Lab Report Spectrophotometric Determination Of An Equilibrium PDF

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Equilibrium lab report for CHEM 120...

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Lab report, hand in within 48 hours

SPECTROPHOTOMETRIC DETERMINATION OF AN EQUILIBRIUM CONSTANT Lab report Name: ID #: TA name: Lab Section: Lab Date: 1. Looking at Table 5 (data sheet), how is Kc related to the initial concentrations ( [Fe3+]o and [SCN– ]o)? Is it what you expected? (Explain briefly) (5 marks) Looking at Table 5 in the results, it is somewhat evident that Kc is actually independent of the initial concentrations of the reactants since they are relatively consistent among the five test tubes. Furthermore, the Kc values are much more relative as opposed to the significant changes in Kc due to temperature variation which holds true because the equilibrium constant is affected by temperature. However, looking at the data for any trends, it appears that Kc drops as the concentration of [SCN–] increases, but this should not be true. So, accounting for relative human error and limited machine capabilities, the value only fluctuated minimally from the average which was 204 (M-1). Thus the results are not exactly as I had expected, but ignoring the small deviations, the results are based on Le Chatelier’s Principle which states that when a chemical system in equilibrium is disturbed there will be a shift in the system to reduce the effect of that disturbance. So, by increasing [SCN–] from tube 1 to 5, there should be no change of Kc since equilibrium will shift to the right of the reaction to increase the concentration of [𝐹𝑒𝑆𝐶𝑁 2+ ] and balance out. 𝐾𝑐 =

[𝐹𝑒𝑆𝐶𝑁 2+ ]𝑒𝑞 [𝐹𝑒 3+] ∗ [SCN − ]

2. To obtain [FeSCN2+]eq in tube 5, you made the assumption that 100% of the ions SCN – had reacted. Now that you know the value for Kc (average in table 5), build an ICE diagram to calculate the true percentage of SCN – reacted in tube 5. Comment on the validity of the assumption. (Make sure you include the ICE diagram in your answer) (10 marks) ICE Table [Fe+3]

+ [SCN-]

⇌ [FeSCN+2]

Initial

0.180 M

0.000199 M

0M

Change

-x

-x

+x

Equilibrium

(0.180 – x) M

(0.000199 – x) M

xM

1

[𝐹𝑒 3+ ] =

(9.00 𝑚𝐿∗0.2000 𝑀) 10.00𝑚𝐿

= 0.180 𝑀[SCN − ] =

(1.00 𝑚𝐿∗0.00199 𝑀) 10.00𝑚𝐿

= 0.000199 𝑀

The average calculated for Kc among the test tubes is 204. 𝑥 [𝐹𝑒𝑆𝐶𝑁 2+ ]𝑒𝑞 → 204 = → 𝐾𝑐 = − 3+ [𝐹𝑒 ]𝑒𝑞 ∗ [SCN ]eq (0.180 − 𝑥)(0.000199 − 𝑥) 204(0.180 − 𝑥)(0.000199 − 𝑥) = 𝑥 → 204𝑥 2 − 0.040596 − 36.72𝑥 − 𝑥 + 0.00730728 = 0 204𝑥 2 − 37.760596𝑥 + 0.00730728 = 0 Using the quadratic formula, 𝑥 =

−𝑏 ± √𝑏2 −4𝑎𝑐 2𝑎

→ x = 0.185 M or 0.000194 M

The 0.185 M is clearly not the correct concentration, so the concentration must be 0.000194 M. Thus, the true percentage of SCN– reacted in tube 5 is: 0.000194𝑀 ∗ 100% = 97.5% %[SCN − ] 𝑟𝑒𝑎𝑐𝑡𝑒𝑑 = 0.000199𝑀 Therefore, the assumption that 100% of the ions of SCN – had reacted because the percentage calculated from the average Kc value (determined from the Kc value of each trial during lab experiment) is extremely high. This result is due to the addition of a high concentration of Fe3+ which leads to SCN– being the limiting reactant in the chemical reaction. A limiting reactant is the substance that must all be consumed when the reaction has reached completion. 3. Use data in table 6 (data sheet) to plot the graph of ln(K c) versus 1/T. From the graph, determine the enthalpy (H) of the reaction, show your work below. Copy and paste the graph below and include with this lab report. (10 marks for the graph, 5 for Enthalpy ) ∆𝐻

Using the Van’t Hoff equation which is 𝑙𝑛𝐾𝑐 = − 𝑅𝑇 +

∆𝑆° 𝑅

, we can calculate the enthalpy of

the reaction by taking the slope of the graph of ln(Kc) versus 1/T and equating it to − Thus, −∆𝐻 =

𝑙𝑛𝐾𝑐 1 𝑇

∗ 𝑅 and plugging in the values into the equation we get:

∆𝐻 = −2020.3 ∗ 8.314 𝐽 𝐾 −1𝑚𝑜𝑙 −1 = −16800𝐽 = −16.8𝑘𝐽

2

∆𝐻 𝑅

.

ln(Kc)

Equilibrium Constant vs. Temperature 5.8 5.7 5.6 5.5 5.4 5.3 5.2 5.1 5 4.9 y = 2020.3x - 1.5881 4.8 R² = 0.9698 4.7 0.00315 0.0032 0.00325 0.0033 0.00335 0.0034 0.00345 0.0035 0.00355 0.0036

1/T

• Please note, for this report you may write out the answers to question 1 and 2 by hand, scan the file and submit as a pdf. You can use the uprint printers to scan and e-mail the file. • Once you are complete, remember to save as .pdf file and submit to MyCourses. • Double check your submission to make sure that you have submitted the correct file. • (Sometimes students accidentally submit the empty template, a corrupt file or the wrong file altogether, double check the file that you have submitted to make sure this doesn’t happen to you.)

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