The concentration of specific ions in solution determines whether the solution is acidic or basic.
Acids and bases can be described as substances that either increase or decrease the
concentration of hydrogen (\(\text{H}^{+}\)) or hydronium (\(\text{H}_{3}\text{O}^{+}\)) ions in
a solution. An acid increases the hydrogen ion concentration in a solution, while a
base decreases the hydrogen ion concentration.
pH is used to measure the concentration of \(\text{H}^{+}\) ions
(\([\text{H}^{+}]\)) and therefore, whether a substance is acidic or basic (alkaline). Solutions
with a pH of less than seven are acidic, while those with a pH greater than seven are basic
(alkaline). The pH scale ranges from \(\text{0}\) to \(\text{14}\) and a pH of \(\text{7}\) is
considered neutral.
The universal indicator changes colour from \(\color{red}{\textbf{red}}\) in
\(\color{red}{\textbf{strongly acidic}}\) solutions through to
\(\color{purple}{\textbf{purple}}\) in \(\color{purple}{\textbf{strongly basic}}\)
solutions.
The term pH was first used by in 1909 by Søren Peter Lauritz Sørensen (a Danish
biochemist). The p stood for potenz and the H for hydrogen. This
translates to power of hydrogen.
pH
pH is a measure of the acidity or alkalinity of a solution.
At the beginning of the chapter we mentioned that we encounter many examples of acids and bases in
our day-to-day lives. The pH of solutions of some household acids and bases are given in Table 9.6.
Fizzy cooldrinks often have very low pH (are acidic).
Drink
pH
Coke
\(\text{2,5}\)
Diet coke
\(\text{3,3}\)
Pepsi
\(\text{2,5}\)
Diet pepsi
\(\text{3,0}\)
Sprite
\(\text{3,2}\)
7 Up
\(\text{3,2}\)
Diet 7 Up
\(\text{3,7}\)
Molecule
Found in
pH
Type
phosphoric acid
fizzy drinks
\(\text{2,15}\)
acid
tartaric acid
wine
\(\text{2,95}\)
acid
citric acid
lemon juice
\(\text{3,14}\)
acid
acetic acid
vinegar
\(\text{4,76}\)
acid
carbonic acid
fizzy drinks
\(\text{6,37}\)
acid
ammonia
cleaning products
\(\text{11,5}\)
base
ammonium hydroxide
cleaning products
\(\text{11,63}\)
base
sodium hydroxide
caustic soda
\(\text{13}\)
base
Table 9.6: The pH of solutions of acids and bases as found in common household items.
pH calculations (ESCPD)
The pH scale is a log scale. Remember from mathematics that a difference of
one on a base \(\text{10}\) log scale (the one on your calculator) is equivalent to a
multiplication by 10. That is:
1 = log(10)
2 = log(100)
3 = log(\(\text{1 000}\))
4 = log(\(\text{10 000}\))
So a change from a pH of 2 to a pH of 6 represents a very large change in the
\(\text{H}^{+}\) concentration.
pH can be calculated using the following equation:
pH = -log[\(\text{H}^{+}\)]
[\(\text{H}^{+}\)] and [\(\text{H}_{3}\text{O}^{+}\)] can be substituted for one another:
pH = -log[\(\text{H}_{3}\text{O}^{+}\)]
The brackets in the above equation are used to show concentration in
\(\text{mol·dm$^{-3}$}\).
Worked example 11: pH calculations
Calculate the pH of a solution where the concentration of hydrogen ions is
The number of moles of \(\text{CH}_{3}\text{COO}^{-}\) is found to be \(\text{0,001}\)
\(\text{mol}\). Calculate the pH of the solution.
Determine the number of moles of hydronium ions in the solution
According to the balanced equation for this reaction, the mole ratio of
\(\text{CH}_{3}\text{COO}^{-}\) ions to \(\text{H}_{3}\text{O}^{+}\) ions is
\(\text{1}\):\(\text{1}\), therefore the number of moles of these two ions in the
solution will be the same.
So, \(n(\text{H}_{3}\text{O}^{+}) = \text{0,001} \text{ mol}\).
Determine the concentration of hydronium ions in the solution
Understanding pH is very important. In living organisms, it is necessary to maintain a constant pH
in the optimal range for that organism, so that chemical reactions can occur.
pH
\(\text{1}\)
\(\text{6}\)
\(\text{7}\)
\(\text{8}\)
\(\text{13}\)
\([\text{H}^{+}]\)
\(\text{1} \times \text{10}^{-\text{1}}\)
\(\text{1} \times \text{10}^{-\text{6}}\)
\(\text{1} \times \text{10}^{-\text{7}}\)
\(\text{1} \times \text{10}^{-\text{8}}\)
\(\text{1} \times \text{10}^{-\text{13}}\)
\([\text{OH}^{-}]\)
\(\text{1} \times \text{10}^{-\text{13}}\)
\(\text{1} \times \text{10}^{-\text{8}}\)
\(\text{1} \times \text{10}^{-\text{7}}\)
\(\text{1} \times \text{10}^{-\text{6}}\)
\(\text{1} \times \text{10}^{-\text{1}}\)
Solution
\(\color{red}{\textbf{strongly acidic}}\)
\(\color{red}{\text{weakly acidic}}\)
neutral
\(\color{blue}{\text{weakly basic}}\)
\(\color{blue}{\textbf{strongly basic}}\)
Table 9.7: The concentration of \([\text{H}^{+}]\) and \([\text{OH}^{-}]\) ions in
solutions with different pH.
A build up of acid in the human body can be very dangerous. Lactic acidosis
is a condition caused by the buildup of lactic acid in the body. It leads to acidification
of the blood (acidosis) and can make a person very ill. Some of the symptoms of lactic
acidosis are deep and rapid breathing, vomiting, and abdominal pain. In the fight against
HIV, lactic acidosis is a problem. One of the antiretrovirals (ARV's) that is used in
anti-HIV treatment is Stavudine (also known as Zerit or d4T). One of the side effects of
Stavudine is lactic acidosis, particularly in overweight women. If it is not treated
quickly, it can result in death.
Litmus paper can be used as a pH indicator. It is sold in strips. Purple litmus paper will
become red in acidic conditions and blue in basic conditions. Blue litmus paper is used to
detect acidic conditions, while red litmus paper is used to detect basic conditions.
In agriculture, it is important for farmers to know the pH of their soils so that they are able to
plant the right kinds of crops. The pH of soils can vary depending on a number of factors, such
as rainwater, the kinds of rocks and materials from which the soil was formed and also human
influences such as pollution and fertilisers. The pH of rain water can also vary, and this too
has an effect on agriculture, buildings, water courses, animals and plants. Rainwater is
naturally acidic because carbon dioxide in the atmosphere combines with water to form carbonic
acid. Unpolluted rainwater has a pH of approximately \(\text{5,6}\). However, human activities
can alter the acidity of rain and this can cause serious problems such as acid rain.
Calculating pH
Textbook Exercise 9.5
Calculate the pH of each of the following solutions:
(Tip: \([\text{H}_{3}\text{O}^{+}][\text{OH}^{-}]\) = \(\text{1} \times
\text{10}^{-\text{14}}\) can be used to determine
\([\text{H}_{3}\text{O}^{+}]\))
A \(\text{KOH}\) solution with a \(\text{0,2}\)
\(\text{mol·dm$^{-3}$}\) concentration of \(\text{OH}^{-}\).
Remember that \([\text{OH}^{-}][\text{H}_{3}\text{O}^{+}]\) =
\(\text{1} \times \text{10}^{-\text{14}}\).
An aqueous solution with a \(\text{1,83} \times
\text{10}^{-\text{7}}\) \(\text{mol·dm$^{-3}$}\)
concentration of
\(\text{HCl}\) molecules at equilibrium (\(\text{K}_{\text{a}}\) =
\(\text{1,3} \times \text{10}^{\text{6}}\))
To determine the concentration of \(\text{H}_{3}\text{O}^{+}\) ions
in solution we must first write the expression for
\(\text{K}_{\text{a}}\), and to do that we need the balanced
equation:
In a typical sample of seawater the concentration the hydronium
(\(\text{H}_{3}\text{O}^{+}\)) ions is \(\text{1} \times
\text{10}^{-\text{8}}\) \(\text{mol·dm$^{-3}$}\), while the
concentration of the hydroxide (\(\text{OH}^{-}\)) ions is \(\text{1}
\times \text{10}^{-\text{6}}\) \(\text{mol·dm$^{-3}$}\).
Is the seawater acidic or basic?
The concentration of \(\text{H}_{3}\text{O}^{+}\) is lower than that
of \(\text{OH}^{-}\). The seawater is therefore basic.
This type of reaction (the transfer of a proton between identical molecules) is known as
auto-protolysis. This reaction is also known as the
auto-ionisation of water and the ions formed are a conjugate acid and base pair
of water:
Auto-protolysis and auto-ionisation of water
Auto-protolysis is the transfer of a proton between two of the same molecules. The
auto-ionisation of water is one example of auto-protolysis.
\(\text{K}_{\text{w}}\) is the equilibrium constant for this process:
\(\text{K}_{\text{w}}\) = \(\text{1} \times \text{10}^{-\text{14}}\) at \(\text{25}\)
\(\text{℃}\)
Salt hydrolysis (ESCPG)
Does neutralisation mean that the pH of the solution is \(\text{7}\)? No. At the equivalence point
of a reaction, the pH of the solution need not be \(\text{7}\). This is because of the
interaction of the salt (formed by the reaction) and water.
At the equivalence point of an acid-base neutralisation reaction there is salt and water. The ions
of water interact with the salt present and form a small quantity of excess hydronium ions
(\(\text{H}_{3}\text{O}^{+}\)) or hydroxide ions (\(\text{OH}^{-}\)). This leads to pH values
that are not equal to \(\text{7}\).
A simple rule for determining the likely pH of a solution is as follows:
A neutralisation reaction does not imply that the pH is neutral
(\(\text{7}\)).
A \(\color{red}{\textbf{strong }\text{acid}} + \color{blue}{\textbf{strong
}\text{base}}\) form a \(\color{darkgreen}{\textbf{neutral}}\) salt and water
\(\color{darkgreen}{\textbf{solution}}\):
A \(\color{red}{\textbf{weak }\text{acid}} + \color{blue}{\textbf{strong
}\text{base}}\) form a \(\color{blue}{\textbf{weak basic}}\) salt and water
\(\color{blue}{\textbf{solution}}\):
A \(\color{red}{\textbf{strong }\text{acid}} + \color{blue}{\textbf{weak
}\text{base}}\) form a \(\color{red}{\textbf{weak acidic}}\) salt
\(\color{red}{\textbf{solution}}\):
Table 9.8: The approximate pH of neutralisation reaction solutions based on the
strength of the acid and base used.
Indicators (ESCPH)
A titration is a process for determining, with precision, the concentration of a solution with
unknown concentration. The theory behind titrations will be discussed later in this chapter. An
indicator is used to show the scientist carrying out the reaction exactly when the reaction has
reached completion.
This experiment looks at the change in colour of an indicator during an acid-base reaction.
It is effectively a very rough titration experiment. Principles that can be applied to
titrations, such as adding a small volume of acid, then swirling, can be applied here as
well. It is important that the learners understand that the pH range that the indicator
changes colour in is not always around \(\text{7}\).
Learners are working with a strong acid and a strong base in this reaction. Concentrated,
strong acids and bases can cause serious burns. Please remind the learners to be careful
and wear the appropriate safety equipment when handling all chemicals, especially
concentrated acids and bases. The safety equipment includes gloves, safety glasses, and
protective clothing.
Indicators
Aim
To investigate the use of an indicator in an acid-base reaction.
Apparatus and materials
one volumetric flask, one conical flask, one pipette, a piece of white
paper or a white tile
A \(\text{1}\) \(\text{mol·dm$^{-3}$}\) solution of sodium hydroxide
(\(\text{NaOH}\)), a \(\text{1}\) \(\text{mol·dm$^{-3}$}\) solution
of
hydrochloric (\(\text{HCl}\)), an indicator
Method
Measure \(\text{20}\) \(\text{ml}\) of the sodium hydroxide solution into a
conical flask. Add a few drops of the indicator.
In this experiment colour change is very important. So place the conical
flask on a piece of white paper or a white tile to make any colour
change easier to observe.
Slowly add \(\text{10}\) \(\text{ml}\) of hydrochloric acid. If there is a
colour change, then stop. If there is no colour change add another
\(\text{5}\) \(\text{ml}\). Continue adding \(\text{5}\) \(\text{ml}\)
increments until you notice a colour change.
Observations and discussion
The solution changes colour after a certain amount of hydrochloric acid is added. This
is because the solution now contains more acid than base and has therefore become
acidic. It can be concluded that the indicator is one colour in a basic solution and
a different colour in an acidic solution.
Indicators are chemical compounds that change colour depending on whether they are in an acidic or
a basic solution. A titration requires an indicator that will respond to the change in pH with a
sensitive and quick colour change. Typical indicators used in titrations are given in Table 9.9.
Titration type
Preferred
indicator
Colour of
acid
Colour of
end-point
Colour of
base
pH range
strong acid +
strong base
bromothymol
blue
yellow
green
blue
\(\text{6,0}\) - \(\text{7,6}\)
weak acid +
strong base
phenolphthalein
colourless
faint pink
pink
\(\text{8,3}\) - \(\text{10,0}\)
strong acid +
weak base
bromocresol
green
yellow
green
blue
\(\text{3,8}\) - \(\text{5,4}\)
Table 9.9: Some typical indicators for typcial titrations.
Revise Grade 11 Acids and Bases for more information on plants and foods that can be used as
indicators.
Notice that the pH range for colour change of the indicator used should match the approximate
pH expected for that type of titration (see Table 9.8):
solution
pH
pH range
neutral
\(\text{7}\)
\(\text{6,0}\) - \(\text{7,6}\)
weak
basic
\(\text{9}\)
\(\text{8,3}\) - \(\text{10,0}\)
weak
acidic
\(\text{5}\)
\(\text{3,8}\) - \(\text{5,4}\)
Figure 9.3: Some typical indicators for typcial titrations: bromothymol blue (left),
phenolphthalein (centre), and bromocresol green (right).
Indicators change colour (Figure
9.3) according to where the \(\text{H}\) is:
So, when an acid is added to aqueous bromothymol blue there will be extra \(\text{H}^{+}\) ions.
The equilibrium will shift (remember le Chatalier's principle) to decrease the number of
\(\text{H}^{+}\) ions. That is, to the left. If sufficient acid is added, the entire solution
will become acidic. This means there will be more HBromothymol blue than
bromothymol blue\(^{-}\) and the solution will become yellow.
