CAPE CHEMISTRY



UNIT 2

CHEMISTRY PRACTICALS

2010 - 2011

Table of Contents

|Page |Title |Module |Skills tested |

|# | | | |

|3 |Standardisation of sodium hydroxide |2 | |

|5 |Calibration of a pipette |2 |ORR, A&I |

|6 |Redox Titration / H2O2 vs KMnO4 |2 |M&M |

|8 |Back Titration |2 | |

|10 |Gravimetric analysis of hydrated salt |2 | |

|11 |Flame tests |2 | |

|12 |Reactions of ethanol |1 |ORR,A&I |

|13 |Paper Chromatography |2 | |

|14 |Theoretical: Functional groups in organic compounds |1 | |

|15 |Water polluting inorganic ions |3 | |

|16 |Distillation of a sample of rum |3 | |

|17 |Salted peanuts “Fuller” |2 |P&D |

|18 |Identity of sample X |3 |M&M, A&I |

|19 |Comparing simple and fractional distillation |2 | |

|20 |Margarine vs Butter |1 |P&D |

MODULE 2

TITLE: STANDARDISATION OF HYDROCHLORIC ACID

SKILLS TESTED: None

AIM: The purpose of this experiment is to determine the concentration of a solution of hydrochloric acid using a standard solution of sodium carbonate.

Introduction: The salt reacts with hydrochloric acid in a 1: 2 molar ratio (1 mole sodium carbonate : 2 moles acid). In this experiment you will prepare a standard solution of sodium carbonate and titrate this solution against a solution of hydrochloric acid of unknown concentration using methyl orange as indicator.

Procedure A: Preparation of the standard solution.

1. Weigh 2.0 g of sodium carbonate to the nearest 0.01g. Dissolve the solid in the minimum amount of water.

2. Record your readings in the table below.

3. Transfer the solid and washings to a volumetric flask. Your teacher will guide you.

4. Add distilled water, swirling at intervals to mix the contents, until the level is within 1 cm of the calibration mark on the neck of the flask.

5. Using a wash bottle, carefully add distilled water to bring the bottom of the meniscus to the mark.

6. Insert the stopper and shake thoroughly (by inversion) to ensure complete mixing.

Procedure B:

1. Rinse and fill the burette with hydrochloric acid solution. Record the initial burette reading.

2. Using a pipette filler, rinse the pipette with some of the sodium carbonate solution and carefully transfer 25cm3 of the solution to a clean conical flask.

3. Add 2 drops of the methyl orange indicator and swirl the flask.

4. Run hydrochloric acid from the burette into the flask, with swirling until an orange colour (with a hint of pink) permanent colouration is seen. Record the burette reading.

5. Repeat the titrations as many times as is necessary to obtain consistent results.

6. Do NOT write a full lab report. Complete the table and write your answers on the dotted lines on THIS paper.

Results and calculations:

|Molar mass of sodium carbonate (g) | |

|Mass of weighing bottle + sodium carbonate (g) | |

|Mass of weighing bottle (g) | |

|Mass of sodium carbonate (g) | |

|Molar concentration of sodium carbonate (mol dm-3) | |

|Burette |Trial |1 |2 |3 |4 | |

|Reading/cm3 | | | | | | |

|(to 2 decimal places) | | | | | | |

|Final | | | | | | |

|Initial | | | | | | |

|Volume used/cm3 | | | | | | |

|average titre/cm3 |(please tick the values used to determine your average) |

Treatment of results:

1. Write a balanced equation for the reaction between sodium carbonate and hydrochloric acid.

………………………………………………………………………………………………………………………

2. Calculate the number of moles of sodium carbonate in 25 cm 3 of the standard solution.

……………………………………………………………………………………………………………………..

3. Using the molar ratio of acid to alkali, calculate the number of moles of hydrochloric acid in the mean volume of acid used.

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

4. Hence, calculate the molar concentration of the hydrochloric acid solution

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

5. Calculate the mass concentration of the hydrochloric acid solution.

………………………………………………………………………………………………………………………

MODULE 2

TITLE: CALIBRATION OF A PIPETTE

Skills tested: ORR, A&I

INTRODUCTION: In this experiment, you will accurately determine the volume of a 10 cm3 pipette which is advertised as being 10cm3+/- 0.06 cm3 . This will be achieved by measuring the mass of water transferred by the pipette, followed by calculating the volume of that mass of water using the known density of water. Without calibration, a systematic error of unknown magnitude may exist.

PROCEDURE

1. Collect distilled water in a beaker and let it stand for about 15 minutes before determining the temperature of the water.

2. Weigh the dry empty beaker on the balance.

3. Using the pipette filler, fill and discharge the pipette into the pre-weighed beaker.

4. Determine the mass of water discharged.

5. Repeat steps 3 – 4 until you have results for 5 trials. N.B. You do not need to empty and dry the beaker between trials.

6. Determine the temperature of the water you have pipetted and take the mean of the two temperatures you have measured. This value will be used as the temperature of the water during the calibration.

7. Write a full lab report.

Calculations

1. Calculate the mean of the five trials.

2. Use the following table to determine the density of water at the temperature of your determination by drawing a calibration curve:

|X |T/oC |26 |28 |30 |32 |34 |36 |

|Y |Density /gcm-3 |0.99678 |0.99623 |0.99565 |0.99503 |0.99438 |0.99369 |

3. Using your answer from question 2 determine the mean volume of water discharged from the pipette.

4. Locate a group in the lab that has a pipette with a volume fairly close to yours. Collect the data from that group and determine the standard deviation of the ten data points.

5. Report the pipette volume and the error

6. Comment on the accuracy and precision of the pipette.

7. Discuss any sources of error in your experiment.

MODULE 2

TITLE: REDOX TITRATION

SKILLS TESTED: M&M, A&I

AIM: To determine the concentration of a solution of hydrogen peroxide

Introduction

Your task is to use 0.020 moldm-3 potassium manganate (VII) solution to find the actual concentration of a solution of hydrogen peroxide, H2O2, which is believed to have partially decomposed.

Procedure

a) Prepare 250 cm3 of a solution of hydrogen peroxide by adding 7.5 cm3 of the stock solution supplied to the volumetric flask and diluting to the required volume.

b) Pipette 25.0 cm3 of this solution into a conical flask.

c) Add 25 cm3 of 1 mol dm-3 sulphuric acid in a measuring cylinder.

d) Titrate the mixture against the potassium manganate (VII) until a permanent pale pink colour

appears in the conical flask.

e) Record the titre volume and repeat until concordant values are obtained.

f) Do not write a full lab report, write your results in the table below and answer the questions that follow.

Results

|Burette |Trial |1 |2 |3 |4 | |

|Reading/cm3 | | | | | | |

|(to 2 decimal places) | | | | | | |

|Final | | | | | | |

|Initial | | | | | | |

|Volume used/cm3 | | | | | | |

|average titre/cm3 |(please tick the values used to determine your average) |

Questions

1. Write the relevant half equations for

i) the reduction of MnO4- ions to Mn2+ in acidic solution

ii) the oxidation of H2O2 to O2 in acidic solution

2. Using the two half equations from question 1, write the full balanced ionic equation.

3. If the stock solution of hydrogen peroxide had a molar concentration of 1.67 mol dm-3, determine the molar concentration of the hydrogen peroxide solution prepared by you.

4. Determine the molar concentration of the hydrogen peroxide solution prepared by you using your titration results.

5. Determine the error in your results.

6. Suggest a reason for the partial decomposition of the hydrogen peroxide.

MODULE 2

TITLE: BACK TITRATION

SKILLS TESTED: None

AIM: To determine the percentages of calcium carbonate and calcium chloride in a mixture of the two.

THEORY:

Sometimes a direct titration would involve a reaction which is too slow and thus an incomplete reaction would occur, therefore a back titration is more suitable. In a back titration, a known excess of the reagent is used to ensure complete reaction and then a second reagent is titrated against the mixture which reacts with the remaining reagent which allows via calculation, the amount of the reagent that was used in the initial reaction.

PROCEDURE:

1. Weigh 1.5 g of the solid mixture of calcium carbonate and calcium chloride and place in a conical flask.

2. Measure 25 cm3 of the 1 mol dm-3 HCl solution and place in the conical flask.

3. Rinse and fill the burette with 0.2 mol dm-3 NaOH solution.

4. Once all effervescence in the conical flask has ceased, add 2 drops of phenolphthalein indicator to the flask and swirl the flask.

5. Titrate the sodium hydroxide against the reaction mixture until the first permanent pale pink colour is seen.

6. Repeat steps 1-5 until consistent results are obtained. (In the interest of time and materials, please try to be as accurate as possible to minimise the number of trials)

7. Do not write a full lab report. Record your readings in the table below and answer the questions that follow.

RESULTS

|Burette |Trial |1 |2 |3 |

|Reading/cm3 | | | | |

|(to 2 decimal places) | | | | |

|Final | | | | |

|Initial | | | | |

|Volume used/cm3 | | | | |

|Average volume |(please tick the values used to determine your average) |

| | |

TREATMENT OF RESULTS

1. Calculate the # of mol of HCl placed in the conical flask.

2. Using the average volume of NaOH used, calculate the # of mol of NaOH used in the titration.

3. Write the balanced equation for the reaction between HCl and NaOH and thus determine the # of mol of HCl remaining after the initial reaction in the conical flask was complete.

4. Write a balanced chemical equation for the reaction between calcium carbonate and HCl.

5. Using your answers from question 1 and question 3, determine the # of mol of HCl that reacted with the calcium carbonate.

6. Determine

a) the # of mol of CaCO3

b) the mass of CaCO3 present in the mixture

7. Hence calculate the percentages of CaCO3 and CaCl2 present in the solid mixture.

MODULE 2

TITLE: GRAVIMETRIC ANALYSIS

SKILLS TESTED: None

AIM: To determine the water of crystallisation in hydrated magnesium sulphate using a volatilisation method.

INTRODUCTION:

Water of crystallisation forms an integral part of the crystalline structure of a stable ionic solid. This water is considered to be one type of essential water and is distinct from water released when compounds are decomposed by heat. In this experiment, a sample of hydrated magnesium sulphate will be converted to the anhydrous salt and the decrease in mass will be used to determine the value of x in the formula MgSO4. xH2O.

PROCEDURE:

1. Weigh an empty crucible and record its mass in a table.

2. Add between 5.00 and 6.00g of hydrated magnesium sulphate to the crucible and record the mass of crucible and solid.

3. Heat the crucible containing the hydrated salt over a Bunsen burner for approximately 10 minutes.

4. Place the crucible and its contents on the heat proof mat and allow it to cool to room temperature.

5. Weigh the crucible and its contents.

6. Reheat the crucible and its contents for 5 minutes and then repeat steps 4 and 5.

7. Write a full lab report.

TREATMENT OF RESULTS

1. Write a chemical equation for the loss of water of crystallisation from one mole of hydrated magnesium sulphate represented by the formula: MgSO4.xH2O

2. Calculate the mass of one mole anhydrous magnesium sulphate.

3. Calculate the number of moles of anhydrous magnesium sulphate formed in the experiment.

4. Calculate: (a) the mass (b) the number of moles of water released from the sample of hydrated salt used in the experiment.

5. Use your answers to 3 and 4(b) to calculate the number of moles of water in one mole of the hydrated salt.

QUESTIONS;

1. Identify two possible sources of error in the experiment.

2. Traditional methods for determining moisture content of solids were time consuming and involved heating in conventional ovens or in vacuum ovens or storing the sample in a dessicator until the material became constant in weight. Suggest ONE method by which the determination could be speeded up.

MODULE 2

TITLE: FLAME TESTS

INTRODUCTION:

Many cations of s and d-block elements may be identified by a flame test. The energy from the burning gas causes promotion of electrons in the ions to higher energy levels. The excited ions then lose energy and undergo electronic transition to a lower energy level, giving off the extra energy in the form of electromagnetic radiation. The frequency of this radiation usually falls in the visible region of the spectrum, so a coloured flame is observed.

PROCEDURE:

1. Clean a nichrome or platinum wire by dipping it in concentrated hydrochloric acid and placing it in a non-luminous Bunsen flame.

2. Continue this cleaning process until no colour at all is produced when the wire is in the flame.

3. Moisten the wire with concentrated hydrochloric acid, dip it in the sample, and hold it in the flame again.

4. Record the colour observed.

5. Carry out a flame test on each of the other samples supplied.

6. Record your results in the table below.

7. Do not write a full lab report.

Results

|Metal ion | | | |

|A |Bromine in 1,1,1-trichloroethane |Decolourisation | |

| | | | |

| | | | |

| | | | |

|B | | | |

| |PCl5 |Dense white fumes | |

| | | | |

| | |Yellow ppt | |

| |I2 in NaOH | | |

|C |Litmus test |Blue to red | |

| | |Ense white fumes | |

| |PCl5 | | |

|D |Litmus test |Blue to red | |

| | | | |

| | | | |

| |Br2(aq) |White ppt | |

| | | | |

QUESTION:

Compound B is a fragrant liquid which when heated with excess concentrated H2SO4 produces a gas, G, which decolourises aqueous bromine. When G reacts with HBr, two isomeric compounds, E and F, with relative molecular mass of 123 are obtained. Write the displayed formula for EACH of the following.

MODULE 3

TITLE: WATER POLLUTING INORGANIC IONS

INTRODUCTION:

Nitrate (V), Phosphate (V), cyanide, and Lead (II) ions may be found in contaminated water. The typical reactions of the NO3 - , PO4 3- , and Pb2+ ions will be carried out in this practical. CN- ions are highly poisonous and will be considered theoretically.

PROCEDURE:

Carry out the following tests. Record all observations and inferences in Tabular form.

1. Nitrate (V) anions:

a) Add Devarda’s alloy followed by sodium hydroxide. Warm mixture and test gas evolved.

b) Mix with iron (II) sulphate crystals, Pour concentrated sulphuric acid slowly down the side of the tube, so that two layers of liquid are formed.

c) Add some copper turnings then add concentrated sulphuric acid. [ TEACHER DEMONSTRATION]

2.Phosphate (V) anions

a) Add aqueous silver nitrate. Test solubility of ppt separately in (i) dil. HNO3 (aq) (ii)NH3 (aq) and discard product immediately!

b) Add aqueous barium chloride. Divide ppt into three portions and test its solubility in (i) dil nitric acid, (ii) dil ethanoic acid, (iii) dil aqueous ammonia.

c) Add an equal volume of ammonium molybdate. Warm gently to accelerate formation of ppt.

d) Add aqueous iron (III) chloride. Test solubility of ppt in (i) dilute nitic acid, (ii) dilute ethanoic acid.

3. Lead (II) cations

a) Add aqueous sodium hydroxide until in excess

b) Add dilute hydrochloric acid. Heat mixture and allow to cool.

c) Add potassium chromate (VI). Divide mixture and add(i) dil nitric acid,(ii) dil ethanoic acid

d) Add aqueous potassium iodide until in excess.

e) Add aqueous sodium sulphide

.

QUESTIONS

(a) A sensitive test for cyanide ions ( concentration limit 1 in 50,000) involves their reaction with polysulphide (S2 2- ) ions to form thiocyanate ions to which a solution of iron (III) chloride is then added. Suggest the colour and formula of the final product.

(b) Cyanide ions react with dil HCl to liberate a poisonous gas which has the smell of bitter almonds. Identify the gas and explain the reaction.

Do not write a full lab report. Tabular form of the result is suggested.

MODULE 3

TITLE: DISTILLATION OF A SAMPLE OF RUM

AIM: To compare the efficiency of separating ethanol from a sample of Unknown Barbados Rum by simple distillation and fractional distillation.

INTRODUCTION: Methods of separation are never 100% efficient. However some methods are more efficient than others. The separation technique which is more suitable is dependent on the type of mixture being separated.

PROCEDURE:

Observe the demonstration for the simple distillation and the fractional distillation of the sample. Please note the distillation rate and if you are careful waft the vapours from each type of distillate for your own determination of alcohol present.

RESULTS

Table of comparison of simple and fractional distillation of a sample of Unknown Barbados Rum

| |Temperature of the lower boiling|Volume of distillation after |Smell of distillate (state |

| |point fraction |demonstration was completed |whether there was a strong or |

| | | |weak smell of alcohol) |

|Simple distillation | | | |

|Fractional distillation | | | |

QUESTIONS:

1. Explain the process of simple distillation and fractional distillation. What areas in the demonstration that could have been improved to obtain a purer distillate?

2. Even industrially not all the water can be removed from an alcohol – water mixture via fractional distillation. Suggest what can be done to remove the last traces of moisture from a distillate obtained from an alcohol-water mixture.

3. Why would there be a difference in the volumes of distillate collected?

4. Based on the strong or weak smell of alcohol of the distillate, which seems to be the more efficient separation technique?

MODULE 2

TITLE: SALTED PEANUTS “FULLER”

Skills tested: P&D

PROBLEM:

As a young boy or girl, you have always enjoyed eating salted peanuts. As you got older and being a science student, you then began to understand the dangers of eating too much salt. In the past, salted peanuts “Fuller” had quoted on its nutritional label that it contained 180 mg of salt per 30 g of salted peanuts. As customers become more health conscious and buying what they deem to be healthier items, manufacturers have no choice but to comply. Salted peanuts “Fuller” are now quoted to have a reduced salt content per 30 g of salted peanuts. You wonder if this is really true or just a marketing gimmick. Plan and design an experiment to determine the amount of salt in the past salted peanuts “Fuller” and present day reduced salt content salted peanuts “Fuller”.

The format of your lab should be as follows

Title

Date

Aim

Hypothesis

Apparatus & Materials

Variables

Procedure

Data Collection (shown in an appropriate format)

Expected results

Discussion (which consists of limitations, precautions and sources of error)

MODULE 3

TITLE: Qualitative analysis of sample X

SKILLS TESTED: M&M, A&I

PROCEDURE:

Carry out the following tests on the sample X. Carefully record all observable changes. Complete the table below. Do not write a full lab report.

|Tests |Observations |Inferences |

|1. Add a heaped spatula load of the solid X to 10 cm3 | | |

|of distilled water in a boiling tube and stir | |………………………………….. |

|thoroughly. Keep the contents of the tube for the | | |

|following tests. | |…………………………………. |

|2. To approximately 1 cm3 of solution, add NaOH(aq) | |Ion(s) possibly present |

|slowly until in excess. | | |

| | |………………………………… |

|3. To approximately 1 cm3 of solution, add NH3 (aq) | |Ion(s) possibly present |

|until in excess. | | |

| | |………………………………… |

|4. To approximately 1 cm3 of solution, add dilute | | |

|hydrochloric acid. | |Ion present ..…………. |

| | | |

| | |Write an ionic equation for the reaction |

| | | |

| | |………………………………… |

|5. Add a spatula load of Devarda’s alloy, then add | |Acid-base nature of the gas evolved |

|NaOH (aq) and warm the mixture. Test gases evolved | | |

|with moist red and blue litmus. | |………………………… |

| | | |

| | |Formula of gas evolved …………. |

| | | |

| | |Ion(s) possibly present |

| | | |

| | |………………………………. |

|6. Add solid or aqueous FeSO4, then pour conc. H2SO4 | |Formula of compound formed |

|slowly down side of tube. | | |

| | |…………………………… |

| | | |

| | |Ion present ……………… |

TREATMENT OF RESULTS

1. What is the likely identity of sample X? [1]

2. Write a balanced equation for the action of strong heat on the solid sample.[2]

3. Suggest two reasons why the solid sample should not be strongly heated in the school laboratory.[2]

4. Comment on the extent to which this sample can contaminate a potable water supply.[2]

MODULE 2

TITLE: Comparing the ethanol content of distillates.

AIM : To determine which of two distillates contains the higher percentage of ethanol.

INTRODUCTION:

The ethanol content of a distillate may be determined by the following procedure:

1. A given volume of the distillate is treated with an excess of potassium dichromate of known concentration by heating the mixture in a water bath for at least 2 hours or in an oven overnight.

2. The sample is then left to cool to room temperature.

3. Any excess dichromate is estimated by adding an excess of potassium iodide to the mixture. The dichromate will oxidize the iodide to iodine.

4. The iodine can then be estimated using a standard solution of sodium thiosulphate.

This procedure was carried out up to the end of step 3 on the same volume of the two samples of alcohol collected by fractional distillation and simple distillation from a previous practical. The iodine mixture obtained from the alcohol sample collected by simple distillation is labeled A; and that from the alcohol sample collected by fractional distillation is labeled B. You are required to determine which distillation method was the more efficient.

PROCEDURE:

Pipette 25cm3 of the iodine mixture into a clean conical flask. Fill the burette with a standard solution of sodium thiosulphate. Titrate the iodine mixture until a pale yellow (straw) colour is obtained. Add 1 cm3 of starch solution and carefully continue the titration until the blue-black colour just disappears. Repeat the titration to obtain consistent results. Record your results in an appropriate format.

TREATMENT OF RESULTS:

The following equations are useful in helping you to arrive at a conclusion re the relative alcohol content of the two alcohol samples: The standard solution of sodium thiosulphate

I2(aq) + 2 S2O32- (aq) → 2I- (aq) + S4O62- (aq)

2I- - 2e → I2 (aq)

Cr2O72- + 14 H+ + 6e → 2 Cr3+ + 7 H2O

3 C2H5OH + K2Cr2O7 + 4H2SO4 → 3CH3CHO + Cr2(SO4)3 + K2SO4 + 7H2O.

The standard solution of sodium thiosulphate contains 0.17 mol dm-3

1. Calculate the number of moles of thiosulphate ions in the titres: (a) from distillate A (b) From distillate B

1. Calculate the quantity of iodine present in 25cm3 of the iodine mixture titrated from:

(a) distillate A (b) distillate B

2. Calculate the molar concentration of iodine in each mixture: A and B.

3. Use the relevant equations to determine: (a) Which sample would have contained the larger quantity of dichromate by the end of step 2 of the first procedure.

4. Hence deduce which sample would have contained the larger quantity of ethanol prior to the first procedure.

MODULE 1

TITLE: MARGARINE VS BUTTER

Skills tested: P&D

PROBLEM

Margarine and butter.

Both can be used as spreads on breads or biscuits or even in cooking to improve flavour. In this modern age where obesity and terms such as “good” cholesterol and “bad” cholesterol are “hot” topics, the question margarine or butter, which is better for your health? was given as an assignment to CAPE chemistry students. Through research, the students determined that the degree of unsaturation in these products plays a role in the health consequences of using these products. Plan and design an experiment to determine the degree of unsaturation in samples of margarine and butter.

The format of your lab should be as follows

Title

Date

Aim

Hypothesis

Apparatus & Materials

Variables

Procedure

Data Collection (shown in an appropriate format)

Expected results

Discussion (which consists of limitations, precautions and sources of error)

-----------------------

E

F

G

Phosphate ions and silver nitrate ( yellow ppt (soluble in both acid and ammonia)

phosphate ions and ammonium molybdate with heating ( yellow ppt

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