Experiment #19
Nichole Morgan
Chem 106, Section 17
1
The purpose for conducting this experiment was to be able to prepare a series of
standard solutions of various concentrations from a stock solution to create a Beer’s law
plot in order to identify the concentration of an unknown and known solution. To do this,
six standard solutions of different concentrations were prepared from varying amounts of
stock solution dispensed from a burette, and their absorbance measured via a
spectrophotometer. From the measured values, a Beer’s law plot was produced illustrating
the relationship between absorbance and concentration (M). It is crucial to measure a
solution of known concentration because from the measured absorbance data a straight
best-fit line can be produced allowing us to accurately determine the concentration of our
unknown and known solutions.
Results:
Actual Concentration of the Known Solution: 7.155*10-6 M
Calculated Concentration of the Known Solution: 6.40*10-6 M
Unknown Solution Number: #3
Calculated Concentration of the Unknown Solution: 1.05*10-5 M
Percent Difference: (7.155*10-6 - 6.40*10-6) = 10.55%
7.155*10-6
The size of the percent difference was 10.55%, which is reasonably accurate.
Because the percent difference is greater than 10% but not too high, it allows us to see that
our data is accurate but not unrealistic. However, our calculated concentration did differ
from the known concentration, this could be due to several reasons such as, not being
precise when identifying the burette reading, improperly preparing the burette for
dispersion, contamination, or technical issues involving the spectrophotometer. To fix such
Experiment #19
Nichole Morgan
Chem 106, Section 17
2
issues in the future, it is imperative that we are more careful, we wash the burette to
prevent contamination or water bubbles, we accurately measure the reading, and we check
the spectrophotometer for error by calibrating it at least twice. Despite some error such as
the uncertainty of our glassware, the burette with an uncertainty of 0.01mL , and the
volumetric flask with an uncertainty of (+-) 0.1%, our R2 value further proves our data was
relatively accurate as our R2 value was 0.9972, very close to 1.00, meaning our data was
plotted extremely close to the determined best-fit line.
Based upon our assay, I believe the accuracy of the concentration of our unknown
sample is relatively close to its actual value, but from our percent error, I believe it to be
about 10% lower than its true concentration, making our calculated value reliable to the
point that we can identify its true value is higher than what was calculated.
An important part of this experiment was properly calibrating the instruments
before collecting data from them. This is critical because if not properly calibrated the
instrument will produce inaccurate results, thus, calibrating an instrument before use
eliminates any source of error. As well, it is also important to calibrate instruments because
it allows for data to be in a consistent range. For this experiment, the instrument that we
calibrated was the spectrophotometer. To do this, we used a blank solution (DI water)
because it does not contain pigments of any sort; it is clear and colorless so it will not
absorb any light. When introduced into the spectrophotometer the reading should be 0.00
since all the wavelengths of light will be transmitted through the liquid. By taking the
reading of the water from then on we are using that as the standard absorbance to compare
other solutions.
Experiment #19
Nichole Morgan
Chem 106, Section 17
3
The information given by the slope of the best-fit line shows an estimate of the
molar absorptivity of the substance, large slope would indicate that the absorbance was
high for higher concentrations, while a small slope would show that the absorbance was
low for all concentrations. The slope can be used to estimate the unknown values of a
standard solution. From the y-intercept, the amount of absorbance can be found given the
concentration is 0. The y-intercept should have a value of 0.00 because the concentration of
DI water is 0.00, the properties of the substance being clear and colorless make it so light is
not absorbed thus, making the absorbance 0.
If the measured absorbance of the unknown is greater than the absorbance of any of
the standard solutions used to make the Beer’s Law plot then the concentration of the
solution cannot be trusted. The high value means the solution is a very concentrated
solution that absorbs more light than a dilute solution. Beer’s law states that there is a
linear relationship between the absorbance of a dilute solution at a specific wavelength and
the concentration of the solution. To alleviate the issue, adding DI water to the unknown
solution to make it more dilute would get a more accurate reading for the Beer’s Law plot.
Failing to rinse the burette with stock solution before dispersion would result in
excess water present in the burette. The excess water would ultimately cause changes in all
data, first the standard solutions would be more dilute and would have lower
concentrations than what was intended. The measured values would all then be lower than
what we want, so our best-fit line and slope would be lower than our expected results.
Finally, the measured absorbance of the unknown solution would be higher than
anticipated, creating a higher percent error and more inaccurate results that we had
anticipated.
Experiment #19
Nichole Morgan
Chem 106, Section 17
4