S. O’Brien
Donald High School
Year 12 Chemistry: Chapter 4:~ Analysing Acids and Bases
4.1 Acid-base Chemistry Revisited
 Acids are proton donors
 Bases are proton acceptors
 Acid-base reactions involve the transfer of a proton from an acid to a base.
This definition was proposed by the Danish chemist Johannes Bronsted and the English
chemist Thomas Lowry in 1923. It is known as the Bronsted-Lowry Theory.
Hydrochloric acid is a strong acid that ionises almost completely in water forming H3O+
and Cl- ions.
Substances such as HCl and Cl- that differ by only one proton (H+) are called a conjugate
acid-base pair, eg H2O / H3O+.
Diprotic acid can donate one or two protons, eg Sulfuric acid H2SO4.
Triprotic acid can donate up to three protons, eg Phosphoric acid H3PO4.
Strong acids, such as hydrochloric acid are acids that readily donate protons. The react
almost completely with water to form ions. Weak acids such as ethanoic (acetic) acid,
ionise only slightly:
Similarly, strong bases are bases that readily accept protons and weak bases accept
protons only to a limited extent.
The concentration of acids and bases can be determined experimentally by volumetric
analysis. Acid-base titrations utilise the neutralisation reaction between acids and bases
to from a salt plus water.
S. O’Brien
Donald High School
4.2 pH
The concentration of H3O+ ions in a solution is referred to as the solution’s acidity.
Acidity is measured using a logarithmic scale, called the pH scale. The definition of pH
is:
Where [H3O+] is the concentration of H3O+ ions measured in mol L-1.
 Neutral solutions have a pH =
at 25C
 Acidic solutions have a pH
 Basic solutions have a pH
4.3 Indicators
An indicator is used during acid-base titration to identify the equivalence point of the
reaction. An acid-base indicator is a substance whose colour depends on the
concentration of H3O+ ions in solution. Indicators are weak acids with their acid from
being one colour and their conjugate base being another.
Indicator
Phenolphthalein
Methyl orange
Bromothylmol blue
Colour of acid form
Colour of base form
pH range
The indicator must be chosen carefully to ensure the end point closely matches the
equivalence point of the reaction.
At the end point, addition of a very small volume of strong acid produces a large change
in pH. This is referred to as a sharp end point.
When a weak base is titrated with a strong acid, or a strong base with a weak acid,
there is a much more gradual change in pH around the end point.
In the case of reactions between weak acids with weak bases the change is so gradual
that this combination cannot be analysed by a simple direct titration.
S. O’Brien
Donald High School
Examples:
The concentration of ethanoic acid (CH3COOH) in a brand of white vinegar was
determined by titration with standard sodium hydroxide solution.
A 25.00 ml aliquot of vinegar was pipetted into a flask and several drops of
phenolphthalein indicator added. Using a burette, 0.995 M sodium hydroxide solution was
slowly added until the indicator turned permanently pink. The volume of sodium
hydroxide solution required to reach this point was 21.56 mL. Calculate the
concentration of acid in the vinegar.
The equation is (STEP 1):
Solution:
Step 2: calculate the amount in mol of the known solution
Step 3: mole ratio
Step 4: calculate the amount in mol of the unknown
Step 5: calculate the concentration of the unknown.
S. O’Brien
Donald High School
Example: A commercial concrete cleaner contains concentrated hydrochloric acid. A
25.00 mL volume of cleaner was diluted to 250.0 mL in a volumetric flask.
A 20.00 mL aliqu0t of 0.448 M sodium carbonate solution was placed in a conical flask.
Methyl orange indicator was added and the solution was titrated with the diluted
cleaner. The methyl orange indicator changed permanently from yellow to pink when
19.84 mL of the cleaner was added. Calculate the concentration of hydrochloric acid in
the concrete cleaner.
Solution:
Step 1: Balanced chemical equation
Step 2: Calculate amount in mol of the known solution
Step 3: Mole ratio
Step 4: Calculate amount of the unknown solution
Step 5: calculate the concentration of the unknown using the dilution factor.
This analytical method is a convenient way to determine the amount of active ingredient
in products such as cleaning agents containing ammonia, antacid tablets or aspirin.
Look at ionic equations (pg 41)
Questions: 4, 5, 10, 11, 6, 7, 16, 17, 18, 22, 28, 9 & 26.
S. O’Brien
Donald High School
4.4 Back titrations
Some acids and bases are so weak that they do not produce a sharp colour change at the
end point of a titration. Back titration is used to overcome this problem.
Back titration has two parts:
 If the substance to be analysed acts as a weak acid, it is mixed with an excess of
strong base. The original amount of strong base is known. All of the weak acid
reacts (in stoichiometric proportions), leaving some of the strong base unused.
 The unused strong base is titrated as normal with a standard solution of a strong
acid. Knowing the original amount of strong base and the amount of strond base
left unused, it is possible to work back to find the amount of weak acid present.
Example: A 1.50 g of lawn fertiliser was boiled with 25.00 mL of 0.9987 M sodium
hydroxide solution. When no further ammonia gas was evolved from the mixture, it was
cooled and titrated with 0.2132 M hydrochloric acid, using phenolphthalein as an
indicator. A titre of 19.78 mL was required. Calculate the percentage of ammonium ions
in the fertiliser.
Step 1: Find the original amount of NaOH used:
n(NaOH) =
Step 2: Find the amount of NaOH that did not react with NH4+ ions.
*write a chemical formula:
n(HCl) =
mole ratio =
n(NaOH) =
Step 3: The amount of NaOH that reacted with NH4+ ions is given by:
n(NaOH) reacting = n(NaOH) used originally - n(NaOH) not reacting with NH4+ ions
S. O’Brien
Donald High School
Step 4: Find the amount of NH4+ in the fertiliser sample.
Write an ionic equations:
Mole ratio:
Step 5: The percentage of NH4+ in the fertiliser sample can now be found:
QUESTIONS: 9 and 26