Integrated Science SME

Key Unit competence: To prepare standard solutions and use
                                               them to determine concentration of other
                                               solutions by method of titration.

Introductory Activity 4
Observe the photo below and attempt the questions that follow:

1. For the bottle of beer 1, and that of Beer 2, we find on the bottle of its
labels 5% alcohol and 6% alcohol respectively.
a) Explain the meaning of 5% and 6% alcohol.
b) 5% alcohol corresponds to which volume of alcohol of the total
volume of 72cl of the bottle of beer 1.
c) Calculate the volume of alcohol corresponding to 6% alcohol of
the bottle of Beer 2 of the total volume of 65cl.
2. On the label of the bottle of fruit juice we find that dilution is 1:5.
Explain this ratio and explain also the purpose of diluting substances.

4.1. Definition of standard solution and primary standard solution
Activity 4.1
1. Explain the concepts of:
a) Solution
b) concentration
2. You are given a basic solution, NaOH (aq), and you are requested
to determine its concentration. How could you proceed to know its
concentration?

A homogeneous mixture is known as a solution.
 An aqueous solution is a solution in which the solvent is water. Its
concentration may be estimated in many ways.

In analytical chemistry, a standard solution is a solution containing a
precisely known concentration of an element or a substance and used to
determine the unknown concentration of other solutions. A known weight of solute is dissolved to make a specific volume. It is prepared using a standard substance, such as a primary standard.

Standardization is the process by which the concentration of a solution is determined by measuring accurately the volume of the solution required to react with an exactly known amount of a primary standard.

A primary standard is defined as a substance or compound used to prepare standard solutions by actually weighing a known mass, dissolving it, and diluting to a definite volume.

Or a substance, which is chemically stable in aqueous solution and its
concentration remains constant with change in time such that it can be used to standardize other solutions.

Some important examples of primary standard are;
• Sodium Carbonate, Na2CO3
• Potassium dichromate, K2Cr2O7
• Benzoic acid, C6H5COOH
• Oxalic acid, H2C2O4
• Iodine, I2

• Sodium oxalate, Na2C2O4
• Butanoic acid, C3H7COOH
• Sodium tetraborate, Borax, Na2B4O7
• Potassium chloride, KCl
• Arsenic(III)oxide, As2O3
• Silver nitrate, AgNO3
Standard solutions are normally used in titrations to determine unknown concentration of another substance.

Application activity 4.1
Differentiate between standard solution and primary standard solution

4.2. Properties of a primary standard solution
Activity 4.2
1. Recall the concept of “primary standard solution”
2. Do research and find out the role and characteristics of a good
primary standard solution.

A good primary standard solution meets the following criteria:
• High level of purity
• High stability
• Be readily soluble in water
• High equivalent weight (to reduce error from mass measurements)
• Not hygroscopic (to reduce changes in mass in humid versus dry
environments)
• Non-toxic
• Inexpensive and readily available
• React instantaneously, stoichiometrically and irreversibly with other
substances i.e. should not have interfering products during titration.
• It should not get affected by carbon dioxide in air

Molar concentrations are the most useful in chemical reaction calculations because they directly relate the moles of solute to the volume of solution. 

The formula for molarity is:

Application activity 4.2
Discuss the properties of a good primary standard solution.

4.3. Preparation of standard solutions
Many of the reagents used in science are in the form of solutions which
need to be purchased or prepared. For many purposes, the exact value of concentration is not critical; in other cases, the concentration of the solution and its method of preparation must be as accurate as possible.

Thus, the preparation of solutions is one of the most fundamental tasks
performed by a chemist.

A solution may be prepared by two methods: dissolution method and
dilution method.

4.3.1. Preparation of standard solution by dissolution of solids
Activity 4.3.1
Some chemicals of high interest used in scientific laboratories, hospitals,
pharmacies, are found in the form of solutions with precise concentration.
1. What is the most common unit of the solution concentration?
2. Using your knowledge in chemistry acquired so far and available
resources identify the requirements and the procedure for the
preparation of such solutions starting with solids.
The most common unit of solution concentration is molarity (M).
The molarity or molar concentration of a solution is defined as the number of moles of solute per one litre (mol/L) of solution.
Molarity, therefore, is a ratio between moles of solute and litres of solution. To prepare laboratory solutions, usually a given volume and molarity are required. To determine molarity, the formula weight or molar mass of the solute is needed.

When preparing a solution starting with a solid, the following steps may be followed:
• Determine the mass in grams of one mole of solute, the molar mass,
Mm.
• Decide volume of solution required, in litres, V.
• Decide molarity of solution required, M.
• Calculate grams of solute required.

In the preparation of solution, glasses, volumetric flask, pipette, droppers, glass rod, measuring cylinder, analytical balance, spatula, beakers, magnetic stirrer and other laboratory devices are used.

The preparation of a solution by dissolution follows the steps illustrated by figure 4.1

The following examples illustrate the calculations for preparing solutions.
Examples:
1. Describe in details how you can prepare the following solution: 50 mL of NaOH, 10%.
Answer:
Preliminary calculations
10% means that in 100 mL of solution only 10 g are pure NaOH. So 50 mL
of NaOH will be prepared by taking the mass of NaOH, dissolving it in water and making up to 50 ml. 

Procedure
• Weigh 5g of NaOH accurately using an analytical balance.
• Carefully dissolve it in a beaker containing about 30mL of distilled water and stir using a glass rod or a magnetic stirrer till you get homogeneous mixture.
• Transfer the solution and the washings in a 50 mL volumetric flask and
mix.
• Add more distilled water until the level is about 2 cm below the
graduation mark on the volumetric flask.
• Top up using distilled water from a dropper until the bottom of the
meniscus is at the level of mark when viewed at eye level.
• Stopper the flask and shake again to homogenise.
• Label your solution: 10% NaOH; 50 mL
• The solution prepared is 10% NaOH
Usually, the minimum label requirements are: (1) Identity of contents (2).
Concentration (3) Your name and (4) date of preparation.

2. Describe in details the preparation of 250 cm3 of a 0.1M Na2CO3
solution.

Answer:
Step 1: Preliminary calculations
Calculate the amount of anhydrous sodium carbonate required to be
dissolved in 250cm3 of solution. i.e:
Molar mass of Na2CO3
= (23 x 2) +12+ (16 x 3) = 106 g/mol
Thus, 1 mole of Na2CO3 has a mass of 106 g.
0.1mole of Na2CO3 solution will have a mass=106g/mol x 0.1mol =10.6 g.
1000 cm3 of 0.1M Na2CO3 solution contains 10.6 g.

Step 2: Procedure
• Weigh 2.65g of sodium carbonate and carefully transfer in a beaker
containing 150mL of distilled water.
• Stir using a glass rod until complete dissolution.
• Using a filter funnel, transfer the solution and the washings, in a 250
mL volumetric flask.

• Using a wash bottle, add more distilled water up to about 1cm of the
graduation mark.
• Using a dropper pipette, complete the volume the graduation mark.
• Stopper the flask and shake to homogenise.
• Label the solution.

Application activity 4.3.1
Give a detailed description, including calculations involved, for the
preparation of the following solutions:
a). 1L of a 0.01M KMnO4solution
b). 250 mL oxalic acid 10g/L.

4.3.2. Preparation of standard solution by dilution

Activity 4.3.2
Some chemicals of high interest used in scientific laboratories, hospitals,
pharmacies, are found in the form of solutions with precise concentration.

Some are obtained by dissolution of solids and others may be obtained
by dilution of stock solution.
1. What is meant by:
a) stock solution? b) dilution?
2. Do research to identify the requirements and the procedure for
the preparation of such solutions starting with solutions or liquids.
Apart from, dissolution method, we can prepare a solution by dilution from the stock solution. This consists of reducing the concentration of a concentrated stock solution to less concentrated

A simple dilution is one in which a unit volume of material of interest (solute) is combined with an appropriate volume of a solvent to achieve the desired concentration.

The dilution factor is the total number of unit volumes in which the solute is dissolved. For example a 1:5 dilution involve combining 1 unit of the solute and 4 unit volumes of the solvent.

Thus, when preparing a solution starting with a solution or liquid reagent: When diluting more concentrated solutions,
• Decide what volume (Vf) and molarity (Mf) the final solution should
be.Volume can be expressed in liters or milliliters.
• Determine molarity (Mi) of starting, more concentrated solution.
• Calculate volume of starting solution (Vi) required using equation:
MiVi= MfVf

Where Mi and Mf are the initial and the final concentrations (or initial
and final molar concentrations or initial and final molarities); Vi and Vf
 are the initial and the final volumes.

Note: Vi must be in the same units as Vf.

The following are the requirements: Graduated pipettes and /or micropipettes, volumetric flasks, Stoppers, Dropping pipettes, Stock solutions, Droppers, Distilled water.

The preparation of a solution by diluting stock solutions follows the
steps illustrated by figure 4.2

Examples:
1. Describe how you can prepare 1L of a 1M H2
SO4 solution from the stock solution with the following specification: density (d) =1.84; percentage (P) by mass: 98%; Mw = 98 g/mol.

Answer:
Preliminary calculations
The molarity of the stock solution may be calculated using the following
relation.

Hence the required volume of concentrated sulphuric acid to be pipetted is 54.35mL.

Procedure:
1) Using a graduated pipette or a burette, measure 54.35mL of
concentrated H2SO4;
2) Using a funnel, carefully transfer the concentrated acid into a 1L
volumetric flask containing about 700mL of distilled water;
3) Put a stopper and gently shake the mixture;
4) Using a wash bottle, add more distilled water until the level is 1cm
below the mark;
5) Using a dropper pipette, make up to 1L dropwise;
6) Stopper the volumetric flask and gently shake.
7) Label the solution and specify the preparation date.

The solution prepared is 1M H2SO4.
2. Calculate the volume of 15M H2SO4 that would be required to prepare
150cm3 of 2MH2SO4.

Application activity 4.3.2
Describe the preparation of 500mL 0.5M HCl solution from a stock
solution of HCl with the following specifications.
• Molar weight: 36.5g/mol
• Percentage of HCl: 36% by mass
• Density: 1.18g/mL

4.4. Titration process focussing on precise measurements
4.4.1. Simple acid-base titrations
Activity 4.4.1
1. Analyse the diagram below and answer the questions that follow

b) Is it important in real life to perform operations illustrated by the
above set up? Explain using at least two examples.
c) What are the necessary requirements for such operations?
2. You are provided with:
S1: solution of HCl (aq) 0.1 M
S2: solution of NaOH (aq)
Materials: Burettes - Indicator (phenolphthalein) - pipettes - Washing
bottles - Conical flasks - Beakers - Retort stands - Funnels

Procedure
a). Using a pipette, transfer 10 mL of S2 into a conical flask.
b). Add three drops of phenolphthalein indicator and titrate it with S1
from the burette.
c). Repeat the titration until you obtain consistent values.
d). Record your readings in the table below:

Titration is a process used in volumetric analysis, in which a solution of one reactant of known concentration, of unknown concentration, the titrant, is added to a known volume of another reactant of unknown concentration, the titrand, until the reaction between the two is complete.
The analyte (titrand) is the specie of interest in a titration i.e, the specie that is subjected to analysis.

If a concentration and volume of the titrant is reacted with the analyte, it is possible to calculate the analyte’s concentration.

The completion of the reaction is signalled by an indicator, i.e a substance that has distinctly different colours in acidic and basic media. E.g phenolphthalein.

The point at which the indicator changes the colour and the titration is
stopped is called end point.

The equivalence point is the point, in a titration, at which the mole ratio of reactants is exactly equal to the mole ratio in the balanced chemical equation. The difference between the end point and equivalent point is called titration error.

Because volume measurements play a key role in titration, it is also known as volumetric analysis. Usually it is the volume of the titrant required to react with a given quantity of an analyte that is precisely determined during a titration.

Titrations can be classified by the type of reaction. Different types of titration reaction include acid-base titrations, redox titrations, complexometric titrations, etc. In our context we will deal with acid-base titrations.

Acid-base titrations
In an acid-base titration, a standard solution of a base is added to a known volume of an acid with unknown concentration or vice versa, until the end point is reached.

At the equivalence point, the acid and the base that have reacted are in
equivalent amounts.

Therefore, naMbVb = nbMaVa,
Where Ma=Molarity of acid; Mb =Molarity of the base
Vb =Volume of base; nb = stoechiometric number of base
Va =Volume of acid; na = stoechiometric number of acid
Indicators and pH-meters can be used to determine the equivalence point.
An acid-base titration needs an indicator which will change colour with the concentrations of hydrogen ions (H+).

Choice of indicators in acid-base titrations
When the technique of acid-base titration is extended to a wide variety of
acidic and alkaline solution, care needs to be taken about the choice of
indicator for any given reaction.

The choice of an inappropriate indicator would lead to incorrect results, and it is therefore extremely important that the indicator is chosen carefully.
The principle on which a choice of indicator is made concerns the strength of the acid or base involved in the reaction. Note that the strength of an acid or base is not to be confused with the concentration of its solution. Example of strong and weak acids and bases and choice of indicator are given in the Table below.

The common equipments used in titration are:
Burette
Pipette
Retort stand and clamp
Dropper pipette
pH-indicator/acid-base indicator
White tile: used to see a colour change in the solution(a white paper can
also be used)
Conical flask (Erlenmeyer flask) or a beaker

How to perform titrations?
It is essential to be familiar with the use of pipette and burettes and how to handle them. The following points are useful in order to correctly perform a titration.
1. The apparatus should be arranged as shown in the activity 4.4.1
2. The burette tap is opened with the left hand and the right hand is
used to shake the conical flask.
3. The equivalence-point is reached when the indicator just changes
permanently the colour.
4. At the end-point, the level of the titrant is read on the burette.
5. The titration is now repeated, three more times are recommended.
Towards the end-point, the titrant is added dropwise to avoid
overshooting.

Notice: Before titration, check if the tip of the burette is filled with the titrant, and doesn’t contain bulb of air. If there is a bulb of air, a quick opening and closing of the tap will expel the air out of the burette.

The results are summarised in a table as shown below:
Burette readings should be written to two decimal places (for burette having precision up to hundredth)

Procedure
• Weigh accurately 1.0 g of sodium hydroxide pellets in a weighing
boat and transfer into a 250cm3 volumetric flask.
• Add a little distilled water and shake to dissolve.
• Make up the solution to the mark with distilled water dropwise.
• Label the solution FA2.
• Pipette 20cm3 (or 25 cm3) of FA2 into a conical flask.
• Add 2-3 drops of phenolphthalein indicator.
• Titrate with FA1 from the burette.
• Record your results in the spaces provided below.

Specimen results
Mass of weighing container + NaOH pellets = 11.30 g
Mass of weighing container alone = 10.30 g
Mass of NaOH pellets = 11.30 g – 10.30 g = 1.00 g
Masses should be recorded to at least two decimal places (Table 11.4)
Volume of pipette used = …………..

Each value or entry in the table must be recorded or written to two decimal places.

Different initial readings should be used. Initial reading in each experiment should be correctly subtracted from the final reading.

Questions:
a). Determine the values used to calculate average volume, and
calculate the average volume of FA1 used.
b). Write the equation of the reaction between NaOH and HCl
c). Calculate the number of moles of FA1 used.
d). Calculate the concentration of FA2 in mol.dm-3.
e). Calculate the concentration of FA2 in g/L.

What compound is represented by the picture above?
What observable changes could be recorded if the substance was heated
and how do you interpret those observations?
When blue crystals of copper(II) is heated it becomes a white powder. This is due to the loss of water molecules attached to it. This water is known as water of crystallisation.

Water of crystallisation is water present in definite proportions in some
crystalline compounds.

Many crystalline salts form hydrates containing 1, 2, 3, or more moles of
water per mole of the compound and the water may be held in the crystal

Skills lab 4
You are asked to investigate the acidity content of aspirin tablets.
Write a procedure including materials, set up and procedure you will use for that purpose.
Your project will be performed either during evenings weekends. You will
request from the school authorities to avail some materials for you.

End unit assessment 4
1. Explain the following terms: a. Standard solution b. Primary standard.
2. Describe the characteristics of a primary standard.
3. A solution is made by dissolving 5.00g of impure sodium hydroxide in
water and making it up to 1dm3 of solution.25cm3 of this solution are neutralized by 30.0cm3 of hydrochloric acid of concentration 0.102M.
Calculate the percentage purity of the sodium hydroxide.
4. You are provided with a white powder containing a mixture of sodium
carbonate and magnesium nitrate and 1M nitric acid. Make a solution
using 10g of the powder and 47.3cm3 of water. You need 0.05dm3
of the acid to fully react with the solution. What is the percentage
composition of magnesium nitrate in the powder?