# AQA paper 1 13/06/2017 required practicals

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#1
So I decided to make some notes for the required practicals that we should know for Tuesday's exam, if I've done any mistakes or missed bits of info out, feel free to comment on it. Hope this little summary helps a few people (I know most of the practical-related stuff will be on paper 3 but some questions do come up in paper 1 and 2).

Required Practical 1: Making a standard solution; performing titrations; indicators

Assume you want to make a standard solution of sodium hydroxide.

1. 1. First, use n = cv to calculate how many moles of sodium hydroxide you need. The volume will usually be 250cm3 (since volumetric flasks come in set volumes), the concentration will be whatever you desire.
2. 2. Next, work out how many grams of sodium hydroxide you need using mass = Mr * mole.
3. 3. Place a weighing bottle on a digital balance and tare it (zero it/reset it).
4. 4. Now, using a clean spatula, transfer the correct amount of sodium hydroxide into the bottle. As you approach the desired mass, start adding it in smaller amounts so you don't add too much.
5. 5. Once you have the correct mass measured out, tip the contents into a clean 250cm3 beaker. It is important that you wash out the weighing bottle with distilled water multiple times and add that to the beaker as well. This ensures you have no solid left in the bottle.
6. 6. Pour some distilled water into the beaker and, using a glass rod, stir thoroughly. Keep adding distilled water until all the solid has dissolved.
7. 7. Once all the solid has dissolved, transfer the solution from the beaker to a 250cm3 graduated volumetric flask. You can do this using a glass funnel.It is important that you wash the beaker, stirring rod and funnel with distilled water multiple times and add that to the volumetric flask as well. This ensures you have not lost any solid.
8. 8. Keep adding distilled water to the volumetric flask. As you approach the 250cm3 mark, add distilled water drop-wise using a clean teat pipette. This will ensure you don't add too much water.
9. 9. Once the bottom of the meniscus is on the 250cm3 mark, close the volumetric flask and invert it multiple times. This will ensure absolutely all the solid has dissolved.

Notes:

• Ø The weaker the concentration of the standard solution, the more accurate the result will be. This is because you're minimising the percentage uncertainty of each reading by maximising the measurement taken.
• Ø If you add too much water to the volumetric flask, you will need to start the whole thing again. This is because you now an unknown volume.

Performing the titration.

1. 1. Using a clean 25cm3 pipette, transfer 25cm3 of your standard solution into a clean conical flask.
3. 3. Next, pour a little bit of acid (the one you're investigating in this titration) into the burette, ensuring the jet is closed.
4. 4. Place a clean beaker under the burette and open the jet. This rinsing of the burette will improve the accuracy of the result since it gets rid of any dirt or other impurities which cling onto the inside of the burette.
5. 5. Once the burette has been rinsed, fill it with acid. Once the jet is completely filled, close it and keep adding acid until you're approaching the 0cm3 mark on the burette.
6. 6. Just like with the volumetric flask, switch to a clean teat pipette once you're close to the mark and add the acid drop-wise. The bottom of the meniscus must lie on the 0cm3 mark.
7. 7. The apparatus is all set up and you may begin.

1. 8. Place the conical flask with the sodium hydroxide under the burette and open the tap. As the acid is flowing, swirl the conical flash gently.
2. 9. As soon as you see a rapid colour change/the colour stops changing, close the jet of the burette. Record your result- this is now your rough titre.
3. 10. Before repeating the experiment, prepare another conical flask with a fresh 25cm3 sample of sodium hydroxide solution. You do not have to fill the burette back up to 0cm3, but it will be easier if you let some drop out until the bottom of the meniscus lies on an integer (whole number).
4. 11. Repeat the experiment until you have 2 concordant results, i.e., results that are within 0.1cm3 of each other. It is important to partially-close the jet until the acid is dropping drop-wise as you approach your rough titre. This will make sure you don't over-add the acid.

Notes:

• Ø If you rinse the burette with distilled water instead, you will dilute the acid. This will decrease the percentage uncertainty but will not give an accurate result if you're trying to get a result that is (close to) the same as the one quoted in the data book.
• Ø If the jet is not filled before you start, your rough titre will be out by a few cm3.
• Ø Having the conical flask against a white background, e.g., a white tile or a white piece of paper, will help you to spot the colour change/absence of colour change faster.

Required Practical 2: Measuring enthalpy changes in the lab

Experiments measuring how much heat is given out by reactants are known as calorimetry (not to be confused with colorimetry) experiments.

Assume you want to find the enthalpy change of solution (usually neutralisation) or dissolution.

1. 1. Add one of the two solutions/solvent into a clean polystyrene cup and start the stopclock. You must use a known volume for this- integer values are generally easier for when you go about processing your results.
2. 2. Place a thermometer in and wait 3 minutes. After 3 minutes, record the temperature.
3. 3. Next, prepare your other solution/solid. Again, the volume/mass must be known. Add it to the cup.
4. 4. Immediately after adding it, place a lid on the cup. The lid must have 2 small holes in it- just big enough to fit a thermometer through and a glass stirring rod.
5. 5. Stir the mixture until the stopclock displays 5 minutes, at which point record the temperature.
6. 6. Stirring at a constant rate, carry on the experiment, taking temperature readings at minute intervals. In general, there is no need to do the experiment for longer than 7 or 8 minutes.

Processing the results:

1. 1. The temperature change isn't as simple as the maximum temperature recorded minus the initial temperature. This is because thermometers don't work instantaneously- they take a few seconds to get adjusted to the temperature change. While the thermometer is adjusting, heat is being lost to the surroundings. This is why you can't just use your peak temperature as the highest one achieved, since it is likely to be slightly lower than what is really should be.
2. 2. Plot a graph of time (x-axis) versus temperature (y-axis). You'll notice there are readings for 0 to 3 minutes and from 5 onwards- there is no temperature reading at the 4th minute.
3. 3. Draw a straight line of best fit through both sets of results. One line for the first 3 readings (which should be almost perfectly horizontal since the initial temperature should have been at room temperature). The other line for the readings after 5 minutes.
4. 4. It is important you extrapolate both lines by 1 minute.
5. 5. Finally, draw a vertical line from 4 minutes on the x-axis to find the real maximum temperature reached.
6. 6. The graph you have produced is known as a cooling graph.

Notes:

• Ø The more insulated the cup is, the more accurate the results will be. This is because you're minimising loss of heat to the surroundings.
• Ø Stirring at the same rate throughout ensures consistency.
• Ø If you're dissolving a solid in a solvent, it's important to calculate whether the solvent was in excess. If it wasn't, you must take this into account. You can only use the number of reacted moles as your n when you do q/n.
• Ø If you're mixing two solutions together, it's important that you allow them to reach the same (room) temperature before beginning.
• Ø If possible, use a thermometer with the highest precision you can. This reduces percentage uncertainty.

Assume you want to find the enthalpy change of combustion.

1. 1. Set up the apparatus- the beaker with the water should be on top of a gauze on a tripod, above a spirit burner. A thermometer should also be in the solution.
2. 2. Take the spirit burner and place it on a digital scale (the lid can be on or off, doesn't matter as long as you repeat the second weighing just like the first). Record the weight.
3. 3. Place the spirit burner under the beaker and light it.
4. 4. Record the temperature at minute intervals for as long as you desire- generally no need to carry on for longer than 7 or 8 minutes.
5. 5. Record the water temperature once you feel like you want to stop and extinguish the spirit burner by placing the lid back on top.
6. 6. Weigh the spirit burner on a digital scale and record the weight.

Notes:

• Ø Due to the nature of the experiment, a lot of heat will be lost to the surroundings.
• Ø Using the initial and final mass of the spirit burner, you can calculate how many moles were used (the fuel will generally be a simple alcohol, such as ethanol, whose Mr you can work out).
• Ø Incomplete combustion of the fuel may also add to the inaccuracy of the result.
• Ø The evaporation of the volatile fuel could also account for the inaccuracy.

• Ø In general, you should assume that volume of solution = mass of solution, unless the question states that the density is not 1.00.

Required Practical 4: Tests for various ions

Assume you want to identify positive metal ions.

1. 1. First, dip a nichrome wire (nickel-chromium) into concentrated HCl. This will clean it.
2. 2. Next, dip the cleaned wire into the compound you're investigating.
3. 3. After lighting a Bunsen burner, hold the wire loop in the clear blue part of the flame and observe the colour formed.

Notes:

• Ø Calcium is brick red.
• Ø Strontium is red.
• Ø Lithium is red.
• Ø Barium is pale green.
• Ø Potassium is lilac.
• Ø Sodium is yellow.
• Ø Magnesium is brilliant white.

Assume you want to identify ammonium ions.

1. 1. Place the unknown solution into a boiling tube. Next, suspend a piece of damp litmus paper on the top and add a hydroxide solution to the tube (such as sodium hydroxide).
2. 2. Shake the tube very gently.

Notes:

• Ø The litmus paper will turn blue due to ammonia gas (which is alkaline).

Assume you want to identify sulphate ions.

1. 1. Pour the unknown solution into a boiling tube and add a little bit of dilute HCl to it. This acid will remove any ions, e.g., carbonates, which could interfere with the test.
2. 2. Next, pour barium chloride solution into it.

Notes:

• Ø A white precipitate of barium sulphate will form.
• Ø You cannot use sulphuric acid instead of hydrochloric. This is because you'd be supplying sulphate ions which would be a false positive result.

Assume you want to identify hydroxide ions.

1. 1. Use a pH indicator.

Notes:

• Ø It will be alkaline.

Assume you want to identify halide ions.

1. Pour the unknown solution into a boiling tube and add a little bit of dilute nitric acid. This will remove any ions which could interfere with the test.

1. 2. Add silver nitrate solution.

Notes:

• Ø If fluoride ions are present, no precipitate will form.
• Ø If chloride ions are present, a white precipitate will form. You may confirm this by dissolving it in dilute ammonia.
• Ø If bromide ions are present, a cream precipitate will form. You may confirm this by dissolving it in concentrated ammonia.
• Ø If iodide ions are present, a yellow precipitate will form. You may confirm this by adding concentrated ammonia- it will not dissolve.
• Ø You may not use hydrochloric, hydrobromic or hydroiodoic acid to eliminate ions as they will give false positive results.

Assume you want to identify carbonate ions.

1. 1. Place your sample in a conical flash and add any acid. The flask must contain a bung on top to stop gas from escaping.
2. 2. The bung should have a delivery tube whose end is submerged in a separate boiling tube containing limewater.

Notes:

• Ø The limewater will turn milky.

Required Practical 8: Electrochemical cells

Assume you want to set up an electrochemical cell.

1. 1. Set up 2 separate, clean beakers.
2. 2. To each beaker, add a 1.00moldm-3 solution of the metal solutions you're investigating.
3. 3. Next, prepare your 2 electrodes. Take each electrode and scrub it with sandpaper. This will help to clean the surface from the oxidative layer. Once you've scrubbed it, purify the surface with propanone. It is very volatile, so it will evaporate quickly. Cleaning the electrode surface is important, as that's the reaction surface.
4. 4. Place each cleaned electrode into the corresponding solution.
5. 5. Now, take a U-shaped glass tube (not sure if it has a specific name) and fill it completely with an inert solution capable of ion transfer (this will be the salt bridge- the solution can be potassium nitrate for example).
6. 6. Stuff cotton wool into both ends of the tube in order to plug it.
7. 7. Turn the tube upside down and submerge both ends into the glass beakers.
8. 8. Finally, using crocodile clips and wires, connect the 2 electrodes to a voltmeter and turn it on.

Notes:

• Ø The salt bridge solution must be inert, otherwise its ions will interfere with either of the 2 solutions.
• Ø Both solutions must be 1M, unless you're investigating the effects of different concentrations on the EMF.
• Ø The voltage reading may be negative if you connected the wires the wrong way round. Just ignore the minus.

Required Practical 9: pH curves

Assume you want to draw pH curves for a titration reaction.

1. 1. Set up all apparatus ready for a titration and perform a rough titration.
2. 2. Wash the tip of the pH probe with distilled water multiple times. This is done to calibrate it.
3. 3. Set up all apparatus ready for a repeat titration. Place a pH probe into the beaker and record the initial pH.
4. 4. Add the acid/alkali from the burette in 2cm3 increments, recording the pH each time.
5. 5. Once you're ~2cm3 from your rough titre, decrease your increments to 0.2cm3.
6. 6. In the same manner, increase your increments back to 2cm3 once you're ~2cm3 past your rough titre.
7. 7. Carry on the experiment until you've added approximately 40cm3 of acid/alkali from the burette (assuming the equivalence point occurs before that).

Processing results:

1. 1. Draw a graph of volume of acid/alkali added (x-axis) against pH (y-axis).
2. 2. The curve should be sigmoidal.
3. 3. In order to calculate the neutralisation point, draw a vertical line down from the equivalence point. This will tell you how much acid/alkali reacted with your standard solution.
4. 4. You can then calculate your results just like in an AS titration.

Notes:

• Ø Best effort should be done to stir the conical flask, even if the pH probe is present.
• Ø If the pH probe isn't washed with distilled water, it may affect the first few readings.
• Ø Since the pH probe may not be 100% accurate, it is advisable to plot a calibration curve before the titration. To do this, use a washed pH probe to measure the pH of solutions with known pHs. You will then be able to draw a calibration graph and see how much the pH probe is out for each reading. Then, you apply this to the results you got from the titration.

Required Practical 11: Identifying metal ions

Assume you want to identify metal ions.

1. 1. Set up 3 different test tubes and pour a bit of your unknown sample into each one.
2. 2. To the first test tube, add, drop-wise, sodium hydroxide. If a precipitate forms, keep on adding sodium hydroxide, drop-wise. If the precipitate is not dissolving, leave it.
3. 3. To the second test tube, add, drop-wise, ammonia. If a precipitate forms, keep on adding ammonia, drop-wise. If the precipitate is not dissolving, leave it.
4. 4. To the third test tube, add, drop-wise, sodium carbonate.
5. 5. For each test tube, you should keep an eye out for a colour change, precipitate formation, precipitate dissolution and effervescence.
6. 6. Repeat the experiment for each unknown sample you're investigating.

Notes:

• Ø If a precipitate forms with sodium hydroxide or ammonia, it will be blue for copper, green for iron(II), white for aluminium and brown for iron(III).
• Ø If the white precipitate dissolves after continuing adding sodium hydroxide, it will form a colourless solution of [Al(OH)4(H2O)2]-.
• Ø If the blue precipitate dissolves after continuing adding ammonia, it will form a deep blue solution of [Cu(NH3)4(H2O)2]2+.
• Ø If a precipitate forms with sodium carbonate, it will be blue for copper, green for iron(II), white for aluminium (plus effervescence) and brown for iron(III) (plus effervescence).
• Ø If you leave a solution of iron(II) in air for long enough, it will turn brown due to oxygen oxidising it to iron(III).
8
4 years ago
#2
Thanks a lot for this, it's really helpful! Do you have a method for the rates or reaction procedures, and do you know which equations we need to memorise for them?
0
#3
(Original post by voltz)
Thanks a lot for this, it's really helpful! Do you have a method for the rates or reaction procedures, and do you know which equations we need to memorise for them?
Rates is in paper 2, I'll be doing some notes on the other practicals as well if you'd find that useful, sometime soon.
0
4 years ago
#4
Thank you for this! I haven't done a lot of practical revision so having it concisely like this is a great help.
0
4 years ago
#5
Thank you so much!! You're a lifesaver
0
4 years ago
#6
(Original post by Rexx18)
Rates is in paper 2, I'll be doing some notes on the other practicals as well if you'd find that useful, sometime soon.
That'd be very helful thanks.

Do you have any predictions for what may come up?
0
4 years ago
#7
Do you have notes on the rest of the required practicals? These are really useful by the way! Thank u x
0
#8
(Original post by ineedhelpasap98)
Do you have notes on the rest of the required practicals? These are really useful by the way! Thank u x
https://www.thestudentroom.co.uk/sho....php?t=4788630 here's the rest
0
2 years ago
#9
Lmfao my exam is today and I’m trying to osmosis these into my brain
0
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