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Chemistry: equilibrium - misunderstanding something or totally hosed experiment?
Jimmy King
Registered User regular
Registered User regular
I'm leaning towards totally screwed up experiment, for several reasons, but am double checking here in case I am just dumb. Basic experiment - mix Fe3+ and SCN- ions to get Fe(SCN)2+. Do this at various concentrations to calculate Beer's law and Le Chatellier's Principle stuff.
So this lab has 3 parts, we have 6 groups, so the teacher had 2 groups each do each part.
Part 1: my group and another did this. Determine the light wavelength with the greatest absorption for Fe(SCN)2+. My group found that to be 576nm, the other group doing this part landed at 560nm... close-ish, but not actually in line with what I've been seeing online. Then test varying concentrations of the FeSCN2+ created by using a ton of KSCN and a small amount of Fe(NO3)3 to ensure that all Fe3+ reacted and the FeSCN2+ molarity would be the same as the Fe3+ molarity. All of our absorptivity values were very close between the two groups. Average the absorptivity values, plot a graph, the slope of the line is the absorptivity constant. This gave us a line with the equation y = 1154.8x + .0374 and so absorptivity constant of 1154.8.
Part 2: Things go south real quick. Use several different, more balanced, concentrations of the above and find their absorbance. Use the previously made graph (or equation of a line) to use the absorbance to find the equilibrium concentration of FeSCN2+. Use that along with an ICE table to determine the equilibrium concentrations of Fe3+ and SCN-. Use those to determine the equilibrium constant at room temp. I have a feeling things have gone to hell because for some reason the groups doing this part used 440nm light, which is way the hell different from 560-576nm, although it is more in line with what I have been seeing online. So, using one example with an absorbance of .285:
Absorbance: .285
Initial Fe: 5ml of .002M diluted to 20mL = 0.0005M
Initial SCN: 5ml of .002M diluted to 20mL = 0.0005M
FeSCN equilibrium molarity:
.285 = 1154.8x + .0374
x = (.285-.0374)/1154.8 = 2.14x10-4M
So Fe and SCN equilibrium molarity each = .0005 - 2.14x10-4 = 2.86x10-4
Using Kc = [FeSCN]/([Fe][SCN])
Kc = 2.14x10-4/(2.86x10-4 x 2.86x10-4) = 2.63x103
That's just a tiny bit off from values I'm seeing on the internet which suggest it should be more like 170-180. Worse, though, is the fact that these values range from -19,000 to 13,400 over 8 samples. The -19,000 comes from a negative SCN- equilibrium concentration which is obviously not possible. Even throwing that out as an outlier, the numbers are pretty screwed.
I am doing this correctly and some combination of incorrect concentrations and incorrect light wavelengths is off, causing stupid values, right?
So this lab has 3 parts, we have 6 groups, so the teacher had 2 groups each do each part.
Part 1: my group and another did this. Determine the light wavelength with the greatest absorption for Fe(SCN)2+. My group found that to be 576nm, the other group doing this part landed at 560nm... close-ish, but not actually in line with what I've been seeing online. Then test varying concentrations of the FeSCN2+ created by using a ton of KSCN and a small amount of Fe(NO3)3 to ensure that all Fe3+ reacted and the FeSCN2+ molarity would be the same as the Fe3+ molarity. All of our absorptivity values were very close between the two groups. Average the absorptivity values, plot a graph, the slope of the line is the absorptivity constant. This gave us a line with the equation y = 1154.8x + .0374 and so absorptivity constant of 1154.8.
Part 2: Things go south real quick. Use several different, more balanced, concentrations of the above and find their absorbance. Use the previously made graph (or equation of a line) to use the absorbance to find the equilibrium concentration of FeSCN2+. Use that along with an ICE table to determine the equilibrium concentrations of Fe3+ and SCN-. Use those to determine the equilibrium constant at room temp. I have a feeling things have gone to hell because for some reason the groups doing this part used 440nm light, which is way the hell different from 560-576nm, although it is more in line with what I have been seeing online. So, using one example with an absorbance of .285:
Absorbance: .285
Initial Fe: 5ml of .002M diluted to 20mL = 0.0005M
Initial SCN: 5ml of .002M diluted to 20mL = 0.0005M
FeSCN equilibrium molarity:
.285 = 1154.8x + .0374
x = (.285-.0374)/1154.8 = 2.14x10-4M
So Fe and SCN equilibrium molarity each = .0005 - 2.14x10-4 = 2.86x10-4
Using Kc = [FeSCN]/([Fe][SCN])
Kc = 2.14x10-4/(2.86x10-4 x 2.86x10-4) = 2.63x103
That's just a tiny bit off from values I'm seeing on the internet which suggest it should be more like 170-180. Worse, though, is the fact that these values range from -19,000 to 13,400 over 8 samples. The -19,000 comes from a negative SCN- equilibrium concentration which is obviously not possible. Even throwing that out as an outlier, the numbers are pretty screwed.
I am doing this correctly and some combination of incorrect concentrations and incorrect light wavelengths is off, causing stupid values, right?
Jimmy King on
0
Posts
you essentially solved Beers law by varying your concentration.
A=ebc
but once you change light wavelength e changes. So you are stuck with an additional variable you can't solve.