unit 2 lab manual

Table of Contents

Page Title Module # 3 Standardisation of hydrochloric acid 2 5 Calibration of a pipette 2 6 Redox Titration / H2O2 vs KMnO4 2 8 Back Titration 2 10 Gravimetric analysis of hydrated salt 2 11 Flame tests 2 12 Reactions of cyclohexane and cyclohexene 1 14 Paper Chromatography 2 15 Water polluting inorganic ions 3 16 Distillation of a sample of rum 3 17 Salted peanuts “Fuller” 2 17 Antique or Fake 2 17 Margarine vs Butter 1 18 Identity of sample X 3

Skills tested

ORR, A&I M&M, A&I

ORR,A&I

P&D P&D P&D M&M, ORR, A&I

Unit 2 Chemistry Practicals 2010-2011 MODULE 2 TITLE:

page 3 of 19

STANDARDISATION OF HYDROCHLORIC ACID

SKILLS TESTED: None AIM: The purpose of this experiment is to determine the concentration of a solution of h ydrochloric 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)

Unit 2 Chemistry Practicals 2010-2011 Burette Reading/cm3 (to 2 decimal places) Final

Trial

1

page 4 of 19 2

3

4

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 cm3 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. ………………………………………………………………………………………………………………………

Unit 2 Chemistry Practicals 2010-2011 MODULE 2 TITLE:

page 5 of 19

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 mass of water discharged 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

0.9967 8

0.9962 3

0.9956 5

0.9950 3

0.9943 8

0.99369

3. 4. 5. 6. 7.

Using your answer from question 2 determine the mean volume of water discharged from the pipette. Calculate the standard deviation of the five masses from question 1 and convert to volume as well. State the capacity of the pipette in the format mean volume ± standard deviation State the capacity of the pipette as shown on the pipette you have used. Using the mean volume you have calculated and true mean value stated on the pipette, calculate the % error via the formula below % error = difference between calculated mean value and true mean value x 100 true mean value

8. Discuss any sources of error in your experiment.

Unit 2 Chemistry Practicals 2010-2011

page 6 of 19

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 con centration 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 Reading/cm3 (to 2 decimal places) Final

Trial

1

2

3

Initial Volume used/cm3 average titre/cm3


4


page 7 of 19

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 usingyour titration results. 5. Determine the error in your results. 6. Suggest a reason for the partial decomposition of the hydrogen peroxide.

Unit 2 Chemistry Practicals 2010-2011 MODULE 2 TITLE: BACK TITRATION

page 8 of 19

SKILLS TESTED: None AIM:

To determine the percentage of calcium carbonate in an impure sample of calcium carbonate

THEORY: Sometimes a direct titration would involve a reaction wh ich 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 1.0 g of the solid mixture and place in a conical flask. 2. Measure 25 cm3 of the 1 mol dm-3 HCl solution and place in the conical flask. Please stir mixture with glass rod until all effervescence has ceased. Remember to rinse glass rod with small amount of distilled water and allow water to run into the conical flask. 3. Rinse and fill the burette with 0.2 mol dm-3 NaOH solution. 4. 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 twountil 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 Reading/cm3 (to 2 decimal places) Final

Trial

1

2

3

Initial Volume used/cm3 Average volume



page 9 of 19

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 in the solid sample.

Unit 2 Chemistry Practicals 2010-2011 MODULE 2 TITLE: GRAVIMETRIC ANALYSIS

page 10 of 19

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.


page 11 of 19

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 Colour

Treatment of results:

1. Why are different colours observed? 2.

(a) Identify any ion or ions for which there was no colour imparted to the flame. (b) Suggest an explanation for this.

3. Why is the cleaning of the nichrome wire necessary before performing a flame test?


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Reactions of cycloalkanes and cycloalkenes Skills tested: ORR, A&I

Module 1

Introduction

This practical involves the use of cyclohexane and cyclohexene and reactions that would distinguish between the two compounds. Procedure

Tests Observations 3 1. Use a measuring cylinder and place a 2cm of the cyclohexane R onto a crucible lid and while holding the lid with a tongs, ignite the liquid by placing the lid over a flame. 2. Repeat test 1 using cyclohexene S

3. Place 2 cm3 of R in a test-tube and add a few drops of bromine liquid and shake and observe. Retain testtube for test 5.

4. Repeat test 3 using cyclohexene S

5. Place test-tube from test 3 in direct sunlight and note any changes

6. Place 2 cm3 of S in a test-tube and add a few drops of acidified potassium manganate solution.

Unit 2 Chemistry Practicals 2010-2011 Treatment of results

page 13 of 19

1. Based on the results of tests 1 and 2, would this method used to distinguish alkanes from alkenes a reliable test? Give a reason for your answer. 2. Suggest a reason for the role that sunlight played in test 5, give the name of the mechanism for the reaction and draw the fully displayed formula of the first organic compound that would have been formed in this reaction. 3. What type of reaction occurred in test 6? Draw the structure of the molecule that was produced. 4. Write the equation of the reaction that occurred between S and bromine liquid and suggest a name for the resulting product.


MODULE 2 TITLE:

page 14 of 19

PAPER CHROMATOGRAPHY

AIM: To separate the components in a mixture of food colouring dyes . INTRODUCTION One method used to physically separate the components of a mixture is chromatography. For each component a retention factor (R f) value, which characterises it may be calculated. Hence the method may also be used to identify the components. PROCEDURE 1. Using a clean ruler and a pencil mark a horizontal line approximately 1.5 cm from the tapered end of the strip of chromatography paper placed on a clean surface. This line will be the origin in the developed chromatogram. Mark a small stroke in the centre of the line to represent the position the spot should be placed. 2. Use the marked strip to estimate the height to which the developing solvent should be poured into the measuring cylinder. N.B: The solvent should be in contact with the paper but below the level of the origin at the start of the separation . 3. Remove the strip of paper and pour the solvent (either water or an ethanol/water mixture into the measuring cylinder (Chromatography tank), stopper the container and leave to stand for a few minutes. 4. Place the strip of chromatography paper on a clean surface and spot on the mixture to be separated as follows: 5. Use the toothpick and make a dot roughly 1 mm across on the origin. Repeat this 4 more times to ensure the spot of mixture is concentrated. 6. Holding the strip by the edge furthest from the origin, place it into the tank so that the tip is submerged. 7. Replace the cover and leave to stand. 8. Remove the developed chromatogram when the solvent has reached about ¾ to the top of the paper. 9. Use a pencil to mark the furthest position (solvent front) to which the solvent has moved. 10. Calculate the R f value for each component in the ink. 11. Draw an ACCURATE chromatogram showing the results. 12. Do not write a full lab report. QUESTIONS

1. a) From your results suggest the number of components that may be present in the dye b) Using the chromatogram, determine the R f values of each component. 2. What is the principle on which the separation was based? 3. Why would different R f values be obtained with the different solvents? 4. What further procedure may be used to obtain a better separation of the components?

Unit 2 Chemistry Practicals 2010-2011 MODULE 3

page 15 of 19

WATER POLLUTING INORGANIC IONS

Skills tested: None

PROCEDURE: Carry out the following tests, using an approximately 2 cm3 sample of the aqueous ion and complete the table below

. Tests 1. Nitrate (V) anions a) Add a small amount of Devarda’s alloy followed by sodium hydroxide. Warm mixture and test gas evolved with moist red and blue litmus.

Observations

Inferences

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

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

b) Add aqueous barium chloride. Divide the ppt into three portions and test its solubility in (i) dil nitric acid, (ii) dil ethanoic acid, (iii) dil aqueous ammonia. 3. Lead(II) ions a) Add aqueous sodium hydroxide slowly until in excess

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

c) Add aqueous potassium iodide until in excess.

d) Add a small amount of aqueous sodium sulphide solution

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 (SCN-) to which a solution of iron (III) chloride is then added. Write the equation for the reaction of cyanide ions with polysulphide ions. Suggest the colour and formula of the final product. (b) Cyanide ions react with dilute HCl to liberate a poisonous gas which has the smell of bitter almonds. Identify the gas and write the equation for the reaction.


page 16 of 19

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. P lease 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 point fraction

Volume of distillation after demonstration was completed

Smell of distillate (state 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. Why would there be a difference in the volumes of distillate collected? 3. Based on the strong or weak smell of alcohol of the distillate, which seems to be the more efficient separation technique?


page 17 of 19

Original salted peanuts “Fuller” vs reduced salt Fuller

Skills tested: None (practice)

Module 2

PROBLEM:

You are given two packages of salted peanuts. The original version, which has a high salt content and a newer reduced salt content version. Determine whether the newer version does in fact have a reduced salt content compared to the original version. The format of ALL plan and design reports should be as follows Title Date Hypothesis Aim Apparatus and Materials Procedure Variables Expected results Treatment of results Chemical principles (theory which would influence your procedure) Sources of error / precautions / limitations Antique or Fake

Skills tested: P&D

Module 2

You find a coin dated 1410 and you are wondering if the coin can be cashed in for a significant reward based on its age. After doing some research, you realize coins contains copper and those dated before 1600 had a much lower copper content than those dated after 1600. You, a new chemistry graduate decide to find out if this coin is really dated 1410 or if it is a fake. Design an experiment to determine the copper content of the coin.

Margarine vs Butter

Skills tested: P&D

Module 1

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.


page 18 of 19

MODULE 3 TITLE: Qualitative analysis of sample X SKILLS TESTED: M&M, ORR,A&I PROCEDURE:

Carry out the following tests on the sa mple X. Carefully record all observable changes. Complete the table below. Do not write a full lab report. Tests 1. Add a heaped spatula load of the solid X to 10 cm 3 of distilled water in a boiling tube and stir thoroughly. Keep the contents of the tube for the following tests. 2. To approximately 1 cm 3 of solution, add NaOH(aq) slowly until in excess.

Observations

Inferences

………………………………….. …………………………………. Ion(s) possibly present …………………………………

3. To approximately 1 cm 3 of solution, add NH3 (aq) until in excess.

Ion(s) possibly present …………………………………

4. To approximately 1 cm 3 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 NaOH (aq) and warm the mixture. Test gases evolved with moist red and blue litmus.

………………………………… Acid-base nature of the gas evolved ………………………… Formula of gas evolved …………. Ion(s) possibly present ……………………………….

6. Add solid or aqueous FeSO 4, then pour conc. H2SO4 slowly down side of tube.

Formula of compound formed …………………………… 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]


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unit 2 lab manual

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