Mr McCondichie

Researching Chemistry

Experiment 1: Volumetric Analysis

U2PPA2: Gravimetric Analysis

Experiment 2A: Preparation of a stand solution of 0.1 mol l^-1 sodium carbonate solution

Experiment 2B: Standardisation of approximately 0.1 mol l^-1 hydrochloric acid

Experiment 2C: Determination of purity of marble by back titration

Experiment 3: Determination of nickel in a nickel(II) salt using EDTA

Experiment 8: Preparation of benzoic acid by hydrolysis of ethyl benzoate

Experiment 10: Colorimetric determination of manganese in steel

Experiment 11: Preparation of cyclohexene from cyclohexanol

Volumetric Anaylsis

Experiment 1A. Preparation of a standard solution of 0.1 mol l^-1 oxalic acid

Introduction

A standard solution is one of accurately known concentration and can be prepared directly from a primary standard which, in this case, is hydrated oxalic acid, (COOH)₂.2H₂O (RFM = 126.1).

To prepare 250 cm³ of 0.1 mol l^-1 oxalic acid solution, the mass of hydrated oxalic acid required can be calculated as 0.1 x 0.250 x 126.1 = 3.15 g.

Requirements

• balance (accurate to 0.01g)
• weighing bottle
• 250cm³ beaker
• 250cm³ standard flask
• wash bottle
• dropper
• glass stirring rod
• filter funnel
• oxalic acid AnalaR, (COOH)₂.2H₂O
• deionised water

Hazcon

Wear eye protection and if any chemical splashes on the skin, wash it off immediately. Oxalic acid is harmful if ingested and irritates the eyes and skin. Wear gloves.

Procedure

1. Transfer approximately 3.2 g of oxalic acid crystals to the weighing bottle and weigh accurately.
2. Pour the oxalic acid crystals into a clean beaker containing about 50 cm³ of deionised water and reweigh accurately the weighing bottle and any remaining crystals. Stir the solution until all the oxalic acid dissolves and then transfer it to a 250 cm³ standard flask.
3. Rinse the beaker several times with deionised water and add all the rinsings to the flask.
4. Make up the solution to the graduation mark with deionised water.
5. Stopper the flask and invert it several times to ensure the contents are completely mixed.
6. Calculate the concentration of the oxalic acid solution using the exact mass of the oxalic acid transferred to the beaker in step 2.

Experiment 1B. Standardisation of approximately 0.1 mol l^-1 sodium hydroxide

Introductions

Sodium hydroxide is not a primary standard and so a standard solution of it cannot be prepared directly from the solid. However, a solution of approximate concentration can be prepared and its exact concentration determined by titrating it against an acid of accurately known concentration using a suitable indicator. In this experiment, a sodium hydroxide solution is standardised against the 0.1 mol l^-1 oxalic acid solution prepared in Experiment 1A. The stoichiometric equation for the titration reaction is:

(COOH)₂ + 2NaOH → 2H₂O + (COONa)₂

Requirements

• 50cm³ burette
• 10cm³ pipette
• 100cm³ beakers
• 100cm³ conical flasks
• wash bottle
• pipette filler
• white tile
• filter funnel
• standardised oxalic acid solution (approx. 0.1 mol l^-1)
• sodium hydroxide solution (approx. 0.1 mol l^-1)
• phenolphthalein indicator
• deionised water

Hazcon

• Wear eye protection and if any chemical splashes on the skin, wash it off immediately.
• 0.1 mol l^-1 oxalic acid irritates the eyes and skin
• 0.1 mol l^-1 sodium hydroxide is corrosive to the eyes and skin.
• Phenolphthalein indicator solution is highly flammable and irritating to the eyes because of its ethanol content.

Procedure

1. Rinse the 10 cm³ pipette with a little of the oxalic acid solution and pipette 10 cm³ of it into a conical flask.
2. Add two or three drops of phenolphthalein indicator to the oxalic acid solution in the flask.
3. Rinse the 50 cm³ burette, including the tip, with the sodium hydroxide solution and fill it with the same solution. Titrate the oxalic acid solution with the sodium hydroxide solution from the burette until the end-point is reached. This is indicated by the appearance of a pink colour.
4. Repeat the titrations until two concordant results are obtained.
5. Calculate the concentration of the sodium hydroxide solution.

Experiment 1c. Determination of the ethanoic acid content of white vinegar

Introduction

Vinegar is a dilute solution of ethanoic acid and the aim of this experiment is to determine the concentration of ethanoic acid in a given sample of white vinegar by titration against the sodium hydroxide solution standardised in Experiment 1B. The stoichiometric equation for the titration reaction is:

CH₃COOH + NaOH → H₂O + CH₃COONa

Requirements

• 50cm³ burette
• 25cm³ pipette
• 100cm³ beakers
• 100cm³ conical flasks
• 250cm³ standard flask
• wash bottle
• pipette filler
• dropper
• white tile
• filter funnel

Hazcon

• Wear eye protection and if any chemical splashes on the skin, wash it off immediately.
• Vinegar irritates the eyes and skin.
• 0.1 mol l^-1 sodium hydroxide is corrosive to the eyes and skin.
• Phenolphthalein indicator solution is highly flammable and irritating to the eyes because of its ethanol content.

Procedure

1. Rinse the 25 cm³ pipette with a little of the vinegar.
2. Dilute the sample of vinegar by pipetting 25 cm³ of it into a clean 250 cm³ standard flask and making it up to the graduation mark with deionised water.
3. Stopper the standard flask and invert it several times to ensure the contents are thoroughly mixed.
4. Rinse the 25 cm³ pipette with a little of the diluted vinegar and pipette 25 cm³ of it into a conical flask. Add two or three drops of phenolphthalein indicator to the diluted vinegar in the conical flask.
5. Rinse the 50 cm³ burette, including the tip, with the sodium hydroxide solution and fill it with the same solution.
6. Titrate the diluted vinegar solution with the sodium hydroxide solution from the burette until the end-point is reached. This is indicated by the appearance of a pink colour.
7. Repeat the titrations until two concordant results are obtained.
8. Calculate the concentration of the ethanoic acid in the diluted vinegar and hence in the undiluted vinegar.

U2PPA2: Gravimetric Analysis

Gravimetric and Volumetric Quiz

Experiment 10: Colorimetric determination of manganese in steel

Introduction

Colorimetry is an analytical technique used to determine the concentrations of coloured substances in solution. It relies on the fact that a coloured substance absorbs light of a colour complementary to its own and the amount of light it absorbs (absorbance) is proportional to its concentration.

Colorimetry is particularly suited to the determination of manganese in steel because the manganese can be converted into permanganate ions, which are coloured. The conversion is achieved in two stages. Using nitric acid, the managanese is first oxidised to manganese(II) ions, which are then oxidised to permanganate ions by the more powerful oxidising agent, potassium periodate.

Requirements

Equipment	Chemicals
standard flasks (50 cm3 and 100 cm3)	steel paper clips
50 cm3 burette	acidified potassium permanganate
green filter	2 mol l-1 nitric acid
optically matched cuvettes	85% phosphoric acid
balance (accurate to 0.001 g)	acidified potassium periodate
glass beakers (50 cm3 and 250 cm3)	5 g potassium periodate per 100 cm3 of 2 mol l-1 nitric acid
Bunsen burner, heating mat and tripod	potassium persulfate
clock glass	propanone
filter funnel	deionised water
tweezers	anti-bumping granules
wash bottle
dropper
wire cutters
standardised 0.0010 mol l-1 colorimeter
measuring cylinders (50 cm3 and 10 cm3)

Hazcon

Wear eye protection and if any chemical splashes on your skin wash it off immediately.

The acidified 0.0010 mol l^-1 potassium permanganate is harmful if ingested and irritates the eyes and skin. Wear gloves.

Both 2 mol l^-1 nitric acid and its vapour are corrosive and toxic, causing severe burns to the eyes, digestive and respiratory systems. Wear gloves.

85% phosphoric acid is corrosive: it burns and irritates the eyes and skin. It is a systemic irritant if inhaled and if swallowed causes serious internal injury. Wear gloves.

Acidified potassium periodate solution is harmful if swallowed and is an irritant to the eyes, skin and respiratory system. It is also corrosive. Wear gloves.

Potassium persulfate is harmful if swallowed or inhaled as a dust. It irritates the eyes, skin and respiratory system, causing dermatitis and possible allergic reactions. Wear gloves.

Propanone is volatile and highly flammable, and is harmful if swallowed. The vapour irritates the eyes, skin and lungs, and is narcotic in high concentrations. Wear gloves.

Procedure

Part A - Calibration graph

1. Rinse the burette, including the tip, with 0.0010 mol l^-1 acidified potassium permanganate and fill it with the same solution.
2. Run 2 cm3 of the permanganate solution into a 50 cm³ standard flask and make up to the graduation mark with deionised water.
3. Stopper the flask and invert it several times to ensure the contents are completely mixed.
4. Rinse a cuvette with some of the solution and fill it.
5. Using a colorimeter (fitted with a green filter) measure the absorbance of the solution in the cuvette. If you have more than one green filter, choose the one that gives maximum absorbance.
6. Repeat steps 2 to 5 with 4, 6, 8, 10, 12 and 14 cm3 of the permanganate stock solution in the burette.
7. Plot a calibration graph of 'absorbance' against 'concentration of potassium permanganate'. Your practitioner will provide you with the accurate concentration of the acidified potassium permanganate stock solution.

Part B - Conversion of manganese to permanganate

1. Degrease a steel paper clip by swirling it with a little propanone in a beaker. Using tweezers remove the paper clip and leave it to dry for a minute or so on a paper towel.
2. Cut the paper clip into small pieces.
3. Weigh accurately about 0.2 g of the paper clip pieces and transfer them to a 250 ³ glass beaker.
4. Add approximately 40 cm³ of 2 mol l^-1 nitric acid to the beaker and cover it with a clock glass. Heat the mixture cautiously, in a fume cupboard, until the reaction starts. Continue heating gently to maintain the reaction, but remove the source of heat if the reaction becomes too vigorous.
5. Once the steel has reacted, allow the solution to cool a little. Add a couple of anti-bumping granules and then boil the solution until no more brown fumes are given off.
6. Once this solution has cooled considerably - no more than 'hand hot' - add about 5 cm³ of 85% phosphoric acid, approximately 0.2 g of potassium persulfate and a couple of fresh anti-bumping granules. Boil the mixture for about 5 minutes.
7. To this solution, add approximately 15 cm³ of acidified potassium periodate solution plus a couple of fresh anti-bumping granules and then gently boil the mixture. The solution will start to turn pink. Continue gently boiling until the intensity of the pink colour remains constant. This should take about 5 minutes.
8. Allow the pink solution to cool to room temperature and then transfer it to a 100 cm³ standard flask, leaving the anti-bumping granules in the beaker.
9. Rinse the beaker several times with a little deionised water and add the rinsings (but not the anti-bumping granules) to the flask.
10. Make up the solution to the graduation mark with deionised water.
11. Stopper the flask and invert it several times to ensure the contents are completely mixed.
12. Using a colorimeter fitted with the appropriate green filter, measure the absorbance of the solution.
13. Use your calibration graph to convert the absorbance to a permanganate concentration and then calculate the percentage by mass of manganese in the steel paper clip.