Acids, Bases, and Salts
Chapters 14 and 15
Some Properties of Acids
1. The word acid comes from the Latin
word acere, which means "sour." All acids
taste sour.
2. In 1663, Robert Boyle wrote that acids
would make a blue vegetable dye called
"litmus" turn red.
3. Acids react with bases (they destroy the
chemical properties of bases).
4. Acids conduct an electric current.
5. Upon chemically reacting with an active
metal, acids will evolve hydrogen gas (H2).
Some Properties of Bases
1. The word "base" has a more complex history
and its name is not related to taste. All bases
taste bitter.
2. Bases are substances which will restore the
original blue color of litmus after having been
reddened by an
Thisacid.
is because they dissolve the fatty
3. Bases destroyacids
theand
chemical
oils fromproperties
your skin andofthis
acids (will react
cutswith
downacids)
on the friction between your
fingers
you rub them
together.
4. Bases will conduct
anaselectric
current.
5. Bases feel “slippery” (soap, bleach) on your
skin.
Some Properties of Salts
1. A salt is the combination of an anion (- ion)
and a cation (+ ion).
2. Salts are products of the reaction between
acids and bases.
3. Solid salts are usually crystalline.
4. If a salt dissolves in water solution, it usually
dissociates into the anions and cations that
make up the salt (amount depends on Ksp)
The Acid Base Theory
The three main theories regarding acids and
bases are… (from least to most inclusive):
1. Arrhenius
3. Lewis
2. Brønsted-Lowry
Arrhenius Theory – late 1890s
DEFINITIONS:
Acid - any substance which donates hydrogen
ions (H+) to water (produces hydronium ions,
H3O+):
HA → H+ + A¯
Base - any substance which produces hydroxide
ions (OH¯) in water.
XOH → X+ + OH¯
When acids and bases react, they neutralize
each other, forming water and a salt:
HA + XOH → H2O + XA
Problems with Arrhenius Theory
The theory did not explain why ammonia (NH3)
was a base (… OH- was the only base so far)
 The theory only considers water as a solvent.
We know that an acid added to benzene will
not dissociate. Solvents are crucial to acid
definition.
 The end result of mixing certain acids and
bases can be a slightly acidic or basic solution.
Arrhenius had no explanation for this
phenomenon (degrees of acidity).

Brønsted – Lowry Theory –
Early 1920s
Two chemists, independent of one another,
proposed a new definition of an acid and a
base:
 An acid is a substance from which a proton
can be removed (or donates protons).
 A base is a substance that can remove a
proton from an acid (or a proton acceptor).
*This definition does not require acids and
bases to be in aqueous solutions.

Reactions Based on Bronsted - Lowry
Which are the acids and bases?:
HCl + H2O → H3O+ + Cl¯
 HCl - this is an acid, because it has a proton
available to be transferred (it can give a proton).
 H2O - this is a base, since it gets the proton that
the acid lost (it has the capacity to accept a
proton).

Conjugate acid-base pairs
Example:
HCl + H2O → H3O+ + Cl¯
Notice that each pair (HCl and Cl¯ as
well as H2O and H3O+ differ by one
proton (H+). These pairs are called
conjugate pairs.
Example: HNO3 + H2O → H3O+ + NO3¯
 The acids are HNO3 and H3O+ and the
bases are H2O and NO3¯. What are the
pairs?

Bases and Conjugate Acid
Base
Name
C2H3O2- Acetate ion
Conjugate
acid
Name
CH3COOH
Acetic acid
NH3
Ammonia
NH4+
Ammonium
H2PO4-
Dihydrogen
phosphate ion
H3PO4
HSO4-
Hydrogen sulfate
ion
H2SO4
Phosphoric
acid
Sulfuric acid
OHNO3-
Hydroxide ion
Nitrate ion
H20
HNO3
water
Nitric acid
H2O
water
H30+
Hydronium ion
Lewis Theory –Early 1920s
Remember drawing Lewis Dot Structures for
ionic and covalent compounds?
 Lewis Theory focuses on the nature of
electrons rather than proton transfer.
DEFINITIONS:
 An acid is an electron pair acceptor and a
base as an electron pair donor.
 Lewis Theory is much more general and will
include also reactions that do not involve
hydrogen or hydrogen ions… this will include
formation of coordinate complex ions.

Lewis acid-base reaction:
BF3 accepts an electron pair from ammonia to
complete octet:
A Lewis acid must have an empty orbital to accept an
electron pair –all cations are L.A.s, and ions/atoms
with incomplete octets.
A Lewis base must have a pair of unshared electrons
that can be donated. Typical Lewis bases are OH-,
H2O, NH3, Cl-, CN-… due to lone pair electrons.
Lewis AB reactions and Formation of
Coordinate Complexes
The metal ion is a Lewis acid and the ligands
coordinated to the ion are Lewis bases.
Autoionization of Water
What are Strong Acids and Bases?

Strong acids are
those that ionized
completely in
water... the H+ ions
are not held strongly
to the conjugate
base.
•The dissociation of a strong base also looks like
the diagram at the right in that it dissociates
into positive and negative ions.
7 Strong Acids
HNO3 - nitric acid
HCl - hydrochloric acid
HBr - hydrobromic acid
HI - hydroiodic acid
H2SO4- sulfuric acid
HClO4 - perchloric acid
HClO3 - chloric acid (wanna be)
•Strong acids are assumed to ionize completely (100%)
in water. They exist as H3O+ ions in water. This is known
as “the leveling effect”. Water has a greater affinity for H+
than the conjugate bases do.
ANIMATION LINKS

Acid ionization equilibrium demo
What are Weak Acids and Bases?
Some acids and
bases ionize only
slightly in water and
are considered
weak.
 A weak acid has a
high affinity for its
H+.
 The most important
weak base is
ammonia.

Balance of ions in solutions
Acidic
Solution
Neutral
Solution
Strong Bases
LiOH - lithium hydroxide
NaOH - sodium hydroxide
KOH - potassium hydroxide
RbOH - rubidium hydroxide
CsOH - cesium hydroxide
Ba(OH)2 - barium hydroxide
Sr(OH)2 - strontium hydroxide
Ca(OH)2 - calcium hydroxide
Uncommon
in labs because
too expensive
GROUP 1
hydroxides
Some
GROUP 2
hydroxides
•Strong bases also ionize completely in water, except
for Sr(OH)2 and Ca(OH)2 which are only slightly soluble
(remember Mg(OH) is insoluble, K =2x10-11).
Polyprotic Acids
Polyprotic acids dissociate in a stepwise fashion
with different Ka values for each step… In the
second and subsequent ionizations the acids
are always weak, whether or not the original is
a strong or weak acid.
 For most of these acids (ex. H3PO4), the first
dissociation contributes the significant amount
of H+ for pH calculations, and the rest are
negligible (except for H2SO4 where second
ionization is significant).

Naming Acids -REVIEW
-ide ending (elements): “hydro____ic acid”
ex. chloride (HCl): hydrochloric acid
-ate ending (polyatomics): “______ic acid”
ex. chlorate (HClO3): chloric acid
-ite ending(polyatomics): “______ous acid”
ex. chlorite (HClO2): chlorous acid
Net Ionic Equations -REVIEW
For aqueous acid-base reactions reactions, it is
common to write equations in the net ionic form.
Standard form:
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
Ionic form:
Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq)  Na+(aq) +
Cl-(aq) + H2O(l)
Net ionic form:
OH-(aq) + H+ (aq)  H2O(l)
(No spectator ions are included)

Things to remember when
writing Net Ionic Equations




Binary Acids: HCl, HBr, and HI are strong: all other
binary acids and HCN are weak. Strong acids are
written in ionic form; weak acids are written in
molecular form.
Ternary Acids: If the number of oxygen atoms in an
inorganic acid molecule exceeds the number of
hydrogen atoms by two or more, the acid is strong
(complete dissociation). We will consider all organic
acids as weak.
Strong: HClO3, HClO4, H2SO4, HNO3
Weak: HClO, H3AsO4, H2CO3, H4SiO4, HNO2
Polyprotic Acids: (acids that contain more than
one ionizable hydrogen atom. Ex: H2SO4,
H3PO4, H2CO3).
 Bases: Hydroxides of Group 1 and 2 elements
(except Be(OH)2 and Mg(OH)2) are strong
bases. All others including ammonia,
hydroxlamine, and organic bases are weak.
 Salts: Salts are written in ionic form if soluble,
and in undissociated form if insoluble. *Know
the solubility rules.
 Oxides: Oxides are always written in molecular
or undissociated form (ex: MgO).
 Gases: Gases are always written in molecular
form (ex: SO2).

Practice Net Ionic Equations
1. AgNO3 (aq) + H2SO4 (aq) 
2. H4SiO4 (aq) + NaOH (aq) 
3. HBr (aq) + KOH (aq) 
1. Ag+ + HSO4- → AgHSO4(s)
2. H4SiO4 + OH- → H3SiO4- + H2O
3. H+ + OH- → H2O
How do we determine the strength (or pH) of a weak acid?
Relationship between Ka and Kb
K a x K b = Kw
 For any acid and it’s conjugate base, this
relationship can be used to determine Ka or Kb.
 Ex: NH3 + H2O ↔ NH4+ + OHNH4+ + H2O ↔ NH3 + H3O+
Kb(NH3)=[NH4+][OH-]
Ka(NH4+)=[NH3][H3O+]
[NH3]
[NH4+]

Therefore, Ka x Kb = [OH-][H3O+]= Kw
Weak acids and bases will have Ka or Kb values less than one,
but greater than water dissociation, Kw (essentially Ka and Kb of H2O)
Soren Sorenson
(1868-1939)
invented the pH
scale while creating
a waypH
to test
the (potential hydrogen scale)
 The
scale
of beer. Beer
+)
aacidity
measure
of
hydronium
ion
(H
O
3
has a pH of about
concentration.
4.5.
pH Scale
is
Hydronium ion concentration indicates
acidity. Each increase in pH # means a 10fold decrease in [H+].
 The higher the [H3O+], the higher the
acidity.

pH scale and [H+]
Calculating pH
The concentration (M or mol/L) of H3O+ is
expressed in powers of 10, from 10-14 to
100.
 Scientists use pH, which is the negative log
of [H3O+].

pH = -log[H3O+]

Note: The significant figures for logarithmic
numbers are given after the decimal, and the
numbers preceding the decimal give the exponent.
Understanding pH and [H+]
Try filling in this table in your notes:
x
vs.
-log x
[H+]
pH
-1
Estimate the pH 10
of a 3M HCl
1
0
solution.
0.1
1
Estimate the pH
0.001
3
of a 3 x10-5 M
HCl solution.
5
10-5
7
10-7
14
10-14

Calculating pH of a strong acid:
Ex: Given a solution of 0.50 M HCl, what
is the pH?
Step 1: Find [H3O+] in mol/L
0.50 mol/L = 5.0 x 10-1 mol/L
Step 2: Place value in equation and solve.
pH = -log[5.0 x 10-1] = 0.30
Practice pH Calculations

1.
2.
3.

1.
2.
Find pH of the following solutions if
[H3O+] is (try estimating first!):
1.00 x 10-3
6.59 x 10-6
9.47 x 10-10
Find [H3O+] if the pH is:
6.678
3. 10.0
2.533
4. 2.56
pOH
You can calculate the pH of a solution if you
know the concentration of hydronium ion.
[OH-]
 If we use the ion product constant of water
we can derive this equation:
[pH][pOH] = 1.00 x 10-14
 Working with this equation leads to:
pH + pOH = 14

Calculating pH of a strong base:
Ex: Find the pH of a solution with an [NaOH]
of 1.0 x 10-8.
Step 1: Solve for [H3O+] in equation:
[H3O+] = 10-14
[OH-]
Step 2: Place values in:
[H3O+] = 10-14
= 10-6 M
[1.0 x 10-8]

Step 3: Solve for pH by placing [H3O+]
in pH = –log[H3O+]
pH = -log(1.0 x 10-6)
pH = 6.0
Practice pH Calculations Using
pOH

1.
2.
3.
4.
Find the pH of the following solutions
with [OH] of:
1.00 x 10-4
2.64 x 10-13
5.67 x 10-2
3.45 x 10-11
Calculating pH, pOH, [H+] and [OH-]

If one of the these values is known, all others
can be found using the following relationships:
pH + pOH = 14
pH
pH=-log[H+]
[H+]=10-pH
pOH
pOH=-log[OH-]
[H+]
[OH-]=10-pOH
[OH-]
[H+] [OH-] = Kw
Calculating pH of a weak acid:
Ex: Find the pH of a 1.00 M HF solution (Ka=7.2 x10-4):
HF(aq) ↔ H+(aq) + F-(aq)
Ka= [H+][F-]
[HF]
Use ICE box method:
HF(aq) ↔ H+(aq) + F-(aq)
I
1.00M
0
0
C
-x
+x
+x
E
1.00 -x
x
x
Ka= [H+][F-] = (x)(x) ≈ x2
[HF]
(1.00-x)
(1.00)
x2 = Ka(1.00) = (7.2 x10-4)(1.00)
x ≈ 2.7 x10-2
CHECK: 5% rule is valid since
2.7 x10-2/1.00 = 2.7%
pH = -log[H+] = 1.57
Try AP 2005 practice problem!!
Acid-base properties of salts


Salt hydrolysis is the reaction in which a salt
dissolved in water produces an acid or basic
solution (opposite of a neutralization
reaction):
Ex1:
AlCl3 + H2O → Al(OH)3 + 3HCl
weak base

strong acid
A salt that contains the conjugate base of a
strong acid will produce a slightly acidic
solution when dissolved in water.
Strength of Acid-base pairs
Strong acids yield WEAK conjugate bases in
water… they have a low affinity for H+
 Weak acids yield STRONG conjugate bases
 Strong bases yield WEAK conjugate acids
 Weak bases yield STRONG conjugate acids


This is important when we determine the
acidic/basic properties of salts….
Acid-base properties of salts

Ex2:
NaC2H3O2 + H2O → NaOH + HC2H3O2
strong base weak acid
A
salt that contains the conjugate base
of a weak acid will produce a slightly
basic solution when dissolved in water.
 What will happen when NaCl dissolves
in water?
Relative
strength of
conjugate A/B
pairs. Reactions
that are most
likely to occur
will be between
top left SA and
bottom right
SB.
Which acids are stronger
than water? Which bases are
stronger than water?
Percent dissociation
Percent dissociation = amount dissociated x 100%
initial concentration
Practice Problem: find Ka given %
dissociation of a weak acid:
Find [H+] first using % dissociation
 Use formula Ka=x2/[acid]0 - x to find Ka

 Make up your own problem from a % dissociation value…
The effect of structure on
acid-base properties


Relative acidity of oxyacids (hydrogen is
attached to oxygen, and acid contains one
other element):
Acid strength increases as the O-H bond is
weakened (or with increasing # of oxygen
atoms on central atom).
HClO4>HClO3>HClO2>HClO
Bond strength of H-F is too strong (F is so
small) to make it a strong acid (will not
dissociate easily).
HI>HBr>HCl>>HF
BUFFERS
A buffer solution is one which resists
changes in pH when small quantities of
an acid or a base are added to it.
 How do buffer solutions work?
A buffer solution has to contain ions or
molecules which will remove any
hydrogen ions or hydroxide ions that
you might add to it -otherwise the pH
will change.
 buffer demo

ex: HF and NaF
ex: NH3 and NH4Cl
Calculating pH of a buffer

Taking the log of the Ka equation (rearranged
to solve for H+)…
…. yields the Henderson-Hasselbalch equation:
This is used to find the pH of a buffer solution
H-H Special case…

What happens with equimolar amounts of
acid and conjugate base ion?

This situation simplifies to pH = pKa, since
log 1=0.
Preparing buffers:
1) A buffer can be prepared by adding a
common ion (the conjugate base) to a
weak acid.
ex: CH3COOH + NaCH3COO
2) Alternatively, a buffer can also be
prepared by partially neutralizing a weak
acid with a strong base to produce the
conjugate base anion.
ex: CH3COOH + OH- → CH3COO- + H2O
Practice Problem
A buffer problem is simply a weak acid
equilibrium problem that involves a common
ion..
 Ex: Calculate the [H+], pH and percent
dissociation of HF in a buffer solution that
contains 1.0 M HF (Ka= 7.2 x10-4) and 1.0 M
NaF.
 ANS:
HF ↔ H+ + FKa= 7.2 x10-4 = [H+][F-] = x(1.0 + x) ≈ x(1.0)
[HF]
1.0 – x
1.0

Solving for x = 7.2 x10-4
[H+] = 7.2 x10-4 and pH = 3.14
% dissociation is 7.2 x10-4 x100% = 0.072%
1.0 M
Buffer Capacity



The buffering capacity of a solution
represents the amount of H+ or OH- the
buffer can absorb without significant changes
in pH.
The pH of a buffered solution is determined
by the ratio [A-]/[HA]. The capacity of a
buffered solution is determined by the
magnitudes of [HA] and [A-].
Therefore, a good buffer solution will have
relatively high concentrations of BOTH
components.
Practice Problem: Buffer Preparation

A chemist needs a solution buffered at pH 4.30 and can choose from the
following acids and their sodium salts:
a) CH2ClCOOH, cloroacetic acid (Ka = 1.35x10-3)
b) CH3CH2COOH, propanoic acid (Ka = 1.3x10-5)
c) C6H5COOH, benzoic acid (Ka = 6.4x10-5)
d) HOCl, hypochlorous acid (Ka = 3.5x10-8)
Calculate the ratio [HA]/[A-] required for each system to yield a pH of 4.30.
Which will work best?
A pH of 4.30 corresponds to [H+] = 5.0x10-5 M, and [H+] = Ka[HA]/[A-]
Using Ka values above, the [HA]/[A-] ratio for each of the salts are:
a) cloroacetic acid:
3.7x10-2
b) propanoic acid:
3.8
best since [HA]/[A-]
c) benzoic acid:
0.78
ratio is closest to 1
d) hypochlorous acid: 1.4 x103
Indicators
pH and titration curves –SA and SB
Adding SA to SB
Adding SB to SA
In what way are the curves different? Similar?
Comparing WA and SA added
to SB….
What are the differences? Similarities?
Adding Acid to Base
Adding WA to SB
Adding SA to WB
In what way are the curves different? Similar?
Adding Base to Acid
Adding WB to SA
Adding SB to WA
Finding Ka from titration curve
More complicated titration curves…
Adding a weak acid to weak base… Adding H2CO3 to strong base…
usually too difficult to measure
two equivalence points seen in the
endpoint
polyprotic acid titration
Neutralization Reaction
Acid-Base Titration
Make up a weak acid name,
formula and Ka….
What will the titration curve look
like?