Practical - Experiment 34 Report - An Equilibrium Constant PDF

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Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant Destiny Cambero CHEM 112 – Farnum T/TH 11:30 AM 09/27/18

Experiment: 34 An Equilibrium Constant Conclusion: For this experiment, we were asked to determine the equilibrium constant of a mixture of Fe3+ (aq) and SCN-. The various concentrations of this mixture that we prepared was plotted on a graph and from its slope we could determine our equilibrium constant using the Beer’s Law equation. In conclusion my average Kc was, my standard deviation of my Kc was, and my % relative standard deviation was

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant

Experiment: An Equilibrium Constant – Abstract The purpose of this experiment was to use a spectrophotometer to determine the equilibrium constant of a chemical system and to then use graphing and data analysis to evaluate our collected data. By doing so we can determine the equilibrium constant for a soluble equilibrium. A spectrophotometer is an instrument used to measure the amount of light that is transmitted through a sample. The spectrophotometric method uses the interaction of electromagnetic radiation with matter. The spectrophotometer measures the transmitted light with a photosensitive detector at visible wavelengths. The wavelength at which the molecules has the most absorption of radiation is used as the set wavelength for the analysis. The amount of visible light transmitted can be affected by many factors such as the concentration of the substance, the thickness of the sample, the cuvet, and the molar absorptivity. The most important factors are transmittance and absorbance. Transmittance is the amount of light that passes through the sample, it is used because it is linear and easy to read. Absorbance on the other hand is most commonly used in calculations because absorbance is directly related to the concentration of a substance. The higher the absorbance means the higher the concentration, the lower the absorbance the lower the concentration. Therefore, the absorbance value is directly proportional to the molar concentration of the absorbing substance. The major results in this experiment was the average equilibrium constant, the standard deviation, and percent relative standard deviation.

Experiment: An Equilibrium Constant - Introduction The goal of this experiment was to use a spectrophotometer to determine the equilibrium constant of a chemical equation and to use graphing techniques and data analysis to evaluate our data. From this data, we would then be able to determine the equilibrium constant for a soluble equilibrium. The method we used for this experiment is the spectrophotometric method which involves the interaction of electromagnetic radiation with matter. The most commonly used wavelengths vary from 400-700nm, which is within the visible spectra. The instrument used to determine this is a spectrophotometer, this measures the amount of transmitted light of a sample. These light intensities are measured with a photosensitive detector. The transmittance is the amount of light that is passed through a sample and is inversely related to the absorbance. The

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant higher the transmittance the fewer ions or “contaminants” are found in the solution leading to a lower absorbance. Whereas a higher absorbance or concentration leads to a lower transmittance because the more ions that are present in a solution, the lower the chances are of light being able to pass through the solution. This implies that the absorbance and concentration of a solution are directly related, when plotted the slope can be used to identify the molar concentration. The equilibrium constant Kc, is used to express the position of equilibrium for a chemical system, it represents the completeness of a chemical reaction. It is a reaction that is reversible and at a state of equilibrium the rate of the forward and the rate of the reverse reaction will be the same, Ratefor=Raterev. This is also the case for both the reactant and product concentrations. A reversible reaction can be represented as aA + bB

expression

❑ ❑

xX + yY, and the Kc is determined by the mass

[ A ]a [ B ]b =K c . The mass expression is equal to the equilibrium constant when a [ C]c [ D ]d

dynamic equilibrium is established between reactants and products. It is used to determine the concentrations at equilibrium and whether a reaction favors the products or reactants at equilibrium. In this experiment, the reversible reaction between Fe3+and SCN-(thiocyanate) ions reaches equilibrium in aqueous solutions, Fe3+ (aq) + SCN- (aq)

❑ ❑

FeSCN2+ (aq). However,

to determine equilibrium, we must know at least on concentration value at equilibrium in order to determine the values of the other concentrations present. This shows us how the concentrations are related to one another and that the rate of a reaction has a direct correlation with the products of the molar concentrations in which each concentration has a power equal to the stoichiometric coefficient of that species. The mathematical relationship that relates absorbance to concertation is the Beer-Lambert Law, otherwise known as beer’s law which can be expressed as A= a*b*c. In this equation, “a” is the Absorptivity, “b” is the path length (cm), and “c” is the concentration. This equation explains that absorbance is proportional to the concentration of the sample that absorbs the light. In order to determine the concentration, we need to first find the molar absorptivity and the path length of the cuvet containing the sample, however for our experiment we can utilize a calibration curve to determine this. So, from our various known concentrations that we prepared of FeSCN2+ and from the absorbance data we collected at the max absorbance, we can plot the absorbance vs. the

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant concentration to get a calibration curve. An I.C.E table can then be utilized to determine the equilibrium concentrations of the remaining reactants; this information will also be useful in determining the equilibrium constant K. The mathematical equation for constant equilibrium was

Kc =

FeSCN ¿ 2+¿ ¿¿ ¿ 3+¿ ¿ Fe ¿ −¿ ¿ SCN ¿ ¿ ¿

.

Since FeSCN2+ doesn’t exist as a solid, we can use Le Châtelier's Principle to prepare a known concentration. This principle states that a nearly complete reaction moving in the forward direction can be reached when a large amount of reactant is added to a small amount of another. By doing this we have a limiting quantity of Fe3+ and larger amounts of SCN-, the limiting reagent in this case would be converted to product. Therefore, we can assume that the FeSCN2+ will be equal to the amount of Fe3+.

Methods For this experiment, we prepared 6 volumetric flasks of 0,1,2,3,4, and 5ml of 0.001M NaSCN into separate 25ml volumetric flasks. We then added We prepared the calorimeter by obtaining two Styrofoam coffee cups to prepare the apparatus followed by a Styrofoam lid, stirrer, and a digital thermometer using a Pasco Capstone system to acquire our data. The reason we use Styrofoam is because it is a good insulator and it insulates the inside from the outside. If we had used a different material or possibly a single cup instead of two, then there is a higher chance that there will not be enough insulation and some of the heat will be lost. So, for this experiment we are to assume that our coffee cup calorimeter is a “perfect” insulator. Since we used the digital thermometer we did not need to observe the change in temperature every few seconds, instead we let the system run continuously and it recorded our data over time. I believe that by using this method we can get more accurate readings giving us better results because of the fact that it removes any personal error. For the first part of Part B I

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant took 50ml of 0.9365M NaOH and measured its temperature alone in the coffee cup calorimeter for about three minutes continuously. After three minutes I added 50ml of 1.1M HCl to the cup with the NaOH and measured the change in temperature for an additional 5 minutes. I then repeated this process for a second trial, and then an additional two trials but with a mixture of NaOH and HNO3. After we plotted our data on graphs. For part C, I collected ~5g of my unknown salt and 20ml of DI water. With the digital thermometer, I measured the initial temperature of the of the DI water by itself in the calorimeter for three minutes. After the three minutes, I added my unknown salt and continuously measured the change in temperature of the reaction for another 5 minutes. I did this for two trials with my salt. After collecting the data, we made additional graphs for this data. The procedures for these experiments were both produced by Beran, J.A. Laboratory Manual for the Principles of General Chemistry. (10th ed.); John Wiley & Sons, Inc. USA, 2015; pp 214-223.

Data Table 1: A Set of Standard Solutions to Establish a Standardization Curve Molar Concentration of Fe(NO3)3 (M)

0.2

Molar Concentration of NaSCN (M)

0.001

Standard Solutions

Blank

Volume of 0 NaSCN (mL)

1

2

3

4

5

1

2

3

4

5

Moles of SCN- (mol)

0

1•10-6

2•10-6

3•10-6

4•10-6

5•10-6

[SCN-] (25.0mL)

0

4.00•10-5

8.00•10-5

1.2•10-4

1.6•10-4

2.00•10-4

[FeNCS2+]

0

4.00•10-5

8.00•10-5

1.2•10-4

1.6•10-4

2.00•10-4

%T

100

90

81

66

46

31

Absorbance

0

.05

.09

.18

.34

.51

(mol/L)

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant Table 2: Absorbance for the Set of Test Solutions Molar Concentration of Fe(NO3)3 (M)

0.002

Molar Concentration of NaSCN (M)

0.002

Test Solutions

6

7

8

9

10

x

Volume of Fe(NO3)3 (mL)

5

5

5

5

5

x

Moles of Fe3+ (mol)

1•10-5

1•10-5

1•10-5

1•10-5

1•10-5

x

2

3

4

5

x

Volume of 1 NaSCN (mL) Moles of SCN-, initial (mol)

2•10-6

4•10-6

6•10-6

8•10-6

1•10-5

x

%T

100

76

61

43

35

x

Absorbance

0

0.12

0.22

0.36

0.46

x

The tables above show the data collected for the standards of the standardization curve and the absorbance for a set of test solutions. Graph Graph explanation

Results Table 3: Calculations of Kc 6 [FeNCS2+], equilibrium, from calibration

7

8

9

10

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant curve (mol/L) Moles FeNCS2+ at equilibrium (10ml) (mol) [Fe3+], equilibrium Moles Fe3+ reacted (mol) Moles Fe3+, equilibrium (mol) [Fe3+], equilibrium unreacted (10ml)(mol/L)

[SCM-], equilibrium Moles SCN-, reacted (mol) Moles SCN-, equilibrium (mol) [SCN-], equilibrium (unreacted)(10ml) (mol/L) Kc Average Kc Standard deviation of Kc Relative Standard Deviation of Kc The tables above represent the calculated results to find the equilibrium constant. Discussion In this experiment, we were expected to determine the enthalpy of neutralization for an acid-base reaction and the enthalpy of solution for the dissolution of a salt. We know that the specific heat is what causes the temperature of a gram of a substance to rise 1°C. We use a calorimeter, which is an apparatus that allows us to measure the flow of energy through a system and its surroundings, to determine this because we know the specific heat capacity of water. The calorimeter is the best apparatus to use because it acts as a perfect insulator preventing heat from escaping into the surroundings, where in this instance the energy is equal to the heat. For part B, we determined the enthalpy of the neutralization reaction with the known values of our change in temperature and the molarity of our acids and bases. My chemical reactions data for my HCl acid was that my average ∆Hn was -63.835 kj/mol, my standard deviation was 0.6293, my relative standard error was 0.9858, and my percent error was 14.40%. My standard deviation and relative was okay because they were under one, nut my percent error is way too high. It is abouve 10% which means that my data was not accurate and precise. Anything below 10% is considered good data. For my HNO3 my average ∆Hn was 68.745 kj/mol, my standard deviation was 1.2657, my relative standard deviation was 1.8412, and my percent error was 20.82%, these results are also above 10%, which means that these results are not accurate as well. Overall, I didn’t have good quality data for my experiment indicating that my errors had an impact on my results. I had could have possibly lacked technique or maybe it’s possible that my lack of consistency with the mixing impacted my results. For part C, we determined the enthalpy od solution for the dissolution of a salt. Lattice energy and hydration energy are very important factors in this experiment. Lattice energy is the energy required to break the ionic bonds within a solid and our hydration energy has to do with the ions

Destiny Cambero CHEM 112 Exp: 34 An Equilibrium Constant that bond with water molecules that creates ion-dipole forces. My unknown was 21and its heat capacity was 0.903J/g °C. This specific heat capacity tells me that my unknown salt was Na2SO4. The average total ∆Hs per mole of the unknown salt was 0.903J/mol, my standard deviation was 372.3, my relative standard deviation was 4.78%, and my percent error was 297.96%. We see a high percent error because there is a low ∆H value and low temperature change, even the smallest errors could affect the results. It is so high because even though we assume our calorimeter is “perfectly insulated”, we know that it is indeed not a perfect insulator. Conclusion Errors Errors that occurred throughout my experiment was that the system was not working correctly. I had to redo my dilutions three times to confirm that I was getting accurate readings. For part B, my trial one (#6) would not change. At first it was possibly my ratios of the mixtures I was making, but after performing a second, and yet a third trial it was apparent that the spectrophotometer was not working correctly. Half the class had one brand of spectrophotometer while the other half had a different brand. The brand that my half of the class had seemed to have a lot of issues while the other half of the class finished quickly. Therefore, I assume that my data isn’t very accurate. A second issue I had was that when filling the volumetric flasks, I filled one over the line for the indicated measurement. It being more dilute would lead to more transmittance and less absorbance, this would affect the data, especially when plotting on a graph. Errors that could have occurred could have been that I didn’t wipe the sides of the cuvets or inserted the cuvet facing the frosted side instead of the clear side. Both of these errors would result in a lower transmittance and a higher absorbance. A second factor could have been that I didn’t calibrate it or if I calibrated with a solution that wasn’t clear like water. You want to use water or a clear solution to calibrate to what would be an accurate representation off 100% transmittance and 0 absorbance. If you use a concentrated solution to calibrate then the results won’t be accurate....