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Essay: Experiment using Le Chatelier's Principle

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  • Published: 29 October 2015*
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  • Words: 987 (approx)
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The purpose of this lab is to observe the changes in a chemical reaction and how factors affect the system like increase concentrations, temperature, pressure or common ions affect the outcome of the overall reaction and its equilibrium. The Le Chatelier’s Principle is used often in order to determine the direction of a reaction, like a balance when disturbed it shifts to more over to one side than the other; in this case if one side of the reaction is stressed the other side will react to counteract for the change. This often means that more or less products will form or reactants will be more present in the end. The five stages performed in this lab with varying chemicals each with stressed the system in various ways. The first reaction is that of 0.2M of cobalt dichloride and concentrated hydrochloric acid, the reaction produced cobalt tetrachloride and hydrogen gas or (CoCl2 + 2HCl = CoCl4 + H2) in which produced a royal blue or aqua marine color. Afterward, half of the contents of the test tube was poured into another and then was mixed with 7ml of water to produced pink Hexaaquacobolt(II), as such this equation is correct : ( CoCl4-2 + 6H2O = Co(H2O)6+2 + 2Cl2). This equation is reversible meaning if hydrochloric acid was mixed with Hexaaquacobolt(II) then the reaction would turn back to blue Cobalt tetrachloride and vice versa. The third test tube reacted with 0.1M AgNO3 or Silver nitrate to produce non-soluable Silver chloride, Cobalt(II) Nitrate and chlorine, the equation shows: (CoCl4 (aq)+ 2 AgNO3(aq) = Co(NO3)2(aq) + 4Cl(aq) + 2AgCl(aq) ). By adding silver Nitrate to the reaction, it removes a chlorine molecule from the equation and thus the reaction shifts to compensate for the lost chlorine molecule; it does that by doing a reverse reaction.
The second stage is a reaction in the change concentration of H+ ions in sodium hydroxide and hydrochloric acid. The stage called ten millimeters of water along with four drops of 0.6M hydrochloric acid in a beaker and same procedure as for the 0.6M sodium hydroxide. These diluted solutions were mixed with another solution containing 1ml of water, 4 drops of acid-base indicator and the diluted acid. The solution turns from orange to red because the indicator reacts to the change of hydrogen ion concentrations by changing color from orange to red. After that reaction, the diluted sodium hydroxide was added and the color of the solution returned to orange and this is because the ion OH- steals the proton from hydrogen, thus becomes water. The equilibrium equation is (HIn = H+ + In-) or (HOH = H+ +OH-), the ‘In’ indicates the indicator.
The third stage is the solubility of calcium hydroxide, which was created by using mixing calcium nitrate and sodium hydroxide; Ex: Ca(NO3)2 + 2NaOH = Ca(OH)2 + 2NaNO3. The calcium hydroxide separated to their individual ions when suspended into water and when the hydrochloric acid was added so the equation would be: Ca(OH)2(aq) ‘ Ca2+(aq) + 2OH-(aq) then HCl was added to make ; Ca2+(aq) + 2OH-(aq) + 2HCl ‘ CaCl2 + 4H2O. Calcium hydroxide and salt is produced and when 10ml of sodium hydroxide is added, Ex: CaCl2 + 2NaOH ‘ 2NaCl + Ca(OH)2. The Le Chatelier’s principle explains that when the same ion is added into a solution of dissociates, it would shift to the precipitate of the disassociated solids and this effect is commonly referred to as the common-ion effect. This happened as the solution’s equilibrium shifts to balance the system to achieve equilibrium state. When the CaCl2 is added to the NaOH, the solubility of Ca(OH)2 will be reduced half as much as CaCl2. the equilibrium will shift to the left decreasing the solubility of Ca(OH)2. The solubility will be higher in NaOH since NaOH produces 1 mole of OH- ion rather than the 2OH- ions in equilibrium.
The fourth stage is observing the effect of change of temperature of a reaction of Cobalt tetrachloride & Hexaaquacobolt(II). When heated in a hot water bath, Pink Hexaaquacobolt(II) begins to lose water and reverts back to CoCl4, while the Cobalt tetrachloride is cooled in ice bath it reverts back to Hexaaquacobolt(II). Since the Cobalt Tetrachloride solution is blue at a high temperature, it can be concluded that the equilibrium shifts to the left with higher temperature and to the right with lower, and by Le Chatelier’s principle, this would mean that the forward reaction is exothermic. The reverse reaction the heating Hexaaquacobolt(II) produced Cobalt tetrachloride then the reaction is exothermic. This equation illustrates that reaction: Heat + Co(H2O)6+2 + 4Cl = CoCl4-2 + 6H2O – Heat. The Fifth stage is involves a reaction of ferric nitrate with potassium thiocyanate in 4 test tubes but later only proceeds with test tube 2 & 4. The key concept of this stage is influencing the concentration of FeSCN2+ in each test tube, so the starting concentration of Fe3+ would be greater than the starting concentration of SCN’ ions. When the Fe3+ concentration become the excess, the equilibrium shifted to the product side until all ions of SCN’ converted to FeSCN32+. This is observed as iron thiocyanate appears to be a deep red/dark crimson color versus the pale yellow of ferric nitrate. The reaction reached equilibrium after the concentration of Fe(SCN)32+ in the solution was equal to the same concentration of SCN’ ions in the solution, meaning the color ceases to change any further. Test tube 2 continued by adding it potassium hydroxide to which in turn produce iron nitrate that appears to be a for a lack of better word deep urine-like color, according to this equation : Fe(SCN)3 + 3KOH = 3KSCN + Fe(OH)3. The 4th test tube contains Fe(SCN)3 but after five drops of 0.1M silver nitrate was added, the reaction produce the precipitate Silver thiocyanate(AgSCN) & Ferric nitrate: Fe(NO3)3 that appear as yellowish. Then more KSCN was added to produce more Fe(SCN)3.

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