Text – Chapter 8
Previous knowledge – Naming Acids and Bases (Gr. 11)
Acid and Bases
Acid
Base












Sour
Electrolytes
Gritty feel
pH – 0 - 6.9
Blue litmus – red
React with bases to form a
salt and water
 Put H+ into solution
 Made by reaction of oxides
and water and binary
covalent molecules and water
Bitter
Electrolytes
Slippery feel
pH – 7.1 – 14
Red litmus – blue
React with acids to form a salt
and water
 Put OH- into solution
 Made by metallic oxides and
water
Acid and Base Names
 Acids
 Contains one or more hydrogen atoms
 General formula

HnX
HCl
H2SO4
 H – hydrogen atom
 n – number of hydrogen atoms (subscript)
 X – monoatomic or polyatomic anion
Acid and Base Names
 When the name of the anion ends in “ide” (X), the acid
is a binary acid, and the prefix is “hydro” and the
ending is “ic”
 When there is a polyatomic ion, that makes up (X), the
acid is a ternary acid. If the ion ends in “ite”, the
ending for the acid is “ous”
 When the polyatomic ion ends in “ate”, the ending for
the acid is “ic”
Acid and Base Names
 Hint – If the name of the anion is “ate”, and the acid is
“ic”, one less oxygen, the acid is “ous”, one more less
oxygen, the acid is prefix “hypo” and ending is “ous”
 If there is one more oxygen than the “ate” polyatomic
ion, the name is, prefix “per” and ending “ic”
 Some organic acids, you just have to memorize the
name. Ex. Ethanoic Acid – CH3COOH
Acid and Base Names
 HCl
Cl- - chloride
Binary – hydro – stem - ic
 H2SO4
SO4 -2 - Sulfate
Ternary – stem - ic
 H2SO3
SO3-2 - sulfite
Ternary – stem - ous
 HCN
 HClO3
 HClO4
 HClO
 H3PO4
Hydrochloric acid
Sulfuric Acid
Sulfurous Acid
Acid and Base Names
 Bases
 Named the same as ionic compounds
 Some you just have to memorize (ie. Ammonia)
Positive ion – cation (+)
NaOH
Magnesium
hydroxide
Aluminum
hydroxide
Negative ion – anion (-)
Sodium hydroxide
Acid – Base Theories
 Arrhenius Acids and Bases
 Hydrogen containing compounds that ionize (ions –
“wanderers” to produce H+ are acids
 Hydroxide containing compounds that produce OH-
ions in solution are called bases
 Not all substances that contain hydrogen atoms or
hydroxide will be acidic or basic. It depends on
electronegativity and polarity between the acidic /
basic unit and the bonded atoms.
Acid – Base Theories
 Bronsted – Lowry Acid and Bases
 Acid – hydrogen ion donor
 Base – hydrogen ion acceptor
 Substance that accepts the hydrogen is the conjugate acid
 Substance that donates the hydrogen is the conjugate
base
 Used for those exceptions that cannot be explained by
Arrhenius
 Truer theory as a “naked” hydrogen ion is very unlikely and
unstable. Hydronium ion is most likely
Acid – Base Theories
 Water can behave as both a conjugate base and acid
(can accept and donate)
 Called Amphoteric substance (behave as an acid and a
base)
Strengths of Acids and Bases
 Based upon structure
 Greater the EN difference, the greater the ionization and
dissociation, means the more “product” is formed, and
more H+ or OH- goes into solution
 Therefore, stronger acid and base (Keq greater than 1)
 First ionization is the strongest
 Second and subsequent ionizations are weaker. (p.607)
Acid – Base Theories
 The number of hydrogens will determine whether it is
monoprotic, diprotic or triprotic acids.
 Structure will determine the strength of the acid.
 Rule – the greater the EN difference, the greater the
polarity, the greater the dissociation (ionization) and
strength of the acid or base.
 Rule – For ternary acids, if the Oxygens out number the
hydrogens by more than 2, the acid will be strong
 Greater net pull, according to the first rule.
Strengths of Acids and Bases
 Keq for an acid is called Ka, and is the measure of how
much of the acid ionizes (H+ or H3O+ and X) and how
much stays together (HnX)
HX
H2XOy
H+ + XH+ + H XOy-
 Keq for a base is called Kb, and is the measure of how
much OH- and + ion is in solution and how much
stays together.
M(OH)z
M+ + Z( OH-)
Strengths of Acids and Bases
 The amount that dissociates, or ionizes, is in
equilibrium with the acid that stays as a “whole”
 If the Keq is greater than 1, it favours “product”
 For Acids and Bases, the Keq greater than 1, means it
dissociates 100%, and all of the reactant (acid or the
base) goes to ions.
Strengths of Acids and
Bases
+
-2
 For example
H2SO4
H
HSO4
2.0
0
0
Therefore,
the
 H SO Ka = 1.00 x 10
the
 H concentration
SO
+ H SO
2.0H-2.0
0 + 2.0of 0+2.0
hydronium
ion
is
 Therefore,
if the concentration
of
the
acid
0
2.0
2.0is 2.0 M, the
concentrations of ions will be:
2.0 M.
I
2
2
3
4
4
C
E
+
4
-
Strengths of Acids and Bases
 Therefore, when a strong acid or base dissociates,
100% turns into ions.
 What would be the OH- for a sodium hydroxide
solution with a molarity of 0.5 M?
Solution – sodium hydroxide is a strong base, therefore, it
dissociates 100%
NaOH
Na+ + OH-
0.5 M
O.5 M
0.5 M
Strengths of Acids and Bases
Strengths of Acids and Bases
 What about weak acids and bases?
 We need to solve for the concentrations of ions, using Ka
or Kb values because they do not dissociate 100%
 Ka or Kb values are less than 1, favouring the acid or base to
stay together and little ionize.
 They may only dissociate 50% or 10%, leaving the majority of
the acid or base as a “whole” and very little in ion form
Strengths of Acids and Bases
 Example
What is the [H+] concentration in a 1.0 M solution of
carbonic acid? (Ka = 4.3 x 10-7)
Hint – Look at the Ka value. It is less than 1. Therefore, it
Note
– Indissociate
your text,100%
the Keq
is written
using
will not
andexpression
is a weak acid.
We need
to use
Bronsted
/Lowry.
You can the
useconcentration
either, as long of
as the
you[H+].
equilibrium
to determine
remember that there are some exceptions in which the
ionization cannot be shown using Arrhenius. Also, the Keq
expression uses water, then omits it. Keep in mind, water
always has a concentration of 1 M. Concentration of
water?!!!!
Strengths of Acids and Bases
 Why are we finding the H+ or OH-?
pH
power of hydrogen or
potential of hydrogen
Hydrogen Ions and pH
 Based upon water
 Highly polar
 Made up of hydronium ion and hydroxide ion
 Self Ionization of water (Kw)
Hydrogen Ions and pH
 Each ion has a value of 1 x 10 -7 M, which, when multiplied
together forms the Keq or Kw, which is 1 x 10 -14
 Both are equal to each other in terms of their
concentrations, and therefore form a neutral substance
(pH = 7)
 If the concentration of H+ is greater than 1 x 10 -7 M, the
solution is acidic (ie. More H+ and less OH-), since all
solutions are in water.
 If the concentration of H+ is less than 1 x 10 -7 M, the
solution is basic (ie. Less H+ and more OH-), since all
solutions are in water.
Hydrogen Ions and pH
 pH = potential hydrogen
 pH scale is based upon the [H+] found in a solution.
 If the [H+] is greater than 1 x 10 -7 M, the solution will be
acidic and the pH will be less than 7.
 If the [H+] is less than 1 x 10 -7 M, the solution will be basic
and the pH will be greater than 7.
 pH is the negative log of the hydrogen ion concentration.
pH Scale
Measuring pH – Acid Base
Indicators
 Indicator – usually a weak acid that accepts hydrogen ions,
and in doing so, changes its chemical structure, which
facilitates a colour change.
 Litmus – changes from red to blue at pH of 7, or 1 x 10 -7 M
 Bromothymol Blue – is yellow below 1 x 10 -7 M, green at 1 x 10
-7 M,
and blue over 1 x 10 -7 M
 Phenopthlalein changes color at a pH of 7-9 (the hydrogen
ion concentration of 1 x 10 -7 M to 1 x 10 -9 M)
Measuring pH – Acid Base
Indicators
OH-
HIn (aq)
Acid Form
Color #1
H+
H+(aq) + In-(aq)
Base Form
Color #2
The change is caused by the removal of a hydronium
ion to form the “Base form” and the addition of a
hydronium ion for the acid form.
Measuring pH – Acid Base
Indicators
 Indicator papers are impregnated with the indicator
solution and when exposed to the hydrogen ion,
change color, depending on the concentration
 Universal Indicators show all pH levels. What do you
think they are made up of?
Neutralization Reactions and
Titration
 If we have high levels of acid in our stomach, we take
an antacid (base) to control it.
 In our small intestine, the acidic chyme, is neutralized
by the bile (basic) to make sure we do not ulcerate our
intestine
 Why does it neutralize?
 We make a salt and water, that does not necessarily work
out to a pH of 7 (more later)
 We could bring the solution to a pH of 7.
Neutralization Reactions and
Titration
[H+] = [OH-]
1 x 10-7 M = 1 x 10-7 M
Ex. If you react a strong acid and a strong base, the ions
in solution will cancel each other out, producing a
neutral solution. The products are always, regardless
of the product pH, a salt and water.
HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
Neutralization Reactions and
Titration
 The point where the number of moles of hydronium
ion equals the number of moles of hydroxide ion, is
called the equivalence point. (not necessarily pH =7)
 Salt
 The compound formed by the cation of the base
bonding with the anion of the acid.
Neutralization Reactions and
Titration
 Titration
 The process of adding a known amount of solution of
known concentration to determine the concentration of
another solution
 When the color changes, this is called the end point,
which is the point of neutralization
 There is only salt and water at this point, if both were
strong acids and bases.
Neutralization Reactions and
Titration
 Animation of Titration
 Lab Primer
 The idea is to calculate the concentration or number of
moles for an unknown solution, using the equivalence
point to determine neutralization
Neutralization Reactions and
Titration
 Titration Curves
Salts
 Salts are the combination of the cation from the base
and the anion from the acid and are the products of
neutralization reactions
 Salts can be acidic, basic or neutral, depending on the
strength of the acid and base that formed it.
 Buffers are an equilibrium condition which consists of
the weak acid and conjugate base (salt) in solution, or
a weak base and its conjugate acid (salt) keeping pH
stable
Salts
 Generally, if a strong acid reacts with a strong base, the
resulting salt will be neutral (pH=7) (ie. The
equivalence point is 7)
 For salts formed from weak acids with a strong base, or
weak bases with a strong acid, the salt will not be
neutral
 This is called by salt hydrolysis, as the cations or
anions from a dissociated salt remove or add hydrogen
ions to water, creating either H+, or OH- in solution
Salts
 Another way to determine the acidity or basicity of a
salt is to look at the net ionic equation and
remembering that strong bases and acids dissociate
100%, while weak acids and bases do not.
 In general
 Acidic salts produce positive ions that release protons
into water
 Basic solutions produce negative ions that attract
protons from water
Salts
Strong Acid +
Strong Base
Neutral solution
Strong Acid + Weak Base
Acidic Solution
Weak Acid + Strong Base
Basic Solution
Buffers
 Buffer
 A substance in which the pH remains constant, when
small amounts of acid or base is added.
 It contains the weak acid and one of its salts (anion –
conjugate base) or the weak base and one of its salts
(cation – conjugate acid)
 Pure water is not a buffer, as when you add acid (H+) or
base (OH-), the concentrations increase, changing the
pH
Buffers
 A buffer is like a sponge
 When hydrogen ions are added, they are absorbed by
the negative ion, forming the “whole” weak acid, that
does not dissociate 100%, lowering the acidity, and
raising the pH to neutral
 When hydroxide is added, they react with the acid to
form the negative ion and water, lowering the basicity,
and lowering the pH to neutral
Buffers
 Animation of Buffering
 Common buffers, keep the pH at a stable level
 Ethanoic acid
 Carbonic acid
 Ammonia
maintain pH 4.76
maintain pH 6.5 (blood)
maintain pH 9