10-Phenanthroline Reaction Lab

When the pH is not at its optimum, the differing pH's will disrupt the bonding between the R groups of the amino acid causing its structure and the shape of the activation site to change

The ending result of this experiment confirms that as five test tubes are lined up with the varying level of absorbance, different results in the level of absorbance will appear as well, this is visible in above table. Thus, this is due to the varying amount of water in the solution. The blank sample had a 0.30 in its level of absorbance.

The Beers Law calibration experiment used many concentrations of crystal violet solutions. Each of these solutions were test and analyzed in order to determine the absorbance of each concentration The results were than graphed and produced a slope of 1.00E05 with an intercept of -2.21E-02.

The values of color absorbance are effective because color absorbance has a linear relationship with concentration values, which in turn, allows us to easily find concentration values for many solutions. Beer’s law describes this phenomenon since the absorbance is directly proportional to concentration. We observed that as the color absorbance increased, the concentration of the FeSCN2+ complex ion increased. This is because as the FeSCN2+ concentration increases, the blood-red color becomes darker due to more presence of the blood-red FeSCN2+ ion. Therefore, the color absorbance increases because there is more blue color absorbed by the darker red color. We then graphed the absorbance and concentration values and created a line of best fit. Using the line of best fit, we were able to predict the equilibrium concentrations of the FeSCN2+ solutions and find the change required to reach equilibrium. Since we already knew the initial concentration of FeSCN2+ and since we already found the equilibrium concentration of FeSCN2+, we can calculate the change in equilibrium. Using this data, we were able to calculate the equilibrium concentration of all of the species in this lab, since we already knew the change from the initial concentration to the equilibrium change. Q is less than K because there was no initial concentration of FeSCN2+, but after the system reached

The purpose of this lab was to figure out what the formation constant (Kf) was for FeSCN2+. To answer this guiding question, Le Chatlier’s principle, volume, absorbance, and concentration all effect outcomes of equilibrium. For this reason, all three components will be collected in order to solve for the average formation constant. Equilibrium is the rate at which the forward reactions and reverse reactions become equal at a constant rate. Equilibrium constant, denoted by K, is the relationship between the concentrations of product over the concentrations of reactants raised to their specific coefficients. In the balanced equation of Fe3+ (aq) + SCN-(aq) ⇌[Fe(SCN)]2+(aq), an excess of Fe3+ is used to move the reaction into completion. In other

For this experiment we are looking to determine the equilibrium constant KC for this reaction, KC = [FeSCN2+]/([Fe3+]•[SCN-]. Using a spectrophotometer, we are going to determine the absorbance of FeSCN2+ with different concentration measurements.

For stock Solution A, the luminol precipitate was dissolved into a 25mL Erlenmeyer flask containing sodium hydroxide (2mL, 3M in H2O). Solution A used a 50 mL beaker to dilute stock solution A (1 mL) in water (9 mL). Stock Solution B called for an previously prepared potassium ferricyanide solution (4 mL, 3% by weight in H2O) and a hydrogen peroxide solution (4 mL, 3% by weight in H2O) in a 25 mL Erlenmeyer flask. Solution B diluted stock solution B (4 mL) with water (16 mL) and placed into a 150 mL beaker. This group of solutions was then moved into a dark room and solution B was poured into the 150 mL beaker containing diluted solution A. Observed and recorded was a bright emission of blue light.

Beer’s law states that absorbance (A) is equal to the molar absorptivity (a) times the pathlength of the cuvette (l) times the concentration of the species in question (c) for an overall equation of A=acl. As more and more Fe3+ was added, the color became a deeper and deeper orange because more FeSCN2+ was being produced. Concentration is proportional to absorbance, so in theory the absorbance of the solution should increase with each 1 mL addition of Fe3+.2 Since the pathlength of the cuvette (1 cm) and the molar absorptivity remain constant, only the concentration of the iron causes variation in the absorption. Once the

Tube three contained 1.0 mL of Buffer, pH 6.0 0.1 M NaPO4 including 0.4 mL of 0.006 M catechol, with 600 uM of substrate concentration and 1.6 mL of distilled water. Tube four contained 1.0 mL of Buffer, pH 6.0 0.1 M NaPO4 including 0.8 mL of 0.006 M catechol, with 1200 uM of substrate concentration and 1.2 mL of distilled water. And finally, Tube five contained 1.0 mL of Buffer, pH 6.0 0.1 M NaPO4 including 1.6 mL of 0.006 M catechol, with 2400 uM of substrate concentration and 0.4 mL of distilled water. All five of the test tubes also contained 1.0 mL of enzyme, which would be added right before testing each tube for absorbance. After preparing all of the test tubes, each tube was inserted in the spectrophotometer one-by-one and recorded its change in absorbance for five minutes.

Iron cations combine with thiocyanate (SCN-) anions to form thiocyanatoiron (FeSCN2+). Looking at the form a[A] + b[B] ⇌ c[C], the formula to find the equilibrium constant is K_c= ([A]^a [B]^b)/[C]^c .2 For this specific reaction, the formula is K_c=[FeSCN^(2+) ]/([Fe^(3+) ][SCN^-]).

The tubes sat at room temperature for 10 minutes, and they were inverted and shaken every 2 minutes in order to mix the enzymes and chelating agents. They sat at room temperature for 10 minutes while being mixed every 2

During the first week of experiments, a standard curve was prepared to aid in the determination of the concentration of Fe(SCN)2+ at equilibrium. With the use of Le Chateliers principle, which states that if a change occurs while a system is at equilibrium, the species will react to offset the change so that it does maintain equilibrium. During the experiment in the second week, several solutions with varying initial concentrations of iron(III) and thiocyanate (SCN-) were prepared. With the use of the standard curve from Beer’s Law and “ICE” tables, the team was able to determine the equilibrium constant for each of the

Determing the Chemical Equilibrium for FeSCN2+ by Observing Colour and Concentration Purpose: The purpose of this lab is to determine the equilibrium constant value of Iron (III) Nitrate with Potassium Thiocynate by observing colour changes when various substances are mixed with the Iron (III) Thiocynate and measuring the depths of 2 solutions after adding or removing some of the solutions until the colours match. Hypothesis: Part 1 If Iron (III) Thiocynate is mixed with distilled water in Test Tube 1, Iron (III) Nitrate in Test Tube 2, Iron (III) Thiocynate in Test Tube 3, distilled water in Test Tube 4 with Test Tube 4 in an ice bath, and distilled water in Test Tube 5 with Test Tube 5 in a hot-water bath, then the order from most intense

Incorporation of assay controls included setting up a spectrophotomer and running the chart recorder with a full-scale deflection before the start of the assay. The set recorder had a corresponding value of 1 for the change in the absorbance. Therefore, prior testing was done to observe whether a change occurred in the readings. This helped to indicate that the results were valid, as they could have been affected by a fault during the setting up of the spectrophotometer. On the other hand this was considered as one of the controls for the experiment. Nevertheless, a new cuvette had to be used for each assay.

The quantity of light absorbed is inversely proportional to the logarithm of the amount of light absorbed by the substance. The amount of light absorbed by a substance is directly proportional to the concentration of the substance (As the concentration increases, the absorbance also increases provided that the wavelength is constant). [1]