unit 5 practical Flashcards
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 tips to look out for when doing double titration?
A
- double check the equations when doing your calculations
7
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
8
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
9
Q
what are some common mistakes when doing redox titrations?
A
- forgetting to add acid if required
10
Q
what are tips to look out for when doing redox titrations?
A
- add water from a kettle if your titration has to be carried out hot
11
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
12
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
13
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
14
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)
15
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
16
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
17
Q
what are some common mistakes when using electrochemical cells?
A
- forgetting to soak your salt bridge
- putting the cells the wrong way round
18
Q
what are some common mistakes when doing melting point experiments?
A
-forgetting to seal the base of the tube
19
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)
20
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
21
Q
what are some common mistakes when doing esterification?
A
- forgetting to add the concentrated acid
- forgetting to distil off the ester
22
Q
what are tips to look out for when doing esterification?
A
- use the acid anhydride for a larger yield
23
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
24
Q
what are some common mistakes when doing organic synthesis?
A
- not refluxing for long enough
25
what are tips to look out for when doing organic synthesis?
- follow the instructions exactly
- record all distillation temperatures
26
what is a standard solution?
- one whose concentration and volume is known exactly
- this can then be used to calculate the concentration of other solutions
27
how to make a standard solution?
- 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:
1. weigh out your solid (see weighing by difference)
2. carefully add a small amount of distilled water to the beaker, use a glass rod to stir until all the solid has dissolved
3. use a funnel to transfer the liquid into the volumetric flask, pouring down the glass rod
4. rinse out the beaker and funnel several times with distilled water. wash the glass rod. ensure that all washings go into the volumetric flask
5. top up to the mark on the volumetric flask
6. put the stopper in firmly and thoroughly mix the solution
7. label the flask with your name
28
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
29
what is the test for ammonia (gas)?
- test with red litmus paper
- +ve result = litmus paper turns blue
30
what is the test for carbon dioxide?
- bubble through limewater
- +ve result = limewater turns milky as calcium carbonate is made
31
what is the test for hydrogen?
- test with a lighted splint
- +ve result = gas burns with a squeaky pop
32
what is the test for oxygen?
- test with a glowing splint
- +ve result = splint relights
33
what is the test for ammonium NH4+?
- heat with NaOH, test vapours with red litmus
- +ve result = litmus turns blue (from evolved ammonia gas)
34
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
35
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
36
what are the tests for copper (II) Cu2+?
- flame test
- +ve result = blue-green flame
- add NaOH
- +ve result blue ppt
37
what is the test for iron (II) Fe2+?
- add NaOH
- +ve result = dark green ppt
38
what is the test for iron (III) Fe3+?
- add NaOH
- +ve result = brown ppt
39
what is the test for potassium K+?
- flame test
- +ve result = lilac flame
40
what is the test for sodium Na+?
- flame test
- +ve result = yellow-orange flame
41
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
42
what is the test for aluminium Al3+?
- add NaOH
- +ve result = white pot that is soluble in excess hydroxide
43
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
44
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
45
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
46
what are the tests for lead (II) Pb2+?
- add NaOH
- +ve result = a white ppt forms which is soluble in excess NaOH but insoluble in excess NH3
- add KI
- +ve result = yellow ppt forms
47
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
48
what is the test for bromide Br-?
- add nitric acid + silver nitrate
- +ve result = cream ppt (of silver bromide)
49
what is the test for carbonate CO3 2-?
- add acid
- +ve result = fizz of co2 produced (can be tested with lime water)
50
what is the test for chloride Cl-?
- add nitric acid + silver nitrate
- +ve result = white ppt of silver chloride produced
51
what is the test for iodide I-?
- add nitric acid + silver nitrate
- +ve result = pale yellow ppt
52
what is the test for sulphates SO4 2-?
- add solution of barium chloride
- +ve result = white ppt of barium sulphate
53
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
54
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
55
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
56
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
57
mostly all group 1s are soluble
58
all nitrates are soluble
59
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
60
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
61
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
62
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
63
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
64
why is CuI (+1) a white ppt?
Cu+ = d^10 = no free d orbitals to promote electrons
65
explain the observation made in test 5. include an equation in your answer [2] (addition sodium thiosulfate solution to CuI)
2Cu2+ + 4I- —> 2CuI + I2
2S2O3 2- + I2 —> S4O6 2+ + 2I
(iodine is used up so colourless)
66
state the feature required for a molecule to show optical activity [1]
chiral centre
67
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
68
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:
- iron (III) chloride - purple solution
- bromine water - white ppt
- iodoform - iodine + sodium hydroxide - yellow ppt of CHI3
69
what are the reagents when using iodoform?
iodine + sodium hydroxide
70
what is the end product when using the iodoform test?
- yellow ppt of CHI3
71
is PbCl2 soluble or insoluble?
insoluble
72
what are the only two soluble compounds of lead
- Pb(NO3)2
- Pb(CH3COO)2
73
- 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
- use nitric acid instead
- as Pb(NO3)2 is soluble
74
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
75
why are the 4 ions that redissolve in excess NaOH the only ones that do?
because they are amphoteric
76
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-
77
what is the enthalpy change equation?
∆H = -mc∆T / n
∆H = joules
m = g
78
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
79
% error equation:
(uncertainty / measurement obtain) x100
(if two measurements, times uncertainty by 2)
80
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)
81
burettes measure to +- 0.05
82
the error = 1/2 of the smallest increment you have on the instrument
83
to what d.p do you read titrations to?
2 d.p
84
when weighing solid, for % error calculation, uncertainty x2
85
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
86
- in the practical:
- making standard solution
- MAKE SURE MIXED
87
what unit should it be when measuring for % error?
g
88
500mg = ___g
0.5g
89
what has occurred if a compound of CHI3 has been formed?
iodoform test
90
what is produced when tollens or fehling is added to an aldehyde?
carboxylic acid
(aldehyde has been oxidised)
91
what is observed when Ba(NO3)2 is added to Pb(NO3)2?
no reaction observed - solution remains colourless
92
what is observed when Ba(NO3)2 is added to MgSO4?
a white ppt is produced
93
what causes the white precipitate in the reaction between Ba(NO3)2 and MgSO4?
- the sulfate ions
- BaSO4 is a white ppt
94
what is observed when Ba(NO3)2 is added to KI?
- no reaction observed - solution remains colourless
95
what is observed when Ba(NO3)2 is added to Na2CO3?
- a white ppt is produced
96
what causes the white precipitate in the reaction between Ba(NO3)2 and Na2CO3?
- the carbonate ions
- BaCO3 is a white precipitate
97
what is observed when Pb(NO3)2 is added to MgSO4?
a white precipitate is produced
98
what causes the white precipitate in the reaction between Pb(NO3)2 and MgSO4?
- the sulfate ions
- PbSO4 is a white precipitate
99
what is observed when Pb(NO3)2 is added to KI?
- a yellow precipitate is produced
100
what causes the yellow precipitate in the reaction between Pb(NO3)2 and KI?
- the iodide ions
- PbI2 is a yellow precipitate
101
what is observed when Pb(NO3)2 is added to Na2CO3?
- a white ppt is produced
102
what causes the white precipitate in the reaction between Pb(NO3)2 and Na2CO3?
- the carbonate ions
- PbCO3 is a white precipitate
103
what is observed when MgSO4 is added to KI?
- no reaction observed - solution remains colourless
104
what is observed when MgSO4 is added to Na2CO3?
- a white ppt is produced
105
what causes the white precipitate in the reaction between MgSO4 and Na2CO3?
- the carbonate ions
- MgCO3 is a white precipitate
106
what is observed when MgSO4 is added to Zn(NO3)2?
- no reaction observed - solution remains colourless
107
what is observed when KI is added to Zn(NO)3)2?
no reaction observed - solution remains colourless
108
what is observed when KI is added to Na2CO3?
no reaction observed - solution remains colourless
109
what is observed when Zn(CO3)2 is added to Na2CO3?
- a white precipitate is produced
110
what causes the white precipitate in the reaction between Zn(CO3)2 and Na2CO3?
- the carbonate ions
- Zn(CO3)2 is a white ppt
111
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
112
what are the hazards associated with HCl, NaOH and phenolphthalein?
- HCl = irritant
- NaOH = irritant
- phenolphthalein = flammable
113
what is the chemical equation for the reaction between NaOH and HCl?
HCl + NaOH —> NaCl + H2O
114
what type of reaction takes place between HCl and NaOH?
neutralisation
115
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
116
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
117
why are titrations usually carried out on a white tile?
- allows the point of colour change to be easily identified
118
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
119
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
120
why are burettes and pipettes always used in titrations?
- burettes and pipettes measure the volumes of solutions very precisely
121
what is a standard solution?
a solution with a known concentration
122
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
123
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
124
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
125
when making up a standard solution, why is it important to not add more than 150cm^3 of deionised water when 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
126
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
127
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
128
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
129
what is the colour change of methyl orange?
- red in acid
- yellow in alkali
130
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
131
what is the chemical equation for the reaction between Na2CO3 and HCl?
- 2HCl + Na2CO3 —> 2NaCl + H2O + CO2
132
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
133
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
134
what reagents can be used to distinguish between aldehydes and ketones?
- tollens’ reagent
- fehling’s reagento
135
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
136
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
137
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
138
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
139
how is litmus paper used to test for acids?
- blue litmus paper will turn red when exposed to an acid
140
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
141
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
142
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
143
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
144
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
145
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
146
how could you reduce the uncertainty in the mass measurement?
- use a digital balance with a greater resolution
- use a larger mass
147
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
148
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
149
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
150
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
151
- 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 chemistry 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
152
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
153
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
154
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
155
- 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)
156
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
157
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
158
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
159
why did the student repeat the titration? [1]
to obtain concordant results (to increase reliability?)
160
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
161
state why no indicator is required in this titration [1]
- potassium manganate (VII) has a deep (purple) colour which disappears on reactio. and when the end-point is reached a single drop more turns the solution pink
neutral answer - it changes colour by itself
162
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)
163
- 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
164
- 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
