Through the experimentation of reactions between iron (III) nitrate and potassium thiocyanate, as well as cobalt (II) chloride hexahydrate in water, equilibrium systems disturbed by stress (changing the amount of reagents and temperature) will shift in order to minimize the stress. Therefore, if the concentration of reactant increases, the rate of the forward reaction will increase and equilibrium will reestablish when the concentration of products increase (vice versa). This can be observed through color changes of the solutions.
Equilibrium Constant Equations for Equilibrium System 1
Fe3+ + SCN- ⇌ FeSCN2+
Yellow Orange + Colorless ⇌ Blood Red Orange
Keq = [FeSCN2+] / [Fe3+][SCN-]
Table 1: Equilibrium System 1 Part A
Data Table studying the effects of concentration within a complex-ion equilibrium reaction between iron (III) nitrate and potassium thiocyanate.
Substance Added
Color of Substance
Resulting Color of Solution
Original - Potassium
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Thiocyanate Solution Clear and colorless n/a Iron (III) Nitrate Clear and orange orange Potassium Thiocyanate Crystals White opaque Dark orange and clear, but changes to blood red and translucent Potassium Nitrate Crystals White opaque spheres White to dark orange to red (solute) No change (solution) Sodium Phosphate Monobasic Crystals Small white crystals Created an orange area in the center of the solution, with swirling it went from red to orange to clear Iron (III) Nitrate Clear and orange Blood red, maroon, brown Potassium Thiocyanate Crystals White opaque Clear precipitate, light yellow solution (continued to add Iron (III) Nitrate till light orange for part B) As seen through Table 1 displaying the effects of concentration within an equilibrium system, the color of the solution changes as the amount of reagent or product varies.
The original solution containing potassium thiocyanate is clear and colorless. However, when iron (III) nitrate or potassium thiocyanate is added to the solution, the overall color becomes darker and more concentrated to orange then to blood red. After a short time, the solution achieves equilibrium, but not at equal concentrations. By adding iron (III) nitrate or potassium thiocyanate, the amount of reagent increases, therefore the forward reaction increases in order to generate more product and increase the concentration of product. In the equilibrium constant equation, when increasing the amount of SCN-, the denominator increases, therefore Q < Keq. In order to reestablish equilibrium, concentration of products need to increase and the concentration of reactants have to decrease through the consummation of reactants and production of
products. The solution experiences no change in color when potassium nitrate crystals are added. As expected, KNO3 is composed of spectator ions during this reaction, meaning it is not within the Keq equation. Therefore, the equation remains the same and the solution remains with no visible change. Adding sodium phosphate monobasic crystals into the solution causes it to go from blood red to clear. This displays the opposite reaction as to when iron (III) nitrate or potassium thiocyanate (reagents) are added to the initial clear and colorless solution. Through an exothermic reaction, H2PO4- in sodium phosphate monobasic reacts with the remaining Fe3+ reactant (as shown below) to produce FeH2PO42+ and heat. We know that heat is produced because a change in Keq indicates a change in temperature. For exothermic reactions, an increase and temperature correlates to a lower Keq and a decrease in temperature correlates to a higher Keq. For endothermic reactions however, a rise in temperature indicates a higher Keq and a loss of temperature indicates a lower Keq. A + heat ⇌ B (for endothermic) A ⇌ B + heat (for exothermic) Fe3+ + H2PO4- → FeH2PO42+ + heat The overall color changes from clear to blood red / maroon / brown with the addition of iron (II) nitrate. With potassium thiocyanate crystals, a clear precipitate forms and the solution becomes light yellow. By adding iron (III) nitrate or potassium thiocyanate, the amount of reactant increases, therefore the forward reaction increases in order to create more product and raise the concentration of it. In the equilibrium constant equation, when increasing the amount of SCN-, the denominator increases, therefore Q < Keq. To reestablish equilibrium, the concentration of product increases and the concentration of reactant decreases through the consummation of reactants for the manufacturing of products. To support Le Chatelier’s Principle, varying concentration within solutions demonstrate that equilibrium systems disturbed by stress will act in a way to reduce it. Table 2: Equilibrium System 1 Part B Data Table studying the effects of temperature within a complex-ion equilibrium reaction between iron (III) nitrate and potassium thiocyanate. Test Tube Initial Color Resulting Color A - Control Light orange, clear Light orange, clear B - Ice Bath Light orange, clear Darker, still clear, more vibrant C - Hot Bath Light orange, clear Translucent, cloudy and fuzzy, light yellow To study the effects of temperature in an equilibrium system, three test tubes were used: a control, an ice bath, and a hot bath. As seen through Table 2, the solution in the ice bath changed from light orange to a darker, more vibrant red-orange. Under cool conditions, the solution has higher levels of red concentration, indicating a loss of heat and a favorable forward exothermic reaction. When you cool the solution down, you remove heat from the products so Keq at 0 degrees celsius is smaller than Keq at 25 degrees celsius. The rate of reverse reaction and rate of forward reaction both slow down but the reverse reaction does so to a greater extent. As shown in table 2, the solution in the hot bath changed from light orange to a cloudy, light yellow. Under high temperatures, the solution has lower levels of red concentration, indicating a gain of heat and a favorable reverse reaction. When you raise the temperature of the solution, energy transfers causing a gain in heat so Keq at 70 degrees celsius is bigger than Keq at 25 degrees celsius. The solutions must maintain an equilibrium state of equal rates, but not equal concentrations. So, if one side increases or decreases, the other side will do the same to reduce stress to some degree. Table 3: Equilibrium System 2 Data Table studying the effects of temperature within a cobalt complex-ion when calcium chloride, hydrochloric acid, and silver nitrate are added. Test Tube Temperature Stance Color A - Control Room temperature Dark blue, clear B - Control w/ Calcium Chloride Ice bath Initial: Blue liquid, white solid Result: slightly lighter blue, white solid did not dissolve B - Control w/ Calcium Chloride Hot bath Slight change: white solid, more vibrant blue B - Control w/ Calcium Chloride and Hydrochloric Acid Ice bath HCl Initial: clear and colorless; caused solution to become a lighter blue Result: became a lighter blue B - Control w/ Calcium Chloride and Hydrochloric Acid Hot bath No change in color C - Control w/ Silver Nitrate Ice bath Clear light brown Made solution pink/purple to lavender C - Control w/ Silver Nitrate Hot bath Immediately white specks (precipitate) forms, color changes back to blue Furthermore, according to Table 3, concerning the effects of temperature within a cobalt complex-ion when reacting to ions of the compounds calcium chloride, hydrochloric acid, and silver nitrate, Le Chatelier’s Principle can be supported. The equation below summarizes the reaction and the initial color of the solution is dark blue / purple: CoCl42- + 6H2O ⇌ Co(H2O)62+ + 4Cl- Blue Pink Testing the control with calcium chloride in an ice bath and a hot bath experiences different color changes. The solution in the ice bath becomes a lighter blue, while the one in the hot bath becomes more of a vibrant blue. In this case, a decrease in temperature lowers the concentration and a rise in temperature gains concentration. A lighter blue occurs when the product side is favored, having more Co(H2O)62+ (pink) than CoCl42- (blue). A vibrant blue is caused by favoring the reactant side because more blue is present than pink. Adding hydrochloric acid to the control with calcium chloride in the ice bath causes the solution to become a light blue due to high concentration of products, while adding it to the hot bath experiences no visible change. This is due to an inconsistency: we failed to heat the test tube for long enough. Adding silver nitrate to the control in an ice bath changes the solution to light pink / lavender because large quantities of Cl- ions, part of the products side, are lost. Equilibrium needs to be restored by making more products, increasing the concentration of pink. Having it in the hot bath however, changes the color back to blue because more reactants are present and an increase in temperature within an exothermic reaction will have a decrease in Keq. The cold and hot baths used in both exothermic equilibrium systems from every experiment will have different initial temperatures. The experiment (testing concentration in equilibrium system 1) has a range of 65 to 70 degrees celsius for the hot bath and does not specify how many ice cubes are to be used in the ice bath. A similar inconsistency goes for time in part B when testing the effect of temperature as well. The instructions provide a broad range of time such as “3 to 5 minutes” and “2 to 3 minutes” instead of a specific time like 45 seconds, for example. Therefore, each group may result in very similar outcomes, but with slight differences in color and physical appearances in general. Our observations could very well reflect inaccuracy as to how the results were supposed to be. As said by Le Chatelier’s Principle and further proved through experimentation and observation, it can be concluded that equilibrium systems experiencing stress (altering the amount of reagents, temperature, and pressure) will react in a way to reduce the stress. This is greatly proven by the way iron (III) nitrate and potassium thiocyanate react when exposed to different concentrations and temperatures, as well as cobalt (II) chloride hexahydrate reacting with calcium chloride, hydrochloric acid, and silver nitrate while undergoing modified temperatures.
The unknown bacterium that was handed out by the professor labeled “E19” was an irregular and raised shaped bacteria with a smooth texture and it had a white creamy color. The slant growth pattern was filiform and there was a turbid growth in the broth. After all the tests were complete and the results were compared the unknown bacterium was defined as Shigella sonnei. The results that narrowed it down the most were the gram stain, the lactose fermentation test, the citrate utilization test and the indole test. The results for each of the tests performed are listed in Table 1.1 below.
Compress the safety bulb, hold it firmly against the end of the pipette. Then release the bulb and allow it to draw the liquid into the pipette.
I did accomplish the purpose of the lab. First, I determined the percentage of water in alum hydrate, and the percentage of water in an unknown hydrate. The results are reasonable because they are close to the example results. Second, I calculated the water of crystallization of an unknown hydrate. Furthermore, I developed the laboratory skills for analyzing a hydrate.
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== == I completed a table to show my results, here is the table: Table 1. Results of different changes of substances Part A Copper (II) Sulfate and Water Reactant description Water (reactant): Color: Colorless Transparency:
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