8.1: Understanding Acids and Bases
According to Arrhenius, ionic compounds separate into ions when they are liquid or in
solution.
Acids
In their pure form (liquid/solids), acids are molecular compounds that contain hydrogen.
They cannot conduct electricity and are termed non-electrolytes.
However, in solution, acids ionize to release hydrogen ion (H+) (aq).
Ionization ---- Any process by which a neutral atom or molecule is converted into an
ion.
HCl(aq) ------ H+ (aq+
Cl- (aq)
We can define an acid as a hydrogen containing compound that ionizes in water to
produce hydrogen ions.
Acids
Electrolyte
Solid
No – behaves as a
molecular
compound
Liquid
No – behaves as a
molecular
compound
Aqueous
Yes – behaves as an
ionic compound in
solution
Bases
Most bases are ionic compounds that contain a hydroxide (OH-), which is released in
solution. We can define a base as an ionic hydroxide that releases mobile hydroxide ions
in solution.
NaOH ----- Na+ (aq) + OH- (aq)
Ba(OH)2(s) ---- Ba2+ (a) + 2OH- (aq)
Bases
Electrolyte
Solid
No
Liquid
Yes
Aqueous
Yes
Properties of acids and bases
Identifying acids in some products is fairly easy. Identifying bases is more difficult.
Acids and bases have different properties however, that enables you to distinguish
between them.
Property
Taste
Electrical
Conductivity in
Solution
1
Feel of Solution
Acids
Taste sour
Conduct electricity
Have no
characteristic feel
Bases
Taste bitter
Conduct electricity
Feel slippery
Reaction with
litmus paper
Reaction with active
metals
Reaction with
carbonate
compounds
Acids
Turns blue litmus
paper red
Produce hydrogen
gas
Produce carbon
dioxide gas
Bases
Turn red litmus
paper blue
Do not react
Do not react
Acid-Base Indicators
Indicators
Acid Colour
Orange IV
2,4-Dinitrophenol
Congo red
Methyl orange
Litmus
Bromothymol blue
Brilliant yellow
Phenolphthalein
Range
red
colourless
blue
red
red
yellow
yellow
colourless
1.4 - 2.6
2.8 - 4.0
3.1 - 4.9
3.2 - 4.4
5.0 - 7.0
6.0 - 7.6
6.6 - 7.9
8.3 - 10.0
Base Colour
yellow
yellow
red
yellow
blue
blue
orange
dark pink
The Arrhenius theory was a major advance in understanding chemical substances and
solutions.
Strong and Weak Acids/Arrhenius Concepts
Strong Acid
An acid that ionizes almost completely in H2O (>99%) to form aqueous hydrogen ions.
Therefore, high percentage ionization, high conductivity.
HNO, (aq) ------ H+ (aq)
+
NO3- (aq)
Weak Acid
An acid that ionizes only partially (<59%) in water to form aqueous hydrogen ions.
Therefore, low percentage ionization, low conductivity.
Strong Base
2
A strong base is the base that dissociates totally. All bases are strong
The Arrhenius Concepts of Acids and Bases
Svante Arrhenius noticed a pattern of the dissociations of acids and bases in water.
HBr (aq) ----------- H+ (aq) + Br- (aq)
H2SO4 (aq) ------- H+ (aq) + HSO4- (aq)
HCLO4 (aq) ------ H+ (aq) + CLO4- (aq)
LiOH (aq) ------- Li+ (aq) + OH- (aq)
KOH (aq) -------- K+ (aq) + OH- (aq)
Ba(OH)2 (aq) -- Ba+2 (aq) + OH- (aq)
8.2 Calculating pH and Hydrogen Ion Concentration [H+]
The relationship between the [H+] and the pH is easy to calculate if the concentration is
only a power of ten.
[H+] = 1.0 × 10-3 mol/L
pH = 3
[H+] = 10 -7 mol/L
pH = 7
However, this method cannot be used if we are given a value such as 2.7 × 10-3 mol/L.
To answer this question, we need to use logarithms.
Log --- The log of a number is the exponent of that number when it is written in
exponential form.
100 = 102 or log10(10-2) = 2
0.01 = 10-3 or log 10(10-3) = 3
Generally:
If y = 10x
The, log10(y) = x
Therefore,
pH = -log[H+] and pOH = -log[OH-]
1. An antacid solution has a hydrogen ion [ ] of 4.7 × 10-11 mol/L. What is its pH?
pH
= -log [H=]
=-log (4.7 × 10-11)
3
= 10.33
2. Determine the hydrogen ion [ ] of a solution that has a pH of 10.33.
[H+] = 10-pH
= 10-10.33
[H+] = 4.7 × 10-11
8.4: Revision of Arrhnius’ Definitions
Although Arrhenius’ definition could be used to predict the existence of most acids and
bases, there are some compounds that it would predict to be neutral but they are not.
The new theory tells us the following:
It is highly unlikely that the particle we call an H+ ion actually exists in an acid solution.
Instead, it would be more likely to bond with the polar water molecule. The result would
be a hydrated proton.
H3O+ (hydronium ion)
Therefore, we can now explain the formation of acids as a reaction with water forming
hydronium ions.
HCl + H2O ------ H2O+ + ClNH3 + H2O ----- NH4+ + OHNa2CO3 ------- 2Na+ + CO3-2
CO3-2 is responsible for the basic character
CO3-2
+
H2O (1)( ------ OH-
+
HCO3-
Bronsted-Lowry Concept
In the Bronsted Theory of an acid and base, acids and bases are defined as follows:
Acid: A substance that donate a proton (H+, a hydrogen atom that has lost its electron)
Base: A substance that accepts a proton from another substance.
The following reaction is thus an acid-base reaction according to Bronsted.
HCl + H2O ----------
H3O+ + Cl-
1. The HCl is an acid because it donates its proton to H2O to form Cl2. H2O is a base because it accepts that proton to form H3O+
The acid on the right of the equation is related to the base on the left, they are said to be
conjugate.
4
Therefore:
HCl --------- ClAcid

H2O ---------- H3O+
Conjugate acid
Base
Conjugate base
Bronsted-Lowry base is specific to each reaction. A base in one reaction can be an
acid in another. These substances are called amphiprotic.
HCO3- (aq)
HCO3- (aq)
+
+
H3O+ (aq) --
OH- (aq) ---
H2CO3- (aq)
CO32- (aq)
+ H2O (1)
+ H2O (1)
Acid-Base Titration
Titration is a common laboratory technique used to determine the concentration of
substances in solution. It is a procedure that involves carefully adding a controlled and
measured volume of a solution from a buret into a measured volume of a sample solution.
Titrant: ------ the solution of unknown that is n the flask of known concentration.
Standard Solution: ------ The solution that is in the flask of known concentration.
The titrant is added drop by drop until the reaction is judged to be complete. The point at
which the indicator changes colour is known as the endpoint. This is when chemically
equivalent amounts of reactants (determined by the mole ratio in a balanced chemical
equation) have been combined.
Acid-Base Titration
Suppose 1.09 g of anhydrous sodium carbonate is dissolved to make 100.0 mL of a
standard solution. Samples (10.00 mL) of this standard solution are then taken and
titrated with the HCl product, which has been diluted by a factor of 10. The titration
evidence collected is shown below.
Trial
1
2
3
Final buret reading (mL)
13.3
26.0
38.8
13.4
Initial buret reading (mL)
0.2
13.3
26.0
0.6
Volume of HCl added (mL)
13.1
12.7
12.8
12.8
5
4
An acid rain sample containing sulfurous acid was analyzed in a laboratory using a
titration with a standard solution of sodium hydroxide. Use the evidence given below to
determine the concentration of the sulfurous acid.
6