
4. _________________________________________
5. _________________________________________
6. _________________________________________
7. _________________________________________


write names and formulae of some common household acids and bases
outline some of the uses and commercial names of common household acids and bases
o common household acids include:
most fruits, carbonated beverages,
teas, and batteries
o common household bases include:
baking soda, ammonia, soap, and
antacids
o more examples on pages 112 – 114 in your textbook
(you should be familiar with these)
2

write balanced equations representing the neutralization of acids and bases in solution
o Neutralization reactions always produce ______________ and a _____________.
o Examples of neutralization reactions:
1.
HCl(aq) +
NaOH(aq) →
2.
HCl(aq) +
Ca(OH)2(aq) →
3.
H4P2O7(aq) +
NaOH(aq) →
(D2) identify various models for representing acids and bases

define Arrhenius acids and bases
o The Arrhenius definition of acids and bases is one of the oldest classifications. It
was developed by Svante Arrhenius during his work with ions in aqueous solution
in 1884. Arrhenius receiving the Nobel Prize in Chemistry in 1903.
o An Arrhenius acid is a substance that increases the concentration of H+ ions
when added to water. Arrhenius acids always begin with hydrogen.
HCl is an example of an Arrhenius acid.
o An Arrhenius base is a substance that increases the concentration of OH- ions
when added to water. Arrhenius bases always end with hydroxide.
NaOH is an example of an Arrhenius base.
Note: Arrhenius acids are frequently referred to as proton, hydrogen ion, or
hydronium ion (H3O+) donors, depending on whether we are trying to
emphasize the species liberated by the acid (proton or hydrogen ion) or the
species present in solution (hydronium ion). To represent the formation of
the hydronium ion, we must include H2O in the chemical equation.
3

define Brönsted-Lowry acids and bases
o When Svante Arrhenius was developing his theory of acids and bases, equilibrium
reactions were not considered. However, equilibrium reactions need to be taken
into account, meaning that another definition of acids and bases is required.
o The Brönsted-Lowry Theory is a more general theory than that of Arrhenius.
o A Brönsted-Lowry acid is a substance that donates a proton.
o A Brönsted-Lowry base is a substance that accepts a proton.
(D3) analyse balanced equations representing the reaction of acids or bases with water

identify Brönsted-Lowry acids and bases in an equation
Identify the reactant acting as an acid and the reactant acting as a base.
NH3 + H2O → NH4+ + OH-
CH3COOH + H2O → CH3COO- + H3O+
(D6) identify chemical species that are amphiprotic


define amphiprotic, and
describe situations in which H2O would act as an acid or base
o In the above examples, H2O acts as a Brönsted-Lowry acid in one situation and a
Brönsted-Lowry base in another situation. H2O is an example of an amphiprotic
substance. Amphiprotic means that the substance can either lose or gain a proton.
o Other amphiprotic substances include H2PO4-, HS-, HCO3-.
o All the amphiprotic species above are polyprotic acids that have lost at least one
proton (polyprotic acids are acids that can lose more than one proton).
o In general, a substance will be amphiprotic if it:
 has a negative charge
 has an easily removable hydrogen
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(D3) analyse balanced equations representing the reaction of acids or bases with water

define conjugate acid-base pair
o A conjugated acid – base pair is a pair of chemical species that differ by only
one proton.
o The conjugated acid is the species with the extra proton.
o The conjugated base is the species without the extra proton.


identify the conjugate of a given acid or base, and
show that in any Brönsted-Lowry acid-base equation there are two conjugate pairs
present
Identify the conjugated pair, the conjugated acid, and the conjugated base.
NH3
CH3COOH
+
H2O
→
NH4+
+
H2O
→
CH3COO-
+
OH-
+
H3O+
What is the conjugate acid of NH3?
What is the conjugate base of H3O?
Brönsted-Lowry acid-base equilibrium general form:

identify an H3O+ ion as a protonated H2O molecule that can be represented in shortened
form as H+
o Refer back up to outcome (D2).
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(D4) classify solution of acid or base as weak or strong, with reference to electrical conductivity

relate electrical conductivity in a solution to the total concentration of ions in the solution
o Recall that electrical conductivity in a solution depends on whether or not ions are
present in the solution. An ionic substance will conduct electricity where as a
molecular substance will not.
o Now, let’s take this one step further. The greater the concentration of ions present
in a solution, the greater the electrical conductivity.
o A substance that is more soluble, or completely ionized in solution, will be a good
electrical conductor.
o A substance that is less soluble, or not completely ionized in solution, will be a
weak electrical conductor.


define and give examples of strong acids and bases, and weak acids and bases, and
write equations to show what happens when strong and weak acids and bases are
dissolved in water
o A strong acid or base is 100% ionized in solution.
o A weak acid or base is less than 100% ionized in solution.
Implications: Do equilibrium reactions involved strong or weak acids and bases? Why?
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Are the terms concentrated and dilute synonymous with strong and weak? Why?
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(D5) analyse the equilibria that exist in weak acid or weak base systems



compare the relative strengths of acids or bases by using a table of relative acid strengths
predict whether products or reactants are favoured in an acid-base equilibrium by
comparing the strength of the two acids (or two bases)
compare the relative concentrations of H3O+ (or OH-) between two acids (or between
two bases) using their relative positions on an acid strength table
o In your Data Booklets:
 “Relative Strength of Brönsted-Lowry Acids and Bases”
o Strong Acids:
 Notice the top 6 strongest acids on the left side of the table. The arrow for
these acids is a one – way arrow, meaning that they dissociate 100%
producing H+ and an anion.
 Notice that the 7th acid is H3O+ and has a Ka of 1.0 (more on this later).
H3O+ is the result of adding any strong acid to water.
o Strong Bases:
 Notice the bottom 2 strongest bases on the right side of the table. The
arrow for these bases is also a one – way arrow, meaning that they
dissociate 100% producing OH- and a cation.
 Notice the 3rd base from the bottom is OH- and has a ka of 1.0 x 10-14.
OH- is the result of adding any strong base to water.
o Weak Acids and Bases:
 The acids are found on the left side of the table between the hydronium
ion and water. The bases are found on the right.
 The weak acids are always separated by an equilibrium arrow from their
conjugated base on the right.
Note: Conjugates of strong acids and bases never act as bases or acids.
o Levelling Effect:
 Strong acids dissociate completely to form H3O+
 Strong bases dissociate completely to form OH Therefore, the strongest acid that can exist is really H3O+, and the
strongest base possible is OH-.
 The top six acids are levelled to produce H3O+ and the bottom two bases
are levelled to produce OH-.