Determination of the formula of Hydrated Iron (II) Sulphate Crystals Essay Sample

Determination of the formula of Hydrated Iron (II) Sulphate Crystals Pages
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Analysis

Method one.

In order to find out the ratio of H2O to FeSO4, the mass of water needs to be calculated, which is done by weighing the iron sulphate before and after the heating. This will tell us the weight that was lost and thus the weight of the water.

Grams (g)

Empty crucible 12.01

Crucible and Iron (II) sulphate crystals 13.48

Mass of Iron (II) sulphate crystals (2-1) (13.48-12.01) 1.47

Crucible and Anhydrous Iron (II) sulphate 12.82

Mass of Anhydrous Iron (II) sulphate (5-1) (12.82-12.01) 0.81

Mass of water (13.48-12.82) 0.66

Repeat

7. Crucible and Anhydrous Iron (II) sulphate (after re-heating) 12.80

8. Mass of Anhydrous Iron (II) sulphate (after re-heating) (7-1) (12.80-12.01) 0.79

9. Mass of water lost during repeat (4-7) (12.82-12.80) 0.02

10. Mass of water lost altogether (2-7) (13.48-12.80) 0.68

Finding the molar ratio between FeSO4 and H2O requires finding the empirical formula using their masses. To find the number of moles of each of the compounds, the equation

moles= mass/molar mass is used.

Finding the Atomic masses of FeSO4 and H2O:

Atomic Masses Fe= 55.8 S=32.1 O4= 16×4

H2=1×2 O=16

Molar mass of FeSO4= 55.8 + 32.1 + (16×4)

= 151.9

Molar mass of H2O = (1×2) + 16

=18

FeSO4 H2O

Mass 0.79g 0.68

Moles 0.79/151.9= 0.005 0.68/18=0.038

Finding the Ratio: 0.005/0.005 0.038/0.005

=1 =7.4

Ratio of FeSO4 to H2O is: 1:7

Therefore, the formula for Iron(II) Sulphate in this experiment is;

FeSO4.7H2O

Thus, the overall equation for the reaction was:

FeSO4.7H2O(s) -> FeSO4(s) + 7H2O(l)

Observations and comments

As heated, the pale green crystals became whiter and joined together to form one block, which became grey around the sides of the crucible. This was because this part of the FeSO4 closest to the heat and so was beginning to be decomposed.

When the ‘anhydrous’ iron sulphate taken out of the crucible after procedure completed; red on the underneath of the iron sulphate. This suggests that some of the Fe2+ had been decomposed to iron oxide and oxidised to fe3+. When the substance turns red, this is what has taken place:

Fe(II)SO4.7H20(s) -> FeSO4(s) + H2O(l)

2Fe(II)SO4(s) -> Fe2(III)O3(s) + SO2(g) + SO3(g)

The Fe2O3 is a red solid.

Method two

Weight of FeSO4.xH2O: 2.92g

Results of Titre: Rough Titre 1 Titre 2 Titre 3

First reading (cm) 0.00 0.00 0.00 0.00

Second reading (cm) 28.90 28.45 28.35 28.40

Titre (cm) 21.10 21.55 21.65 21.60

Chosen titres to find average titre from are: 21.55cm, 21.65 cm and 21.60cm

To find the average titre of KMnO4 used to oxidise the Fe2+ ions, add the three titres together that are within 0.10cm difference of each other and divide by three:

Average titre:= (21.55 + 21.65 + 21.60)/ 3

= 21.60cm

Therefore, it requires an average volume of 21.60cm of 0.01 mol dm- KMnO4 to oxidise all the Fe2+ ions in 25cm of FeSO4.xH2O (dissolved in 1mol dm- concentration of H2SO4).

To find the ratio of FeSO4 to H2O requires firstly to calculate the number of moles of KMnO4 required to oxidise the Fe2+ ions. In order to calculate this, the equation

no. of moles= (volume in cm3/1000) x concentration

is used.

No. of moles of KMnO4 = (21.60/1000) x 0.01

= 0.0002 moles

Redox half equations that occur during the reaction:

MnO4-+ 5e- + 8H+ -> Mn 2+ + 4H2O

Fe2+ -> Fe3+ + e-

Full balanced ionic half equation:

MnO4-(aq) + 5Fe2+(s) + 8H+ -> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)

There are 5 Fe2+ ions to every one MnO4- ion in order to balance the number of electrons required by for each Mn2+ to form. The full ionic equation shows that to every 1 mole of MnO4- ions, there were 5 moles of Fe2+ ions. Therefore, to calculate the number of moles of Fe2+ ions; multiply the number of moles of KMnO4 by 5:

No. of moles of Fe2+ ions = 0.0002 x 5

= 0.001moles

However, that is only the number of moles of Fe2+ ions in 25cm but that is only a 10th of the number of ions in the total 250cm (250/25=10). To find the number of moles of Fe2+ ions in the whole of the solution multiply the number of moles of Fe2+ in 25cm by 10.

No. of moles of Fe2+ in 250cm= 0.0010825 x 10

= 0.010 moles

0.01 moles of Fe2+ means the same as saying there is 0.01 moles of FeSO4. Therefore, in order to find the mass of the FeSO4, you can use the equation;

mass= no. of moles x molar mass.

Mass of FeSO4= 0.01 x 151.9

= 1.52g

As the total mass of the FeSO4.xH2O was 2.92g in order to find the mass of water, you can subtract the weight of FeSO4:

Weight of water = 2.92-1.52

= 1.40g

Now that I have the mass of both the FeSO4 and H2O, I can find the ratio of the two by finding the number of moles of both.

FeSO4 H2O

Mass (g) 1.52 1.40

Moles= mass/molar mass 1.52/151.9 1.40/18

= 0.01 = 0.08

Ratio= both molar values divided by the smaller molar value of the two

0.01/0.01 0.08/0.01

=1 = 8

1:8

Therefore, the formula for Iron(II) sulphate in this experiment is;

FeSO4.8H2O

Observations and comments

Pale green crystals; dissolved into the solution that remained colourless.

As the Fe2+ in solution being titrated it remained colourless and then when went a pale pink indicating that all the fe2+ had been oxidised by the MnO4- ions.

Evaluation

Percentage errors

Method one

In all, I weighed the crucible/and iron sulphate 4 times therefore, to calculate the total measurement percentage error for method one, the percentage error must be found for all four of these and added together. To calculate the percentage measurement error, you use number that is half the smallest calibration/measurement that the measuring device can measure for example, for the top pan balance, it can measure down to 0.01g so half of that is 0.005 and you divide this by the measurement read using that device. This is then multiplied by 100 to get a percentage:

For the top pan balance

Measurement uncertainty percentage= (0.005/mean number)x 100

Crucible (alone) percentage error= (0.005/12.01) x 100

= 0.04%

2. Crucible and Iron (II) sulphate = (0.005/13.48) x 100

= 0.03%

3. Crucible and Anhydrous Iron (II) sulphate = (0.005/12.82) x 100

= 0.04%

4. Repeat, crucible and Anhydrous Iron (II) sulphate= (0.005/12.80) x 100

=0.04%

Total percentage error = 0.04 + 0.03 + 0.04 + 0.04

= 0.15%

The total measurement percentage uncertainty is extremely small so any significant discrepancies in the ratio of iron(II) sulphate to water will not be due to measurement errors. However, to reduce the measurement percentage error even further for method one, a more sensitive top pan balance could be used (although the minor percentage uncertainty for this method would suggest that the measurement device used does not need to be changed ).

Procedural uncertainties and improvements:

There are a number of procedural uncertainties for method one including the fact that the procedure of heating the iron (II) sulphate to drive off the water is only done twice. The problem with this is that because it is not reheated and weighed continuously until it reaches a constant mass, there is the possibility that not all the water is driven off. This would mean that the weight that was calculated as the water is not actually the total amount and therefore the ratio of anhydrous iron sulphate to water would not be completely accurate.

Another procedural uncertainty is the problem with the size of the hydrated iron (II) sulphate crystals used; as the crystals were a relative size, the outside of some of the crystals (depending on how close they were to the heat) would have been dehydrated before the inside of the crystals and it is likely that the inner parts of some of the iron sulphate remained hydrated despite the outside of the substance appearing to be dehydrated (shown by the colour change from green to white). Again, this would mean that the weight calculated for the loss of water was not completely accurate as not all the water was driven off, making the calculated mass of water lower than it really was. Following this procedural error, there is the issue around the shape of the crucible as it was quite deep.

This made judging when to stop heating difficult because although the showing part of the substance appeared as a white solid (suggesting anhydrous form), on taking the substance out, the bottom of the solid was a red colour suggesting that some of the iron sulphate had decomposed to iron (III) oxide. As the bottom part of the iron (II) sulphate was closer to the heat, it lost water first and so, became overheated whilst the top still appeared as the anhydrous form. Uneven heating on top was evident also as the iron sulphate around the edges was a darker grey colour than in the centre. This greatly affects the accuracy of the results as the mass being weighed after heating is no longer just the anhydrous iron (II) sulphate as it should be and so, the lost mass calculated as the water is not necessarily correct.

Another procedural uncertainty is the possibility that the anhydrous iron sulphate combined with moisture in the air following the heating process and so, ultimately was not completely dehydrated thus weighing more than the anhydrous iron sulphate would have.

All of the above procedural uncertainties have a detrimental affect on the accuracy of the results. The most significant of the procedural uncertainties however is the uneven heating that occurred and resulted in the part decomposition of the iron (II) sulphate. This is the most important uncertainty because it meant that whilst some of the iron (II) sulphate remained hydrated, other parts decomposed, which meant that the procedure carried out was very inaccurate, failing to find the mass of just anhydrous iron sulphate as should have been done.

In order to alleviate this uncertainty a number of improvements can be made. Firstly, a shallower dish should be used. If this was used together with a powdered form of the iron (II) sulphate, this would ensure that the heat reached all parts of the iron (II) sulphate equally. In addition, to prevent any chance of the iron (II) sulphate combining with any substance that may have already been on the shallow dish including moisture; the dish should be washed and dried by heating it and then cooled in a desiccator so that it doesn’t pick up any moisture from the air. Then, after heating the iron (II) sulphate; place the shallow dish with the anhydrous iron (II) sulphate back into a desiccator to cool; this would ensure that the anhydrous form did not combine with any moisture in the air on cooling.

Before placing the heated evaporating dish in to the desiccator, allow it to cool for one minute and leave the desiccator lid open slightly for 30 seconds once it the hot dish has been placed in side. This is done in order to ensure that as the heated air cools, a partial vacuum is not formed. The procedure of heating the iron (II) sulphate should be done for 2 minutes and the process should be repeated until the substance is a white solid and has reached a constant mass (will not lose anymore mass). Repeating the process until the substance reaches a constant mass ensures that all the water that has combined with the iron (II) sulphate is driven off and therefore, can be calculated as the accurate mass of water.

Moreover, this method should be repeated at least 3 times. Repeating the procedure increases the reliability of the results if they concur and finding an average between the three masses would improve the accuracy of the results. Also, if there is an anomalous result; repeating the method 3 times allows you to identify it and thus, ignore it whereas, with the procedure I completed; because it was only don’t once it is possible that the mass of water I calculated is an anomaly. Yet, I am unable to identify it as one because I have no other answer to compare it to. Additionally, to increase the accuracy of the results even further; another procedural improvement would be to increase the mass of iron (II) sulphate should be used as this reduces the measurement percentage uncertainty. If 50g of iron (II) oxide was used rather that 12.01, for example, the measurement percentage error would be:

(0.005/50) x 100

=0.01%

Whereas, using 12.01g has a measurement percentage uncertainty of 0.04%.

Method 2:

Measurement uncertainties

Percentage uncertainty from weighing the iron sulphate

(0.005/2.92) x 100

=0.17%

Percentage uncertainty from measuring the burette

(0.05/21.65) x 100

=0.23%

Percentage uncertainty from measuring 25.00cm in the volumetric flask

(0.15/25.00) x 100

=0.60%

Percentage uncertainty for measuring 25cm using the pipette

(0.05/25) x 100

=0.20%

Total measurement percentage error

=1.20%

The total measurement percentage uncertainty for method 2 is considerably larger than it is for method one but it is still a very small percentage uncertainty. The greatest percentage uncertainty was caused by the volumetric flask. The percentage uncertainty of the flask would be improved if a more accurate flask was used, which had a smaller margin of error for example, +/- 0.05cm rather than 0.15cm.

The second largest percentage uncertainty was caused by burette, which only has calibrations every 1cm. This meant that you are required to interpolate between calibrations if the level of KMnO4 is between a calibration; giving rise to possibility of inaccuracy caused by human error. The percentage error would be reduced if a more accurate burette was used, which had more precise calibrations; every 0.5cm for example.

However, as the total measurement percentage uncertainty is much smaller than 5%, it is not considered a significant amount and would not have had a noticeable affect on the accuracy of the results.

Procedural Uncertainties and improvements.

The most significant measurement uncertainty concerning method 2 is the transferral of the Fe2+ ions. Firstly, the ions are transferred from the weighing boat to a container to be mixed with the H2SO4 and water and then the ions are transferred to the volumetric flask. In this interlude, some of the Fe2+ ions would have been lost in the containers and on the mixing rod. However, the amount of Fe2+ ions that are lost is minimised by rinsing all the weighing boat, mixing container and mixing rod with de-ionised water so that many of the fe2+ ions that were left behind are washed into the volumetric flask anyway.

The colour of the KMnO4 also gives rise to the problem of misreading the measurement as it is a deep purple colour, which meant that it was difficult to judge where the meniscus is. Additional to this uncertainty is the fact that an average was found only from 3 repeats and therefore the procedure is not very reliable as all three of the measurements could be inaccurate. However, in comparison to method one which only did the procedure once, method two has greater reliability. The possibility of the human error allowing too much KMnO4 into the iron (II) sulphate solution was also solved by using a trial titre to see the approximate volume of KMnO4 needed to titrate the iron (II) sulphate solution. This ensured that the titre was not missed on the repeats.

There are no anomalous measurements in method two as each of the three volumes differs only by 0.05cm. The slight differences between the volumes are not great enough to classify any of them as anomalous. The rough titre does differ from the other three by about 0.5cm but as it is not included in the calculations, this does not matter.

In order to improve the accuracy of the results; rather than only completing 3 repeats, 12 repeats should be completed and an average titre found between the 12. This would greatly increase the reliability as well as accuracy of the experiment . Secondly, once a model titre has been completed where the colour of the solution that has been titrated is a pale pink colour; this should be used as an exemplar titrate, which I would use to compare it’s colour to all the following titres. This would help to ensure I did not add too much or too little KMnO4. Also, if a powdered form of iron (II) sulphate was used then the container used to dissolved all the iron sulphate with the water and acid would not need to be used.

Using a powder form would mean that the acid, water and iron sulphate could be placed directly into the volumetric flask and therefore, the loss of Fe2+ ions resulting from the transferral between the mixing container and the volumetric flask would not occur thus improving the accuracy of the results. If the iron sulphate was weighed directly in the volumetric flask, this would prevent any Fe2+ ions being lost during transferral process. This could be done by weighing the volumetric flask without the iron sulphate and then taking this weight away from the mass of the volumetric flask containing the iron sulphate; this would tell you the weight of the iron (II) sulphate. In addition to these procedural improvements, a greater mass of iron (II) sulphate should be used as this would decrease the percentage uncertainty of using the top pan balance. Also, if a greater volume of iron sulphate in solution was used in a larger volumetric flask, this would reduce the percentage uncertainty of using a volumetric flask. Thus, improving the accuracy of the results obtained.

Which method is most accurate?

Looking at the percentage measurement uncertainty for the two methods it would appear that method one is a more accurate procedure as it’s total measurement percentage uncertainty is only 0.15% compared to method two’s, which is 1.20%. Yet, the parentage uncertainties are both very small (they are both well under 5% and so, not considered significant uncertainties) and would neither significantly affect the accuracy of the results. Therefore, when deciding on which method is more accurate, just because method one has a lower measurement percentage error does not mean that it is more accurate as it’s procedural uncertainties may be much more significant than method two’s.

In method one, the whole procedure is only completed once so there is only one mass of water found. This makes the method very unreliable as there are no other results to compare to, to ensure that the mass of water calculated is not simply an anomaly. Moreover, finding the average between a number of results increases the accuracy of the overall result but you can not do this with method one. Additionally, as the procedure was not repeated until the iron sulphate reached a constant mass, it is quite likely that not all of the water was driven off; making the weight of water calculated not entirely accurate. Conversely, in method two, 4 titrations were completed and an average was found from 3. Not only does this method allow you to identify an anomaly but it allows you to find an average between a number of titres, which means you are more likely to get closer to the actual answer.

There is the procedural uncertainty with method two surrounding the loss of Fe2+ ions due to transferring them but since the mixing container and weighing boat were rinsed with de-ionised water, the loss of ions should have been at a minimum. However, as the procedure in method one stands, there are a number of procedural uncertainties including the deepness of the crucible and the iron (II)sulphate crystals as well as not knowing how intense the heat should be. All of these issues can lead to the uneven heating of the iron (II) sulphate and ultimately, the part decomposition of the iron (II) sulphate. This means that when weighing what should be only anhydrous iron (II) sulphate, you are not. If this problem can arise when I was conduction the experiment, this means it could arise again suggesting that the method has some significant flaws.

Therefore, comparing the two methods; I believe that method two is the most accurate. This is for the reasons stated above and although method two only repeats the titration three times; method one only obtains one mass of water. Moreover, the procedural uncertainties concerning method one are likely to have a more significant affect on the accuracy of the results than the procedural uncertainty of method two.

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