Intro
to
Acids
&
Bases
UNIT 4
Snowball Fight!
Properties of Acids & Bases
 Acids and bases are special substances with very
distinct properties. It is good to think of acids and
bases as opposites.
Key characteristics of acids
Key characteristics of bases
-
-
-
Sour taste (eg. Lemons, grapefruit,
vinegar, sour milk)
React with active metals such as zinc and
magnesium to produce hydrogen gas
Form electrolytic solutions (conduct
electricity) because they produce ions
Cause certain dyes to change color
(litmus paper turns red)
Neutralized by bases (neutralized means
that the substance no longer has acidic or
basic properties) to form salts
-
Bitter taste
Generally no noticeable reaction with
active metals
Form electrolytic solutions (conduct
electricity) because they produce ions
Cause certain dyes to change color
(litmus paper turns blue)
Slippery feel (eg. soapy feel)
Neutralized by acids to form salts
Important Concepts for this Unit
Dissociation
Ionic
Compounds
Ionization
Covalent
compounds
Dissociation
 Is the breakdown of an ionic compound
 Is the breaking up of a compound into simpler
constituents that are usually able to recombine
under the right conditions
 It is usually reversible

KOH(s) ⇌ K+(aq) + OH-(aq)

Note that bases undergo dissociation.
Review: Dissociation Equations
 There are two important things to notice
about writing dissociation equations:
 Generally
DO NOT include H2O as a reactant. We
know something has been dissolved in water
when we see the (aq) notation. We will make
some exceptions later to this rule
 Ion charges MUST BE included!
Ionization
 Process by which neutral atoms or molecules are
converted to electrically charged ions

Process of dissolving molecular compounds (covalently
bonded) in water to produce ions.
 Ionization is usually irreversible
 Most molecular compounds do not undergo
ionization. However, acids ALWAYS do.
In fact, all acids produce hydrogen ions in a solution.
 H2SO4(g)  2H+(aq) + SO42-(aq)

 So what is actually happening?
 Evidence suggests that the hydrogen ion in acids
actually bonds to a water molecule forming a
hydronium ion, H3O+.
Ex: H2SO4(g)+ 2H2O(l)  2 H3O+(aq) + SO42-(aq)
 You should be comfortable using either method of
representation: one will mean the same as the other.

Definitions
 There are two main theories or definitions of
acids & bases that we will discuss in this
class
 Arrhenius’ Theory
 Bronsted-Lowry Theory
Arrhenius’ Theory of Acids & Bases
 In the 1880’s, Svante Arrhenius determined that in
aqueous solution...

An acid is a source of hydrogen ions (H+) or hydronium ions (H3O+)

Eg. HCl(aq) + H2O(l) ⇌ H3O+ (aq) + Cl- (aq)
H
Cl

H
O
H
–
H+
O
H H
Cl
A base is a source of hydroxide ions (OH-)

Eg. NH3 (aq) +H2O(l) ⇌ NH4+ (aq) + OH-(aq)
H+
H
H
N
H
H
O
H
H
N
H
H
–
O
H
Bronsted and Lowry Theory
 2 chemists working independently,
Johannes Bronsted and Thomas Lowry,
came up with what is now known as the
“Bronsted-Lowry Theory of Acids and
Bases.”
–
–
Acids are proton (H+) donors.
Bases are proton (H+) acceptors
Proton Donation
 How are acids “donors?”
 HCl  H+ + Cl This shows that HCl produces an H+, but to donate
implies that something will receive the H+. So, we can
see the donation with the ionization equation:
HCl + H2O  H3
acid
base
+
O
+
Cl
Conjugate Acid-Base Pairs
 conjugate base: the species that remains after an
acid has given up a proton
 conjugate acid: the species that is formed when a
base gains a proton
HCl + H2O  H3
acid
base
conjugate acid
+
O
+
Cl
conjugate base
base
acid
NH3 + H2O
+
NH4
conjugate acid
+
OH
conjugate base
 Notice that the ammonia has become an ammonium ion by accepting a
H+ from the water.
Amphiprotic/ Amphoteric
 In the two examples water first acts as a base, then as
an acid.
 Any species which can both accept and donate
protons is called amphiprotic. (also known as
amphoteric)

They can act as an acid in one reaction but a base in another.
 Conjugate acid-base pairs differ from each other by
the presence or absence of a single hydrogen ion
(or proton).
 Every acid has a conjugate base, and every base has
a conjugate acid.
 We can now express these equations with a double
arrow, since it represents acid-base equilibrium
Examples
 Example 1: Write the conjugate bases
for the following acids:
 A) HF
 B) H2SO4
Answers
 A) F-
 B) HSO4-
Examples
 Example 2: Write the conjugate acids
for the following bases:
 A) PO43 B) SO42-
Answers
 A) HPO4-2
 B) HSO4-
Example
 Example 3: In the following two reactions which
substance is amphoteric? When is it an acid? A base?
 A) HSO4- + H3O+
H2SO4 + H2O
 B) HSO4- + OHSO42- + H2O
 Answer:
 Forward: HSO4-, A = base, B = acid
 Reverse: H20, A = Base, B = acid