Chemistry Core Practicals

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  • Created by: Garvey
  • Created on: 01-05-17 13:23

1. Measure the molar volume of a gas

Method

1. Place 30cm3 of 1moldm-3 ethanoic acid in a boiling tube

2. Set up apparatus

3. Place approx. 0.05g calcium carbonate into a test tube

4. Weigh the test tube and its contents accurately

5. Tip CaCO3 into boiling tube, quickly replacing the bung

6. Once the reaction is complete, measure the volume of gas collected

7. Re-weigh test tube which contained the CaCO3

8. Repeat 6 more times, increasing mass of CaCO3 by 0.05g each time, but don’t exceed 0.4g

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1. Measure the molar volume of a gas

Variables

  • Independent: Mass of calcium carbonate
  • Dependent: Volume of carbon dioxide evolved
  • Controls: Volume and concentration of ethanoic acid, temperature

 Error

  • Some CO2 escaped before the bung was put on
  • Some CaCO3 may get stuck in the test tube increasing the time before the bung is put on so pour the ethanoic acid into the CaCO3 instead

Results

  • Plot graph of mass CaCO3 (x) against volume of CO2 collected (y) with a line of best fit
  • As the mass of CaCO3 increases, the volume of CO2 evolved increases
  • Calculate moles CaCO3 in 0.25g and hence the volume of one mole of CO2
  • Show that the ethanoic acid was in excess
  • 2CH3COOH + CaCO3 → Ca(CH3COO)2 + CO2 + H2O
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2. Find the conc. of a solution of NaOH

Method

1. Dissolve an accurately known mass (approx. 2.5g) of sulfamic acid in approx. 100cm3 deionised water in a beaker

2. Transfer the solution and washings into a 250cm3 volumetric flask and make up to the mark

3. Prepare titration apparatus – acid solution will go in the burette, NaOH will go in the conical flask

4. Pour a 25cm3 aliquot of NaOH solution (unknown concentration) into a 250cm3 conical flask

5. Add 4 drops of methyl orange indicator

6. Titrate the contents of the conical flask against the prepared acid solution

7. Repeat until you have 2 concordant results (±0.2)

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2. Find the conc. of a solution of NaOH

Variables

  • Independent: Volume of NaOH solution (it will stay the same)
  • Dependent: Volume of acid solution
  • Controls: Concentrations, volume of indicator, temperature

Errors

  • Burette, pipette, volumetric flask (can be calculated)
  • Viewing colour change (subjective, white tile)

Results

  • Yellow solution to redy-orange solution
  • Neutralisation reaction
  • Calculate: total moles sulfamic acid, concentration sulfamic acid, average titre, mole sulfamic acid used, moles sodium hydroxide, concentration sodium hydroxide
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3. Find the conc. of a solution of HCl

Method

1. Wash out the 250cm3 volumetric flask with distilled water

2. Pipette 25cm3 hydrochloric acid solution into the volumetric flask and make up to the mark

3. Prepare titration apparatus – sodium hydroxide solution in the burette, dilute hydrochloric acid solution in the conical flask

4. Pour a 25cm3 aliquot of dilute HCl into the conical flask with 2 drops of phenolphthalein indicator

5. Titrate the contents of the flask against the NaOH solution

6. End point is indicated by a pale pink colour which persists for 5 seconds or more with continued swirling

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3. Find the conc. of a solution of HCl

Variables

  • Independent: Volume of dilute hydrochloric acid solution
  • Dependent: Volume of sodium hydroxide solution
  • Controls: Concentrations, temperature, volume of indicator

Errors

  • Burette, pipette, volumetric flask (can be calculated)
  • Viewing colour change (subjective, white tile)

Results

  • Colourless to pale pink solution
  • Neutralisation reaction
  • Calculations: mean titre, moles of NaOH in titre, moles HCl in titre, moles HCl in 250cm3, concentration of HCl
  • NaOH + HCl → NaCl + H2O
  • 2NaOH + CO2 → Na2CO3 + H2O
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4. Investigate the hydrolysis of halogenoalkanes

Method

1. Set up a water baths by filling a 250ml beaker up to the ¾ mark with 50˚C water

2. Fill 3 test tubes each with 5cm3 ethanol and add 4 drops of 1-iodobutane to the 1st tube, 1-bromobutane to the 2nd and 1-chlorobutane to the 3rd

3. Loosely place a bung on each test tube and place them in the water bath

4. Pour 5cm3 silver nitrate solution into 3 clean test tubes

5. When the halogenoalkane-ethanol solutions have reached the temperature of the water bath solutions, add the silver nitrate solution to one of the test tubes and start the stop clock

6. Measure the time taken for a precipitate to appear

7. Repeat steps 5 and 6 for the other 2 halogenoalkanes

8. Repeat steps 1-7 using 1-bromobutane, 2-bromobutane and 2-bromo-2-methylpropane

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4. Investigate the hydrolysis of halogenoalkanes

Variables

  • Independent: The halogenoalkane
  • Dependent: Time taken for a precipitate to form
  • Controls: Volume and concentration of ethanol, silver nitrate solution and the halogenoalkanes, water bath temperature

 Errors

  • Mixing up test tubes – keep them labelled
  • Measuring cylinder
  • Exact point of precipitation – subjective

Results

  • Time for part 1: 1-chlorobutane > 1-bromobutane > 1-iodobutane
  • Time for part 2: 1-bromobutane > 2-bromobutane > 2-bromo-2-methylpropane
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5. Investigate the oxidation of ethanol

Method

1. Place a few anti-bumping granules in a 50cm3 pear shaped flask

2. Add 10cm3 acidified sodium dichromate to the flask

3. Set up the flask for reflux

4. Mix 2cm3 ethanol with 5cm3 water

5. Place the flask in a beaker of iced water and add the ethanol solution dropwise down the condenser into the flask (this must be done slowly)

6. When all the ethanol has been added, heat the flask gently with a small Bunsen flame so the solution boils for 10mins

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5. Investigate the oxidation of ethanol

7. Perform the following tests:

  • a. Measure the pH of the distillate using universal indicator paper
  • b. Add a few drops of acidified potassium dichromate solution to 1cm3 of the distillate and warm the mixture in a 60°C water bath
  • c. Add a quarter of a spatula of calcium carbonate powder to 1cm3 of the distillate.
  • d. Add a 1cm long length of magnesium ribbon to 1cm3 of the distillate
  • e. Add 1cm3 of Fehling's solution to 1cm3 of the distillate and warm the mixture gently using a water bath
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5. Investigate the oxidation of ethanol

Results

  • Oxidation of ethanol to ethanoic acid
  • a. Low pH so acid
  • b. No change (no oxidation or reduction)
  • c. effervescence as CO2released
  • d. effervescence as H2released (acid)
  • e. No change as fully oxidised already
  • CH3CH2OH + 2[O] → CH3COOH + H2O
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6. Chlorination of 2-methylpropan-2-ol

Method

1. Pour 20cm3 2-methylpropan-2-ol and 70cm3 conc. HCl into a large conical flask with a bung

2. Swirl the contents of the flask, then remove the bung to release the pressure before replacing it and continuing to swirl

3. Continue this process for around 20 minutes when there should be 2 layers in the flask, the upper one being the crude product

4. Add approx. 6g of powdered anhydrous CaCO3 to the flask and swirl until dissolved (to ensure that any unreacted alcohol is in the lower aqueous layer

5. Transfer the reaction mixture into a separating funnel and allow the mixture to settle before running off and discarding the lower layer (retaining the upper layer in the funnel)

6. Add approx. 20cm3 of sodium hydrogencarbonate solution, swirl, and remove the bung at frequent intervals to release the CO2 produced before discarding the lower aq layer

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6. Chlorination of 2-methylpropan-2-ol

7. Repeat the washing with sodium hydrogencarbonate, shake the separating funnel and release the CO2 at frequent intervals

8. Run off and discard the lower layer, ensuring none remains in the tap

9. Run off all the organic layer into a small conical flask and add a spatula full of anhydrous sodium sulfate

10. Add a bung and swirl the contents to mix and then occasionally until the liquid looks completely clear

11. Decant the organic liquid into a 50cm3 pear-shaped flask

12. Set the flask up for distillation

13. Collect the fraction boiling between 50˚C and 52˚C and place the pure product into a labelled sample tube

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6. Chlorination of 2-methylpropan-2-ol

Results

  • Test the distillate by placing a few drops in a test tube. Add 5cm3 ethanol and 1cm3 NaOH(aq) and warm in a water bath. Add excess nitric acid followed by a few drops of silver nitrate solution. A white precipitate will form indicating the presence of chloride ions.
  • (CH3)2CHOH + HCl → (CH3)2CHCl + H2O
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7. Identify unknown organic (l) and inorganic (s)

Method: Part 1 – for organic liquids A, B, C

Start each test by placing 10 drops of each liquid into 3 separate test tubes

1. Add a 1cm depth of bromine water to each test tube and shake

2. Add acidified potassium dichromate to each test tube and warm in a 60˚C water bath for 5 minutes

3. Add 1cm3 of Fehling’s solution to each test tube and heat the resultant in a water bath

4. Add 1cm3 ethanol, 1cm3 dilute sodium hydroxide solution and warm the mixture for 5 minutes before acidifying each mixture with dilute nitric acid and adding 5 drops of silver nitrate solution

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7. Identify unknown organic (l) and inorganic (s)

Results: Part 1

  • Organic liquid A – an alcohol (propan-1-ol)
  • Not decolourised, no layers
  • No layers, green solution
  • Orange solution
  • No layers, dusty yellow
  • Organic liquid B - an alkene (cyclohexene)
  • Decolourised, 2 layers
  • No change
  • Green
  • Brown precipitate
  • Organic liquid C – a halogenoalkane (2-bromobutane)
  • Not decolourised, 2 layers
  • No change
  • Blue
  • Cream precipitate
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7. Identify unknown organic (l) and inorganic (s)

Method: Part 2 – for inorganic solids X, Y, Z

1. Conduct a flame test for each of the 3 solids

2. Dissolve a spatula full of each solid into 3 test tubes containing 10cm3 distilled water, and split into 3 portions

a. Add 5cm3 dilute nitric acid followed by 10 drops of silver nitrate solution and then add dilute ammonia solution

b. Add 5cm3 dilute nitric acid followed by 10 drops barium chloride solution

c. Add 2cm3 chlorine water

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7. Identify unknown organic (l) and inorganic (s)

Results: Part 2

  • Inorganic solid X – calcium bromide
  • Crimson red
  • cream precipitate
  • No change
  • Orange
  • Inorganic solid Y – sodium sulfate
  • Bright orange/yellow
  • No change
  • White precipitate
  • No change
  • Inorganic solid Z – potassium carbonate
  • Lilac
  • Effervescence
  • Effervescence
  • Effervescence
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7. Identify unknown organic (l) and inorganic (s)

Variables

  • Independent: the reagent used to test the substances
  • Dependent: the colour/appearance of the mixture during/after the reaction
  • Controls: temperature (especially part 1, test 2), volumes of the organic liquids and reagents

Errors

  • Colours are subjective
  • Inaccurate measuring
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8. Calc. the ΔH for the thermal decomp. of KHCO

Method

1. Place approx. 3g K2CO3(s) in a test tube and accurately weigh the test tube and its contents

2. Use a burette to dispense 30cm3 of 2moldm-3 hydrochloric acid into a polystyrene cup which is supported in a 250cm3 beaker

3. Measure the temperature of the acid

4. Continue measuring the temperature whilst adding the acid and swirling and record the highest temperature reached

5. Reweigh the empty test tube

6. Repeat steps 1-5 using approx. 3.5g KHCO3 instead of K2CO3, and record the lowest temperature

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8. Calc. the ΔH for the thermal decomp. of KHCO

Variables

  • Independent: The reagent used (i.e. K2CO3 or KHCO3)
  • Dependent: Temperature
  • Controls: Volume and concentration of HCl, use same equipment

Errors

  • Thermometer, mass balance (can be calculated)
  • Human reading error

Results

  • Calculate: energy change for each, enthalpy change for each, enthalpy change for thermal decomposition
  • Show that HCl is in excess
  • Draw energy level diagrams
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9. Determine Ka for a weak acid

Method

1. Set up the datalogger to read the pH, or calibrate the pH meter

2. Pipette 25cm3 of 0.1moldm-3 ethanoic acid solution into a 250cm3 conical flask

3. Fill a burette with sodium hydroxide solution

4. Add 2 or 3 drops phenolphthalein to the conical flask

5. Titrate the ethanoic acid solution with sodium hydroxide solution until the mixture just turn red

6. Pipette a further 25cm3 ethanoic acid solution into the conical flask

7. Record the pH of this solution

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9. Determine Ka for a weak acid

Variables

  • Independent: The volume of ethanoic acid
  • Dependent: The volume of sodium hydroxide solution
  • Controls: Concentration of NaOH and CH3COOH, temperature (this may vary slightly in the classroom)

Errors

  • Colour change is subjective – use a white tile
  • Burette (read from bottom of meniscus), datalogger, pipette (can be measured)

Results

  • Average pH: 4.77 (in theory)
  • Calculate [H+], Ka
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10. Construct EC cells and measure E

Method

1. Clean *****s of zinc and copper using sandpaper

2. Set up a zinc half-cell by pouring 50cm3 zinc sulfate solution into a 100cm3 beaker and standing the ***** of zinc in the beaker

3. Set up a copper half-cell by pouring 50cm3 copper(II) sulfate solution into a 100cm3 beaker and standing the ***** of copper in the beaker

4. Make an electrical connection between the 2 beakers by joining them with a ***** of filter paper that has been dipped in a saturated solution of potassium nitrate

5. Join the 2 metal *****s with a voltmeter, using the connecting wires and crocodile clips

6. Record the electrode potential of the [Zn(s)|Zn2+(aq)] and [Cu2+(aq)|Cu(s)] system (if it is a negative value, reverse the connections)

7. Repeat steps 1-6 using different combinations of metal/metal ion half cells

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10. Construct EC cells and measure E

Variables

  • Independent: Metals used
  • Dependent: Electrode potential value
  • Control: Same salt bridge (same solution and type of filter paper), same voltmeter and wires

Errors

  • Not using standard conditions of 100KPa, 1moldm-3, 298K
  • Impurities may still remain on the metal strips

Results

  • Record electrode potentials for each cell
  • Calculate Ecell­
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11. Amount of Fe in Fe tablet using titration

Method

1. Crush the iron tablets using the pestle and mortar

2. Transfer the crushed tablets to a weighing boat and measure their combined mass

3. Empty the crushed tablets into a small beaker and reweigh the weighing boat

4. Add 100cm3 1.5moldm-3 sulfuric acid into the beaker and stir to dissolve as much as possible

5. Filter the solution into a volumetric flask and add washings (with sulfuric acid). Make up to the mark with distilled water, stopper and shake

6. Pipette 25cm3 of this solution into a conical flask

7. Titrate the iron(II) solution with potassium manganate(VII) solution until the mixture has just turned pink (on standing the pink colour will disappear due to a secondary reaction between the KMnO4 and another ingredient in the tablet)

8. Repeat the titration until concordant results are obtained

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11. Amount of Fe in Fe tablet using titration

Variables

  • Independent: Volume of iron(II) solution
  • Dependent: Volume of potassium manganate(VII) solution
  • Controls: Concentrations, temperature

Errors

  • Human error in judging the end point – colour change is subjective so use a white tile
  • Burette, pipette, mass balance, volumetric flask – can be measured
  • Burette – bring down to eye level and read the bottom of the meniscus

Results

  • Form equation
  • Calculate: average titre, moles MnO4-, moles Fe, mass of Fe in each tablet, percentage of Fe in each tablet
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12. Prepare a transition metal complex

Method

1. Weigh a test tube, add 1.4-1.6g copper(II) sulfate and reweigh the test tube

2. Add 4cm3 water to the test tube using a graduated pipette

3. Stand test tube in a water bath (kettle and 100cm3 beaker) and stir gently to dissolve the copper(II) sulfate

4. Remove solution from water bath

5. In a fume cupboard with gloves, stir 2cm3 concentrated ammonia solution into the copper(II) sulfate solution

6. Pour the contents of the test tube into 6cm3 ethanol that has been pipetted into a beaker. Mix well and cool in an ice bath

7. Filter the crystals using a Büchner funnel and flask, using cold ethanol for the washings and rinsing the crystals

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12. Prepare a transition metal complex

8. Scrape the crystals onto a fresh piece of filter paper and cover with a second piece, patting the paper to dry the crystals

9. Once dry, measure the mass of the crystals

Results

  • Calculate: Moles copper(II) sulfate, theoretical yield, percentage yield
  • Comment on percentage yield
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13.a) Rate of the I -C H reaction using titration

Method

1. Mix 25cm3 1moldm-3 propane(aq) with 25cm3 1moldm-3 sulfuric acid in a beaker

2. Start the stop clock as soon as you add 50cm3 0.02moldm-3 iodine solution, and shake the beaker to mix well

3. Withdraw 10cm3 of the mixture and transfer it to a conical flask

4. Stop the reaction in this sample by adding a spatula of sodium hydrogencarbonate, noting the exact time this happens

5. Titrate the sample with 0.01moldm-3 sodium thiosulfate(VI), until there is a pale yellow colour, then add a few drops of starch indicator (blue) before continuing to titrate until the solution turns colourless

6. Continue to withdraw 10cm3 samples at 5 minute intervals, repeating steps 4-5 each time

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13.a) Rate of the I -C H reaction using titration

Variables

  • Independent: Time the sample was removed
  • Dependent: Concentration of iodine present
  • Controls: Volumes, temperature, concentrations of original reagents

Errors

  • Colour change is subjective
  • Noting the exact times
  • Measurable equipment uncertainty errors

Results

  • (didn’t add starch at right time)
  • Graph of titre against time
  • Order of reaction, rate equation
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13.b) Use clock reaction to determine rate eqn.

Method

1. Measure xcm3 potassium iodide solution into a small beaker on a white tile

2. Add 5cm3 sodium thiosulfate solution

3. Add 10 drops of starch solution to the mixture

4. Start the stop clock when you add xcm3 sodium peroxodisulfate solution

5. Stop the clock when a blue colour appears and note the time

6. Repeat steps 1-5 using the following volumes. Distilled/deionised water is added to make the total volume up to 25cm3

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13.b) Use clock reaction to determine rate eqn.

Mixtures:

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13.b) Use a clock reaction to determine a rate equ

Variables

  • Independent: Volumes of potassium iodide solution and sodium peroxodisulfate solution
  • Dependent: Time taken for the starch to turn blue
  • Controls: Temperature, volume of sodium thiosulfate, total volume

Errors

  • Colour change is subjective – use a white tile and have 1 person judge it
  • Accuracy in measuring the burette – percentage uncertainty can be calculated

Results

  • Calculate concentration of iodide ions and peroxodisulfate ions
  • Plot a graph of rate against concentration
  • Order of reaction, rate equation
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14. Determine the Ea for Br and Br

Method

1. Pipette 10cm3 phenol solution and 10cm3 bromide/bromate solution into 1 boiling tube

2. Add 4 drops of methyl red indicator

3. Pipette 5cm3 sulfuric acid solution into another boiling tube

4. Use a kettle and beaker to prepare a water bath at 75˚C (±1) and stand the 2 boiling tubes in it

5. When the contents of the boiling tube have reached the water temperature, pour the contents of one tube into the other and then back into the empty one, starting the stopwatch at the same time

6. Leave in the water bath and time until the methyl red indicator disappears

7. Repeat at 65˚C, 55˚C, 45˚C, 35˚C, 25˚C, 15˚C, using ice for the lowest temperature

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14. Determine the Ea for Br and Br

Variables

  • Independent: Temperature
  • Dependent: Time taken for colour change
  • Controls: Volume and concentration of phenol, bromide/bromate solution, methyl red, and sulfuric acid, pouring technique

Errors

  • Colour change is subjective – use a white tile and have the same person judge it each time
  • Pipette, thermometer uncertainty can be measured

Results

  • Convert temperatures to kelvin, calculate 1/temp., and take natural logs of the times
  • Plot ln(t) (y axis) against 1/T (x axis)
  • Calculate the activation energy using the Arrhenius equation and the gradient of the graph
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15. Analyse organic and inorganic unknowns

Method: Inorganic ions

1. Flame test

2. Use sodium hydroxide to identify metal cations

3. Use silver nitrate to identify halide ions

4. Test for sulfate ions

5. Test for carbonate ions

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15. Analyse organic and inorganic unknowns

Errors

  • Mixing up the unknown compounds – label well
  • Colours are subjective
  • Measurement uncertainty

Results

  • Compound A formed a blue-green flame and turned the lime water cloudy, suggesting that it is copper carbonate
  • Compound B formed an orange flame and a white precipitate with BaCl, suggesting that it is sodium sulfate
  • Compound C formed a brick red (or orange-red) flame, a white precipitate with NaOH and a cream/yellow precipitate with AgNO3 which dissolved in concentrated ammonia, suggesting that it is calcium bromide
  • Compound D formed a green precipitate with NaOH, and a white precipitate with AgNO3, suggesting that it is iron(II) chloride
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15. Analyse organic and inorganic unknowns

Method: Organic analysis

1. Test for alkenes

2. Test for aldehydes

3. Test for carboxylic acids

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15. Analyse organic and inorganic unknowns

Errors

  • Mixing up the unknown compounds – label well
  • Colours are subjective
  • Smells are subjective
  • Measurement uncertainty

Results

  • Solution X decolourised the bromine water, suggesting that it is an alkene
  • Solution Y formed a silver mirror, suggesting that it is an aldehyde
  • Solution Z formed a fruity smell suggesting that it is a carboxylic acid
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16. Synthesise aspirin from 2-hydroxybenzoic acid

Method

1. Weigh 2g 2-hydroxybenzoic acid and place in pear-shaped flask clamped and suspended in a beaker of water

2. Add 5cm3 ethanoic anhydride and 5 drops concentrated sulfuric acid before attaching a condenser

3. In a fume cupboard, carefully warm the mixture in the water bath using a Bunsen burner and gently swirl until all the solid had dissolved

4. Continue warming for another 10 minutes

5. Remove the flask form the hot water bath and add 10cm3 crushed ice and distilled/ deionised water

6. Stand the flask in iced water until precipitation is complete

7. Filter off the product using a Büchner funnel and suction apparatus

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16. Synthesise aspirin from 2-hydroxybenzoic acid

8. Wash the crystals with the minimum volume of iced water

9. Recrystallise the aspirin in the minimum volume of a mixture of ethanol:water, 1:3

10. Filter and dry

11. Measure the mass of the pure, dry crystals

12. Measure the melting temperature using the melting temperature apparatus

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16. Synthesise aspirin from 2-hydroxybenzoic acid

Variables

  • Independent: Mass of 2-hydroxybenzoic acid
  • Dependent: Mass (and purity) of aspirin
  • Controls: Temperature, volumes of reagents

Errors

  • Mass balance, measuring cylinders, thermometer – can be calculated

Results

  • Mass and melting temperature of dry aspirin obtained
  • Theoretical and percentage yield
  • Data book comparison of melting temperature
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