Practicals Chemistry AQA A-level Rated A+

Practicals Chemistry AQA A-level Rated A+ Required Practical 1 - titration and making a standard solution Making a standard solution 1. weigh the required mass of solid using a sample bottle on a 2 dp balance. 2. transfer the solid into a beaker using distilled water to wash it out- preventing any solid remaining in the sample bottle. 3. reweigh the sample bottle and work out the difference. 4. Add 100cm3 of distilled water to the beaker to dissolve the solid (small volume of water) using a stirring rod to evenly distribute the solution. fer the dissolved solution into a 250cm3 graduated volumetric flask using a funnel, rinsing contents of the beaker and stirring rod with distilled water into the flask. 6. make up the solution to 250cm3 by adding more distilled water, using a dropping pipette for the last few drops. making sure the bottom of the meniscus is touching the line marking 250cm3. 7. invert the volumetric flask several times to ensure that the solution is uniform/evenly distributed. Required Practical 1 Making a solution and acid-base titration Acid-Base Titration (neutralisation)/redox titration 1. Rinse the equipment - with the solutions that they will contain e.g. pipette washed w/ alkali/standard solution, burette rinsed with the known sol conc./acid and the conical flask is rinsed w/ distilled water. 2. Pipette 25cm3 of alkali/standard solution into a conical flask. touch the end of the pipette in the alkali/standard sol once in the conical flask - ensures the correct amount of alkali has been added. 3. Add the known conc. sol into the burette using a funnel- prevents spillages of sol. 4. Add a few drops of indicator to the conical flask (phenolphthalein pink in alkali and colourless in acid) 5. Place a white tile under the conical flask - to accurately observe the colour change. 6. Add the acid to the alkali / known sol to unknown - swirling the conical flask. Add drop wise as the reaction reaches it's end point. 7. repeat until 2 concordant results are obtained. within 0.1 of each other. record the initial titre, final titre- Titre. To find the conc. of the unknown sol. Why is a volumetric pipette more accurate than a measuring cylinder? because it has a smaller uncertainty Why use a conical flask and not a beaker? conical flask is easier to swirl Why only use small drops of indicator? indicators tend to be weak acids - so adding too much will affect the result of the titration. how do you reduce percentage uncertainty in a titration? - replacing any measuring cylinders with pipettes and burettes which have a lower apparatus uncertainty. - increase the vol and conc. of the substance in the conical flask - decrease the concentration of the sol in the burette so a larger volume will be needed. Reducing percentage uncertainty when measuring mass using a more accurate balance using a larger mass Redox titration (not required) only sulphuric acid can be used with manganate other wise a large inaccurate vol of manganate will be added if HCl and poisonous Cl2 gas produced. not enough H+ ions. manganate is self indicating Required Practical 2 Measuring Enthalpy Change- in solution 1. Place a polystyrene cup in a beaker to provide support and insulation 2. Add a volume/mass of one of regents/only regent, using a volumetric pipette, to the polystyrene cup and place the lid over it, which is clamped with a thermometer. 3. make sure the bulb of the thermometer is in the liquid/sol and measure the initial temperature of the regent. 4. if 2 reagents are involved do the same for the second reagent, measuring it's initial temperature. 5. Add the second reagent to the first in the polystyrene cup and place lid back on. 6. Start the stop watch. Measure the temperature at regular intervals e.g. 1 min (until the temp becomes constant). Mixing the solution. 7. Plot a graph of temperature against time 8. extrapolate back to find the temperature before mixing to work out the temperature change. What happens if the reaction is slow? then the temp rise will be difficult to obtain because cooling is happening at the same time. To counteract: measure the temp at regular time intervals and extrapolate the temp/time curve back to when the reactants were added together. Errors in the method Why would the enthalpy change be different to that in the reaction? Heat loss/heat gain from the surroundings - temp change lower/higher Energy can be absorbed by the apparatus. incomplete reaction/ slow reaction. Errors: What does the method assume about the solutions? The solutions have the same specific heat capacity as water. q=mcΔT m = mass of solution (not the solid/excess) if 2 solutions - add them together When water - this is the mass of the water. n = mol of the solution not in XS Measuring Enthalpy Change - Combustion - flame calorimetry 1. measure the mass of the spirit burner containing the alcohol/fuel 2. Measure a volume of water e.g. 50cm3 and add it to a copper container- calorimeter. 3. Measure the initial temperature of the water 4. burn the spirit burner under the calorimeter until a specific temperature has been reached e.g. a temp change of 50 degrees Celsius. stirring the water to keep the temp increase uniform. 5. Once this temp has been reached, extinguish the spirit burner and measure it's final mass- use this to workout the mass lost. 6. use this mass to work out how much mass of alcohol reacted- and then use q=mcΔT. Errors in this reaction: CALORIMETRY COMBUSTION Heat loss to the surroundings- heat loss from evaporation of water Incomplete combustion Evaporation of Fuel after weighing - lowers it's mass heat capacity of the calorimeter is not included Required Practical 3 Investigation how the rate of a reaction changes with temperature Sodium Thiosulphate and HCl = Disappearing cross 1. Measuring cylinders - measure 50cm3 of 0.05 mol of sodium thiosulphate. In another- measure 5cm3 of HCl - place them I seperate boiling tubes. 2. measure the initial temperature of the 2 solutions before you add one to another in a conical flask. 3. Add the sodium thiosulphate to conical flask first and then the HCl. The conical flask should be placed over a white tile with a black cross, start the stopwatch and measure the time it takes for the cross to disappear due to the cloudy suffer being produced. 4. Repeat with 4 more different temperatures using water baths of different temperatures and ice. 5. rate is measures as 1/time Plot a graph of rate against temperature and a graph of lnK AND 1/T (temp) to work out the activation energy for the reaction using the gradient k is proportional to rate therefore lnK = lnrate Required Practical 4 - Identifying cations and anions Cations- Group 2 and NH4+ - Add 10 drops of 0.1 moldm-3 metal ion solution to a test tube - Add 10 drops of 0.6moldm-3 of NaOH mixing well -continue to add dropwise until in excess. Group 2 metal cations = form white ppt/solids strontium and Barium OH - soluble so they will dissolve If you add sulphuric acid instead - mg and ca will dissolve and strontium and barium sulphates will form a white precipitate. NH4+ = e.g. from ammonium chloride Add NaOH dropwise - gently heat the solution - hold damp red litmus paper over the test tube = Litmus paper goes blue due to the production of ammonia gas Required Practical 4 - Identifying cations and anions Anions - Halides, OH-, SO4- and CO32- Halides - nitric acid and silver nitrate Cl- = Dilute ammonia - ppt dissolves Br- = in conc. ammonia - ppt dissolves I- = in conc. and dilute - no change - ppt remains And reactions with H2SO4 (group 7 topic) carbonate - react with HCl - produces effervescence of CO2 / bubble through limewater = turns cloudy. Sulphate - BaCl and HCl - white ppt of BaSO4 Hydroxides - Sodium, potassium and ammonium hydroxides = soluble most others are insoluble. use pH meter 8. Measuring the EMF of a cell Method 1. Clean a piece of copper and zinc using fine grade sandpaper. 2. Degrease the metal using cotton wool and propanone. 3. Place the copper into a 100cm3 beaker with 50cm3 of 1moldm-3 of copper sulphate 4. Place a strip of Zinc into a 100cm3 beaker with 50cm3 of 1moldm-3 of Zinc Sulphate. 5. use a strip of filter paper soaked in potassium nitrate for the salt bridge 6. connect the 2 half cells by connecting the metals to crocodile clips attached to leads joined to the voltmeter. Why Potassium Nitrate a suitable salt bridge? It does not react with the elecrode solutions and electrodes It allows the movement of ions between half cells it maintains a constant charge Why would potassium chloride be unsuitable in the presence of Cu and why would a metal wire be unsuitable? The chloride ions would produce a complex with the Cu ions. A metal wire would set up it's own electrode system. What if there is a half cell with 2 solutions e.g. Fe 2+ and Fe3+ use a platinum electrode because it is inert. The fe2+ sulphate solution and Fe3+ sulphate solution electrolyte should both be at 1moldm-3. 9. Titration Curves - Investigating how pH changes in WA/SB and SA/WB Method 1. Transfer 25cm3 of acid to a conical flask using a volumetric pipette. 2. measure the initial pH of the acid with a pH meter 4. Add the alkali from the burette in small amounts e.g. 2cm3 at a time. Noting the volume added. Stir mixture to equalise the pH 5. measure and record the pH to one d.p. 6. when approaching end point add alkali in smaller volumes, drop by drop until the alkali is in excess. Why do you calibrate the pH meter? Calibrate pH meter by measuring the pH of a known pH buffer solution// distilled water. Because pH meter can lose accurate over time when in storage. How else can you improve accuracy of a acid base titration? pH curves By maintaining a constant temperature. Required Practical 5: Organic Distilling a product from reaction e.g. Primary alcohol - Aldehyde Separating an Organic product from it's reacting mixture. Collect the liquid at it's approximate boiling point. Dilute H2SO4 and Limited Potassium dichromate Dichromate - irritant and Toxic you add the dilute sulphuric acid (10cm3) to a pear-shaped flask. Add 3g of Potassium Dichromate and anti-bumping granules. Shake contents of the flask to mix. Add 1.5 cm3 of the alcohol (propane-1-ol) in drops using a dropping pipette as you shake the flask to mix the contents. Assemble distillation and gently heat and distil he liquid into a conical flask/test tube. - the aldehyde will vaporise at it's boiling point but the alcohol has a higher boiling point so it will not vaporise - Pass into the condenser - The condenser will cool the vapour - condense into a liquid. - condenser - thermometer - bulb should be at the T-junction connecting to the condenser to measure the correct boiling. - Water goes in through the bottom of the condenser, against gravity to allow more efficient cooling and prevents back flow of water. Required Practical 5 Organic Reflux of Primary alcohol to carboxylic acid/ Secondary alcohol to Ketone Reflux also used to form a pure organic Product Excess Potassium Dichromate Conc. acid catalyst H2SO4 1. Dissolve the excess potassium dichromate in a boiling tube 2. Put primary alcohol into a 50cm3 pear shaped flask, add distilled water and bumping granules (to prevent liquid from bumping the sides of the flask and spilling) 3. Add the condenser - has np lid Water in at the bottom and out at the top - against gravity for efficient cooling. 4. Add the small amounts of the Conc. Sulphuric acid, using a dropping pipette into the flask. 5. Whilst the mixture is warm add the potassium dichromate solution down the condenser in drops, using a dropping pipette. The energy released from the reaction - causes the mixture to boil. Add the Sodium Dichromate drop by drop to allow the mixture to continue to boil without the need for external heating. 6. Once all the potassium dichromate is added, use a low bunsen burner flame to keep mixture boiling for 10 minutes- not allowing any vapour to escape. 7. distil any product for further purification. the condenser prevents any vapour from escaping by condensing them back into liquids, returned to the reaction mixture. Heating organic reaction mixtures for long periods of times- they boil - vaporise and nonsense until only the product is left. During reflux what precautions do you take? NEVER SEAL THE CONDENSER - build up of pressure from gas could cause apparatus to break and explode. Anti-bumping granular added in distillation and reflux to prevent uneven boiling and stop the liquid hitting the sides of the flask - smaller bubbles instead of larger bubbles. Required Practical 10: Organic Preparation of a pure organic solid and test of its purity Recrystallisation aspirin - ester Reflux Distill - distallate Purification of the product: 1. Dissolve in the minimum volume of hot solvent: Minimum volume - so the solution is saturated Hot solvent - To dissolve both the product and impurities, solubility increases with increasing temp, and so that on cooling the crystals will reform. 2. Filter the solution hot : To remove any insoluble impurities using a filter funnel and fluted filter paper. 3. Cool the filtered solution by inserting the beaker in ice: Increases the yield of crystals and soluble impurities will remain in solution because they are present in small quantities. 4. Suction filtrate with a Buchner flask to separate out the crystals, the water pump connected to the Buchner flask will reduce the pressure and speeds up the filtrate. 5. wash crystals with distilled water: removes soluble impurities. 6. Dry the crystals between absorbent paper you can use melting point and compare with data booklet to test purity - the lower the mpt the lower the less pure Why would there be loss of yield in this process? Crystals may be lost during filtering or washing. Some product may stay in solution after recrystallisation. Side reactions may occur. If the crystals are not dried properly - the mass may be larger - increases percentage yield. Required Practical 10: Organic Preparing a pure organic liquid Cyclohexene from cyclohexanol // ester Reflux and Distill 1. put the distillate of impure product into a separating funnel 2. Add sodium hydrogen carbonate - to neutralise any acid reactant Shaking and releasing the pressure at intervals from CO2 produced. Add Sodium chloride - to separate the organic layer from the aqueous layer. 3. Allow the layers to separate in the funnel and run and discard the aqueous layer- this is the lower layer and crude organic solution is the upper layer. 4. Transfer upper layer into a small conical flask. 5. Add the drying agent - anhydrous calcium chloride - removes any water until the dry organic solution appears clear. Drying agent - insoluble in the organic liquid and should not react with it. 6. Decant and then redistill to collect pure product / measure the mass of the dry product in a sample weighing sample container. 7. test boiling point Required practical 7: Measuring the initial rate of reaction By initial rate method Iodine Clock Hydrogen peroxide reacts with Iodide ions to form Iodine. Thiosulphate ions react with the iodine, removing it to form more iodide ions. 1. Rinse the burette with potassium iodide (accuracy). 2. Add the potassium iodide to the 50 cm3 burette 3. Transfer 10cm3 of hydrogen peroxide from a burette into a clean, dry beaker. 4. Use a 50cm3 measuring cylinder to measure out 25cm3 of sulphuric acid, add this to a 250cm3 beaker. 5. use a 25cm3 measuring cylinder to add 20cm3 of distilled water to the 250cm3 beaker. 6. use plastic dropping pipette to add 1cm3 of starch to the beaker. 7. Use the burette to add 10cm3 of potassium iodide to the 250cm3 beaker. 8. Finally add 5cm3 of the sodium thiosulphate to the mixture in the 250cm3 beaker. 9. Stir the mixture in the 250cm3 beaker. Pour the 10cm3 of hydrogen peroxide from the beaker into the mixture in the 250cm3 beaker and immediately start the timer. 10. stop the timer when the mixture in the beaker goes blue black. Record the time. 11. Rinse the beaker with distilled water and dry out. 12. Repeat steps 4 more times with different concentrations of potassium iodide. 13. plot initial Rate against concentration to find the order of the reaction. Required practical 7: Measuring the rate of a reaction- Continuous monitoring method 1. Add 50cm3 of 0.8moldm3 of HCl to a conical flask. 2. set up a gas syringe 3. Add 6cm of magnesium strip to the conical flask. 4. Close the flask with the bung joined to the gas syringe and start the timer. 5. swirl the flask. 6. Record the volume of gas produced every 15 secs for 2 and half minutes. 7. Repeat with different concentrations of HCl. 9. Plot volume against time for each concentration of HCl and draw a tangent to time =0 s for each concentration to measure the initial rate of reaction. 10. compare the rate values obtained. Required practical 7: Other ways of measuring rate 1. Colorimeter - using a calibration curve and absorbance over time 2. Quenching - by cooling, neutralising an acid base catalyst, Diluting. 3. Measuring loss of Mass 4. pH