Introduction;
Most chemical reactions are reversible which means that they react both forwards and backwards. A forward reaction is generically A + B -> C + D and a backward reaction is C + D -> A + B. Higher concentrations of A and B at the beginning of the reaction cause it to shift right towards C and D before beginning to slow down as more C and D are created. Eventually, the amount of A and B being formed is the same as the amount of A and B being used, which results in chemical equilibrium, denoted with Keq. From this reaction, the equilibrium constant can be calculated by using the ratio of the product of the products over the product of the reactants. Thus Keq= [C]*[D] / [A] * [B].
Some reversible reactions reach equilibrium faster than others such as that of Iron (III) ion (Fe3+) with thiocyanate ion (SCN-) that forms thiocyanatoiron(III) (FeSCN2+). In this reversible reaction Fe3+ reacts with SCN- to produce FeSCN2+ in water. For this reaction A is the iron (III) ion, B is the thiocyanate ion and C is the thiocyanato iron (III). Its Keq value is equal to
K=([Fe(SCN)^(2-)])/([Fe^(3+) ][SCN^-])
and this may now be used to find the equilibrium constant with known or calculated concentrations.
Methods;
Standard Curve;
Five clean and dry test tubes are obtained and labeled 1-5. Each is filled with exactly 2.50 mL of .200 M Fe(NO3)3 using a burette. Then 0.50 mL of 0.002 M KSCN solution is added to test tube 1. 0.75 mL of 0.00200 M KSCN is added to test tube 2 and so on in increments of 0.25 mL. Finally, enough 0.5 M HNO3 is added to each test tube so that the final volume is equal to 10.0 mL. Each test tube is mixed and then the contents of each are added to a cuvette and tested within a spectrophoto...
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...imary productivity in large regions of the global open ocean. Copper (Cu), on the other hand, is a common anthropogenic contaminant to estuarine and coastal oceans that can act as a toxicant to microorganisms at elevated concentrations. The organic complexation of dissolved iron and copper by largely uncharacterized natural ligands in seawater has proven to be an integral component in the oceanic biogeochemistry of these metals, governing aspects of their solubility, supply and bioavailability in the marine environment. Recent research projects in the Buck lab have examined the distributions, sources and sinks of natural iron- and copper-binding organic ligands in seawater, biological transformations of iron and copper species, and the influence of copper-binding ligands on bioavailability and toxicity of copper in contaminated coastal and estuarine environments.
Use blue LED color for maximum absorbance of Ferroin and scroll down the colorimeter screen to view absorbance at your chosen wavelength. Measure the initial absorbance of the mixture with colorimeter, record it and use this information to determine the molar extinction coefficient for Ferroin. Place the cuvette into 40°C water bath and let it heat up. Remove cuvette from water bath to measure the absorbance of mixture with two minutes of interval in between, putting the cuvette immediately back after the measurement. Be sure to dry the cuvette (ex. with paper towel) before putting into colorimeter. Continue until the absorbance drops below 0.2 or when you have the 10th measurement. For second experiment, repeat procedure using 0.20M sulfuric acid by diluting 0.40M sulfuric acid. For the last experiment, use 0.40M sulfuric acid again but put into 45°C water bath instead of
Abstract: This week we experimentally determined the rate constant k for the reaction 2HCl (aq) +Na2S2O3 (aq) → S (s) + SO2 (aq) + H2O (l) + 2NaCl (aq). In order to do this the average reaction time was recorded in seconds during two trials. The data from the experiment shows this reaction is in the first order overall: rate=.47s-1 [HCl]0 [Na2S2O3]1. These findings seem to be consistent with the expected results
Put 1mL of 0.1M cobalt (II) chloride hexahydrate dissolved in 95% ethanol into a test tube. Then add 1mL of deionized water. Tap the end of the test tube to mix the solution and record the pertinent data in section 2 of the Data Table. Discard the solution in the appropriate container as directed to you by your lab instructor.
The color that was chose to be shined through the sample was purple. The spectrophotometer was set at a wavelength of 400nm to represent the purple color. It was zeroed using the blank meaning the spectrophotometer read zero as absorbance amount. The blank consisted of 5mL of water and 2.5 mL AVM and it was placed in cuvette. A solution with a known concentration of 2.0x10-4 M was used in the spectrometer. For this solution, 5 mL of the solution with 2.5 mL of AMV was placed in the cuvette. The cuvette was placed inside of spectrophotometer and the amount of absorbance was recorded. This procedure that involves a solution with a known concentration was repeated for the concentrations:1.0x10-4 M,5.0x10-5 M,2.0x10-5M, and1.0x10-5M.A unknown solution absorbance was measured by putting 5 mL of unknown solution with 2.5 mL AMV in a cuvette. The cuvette was placed in the spectrophotometer and the amount of absorbance was recorded. The procedure that deals with the unknown solution was repeated 2 more times with the same solution and the same amount of solution and AMV. The average of the three unknown solution was calculated and the concentration of the unknown solution was
Purpose: The purpose of the lab was to perform a series of chemical reactions in order to transform copper within different reactions in order to start and end with solid brown copper.
To complete this lab several chemicals must be measured and transferred to test tubes. First 5.0 mL of 0.200 M Fe(NO3)3 must be diluted to a total volume of 50 mL in a flask. Next 0.0020 M SCN–. This solution is then added to 4 test tubes in 1 mm increments. Each test tube is then put in to
In acid-base titration solution without a known molarity is placed in an Erlenmeyer after it’s volume is measured. An indicator is added to the solution most of the time it is phenolphthalein. The solution with a known concentration is placed in burette with a tap in the end. By opening the tap slightly solution in the burette is poured in to the solution in Erlenmeyer drop by drop. After a while the solution in Erlenmeyer forms a color change. This is the turning point for the solution. At the turning point by the volume consumed in burette the molarity of the other solution can be
A cuvette was filled 3/ 4ths of the way and the absorbance measured in a spectrophotometer. The data was compiled as a class and recorded. The Spectrophotometer was blanked using a test tube of distilled water.
The purpose of the experiment is to identify and understand reactions under kinetic and thermodynamic control. A reaction under kinetic and thermodynamic control can form two different types of products. A reaction under kinetic control is known to be irreversible and the product is formed quickly. A reaction under thermodynamic control is known to require rigorous conditions. It is also reversible. The final product is more stable than the product made by kinetic control. The chart below shows the two types of reaction coordinates:
and the volume of acid. By doing all this we will make sure that it is
The purpose of the experiment is to study the rate of reaction through varying of concentrations of a catalyst or temperatures with a constant pH, and through the data obtained the rate law, constants, and activation energies can be experimentally determined. The rate law determines how the speed of a reaction occurs thus allowing the study of the overall mechanism formation in reactions. In the general form of the rate law it is A + B C or r=k[A]x[B]y. The rate of reaction can be affected by the concentration such as A and B in the previous equation, order of reactions, and the rate constant with each species in an overall chemical reaction. As a result, the rate law must be determined experimentally. In general, in a multi-step reac...
tube. Add 6 mL of 0.1M HCl to the first test tube, then 0.1M KMnO4 and
This is the first reaction in the Harcourt Essen experiment. The iodine is oxidised to produce I2 wh...
This experiment consists of titrating the ferrous ion with permanganate ion to study the oxidation-reduction reaction. The ions react in acidic solution to give ferric ion and a reduced ionic form of manganese. All the reactants and products except permanganate ion are weakly colored, whereas permanganate is a very intensely colored ion. Then a solution of permanganate is removed as long as there is a ferrous ion present to react with it. But as soon as the entire ferrous ion has been oxidized, the next small portion of added permanganate colors the solution. The first appearance of a permanent pink color indicates the endpoint of the experiment. From the titration it will be able to calculate the percentage of iron in the sample from the data.
In this experiment three different equations were used and they are the Stoichiometry of Titration Reaction, Converting mL to L, and Calculating the Molarity of NaOH and HCl (Lab Guide pg. 142 and 143).