1
Q
what are some common mistakes when making a standard solution?
A
- not weighing by difference
- not ensuring all your solid is dissolved before transferring to the flask
- forgetting to rinse your stirring rod/breaker/funnel
- not putting the washings into the volumetric flask
- filling the volumetric flask over the line
- failure to fully mix the solution (not inverting the flask enough times)
2
Q
what are tips to look out for when making a standard solution?
A
- write down all the possible numbers from your weighing in rough
- keep your weighing boat on your desk and dont let anyone touch ut
- only use 100cm^3 of water to dissolve your solid initially
- use a pasteur pipette to add the last few drops of water to your volumetric
- remember the miniscus should sit on the line
- put your palm over the flask’s lid when mixing to ensure you dont lose and solution
3
Q
what are some common mistakes when doing a titration?
A
- not conditioning your equipment with the correct solutions
- not eliminating bubbles from the burette
- filling the burette over your head
- not reading the burette at eye level
- titrating with the funnel in
- using your rough as an actual result
4
Q
what are tips to look out for when doing a titration?
A
- practice using your burette to let out one drop at a time when you are conditioning
- balance any equations before you start so that you’re ready to do your calculation as soon as you have your result
- check you know what the end point will look like before you start titrating
- add half drops near the endpoint
- complete your rough very quickly so that you’re ready can spend more time on your ‘real’ titrations
5
Q
what are some common mistakes when doing double titration?
A
- make sure you add the indicators in the right order - phenolphthalein is always first
6
Q
what are some common mistakes when doing back titration?
A
- don’t forget its alkali in the burette, be especially careful to wear your goggles
- remember the colour changes will be the opposite to usual
7
Q
what are tips to look out for when doing back titration?
A
- if you only take a portion of your original reaction don’t forget to factor this into your calculation
8
Q
what are some common mistakes when doing inorganic tests (precipitation reactions) ?
A
- failing to label all your test tubes at the start
- giving answers as ‘yes’ instead of stating what the observations are
9
Q
what are tips to look out for when doing inorganic tests (precipitation reactions)?
A
- work out which reactions will give precipitates before you start
- remember some substances are ‘sparingly soluble’ so may be inconclusive - it is fine to record this
10
Q
what are some common mistakes when doing gravimetric analysis?
A
- failing to test the filtrate with reagent to see if all the solid has precipitated
- forgetting to rinse the beaker into the filter paper
11
Q
what are tips to look out for when doing gravimetric analysis?
A
- you may have to repeat the precipitation several times but make sure all your solid ends up in the same filter paper (you can add a a second piece of filter paper under the first if needed)
12
Q
what are some common mistakes when measuring rate?
A
- taking too long, this allows the room temperature to change and will affect your results
- not measuring concentrations / volumes exactly
- writing down the tenths and hundredths of seconds from your stopwatch - dont do this
13
Q
what are tips to look out for when measuring rate?
A
- plot your graph for rate as you go along, this way you can immediately repeat any experiments you need to - your points should be really close to your line of best fit
14
Q
what are some common mistakes when using electrochemical cells?
A
- forgetting to soak your salt bridge
- putting the cells the wrong way round
15
Q
what are some common mistakes when doing melting point experiments?
A
-forgetting to seal the base of the tube
16
Q
what are some common mistakes when doing recrystallisation?
A
- using too much solvent
- putting into an ice-bath too quickly (you will only make tiny crystals)
17
Q
what are tips to look out for when doing recrystallisation?
A
- if you cant see any crystals forming, scratch the bottom of the glass with a stirring rod. alternatively, seed the crystals by dipping a glass stir rod in the hot solution. let the solvent dry, then dip the stir rod back into the cooled solution
- make sure the solution is fully cool before filtering
18
Q
what are some common mistakes when doing esterification?
A
- forgetting to add the concentrated acid
- forgetting to distil off the ester
19
Q
what are tips to look out for when doing esterification?
A
- use the acid anhydride for a larger yield
20
Q
what are some common mistakes when doing thermochemistry?
A
- taking the thermometer out to read the temperature
- forgetting to use a lid
- not taking enough readings before and after adding the reagent
21
Q
what are some common mistakes when doing organic synthesis?
A
- not refluxing for long enough
22
Q
what are tips to look out for when doing organic synthesis?
A
- follow the instructions exactly
- record all distillation temperatures
23
Q
what is a standard solution?
A
- one whose concentration and volume is known exactly
- this can then be used to calculate the concentration of other solutions
24
Q
how to make a standard solution?
A
- standard solutions of solids can be prepared by weighing a mass of solid, and dissolving it in a known volume of solvent in a volumetric flask:
- weigh out your solid (see weighing by difference)
- carefully add a small amount of distilled water to the beaker, use a glass rod to stir until all the solid has dissolved
- use a funnel to transfer the liquid into the volumetric flask, pouring down the glass rod
- rinse out the beaker and funnel several times with distilled water. wash the glass rod. ensure that all washings go into the volumetric flask
- top up to the mark on the volumetric flask
- put the stopper in firmly and thoroughly mix the solution
- label the flask with your name
25
how do you weigh by difference?
- put the bottle and top on the balance and press on
- add between 1.1 and 1.4g
- put the lid on and remove the bottle from the balance
- re-zero the balance
- add the full bottle (with the lid) and write down the weight to 3.d.p
- carefully empty the solid into a beaker
- put the bottle containing any residual solid and the lid back on the balance, write down the weight to 3.d.p
- calculate the accurate weight of the solid to 3d.p
results:
- weight of bottle + top + solid = ____ g
- weight of bottle + top + residual solid = ____ g
- accurate weight of solid (1-2) = ____ g
26
what is the test for ammonia (gas)?
- test with red litmus paper
- +ve result = litmus paper turns blue
27
what is the test for carbon dioxide?
- bubble through limewater
- +ve result = limewater turns milky as calcium carbonate is made
28
what is the test for hydrogen?
- test with a lighted splint
- +ve result = gas burns with a squeaky pop
29
what is the test for oxygen?
- test with a glowing splint
- +ve result = splint relights
30
what is the test for ammonium NH4+?
- heat with NaOH, test vapours with red litmus
- +ve result = litmus turns blue (from evolved ammonia gas)
31
what are the tests for magnesium Mg2+?
- flame test
- +ve result = no flame colour
- add NaOH
- +ve result = white ppt that is insoluble in excess
- add NH3 (aq)
- +ve result = white ppt that is insoluble in excess
- add Na2CO3 (aq)
- +ve result = white ppt
32
what are the tests for calcium Ca2+?
- flame test
- +ve result = brick red flame
- add NaOH
- +ve result = white ppt that is insoluble in excess
- add NH3
- +ve result = no reaction
- add Na2CO3
- +ve result = white ppt
33
what are the tests for copper (II) Cu2+?
- add NaOH
- +ve result blue ppt
34
what is the test for iron (II) Fe2+?
- add NaOH
- +ve result = dark green ppt
35
what is the test for iron (III) Fe3+?
- add NaOH
- +ve result = brown ppt
36
what is the test for potassium K+?
- flame test
- +ve result = lilac flame
37
what is the test for sodium Na+?
- flame test
- +ve result = yellow-orange flame
38
what are the tests for zinc (II) Zn+?
- flame test
- +ve result = no flame colour
- add NaOH
- +ve result = white ppt that is soluble in excess
- add NH3
- +ve result = white ppt that is soluble in excess
39
what is the test for aluminium Al3+?
- add NaOH
- +ve result = white ppt that is soluble in excess hydroxide
40
what is the test for chromium (III) Cr3+?
- add NaOH
- +ve result = grey-green ppt forms, soluble in excess NaOH to give a dark green solution, but not soluble in excess NH3
41
what are the tests for strontium Sr2+?
- flame test
- +ve result = crimson flame
- add NaOH
- +ve result = Sr(OH)2 is sparingly soluble - ppt depends on the concentration of the solution
- add Na2CO3
- +ve result = white ppt
42
what is the test for manganese (II) Mn2+?
- add NaOH
- +ve result = off-white ppt (insoluble with excess NaOH or NH3) that rapidly turns brown
43
what are the tests for lead (II) Pb2+?
- add NaOH
- +ve result = a white ppt forms which is soluble in excess NaOH
- add KI
- +ve result = yellow ppt forms
44
what are the tests for barium Ba2+?
- flame test
- +ve result = apple green flame
- add NaOH
- +ve result = doesnt form a ppt
- add Na2CO3
- +ve result = white ppt
45
what is the test for bromide Br-?
- add nitric acid + silver nitrate
- +ve result = cream ppt (of silver bromide)
46
what is the test for carbonate CO3 2-?
- add acid
- +ve result = fizz of co2 produced (can be tested with lime water)
47
what is the test for chloride Cl-?
- add nitric acid + silver nitrate
- +ve result = white ppt of silver chloride produced
48
what is the test for iodide I-?
- add nitric acid + silver nitrate
- +ve result = pale yellow ppt
49
what is the test for sulphates SO4 2-?
- add solution of barium chloride
- +ve result = white ppt of barium sulphate
50
C4H3KO8.nH2O + 3NaOH —> C4Na3KO8.nH2O + 3H2O
- a student carries out a two-part experiment to determine the value of n in hydrated potassium textraoxalate
- in part 1, the student weighs exactly 1.78g of the solid and transfers all of it to a 250cm^3 beaker
- she adds 100cm^3 of deionised water while stirring to ensure that it all disolves
- describe, giving full practical details, how the volume is made up to exactly 250^3 [3]
- pour liquid from beaker —> 250cm^3 volumetric flask
- add washings to volumetric flask
- make up to 250
- invert to ensure mixture is homogenous
51
when making a standard solution, why put it in a beaker first?
- to avoid spilling because the neck of the volumetric flask if thin
- or solid can get stuck on the side
52
when making a standard solution,
why do you dissolve the solid in a small amount of water first?
- to avoid spilling
- to avoid going over 250cm^3
53
when making a standard solution, why do you wash out the beaker and add it to the volumetric flask?
- to ensure all the solid is tranferred
54
mostly all group 1s are soluble
55
all nitrates are soluble
56
what metal’s salt produces:
- a white ppt when reacted with H2SO4
- no visible reaction when reacting with NaOH
- an orange/brown solution when reacting with Cl2
barium
57
what metal’s salt produces:
- no visible reaction when reacted with H2SO4
- a white ppt when reacting with NaOH
- an orange/brown solution when reacting with Cl2
magnesium
58
compound w is a simple inorganic salt containing one d-block metal cation and one anion.
tests were carried out to identify compound W
- test - observation - inference
- add compound w to water - ________ - Fe2+ or Cr3+ ions could be present
- add aq NaOH drop-wise to solution of compound W, then add excess NaOH - ___________ - confirm Cr3+ ions present
- add aq silver nitrate - cream ppt forms - __________
- green solution
- grey-green ppt
- Br-
so compound w = CrBr3
59
if something unknown is insoluble but know it’s either group 1 or group 2, it’s probably group 2 because most group 1s are soluble
60
a solution of an unknown salt containing one s-block metal cation and one anion was tested
- WHAT ARE THE CONCLUSIONS FOR EACH TEST AND OBSERVATION
- solution of unknown salt
- test 1 = add CO3 2- = white ppt formed = conclusion =
- test 2 = add OH- = no change observed = conclusion =
- test 3 = add Cl2 = grey solid/brown solution formed =
- test 4 = add Cu2+ = white ppt in brown solution = formula of white ppt =_________. brown colouration due to=_______
- test 5 = add sodium thiosulfate solution TO RESULTS OF TEST 4 = observation=
- group 2
- barium
- I- present
- CuI
- I2
- colourless
61
why is CuI (+1) a white ppt?
Cu+ = d^10 = no free d orbitals to promote electrons
62
explain the observation made in test 5. include an equation in your answer [2] (addition of sodium thiosulfate solution to CuI)
2Cu2+ + 4I- —> 2CuI + I2
2S2O3 2- + I2 —> S4O6 2+ + 2I
(iodine is used up so colourless)
63
state the feature required for a molecule to show optical activity [1]
chiral centre
64
state two chemical tests that will give a positive result for 5-aminobenzene-1,3-dicarboxylic acid but not ethyl-4-nitrobenzoate
- Na2CO3
- NaNO2 + HCl (HONO) - warm HONO above 5°C
65
state three chemical tests that will give a positive result for 1-(4-hydroxy-3-nitrophenyl)ethanone but not ethyl-4-nitrobenzoate
- tests for phenols:
• neutral iron (III) chloride - purple solution
• bromine water - white ppt
- iodoform - iodine + sodium hydroxide - yellow ppt of CHI3
66
what are the reagents when using iodoform?
iodine + sodium hydroxide
67
what is the end product when using the iodoform test?
- yellow ppt of CHI3
68
is PbCl2 soluble or insoluble?
insoluble
69
what are the only two soluble compounds of lead
- Pb(NO3)2
- Pb(CH3COO)2
70
- a student planned to distinguish between eight compounds:
• magnesium hydroxide
• iron (II) hydroxide
• chromium (III) hydroxide
• lead (II) hydroxide
• magnesium carbonate
• iron (II) carbonate
• chromium (III) carbonate
• lead (II) carbonate
- step 1 : add dilute acid until all the solid has disappeared. record any effervescence
- step 2: add 1cm^3 of NaOH solution to each solution formed in step 1. record any ppt observed
- the student plans to use dilute hydrochloric acid in step 1. his teacher tells him that this is not the correct acid to use. explain why hydrochloric acid should not be used and suggest an appropriate acid to use in its place
- PbCl2 is insoluble so wont be able to do step 2 / HCl would form INSOLUBLE compound with Pb2+
- use nitric acid instead
- as Pb(NO3)2 is soluble
71
what are the 4 ions that redissolve in excess NaOH?
- Cr3+ —> grey-green —> dark green
- Pb2+ —> white ppt —> colourless
- Al3+ —> white ppt —> colourless
- Zn2+ —> white ppt —> colourless
72
why are the 4 ions that redissolve in excess NaOH the only ones that do?
because they are amphoteric
73
you are supplied with four unlabelled aqueous solutions containing the following species:
- carbonate CO3 2-, iodide I-, chlorine Cl2, thiosulfate S2O3 2-
you are also provided with the following reagents:
- dilute H2SO4 and AgNO3
- devise a scheme whereby all four of the unlabelled solutions could be positively identified
you should include observations and ionic equations for any reactions occuring [8]
- add AgNO3 to all 4
- one turns yellow ppt —> I-
- Ag+ (aq) + I- (aq) —> AgI (s)
- add H2SO4 to remaining 3
- fizz —> carbonate
- CO3 2- + 2H+ —> CO2 + H2O
- add I- (identified) to remaining 2
- one goes brown —> Cl2
- 2I- + Cl2 —> 2Cl- + I2
- add I2 to last solution
- brown —> colourless
- 2S2O3 2- + I2 —> S4O6 2- + 2I-
74
what is the enthalpy change equation?
∆H = -mc∆T / n
∆H = joules
m = g
75
if solution C and solution E are reacted together, and form a white ppt that dissolved when extra solution E is added, what does it make solution E to be?
NaOH
76
% error equation:
(uncertainty / measurement obtain) x100
(if two measurements, times uncertainty by 2)
77
titre found = 26.30
the burette used in the titrations has an uncertainty for each reading of +- 0.05cm^3. estimate the maximum % error in the titre [1]
((0.05 x 2) / 26.30) x 100
= 0.38%
(because you find 2 titration values ( start and end)
78
burettes measure to +- 0.05
79
the error = 1/2 of the smallest increment you have on the instrument
80
to what d.p do you read titrations to?
2 d.p
81
when weighing solid, for % error calculation, uncertainty x2
82
should you use the rough results when calculating the mean?
no
(in practical tho, can change rough to number 1 if concordant)
- dont technically need rough column
83
- in the practical:
- if making standard solution
- MAKE SURE MIXED
84
what unit should it be when measuring for % error?
g
85
500mg = ___g
0.5g
86
what has occurred if a compound of CHI3 has been formed?
iodoform test
87
what is produced when tollens or fehling is added to an aldehyde?
carboxylic acid
(aldehyde has been oxidised)
88
in qualitative analysis, why does it not matter exactly how much of each solution is added to the other
- it is qualitative analysis which means the exact measurements are not being recorded
- it only matters that enough of the solution is added for a possible reaction to be observed
89
what are the hazards associated with HCl, NaOH and phenolphthalein?
- HCl = irritant
- NaOH = irritant
- phenolphthalein = flammable
90
what is the chemical equation for the reaction between NaOH and HCl?
HCl + NaOH —> NaCl + H2O
91
what type of reaction takes place between HCl and NaOH?
neutralisation
92
why should the burette be filled below eye level?
so that if any of the acid spills whilst being poured in, it wont splash into your face
93
in salt preparation by titration, why is the titration repeated without the indicator?
- this ensures that the salt produced in the repeated titration will not be contaminated with phenolphthalein
94
why are titrations usually carried out on a white tile?
- allows the point of colour change to be easily identified
95
why is the e.g NaOH placed in a conical flask?
- the conical flask allows the mixture to be swirled without losing any of the contents
96
why must the reaction mixture be swirled during the titration?
- swirling ensures all the reaction particles collide and react
- this helps to give a more accurate end point for the reaction
97
why are burettes and pipettes always used in titrations?
- burettes and pipettes measure the volumes of solutions very precisely
98
what is a standard solution?
a solution with a known concentration
99
what apparatus is required to make up a standard solution?
- weighing bottle or boat
- 250cm^3 volumetric flask
- digital mass balance
- funnel
- 250cm^3 beaker
- glass rod
- pipette
100
describe how to make up a standard solution of anhydrous sodium carbonate
- accurately weigh out approximately 2.75g of anhydrous Na2CO3 into a weighing boat. record the mass
- tip the solid into a 250cm^3 beaker and reweigh the weighing boat
- dissolve the solid in deionised water, stirring with a glass rod. do not add more than 150cm^3 of distilled water
- pour the solution into the 250cm^3 volumetric flask via a funnel
- rinse the beaker and glass rod three times and transfer the washings into the flask
- make the solution up to the mark with distilled water so that the bottom of the meniscus is level with the graduation mark
- add the stopper and shake the mixture thoroughly
101
why must the weighing boat be reweighed after the contents are transferred to a beaker?
- this means that the exact mass of solid transferred can be calculated - as some traces of the solid may have been left behind in the weighing boat
102
when making up a standard solution, why is it important to not add more than 150cm^3 of deionised water when first dissolving the solid?
- this is enough to dissolve it and leaves enough volume left for the washings, allowing the solution to be carefully made up to the 250cm^3 mark
103
why must the volumetric flask and graduation mark be at eye level when adding the final drops?
- this ensures the graduation mark is being viewed at the right angle, allowing the solution to be made up to 250cm^3 more accurately
104
what is the meniscus?
- the curved surface of the liquid within a tube
- when making the solution up to the graduation mark, the bottom of the meniscus should be exactly in line with the graduation mark
105
what is the purpose of ‘washing’ the beaker and glass rod into the solution in the volumetric flask?
- washings ensure there is no solute left behind in the beaker or on the glass rod
- this is important to ensure the concentration of the standard solution is as accurate as possible
106
what is the colour change of methyl orange?
- red in acid
- yellow in alkali
107
describe how to titrate a standard solution with HCl
- fill the burette with the HCl solution and record the initial burette reading
- accurately pipette 25.0cm^3 of the e.g Na2CO3 solution (standard solution) into a conical flask
- add 3 drops of indicator to the flask
- the standard solution is titrated with the acid until, on the addition of one drop of acid, the indicator changes colour
- record the burette reading and repeat until the results are concordant
108
what is the chemical equation for the reaction between Na2CO3 and HCl?
- 2HCl + Na2CO3 —> 2NaCl + H2O + CO2
109
why is the pipette rinsed with (the standard solution) before use and the burette rinsed with (the acid) before use?
- rinsing the equipment with the solutions removes any water which may be in the equipment
- this is important because the water will affect the concentrations of the solutions
- therefore, rinsing ensures a more accurate titration experiment
110
how can aldehydes and ketones be identified from a collection of unknown organic samples?
- add 2,4-DNP to all possible samples
- the aldehydes and ketones will be identified by the positive results of a bright orange/yellow ppt
111
what reagents can be used to distinguish between aldehydes and ketones?
- tollens’ reagent
- fehling’s reagent
112
describe the test for aldehydes using tollens’ reagent
- add the tollens’ reagent to the unknown sample
- a positive results of of a silver mirror forming on the inside of the test tube will indicate the presence of an aldehyde
113
how can you test for methyl ketones?
- iodoform test:
- warm the sample with iodine and sodium hydroxide
- a positive result is indicated by a yellow ppt and an antiseptic smell
114
describe the test for aldehydes using fehling’s reagent
- add the blue fehling’s reagent to the unknown sample
- a positive result of the formation of a brick red ppt will indicate the presence of an aldehyde
115
what can be tested for to distinguish between butanone and diphenylmethanone?
- since butanone has the structure CH3COCH2CH3, it is a methyl ketone
- therefore a test for methyl ketones (iodoform test) will distinguish between the two compounds
116
how is litmus paper used to test for acids?
- blue litmus paper will turn red when exposed to an acid
117
outline how to indirectly determine the enthalpy change of reaction of magnesium oxide and carbon dioxide
- measure 50cm^3 of HCl into the calorimeter using the pipette. place the thermometer into the HCl and leave it to allow the temperature reading to stabilise
- accurately weigh out 0.90g of MgO in a weighing boat. record the mass
- record the temperature of the acid and start the stopwatch. record the temperature of the acid every 30 seconds for 2minutes 30. at 3 minutes, add the MgO to the HCl and mix thoroughly
- when the stopwatch reaches 3 mins 30, record the temperature of the reaction mixture
- record the temperature of the mixture every 30 seconds until the temperature drops for 5 readings
- weigh the weighing boat again. record the mass of MgO added to the calorimeter
- use the data to construct a graph and calculate the enthalpy change of the reaction
- using the values of ∆H, calculate the enthalpy change for the reaction
118
when filling the pipette with HCl, how should the measurement be taken?
- the bottom of the meniscus (curve of the liquid) should be in line with the 25cm^3 mark
- the reading should be taken at eye level to avoid parallax errors
119
how can you use a graph to find an accurate temperature change of a reaction?
- plot the temperature of the solution before reaction and after the reaction
- draw two lines of best fit and extrapolate them to the point that the reaction started
120
why is it generally hard to get accurate results in calorimeter experiments?
- there is always heat lost to the surroundings which means the temperature measurements are not completely accurate
121
how can you prevent heat loss to the surroundings/apparatus?
- use a polystyrene cup to hold the reaction mixture because polystyrene is a good insulator
- place a lid on the reaction mixture
- place the polystyrene cup in a beaker of cotton wool to increase insulation
- avoid large temperature differences between the surroundings and the calorimeter
122
other than preventing heat loss, how can the accuracy of the experiment be improved?
- read the thermometer at eye level to avoid parallax errors
- stir the solution to evenly distribute the temperature
- use a digital thermometer for more accurate and faster readings
123
how could you reduce the uncertainty in the mass measurement?
- use a digital balance with a greater resolution
- use a larger mass
124
give the experimental procedure to investigate the enthalpy change of combustion of methanol
- add 100cm^3 of deionised water to a 250cm^3 conical flask
- place the conical flask in the clamp and position it on the stand at a height so that the spirit burner can fit underneath it
- weigh the spirit burner and lid containing methanol. record the mass
- record the initial temperature of the water
- place the spirit burner under the conical flask and light the wick
- allow the flame to heat the water to around 40°C
- extinguish the flame and record the final temperature of the water
- re-weigh the spirit burner and lid and record the mass. calculate the mass of methanol used
- determine the energy released by methanol and calculate the enthalpy change of combustion
125
how can you prevent heat loss to the surroundings/apparatus when investigating the enthalpy change of combustion of methanol?
- place a lid on the calorimeter
- avoid large temperature differences between the surroundings and the calorimeter
126
other than heat loss, why might the experimental enthalpy of combustion value be less exothermic than the data book value for investigating the enthalpy change of combustion of methanol?
- loss of fuel or water by evaporation
- incomplete combustion of the fuel - this could deposit soot within the burner, affecting the final mass of the fuel
- it is unlikely the reaction takes place under standard conditions - unlike the data book value
127
how can you avoid incomplete combustion when investigating the enthalpy of combustion using calorimetry?
- ensure there is a consistent flow of air to the spirit burner so that the fuel burns in sufficient oxygen
- this will encourage complete combustion to take place
128
- hydrated chromium chloride has a molecular formula of CrCl3H12O6 and a molar mass of 266.6gmol^-1
- it has the somewhat unusual property of existing in a number of octahedral isomers
- these isomers differ in terms of the number of chloride ions and water molecules that are acting as ligands in the complex ion, and the number of chloride ions and water molecules that are not part of the complex ion
- four possible isomers are shown
- a chemist was provided with a solid compound W, which is known to exist as one of the chromium chloride isomers shown above. after the analysis the following results were obtained:
- METHOD 1:
- dissolve a small amount of compound W in water and add aqueous sodium hydroxide, initially drop by drop and then in excess
- green solution formed on dissolving in water. a grey/green gelatinous precipitate was formed on addition of a few drops of aqueous sodium hydroxide. the precipitate dissolved in excess aqueous sodium hydroxide to form a dark green solution
- METHOD 2:
- dissolve 13.3g of compound W in water and add excess aqueous silver nitrate. filter the silver chloride that is precipitated, wash with a small volume of water to remove surface impurities and heat to constant mass
- 7.18g of silver chloride were obtained
- use ALL this information to identify compound W and hence draw a structure for the octahedral complex ion
- in your answer you should make it clear how your explanation is linked to the results and give ionic equations for the reactions taking place [8]
- METHOD 1:
- green solution suggests Cr3+ (aq)
- confirmed by Cr3+ (aq) + 3OH- (aq) —> Cr(OH)3 (s)
- dissolves in excess NaOH (aq)
- Cr(OH)3 (s) + 3OH- (aq) —> [Cr(OH)6]3- (aq)
- METHOD 2:
- Ag+ (aq) + Cl- (aq) —> AgCl (s)
- 13.33g of W = 13.33/266.6 = 0.05 mol
- 7.18g of AgCl = 7.18/143.5 = 0.05 mol
- 1 mol of W contains 1 mol of Cl- ions not coordinately bonded to Cr3+
- therefore compound W is isomer III / [CrCl2(H2O)4]Cl.2H2O
129
state why the egg shell is ground to a very fine powder before reaction with the acid? [1]
to increase the rate of reaction/to increase the surface area of the CaCO3
130
state why the burette was rinsed with the sodium hydroxide solution before filling [1]
- to remove any water (from the burette) which would dilute the sodium hydroxide solution / no impurities in the burette
131
state why the contents of the conical flask were swirled during the titration [1]
- to make sure that reaction (between NaOH (aq) and HCl (aq) is complete / all HCl has reacted / all NaOH has reacted
132
- titration 1 - mass 0.455 - volume NaOH 30.50 cm^3
- titration 2 - mass 0.516 - volume NaOH 16.30
- titration 3 - mass 0.482 - volume NaOH 24.80
- titration 4 - mass 0.535 - volume NaOH 12.90
- titration 5 - mass 0.469 - volume NaOH 22.60
identify the titration that has the largest % error in the volume of NaOH (aq) used and give a reason for your choice [1]
- titration 4
- smallest volume (of NaOH)
133
lowri’s % of calcium carbonate in the powdered eggshell was slightly higher than the actual value. when asked to suggest why, she said ‘i did not dry the eggshell sufficiently before grinding and weighing’
explain whether Lowri’s statement could account for this inaccurate result [1]
- no
- wet sample means that the actual mass of calcium carbonate is smaller than that recorded therefore the % by mass is lower
134
compound A can be prepared in the lab by nitration of methyl benzenecarboxylate using a nitrating mixture of concentrated nitric acid and concentrated sulfuric acid.
stage 3: the nitrating mixture was prepared by adding 3cm^3 of concentrated nitric acid to a dry beaker and cooling by partially immersing it in an ice-water bath. then 3cm^3 of concentrated sulfuric acid was added slowly, with swirling, and the mixture was then allowed to cool for 10 minutes
suggest why the reaction mixture was cooled during stage 3 [1]o
- to prevent more than one nitro group being substituted
135
compound A is very much more soluble in hot ethanol than it is in cold ethanol.
describe how you would purify the sample of compound A [3]
- dissolve solid in MINIMUM volume of HOT ethanol
- filter off any insoluble impurities whilst hot
- allow to cool (to crystallise solid)
- filter and dry
136
why did the student repeat the titration? [1]
to obtain concordant results (to obtain a more reliable value)
137
state why the potassium manganate (VII) solution has to be standardised [1]
- its purity is not high enough (to be a primary standard)
- it absorbs moisture from the air
- concentration of solution changes over times
138
state why no indicator is required in this titration [1]
- potassium manganate (VII) has a deep (purple) colour which disappears on reaction and when the end-point is reached a single drop more turns the solution pink
neutral answer - it changes colour by itself
139
suggest why the iron solution was made up using 5 tablets rather than just one tablet. refer to the mean titre in your answer [1]
- when using only one tablet the mean titre would be less than 5cm^3 and such a small volume has a high percentage error
(must have reference to small volume AND high error)
140
- the solubility of calcium hydroxide in water was found from a solution:
- add the solid to about 1200cm^3 of distilled water
- stir the mixture for ten minutes utes
- filter the mixture
state why the solid was added to DISTILLED water [1]
- to avoid contamination / ensure that all Ca2+ ions came from the solid
141
- the solubility of calcium hydroxide in water was found from a solution:
- add the solid to about 1200cm^3 of distilled water
- stir the mixture for ten minutes utes
- filter the mixture
state why the mixture was stirred for ten minutes [1]
- so that all the calcium hydroxide that could dissolve had dissolved / to produce a standard solution / to ensure homogeneity
142
when making a standard solution, why dont you simply add 1.14g of metal hydroxide to 250cm^3 of water? [1]
- weighing bottle would not have been washed / difficult to dissolve solid in volumetric flask / final volume would not necessarily be 250cm^3
143
name a suitable piece of apparatus for transferring 25.0cm^3 of metal hydroxide solution to a conical flask [1]
pipette
144
state why you add an indicator when titrating? [1]
- to show the end point / when to stop adding acid / when it’s neutralised
145
state what is seen when aqueous chlorine is added to a solution of calcium bromide [1]
red / brown solution
146
- gwyneth uses different tests to identify the four solutions : potassium carbonate, sodium hydroxide, barium chloride and magnesium nitrate
- each test allows her to distinguish between some of the solutions. for each test state the solution(s) that would give a visible change and the observation(s) that would be made:
i) addition of litmus solution [1]
sodium hydroxide solution would turn blue/purple
147
what is a further test that can be done after the addition of silver nitrate to distinguish between sodium bromide and sodium iodide. give the reagent and observations for this further test [2]
- reagent : (dilute) ammonia solution
- observation with sodium bromide: ppt dissolves in part (if concentrated ammonia used, then sodium bromide will dissolve completely)
- observations with sodium iodide : ppt does not change
148
when bromine water is added to a solution of sodium iodide, a reaction occurs. write an equation for this reaction [1]
2NaI + Br2 —> 2NaBr + I2
149
magnesium nitride reacts with water to form magnesium hydroxide and ammonia:
Mg3N2 + 6H2O —> 3Mg(OH)2 + 2NH3
- calculate the minimum mass of magnesium nitride required to form 1.75g of magnesium hydroxide, giving your answer to three significant figures [3]
moles Mg(OH)2 = 1.75 / 58.32(Mr) = 0.0300
moles Mg3N2 = 0.01
mass Mg3N2 = 0.01 x 100.9 = 1.01g
150
the element calcium reacts with water to produce calcium hydroxide and hydrogen gas:
Ca (s) + 2H2O (l) —> Ca(OH)2 (aq) + H2 (g)
- the same student repeated the reaction using the same mass of barium. he noticed that the volume of gas, still at the same temperature and pressure, was less
give the reason why the volume of gas produced was less [1]
fewer MOLES of barium used / barium has a higher Ar
151
the element calcium reacts with water to produce calcium hydroxide and hydrogen gas:
Ca (s) + 2H2O (l) —> Ca(OH)2 (aq) + H2 (g)
- the same student repeated the reaction using the same mass of barium. he noticed that the volume of gas, still at the same temperature and pressure, was less
suggest another difference that the student would observe when barium was used in place of calcium [2]
- reaction faster / more vigorous / less cloudy solution formed with barium
- because ionisation energy of barium is less / electrons lost more easily from barium / barium is lower in the group / barium hydroxide is more soluble
152
starting from a solution of barium hydroxide, describe how he could obtain a pure, dry sample of barium sulfate.
you should include an ionic equation for the reaction [3]
- add sulphuric acid / sodium solution / potassium sulfate solution
- filter
- Ba2+ + SO4 2- —> BaSO4
153
- callum prepares copper (II) sulfate solution from hydrated copper (II) sulfate, CuSO4.5H2O
- he measures a mass of hydrated copper (II) sulfate and uses this to make exactly 250cm^3 of copper (II) sulfate solution of concentration 0.250moldm^-3
describe giving full practical details, how callum should prepare the 250.0cm^3 of copper (II) sulfate solution [5]
- weighing method
- dissolve copper sulfate in a smaller volume of distilled water
- transfer to 250.0cm^3 volumetric flask
- use of funnel
- wash funnel / glass rod / beaker wit distilled water into volumetric flask
- add distilled water up to mark
- shake solution / mix thoroughly
154
a student weighs exactly 1.78g of the solid and transfers all of it to a 250cm^3 beaker. she adds 100^3 of deionised water while stirring to ensure that it all dissolves.
describe, giving full practical details, how the volume is made up to EXACTLY 250^3 [3]
- transfer to 250cm^3 volumetric flask (1)
- use of funnel
- wash funnel / glass rod / beaker with deionised water into volumetric flask
- add distilled water up to mark
- shake solution / mix thoroughly
155
C4H3KO8.nH2O + 3NaOH —> C4Na3KO8.nH2O + 3H2O
- student weighs 1.78g of the solid and transfers to 250cm^3 beaker
- she uses 25.0cm^3 samples of the potassium tetraoxalate
- using sodium hydroxide of concentration 0.100moldm^-3
- mean titre of NaOH used = 21.00cm^3
calculate the relative molecular mass of C4H3KO8.nH2O and hence the value of n. you must show clearly how you obtained your answer [5]
- n(NaOH) = 21.00/1000 x 0.1 =0.0021 2.1x10^-3
- n(C4H3KO8.nH2O) = 2.1x10^-3 / 3 = 7x10^-4
- in 250cm^3 = 7x10^-3
- Mr C4H3KO8.nH2O = 1.78 / 7x10^-3 = 254.29
- nH2O = 36.16
- n = 2
156
metal Y reacts with cold water according to the following equation:
Y (s) + 2H2O (l) —> Y(OH)2 (aq) + H2 (g)
when 2.27g of metal Y are added to 600cm^3 of water, the concentration of the metal hydroxide solution formed on complete reaction is 0.0431moldm^-3. identify metal Y [3]
- n[Y(OH)2] in 600cm^3 = 600/1000 x 0.0431 = 0.02586
- 1:1 ratio
- so n(Y) = 0.02586
- Ar = 2.27/0.02586 = 87.8
- Y = strontium
157
- U = CnH2n-2
- a 3.50g sample of U was treated with 10cm^3 of liquid bromine (an excess) at 20°C, the reaction being carried out in a fume cupboard
- hazard = ________
- risk = _______
- control measure = reaction must be carried out in the fume cupboard. wear eye protection and protective gloves [1]
- hazard = bromine is very toxic OR bromine is corrosive
- risk = bromine produces a vapour which is very toxic IF INHALED
OR bromine causes severe burns to the eyes and skin if spilt onto SKIN or EYES
158
describe a test that would give a positive result for butan-2-ol but not butan-1-ol. give the name of the reagent(s) used and the observation(s) made [2]
reagent: (aqueous) iodine AND (aqueous) sodium hydroxide
- observation: yellow crystalline solid formed
159
- a student was asked to identify four inorganic salts, labelled P, Q, R and S, by their reaction with aqueous sodium hydroxide and by the interactions between their solutions. the student was told that each solution contains:
• a different cation : Mg2+, Ba2+, Pb2+ or Fe2+
• a different anion: I-, Cl-, NO3-, SO4 2-
- the student tested approximately 2cm^3 at a time of each solution with NaOH (aq) and then with a few drops of the other solutions in turn.
- when no observable reaction happened, ‘NOR’ was recorded in the table of results:
NaOH (aq):
- with P = white ppt, insoluble in excess NaOH
- with Q = white ppt, soluble in excess
- with R = green ppt, insoluble in excess
- with S = NOR
S:
- with P = NOR
- with Q = white ppt
- with R = heavy white ppt
R:
- with P = NOR
- with Q = white ppt
Q:
- with P = bright yellow ppt
- identity the cations and anions present in each of the four inorganic salts P, Q, R and S
- give an ionic equation to show the formation of ANY ONE of the precipitates formed [10]
P = Mg2+ and I-
Q = Pb2+ and NO3 -
R = Fe2+ and SO4 2-
S = Ba2+ and Cl-
- e.g Pb2+ (aq) + 2I- (aq) —> PbI2 (s)
160
- the reaction between solid barium hydroxide and solid ammonium thiocyanate is endothermic.
- the student measured a 10.0°C temperature change when using 50g of water
- deduce the maximum temperature change if the experiment were repeated using 100g of water rather than the 50g used in the first experiment, with all other factors kept the same. give a reason for your answer [1]
- 5.0°C
- because the same volume of heat is produced but dissipated through twice the volume
161
- a student carried out the following experiment to determine the % of iron in wire
- step 1: preparation of a solution of iron (II) ions:
- the wire was placed in a conical flask and excess sulfuric acid was added.
- a diagram shows the apparatus used. a piece of rubber tubing with a slit in it lets hydrogen escape but stops any air entering the flask
suggest why it was necessary to prevent any air from entering the flask [1]
prevents the oxidation of Fe2+ to Fe3+
162
name: (CH3)3COH
methylpropan-2-ol
163
name: CH3CH2CH2CH2CN
pentanenitrile
164
CH3CH2CH2CH2CN shows optical isomerism.
give one reaction common to both optical isomers. give the reagent and the structural formula of the organic product formed [2]
- (reduction) LiAlH4 to form CH3CH(CH2NH2)CH2CH3
- (hydrolysis) HCl to form CH3CH(COOH)CH2CH3
- (hydrolysis) NaOH to form CH3CH(COONa)CH2CH3
165
state how an ester is separated from the reaction mixture in esterification [1]
distillation
166
50cm^3 of barium chloride solution was transferred into a 250cm^3 beaker and 50cm^3 of 0.506moldm^-3 sodium carbonate solution (an excess) was added
describe how the student could have confirmed experimentally that all of the barium ions had been precipitated [1]
add further sodium carbonate solution to the filtrate and no precipitate will form
167
write an ionic equation for the reaction of carbonate ions with hydrogen ions (H+) , to form carbon dioxide as one of the products [1]
2H+ + CO3 2- —> H2O + CO2
168
- the oxides and chlorides of two elements, X and Y, which exhibit the following properties:
X:
- properties of oxide:
• white solid of melting temperature 2800°C
• it is insoluble in water but readily dissolves in dilute acid
• addition of aqueous sodium hydroxide to this solution forms a white ppt, which is insoluble in excess aqueous sodium hydroxide
- properties of chlorides:
• white solid with melting temperature of 712°C
• it is readily soluble in water
• its solution gives a white ppt with CO3 2- (aq) but no precipitate with SO4 2- (aq)
use all the information provided to suggest the identity of element X. show your reasoning and include ionic equations to support your answer [4]
- high melting temperature of oxide and chloride —> ionic compounds
- basic oxide (not amphoteric) —> oxide of metal
- insoluble carbonate —> not group 1 metal
- hydroxide and carbonate are both white —> not transition metal
- white ppt with hydroxide which is insoluble in excess —> Mg
- no precipitate with sulfate —> Mg
Mg2+ (aq) + 2OH- (aq) —> Mg(OH)2 (s)
Mg2+ (aq) + CO3 2- (aq) —> MgCO3 (s)
169
- the oxides and chlorides of two elements, X and Y, which exhibit the following properties:
Y:
- properties of oxides:
• white solid of melting temperature 1750°C
• it is insoluble in water and does not react with dilute acid or dilute alkalis
- properties of chlorides:
• colourless liquid with boiling temperature of 58°C
• it reacts vigorously with water to give a white ppt, an acidic solution and misty fumes
• at 60°C and 1 atm, 5.000g of the chloride of Y occupies a volume 805.5cm^3
use all the information provided to suggest the identity of element Y. show ur reasoning and include an equation to support your answer [6]
- no reaction with acid/alkalis —> not a metal oxide
- high melting temperature of oxide —> oxide has giant molecular structure
- low boiling temperature of chloride —> chloride has simple covalent structure
- oxide has giant molecular structure and chloride has simple covalent structure —> group 4 element
- misty fumes and white ppt with water —> Si
chloride Y:
- 333K and 1.01x10^5 Pa
- from PV=nRT
- n = 1.01x10^5 x 805.5x10^-6 / 8.31 x 333 = 0.0294
- Mr = 5.000 / 0.0294 = 170.1
—> chloride is SiCl4 / Y is silicon
SiCl4 + 2H2O —> SiO2 + 4HCl
SiCl4 + 4H2O —> Si(OH)4 + 4HCl
170
outline a suitable laboratory method to investigate the rate of each of the following reaction at constant temperature:
Mg(s) + H2SO4 (aq) —> MgSO4 (aq) + H2 (g) [1]
- use gas syringe to measure increase in gas volume over time
- measure loss in mass over time
171
outline a suitable laboratory method to investigate the rate of each of the following reaction at constant temperature:
CH3COCH3 (aq) + I2 (aq) —> CH2ICOCH3 (aq) + HI (aq) [2]
[1] =
- colorimetry / follow change in colour over time
- sample at intervals, quench and titrate
[2]=
- follow change in colour DUE TO CHANGE IN IODINE CONCENTRATION over time
- sample at intervals, quench and titrate (iodine) AGAINST SODIUM THIOSULFATE SOLUTION
172
write the equation that represents the molar enthalpy change of combustion of propane-1-ol [2]
C3H8O (l) + 4 1/2 O2 (g) —> 3CO2 (g) + 4H2O (l)
173
- phenyl benzoate can be prepared from phenol and benzoyl chloride
- step 1: 28cm^3 of 1 moldm^-3 sodium hydroxide and 1.5g of phenol were added to a 50cm^3 conical flask
suggest the purpose of the sodium hydroxide solution [1]
- reacts with phenol to produce the (nucleophilic) phenoxide ion
- removes the proton from phenol
174
- phenyl benzoate can be prepared from phenol and benzoyl chloride
- step 1: 28cm^3 of 1 moldm^-3 sodium hydroxide and 1.5g of phenol were added to a 50cm^3 conical flask
- step 2: in a fume cupboard, 3.0cm^3 of benzoyl chloride was added to the conical flask. a well-fitting cork was placed on top of the conical flask, and the flask shaken vigorously at intervals for 15minutes
name another reagent that could have been used in place of benzoyl chloride (step 2) to prepare phenyl benzoate [1]
give an advantage of using benzoyl chloride in this reaction [1]
- benzoic anhydride (accept benzoic acid)
- benzoyl chloride more reactive / not an equilibrium process
175
a student is asked to identify the contents of six test tubes containing aqueous solutions of the following six inorganic compounds:
- barium nitrate
- potassium iodide
- aluminium sulfate
- sodium hydroxide
- iron (II) nitrate
- lead (II) nitrate
- one of the solutions has a pale green colour. no two test tubes contain the same compound.
- devise a scheme whereby each of the unlabelled solutions could be positively identified
- no other reagents are available
- you should include the observations for any reaction that enables you to identify each of the solution. explain your reasoning [6QER]
- Fe(NO3)2 identified from pale green solution
- Fe(NO3)2 added to other five solutions
- pale green ppt formed identifies NaOH
- NaOH added to remaining four solutions:
• two solutions give white ppt which are soluble in excess NaOH indicating Pb(NO3)2 and Al2(SO4)3
• two solutions give no observable reaction indicating Ba(NO3)2 and KI
- take one of the solutions which gave a white ppt with NaOH and add in turn to both solutions that gave no observable reaction with NaOH:
• if white ppt formed from Al2(SO4)3 and Ba(NO3)2 both identified
• if bright yellow ppt formed then Pb(NO3)2 and KI both identified
176
- a student was asked to carry out a chemical test on each set of compounds that would allow him to distinguish between the compounds within each set.
- for each set of compounds:
• describe one chemical test the student could use
• give the relevant observations that would allow the student to distinguish between the compounds
• give an equation for one reaction that gives an observable change
SET 1:
A = barium nitrate
B = aluminium nitrate
C = magnesium nitrate [4]
- dropwise addition of NaOH (aq), then in excess
- A = no ppt forms
- B = white ppt forms, dissolves in excess
- C = white ppt forms, does not dissolve in excess
Mg(NO3)2 + 2NaOH —> Mg(OH)2 + 2NaNO3
Al(NO3)3 + 3NaOH —> Al(OH)3 + 3NaNO3
Al(OH)3 + NaOH —> Na[Al(OH)4]
177
- a student was asked to carry out a chemical test on each set of compounds that would allow him to distinguish between the compounds within each set.
- for each set of compounds:
• describe one chemical test the student could use
• give the relevant observations that would allow the student to distinguish between the compounds
• give an equation for one reaction that gives an observable change
SET 2:
D = phenol
E = ethanol
F = cyclohexene [4]
- addition of Br2 (aq) (room temperature)
- D = Br2 decolourised and white ppt forms
- E = no observable change
- F = Br2 decolourised
C6H5OH + 3Br2 —> C6H2Br3OH + 3HBr
C6H10 + Br2 —> C6H10Br2
178
- a student was asked to carry out a chemical test on each set of compounds that would allow him to distinguish between the compounds within each set.
- for each set of compounds:
• describe one chemical test the student could use
• give the relevant observations that would allow the student to distinguish between the compounds
• give an equation for one reaction that gives an observable change
SET 3:
G = chlorobenzene, C6H5Cl
H = (chloromethyl)benzene, C6H5CH2Cl [3]
- addition of NaOH (aq), gentle heat, followed by AgNO3 (aq)
- G = no observable change
- H = white ppt forms
AgNO3 + Cl- —> AgCl + NO3 -
Ag+ + Cl- —> AgCl
179
- a student was asked to carry out a chemical test on each set of compounds that would allow him to distinguish between the compounds within each set.
- for each set of compounds:
• describe one chemical test the student could use
• give the relevant observations that would allow the student to distinguish between the compounds
• give an equation for one reaction that gives an observable change
SET 4:
I = methyl ethanoate, CH3COOCH3
J = ethanamide, CH3CONH2 [3]
- addition of NaOH (aq), gentle heat and damp red litmus paper
- I = no observable change
- J = pungent smelling gas / turns damp red litmus paper blue
CH3CONH2 + NaOH —> CH3COONa + NH3
180
give the equation for the reaction of magnesium hydroxide with hydrochloric acid [1]
Mg(OH)2 + 2HCl —> MgCl2 + 2H2O
181
- a solution was transferred quantitatively into a 250cm^3 volumetric flask and was made up yo the mark with the same hydrochloric acid
describe how the solution was transferred quantitatively into the volumetric flask [1]
state why the solution is transferred in this way [1]
- using funnel
- rinse funnel / glass rod / beaker with HYDROCHLORIC ACID ensuring that washings go into the flask
- everything that was in the beaker hoes into the flask
- no analyte is left in the funnel/in beaker
- no analyte is lost in the process
182
in an experiment to determine ∆H, a student used 30.0cm^3 of 2.00moldm^-3 hydrochloric acid and 2.29g of powdered anhydrous potassium carbonate
outline a method to determine the temperature change during the reaction [2]
- add the HCl to the insulated polystyrene cup
- measure the temperature at regular intervals e.g every 30s
- when the temperature of the acid remains constant, at the next time interval, add the powdered potassium carbonate
- stir the reaction mixture
- continue to record temperature at regular time intervals
- plot a temperature-time graph to determine the temperature change
183
what are the units of ∆H in ∆H=-mc∆T / n?
J mol^-1
(so might need to convert to/from kJmol^-1)
184
2KHCO3 (s) —> K2CO3 (s) + CO2 (g) + H2O (g)
suggest one reason why the enthalpy change for this reaction cannot be determined directly by calorimetry [1]
- the reaction needs large amount of heat so ∆T cannot sensibly be measured
(accept any reference to difficulty of measuring temperature change due to reaction when strong heating is required for the reaction to take place)
*answer will be specific to different equations / no 1 right answer
185
- both propanedioic acid and compound X have the same molecular formula. some of the properties of compound X are given below:
• compound X doesnt show optical or geometric isomerism
• the simplified 1H NMR spectrum of compound X shows three separate peaks in the area ratio 1:1:2
• a solution, containing 0.704g of compound X, on reaction with excess sodium carbonate, gives 83.0cm^3 of a colourless gas at 25°C and 1 atm pressure
• compound X doesnt react with 2,4-DNPH
• compound X decolourises aqueous bromine to form compound Y which had a chiral carbon
use all the information to identify and give the structures of compound X and compound Y. show the chiral carbon in compound Y [6]
- no reaction with 2,4-DNPH —> no carbonyl group
- aqueous bromine decolorises —> X contains a C=C bond
- reacts with sodium carbonate to form colourless gas —> carboxylic acid group present
- n(CO2) = (83.0x10^-3) / 24.5 = 0.00339mol
- n(X) = 0.704/104.04 = 0.00677 mol
- ratio CO2 : X —> 1:2 therefore one carboxylic acid group present
- X doesnt show optical isomerism because none of the carbon atoms is bonded to four different groups / is no chiral carbon
- X doesnt have geometric isomerism because one carbon of the C=C group has two -OH groups attached
- 1H NMR spectrum of compound X would show three peaks —> carboxylic acid -OH, -CH=C, alcohol -OH
186
for ∆H, if temperature decreases, it’s -temperature change
187
∆T = -9.8
give ONE assumption made when finding the value of ∆H from experimental results [1]
- no heat is gained (bc endothermic)
- solid is completely dissolved
- assume that dissolving a salt in water adds no mass
188
Na2S2O3 (s) + 5H2O (l) —> Na2S2O3.5H2O (s)
suggest one reason why the enthalpy change for this reaction cannot be determined directly [1]
- it’s not possible to prevent some of the sodium thiosulfate from dissolving other hydrates / side-products may form
- solid is likely to dissolve in water - when water added to sodium thiosulfate, will dissolve
189
- this question considers two different methods of preparing soluble salts
- in this method, an insoluble solid is added to a suitable solution to form a soluble salt
- one example is the reduction of excess copper (II) carbonate with hydrochloric acid to form copper (II) chloride
CuCO3 (s) + 2HCl (aq) —> CuCl2 (aq) + CO2 (g) + H2O (l)
state why this method uses excess copper (II) carbonate and not excess hydrochloric acid [1]
- excess insoluble reactants can be separated by filtration
- excess acid cannot easily be removed
- if acid in excess, it will contaminate the salt solution
190
- this question considers two different methods of preparing soluble salts
- in this method, an insoluble solid is added to a suitable solution to form a soluble salt
- one example is the reduction of excess copper (II) carbonate with hydrochloric acid to form copper (II) chloride
CuCO3 (s) + 2HCl (aq) —> CuCl2 (aq) + CO2 (g) + H2O (l)
describe the practical steps for making pure, crystalline copper (II) chloride from solid copper (II) carbonate and hydrochloric acid [3]
- add excess CuCO3 to the acid, a little at a time, with stirring
- continue to add CuCO3 until no more fizzing / solid remains
- filter excess unreactive CuCO3
- evaporate the filtrate (to reduce volume) / allow to crystallise
191
in a titration, state why it’s not necessary to fill the burette to exactly zero [1]
volume used from burette is measured by difference
192
- titration of aqueous sodium carbonate with hydrochloric acid to form sodium chloride
the teacher said that the method would not give pure sodium chloride
state how the student should adapt the method in order to get pure solid sodium chloride [1]
- repeat the experiment without the indicator
193
- transition metals display basic or amphoteric behaviour
- describe how you would use sulfuric acid and sodium hydroxide solution to determine whether iron (II) oxide and chromium (III) oxide are basic or amphoteric
- include observations and equations in your answer [6QER]
ACID/BASE:
- Fe2+ ions react with acids
- iron (II) oxide is a basic oxide
- Cr3+ ions react with both acids and bases
- chromium (III) oxide is an amphoteric oxide
RELEVANT EQUATIONS AND OBSERVATIONS:
- iron (II) oxide reacts with sulfuric acid
- FeO + H2SO4 —> FeSO4 + H2O
- pale green solution formed
- iron (II) ions formed react with a base / aqueous sodium hydroxide added dropwise
- Fe2+ (aq) + 2OH- (aq) —> Fe(OH)2 (s)
- green ppt formed
- green ppt of Fe(OH)2 doesnt react / dissolve in excess aqueous sodium hydroxide
- chromium (III) oxide reacts with sulfuric acid
- Cr2O3 + 3H2SO4 —> Cr2(SO4)3 + 3H2O
- violet-blue / green solution formed
- chromium (III) ions reacts with a base / sodium hydroxide added dropwise
- Cr3+ (aq) + 3OH- (aq) —> Cr(OH)3 (s)
- grey-green ppt formed
- chromium (III) hydroxide ppt dissolves in excess aqueous sodium hydroxide
- Cr(OH)3 + 3OH- (aq) —> [Cr(OH)6] 3-
- deep green solution formed
194
if naming homologous series for different functional groups, be specific,
e.g primary amine, secondary alcohol
195
explain how a loss of solution is minimised when transferring it to a volumetric flask [2]
- add water to beaker
- add washings to volumetric flask
196
explain the purpose of inverting the volumetric flask [1]
ensure the conc is the same throughout
197
in calorimetry, why do you take the temperature 4 times before continuing the reaction?
- to allow the temperature to equibilate / ensure constant temperature
198
for ∆H = -mc∆T / n, if one thing in excess, use the moles of the thing that isnt in excess, but otherwise, the n = the solid
199
what is a hydrated salt?
- one that has coordinate bonds that form with ionic lattice
200
what are some improvements to do in a calorimetry practical?
- crush into powder —> larger s.a —> larger ∆T
- insulate —> prevent heat loss / gain
- record many temperatures at the start —> ensure constant —> so temp change only due to reaction
- plot + extrapolate —> gives max value of ∆T
- use larger mass —> lower error
201
what assumption does ∆H = -mc ∆T / n assume?
that adding the solid doesnt add any mass
202
ester + H+ —> ____ + _____
acid hydrolysis:
carboxylic acid + alcohol
203
ester + excess OH - —> ____ + _____
carboxylic acid BUT O- and Na+ (the salt) and alcohol (not salt bc not acidic)
204
if both OH and COOH in a compound, and making an ester, see what the other compound has then can see where joins/where reacts
205
explain why the funnel is removed from the burette before the initial burette reading is taken [1]
- small drops of liquid may fall from the funnel (into the burette) during the titration (leading to a false burette reading)
206
in a redox titration, explain why the addition of boiling water to the conical flask, to maintain a temperature of 70°C, doesnt have an effect on the titration results [1]
- water doesnt react with any of the reagents
- adding water doesnt change the number of moles of reagents added
207
when potassium manganate is used in a titration (and so dont need an indicator) what colour does it turn when the end-point is reached?
purple —> pink
a level chemistry
flashcards
Decks in class (21)
# Cards
-
3.1 redox and standard electrode potential74
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3.2 redox reactions30
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3.3 chemistry of the p-block102
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3.4 chemistry of the d-block transition metals132
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3.5 chemical kinetics53
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3.6 enthalpy changes for solids and solutions53
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3.7 entropy and feasibility of reactions41
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3.8 equilibrium constants26
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3.9 acid-base equilibria62
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4.1 stereoisomerism33
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4.2 aromaticity37
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4.3 alcohols and phenols93
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4.4 aldehydes and ketones72
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4.5 carboxylic acids and their derivatives82
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4.6 amines87
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4.7 amino acids, peptides and proteins51
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4.8 organic synthesis and analysis145
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equations37
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unit 5 practical207
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unit 3127
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unit 4117
