1. If the molarity of the NaOH solution is increased, it would take less drops to observe a colour change. A higher concentration lowers the amount of solution required to produce a reaction due to the presence of more molecules and a higher frequency rate, therefore less of the substance is needed to cause an equilibrium shift.
2. When you add 6.0 M NaOH into the iron (III) thiocyanate ion equilibrium system, the concentration of Fe+3 ion decreases. This causes the equilibrium to favour the reactants, which explains the observed colour change of the solution.
3. If the hydrated cobalt (II) ion complex was refrigerated, the equilibrium would shift and the reactants side would be favoured. When the equilibrium shifts to the left, the colour …show more content…
According to the data, when the temperature was increased, the reaction shifted to the right (products' side) and when the temperature was decreased, the reaction shifted to the left (reactants' side). With this evidence, one can conclude that the forward reaction was endothermic and this supports the Le Chatelier's Principle that states in an endothermic reaction the temperature increases the forward rate.
5. The addition of NaCl into the hydrated cobalt (II) ion equilibrium would increase the concentration of chloride ions and cause the formation of some sodium ions upon dissociation. Which in turn, would cause the equilibrium to shift to the right side in order to use up the excess chloride ions. The shifting of the equilibrium shows that it favours the products and results in a purple colored solution.
6. Cr2O7 2+ (aq) +2OH- (aq) ⇌ 2CrO4 2- (aq) + H2O (l)
7. n
8. Barium dichromate appeared as a fairly clear/orange solution, while the barium chromate was a milky orange colour due to the presence of precipitates. The solubility of barium dichromate was more soluble since the solution was still visible relative to barium chromate in which the solution was cloudy and non-transparent.
9. No precipitation was formed in K2Cr2O7 because there was a lower concentration and the Ba(NO3)2 was unable to produce precipitates. In step 7 and 8,
2. When you add 6.0 M NaOH into the iron (III) thiocyanate ion equilibrium system, the concentration of Fe+3 ion decreases. This causes the equilibrium to favour the reactants, which explains the observed colour change of the solution.
3. If the hydrated cobalt (II) ion complex was refrigerated, the equilibrium would shift and the reactants side would be favoured. When the equilibrium shifts to the left, the colour …show more content…
According to the data, when the temperature was increased, the reaction shifted to the right (products' side) and when the temperature was decreased, the reaction shifted to the left (reactants' side). With this evidence, one can conclude that the forward reaction was endothermic and this supports the Le Chatelier's Principle that states in an endothermic reaction the temperature increases the forward rate.
5. The addition of NaCl into the hydrated cobalt (II) ion equilibrium would increase the concentration of chloride ions and cause the formation of some sodium ions upon dissociation. Which in turn, would cause the equilibrium to shift to the right side in order to use up the excess chloride ions. The shifting of the equilibrium shows that it favours the products and results in a purple colored solution.
6. Cr2O7 2+ (aq) +2OH- (aq) ⇌ 2CrO4 2- (aq) + H2O (l)
7. n
8. Barium dichromate appeared as a fairly clear/orange solution, while the barium chromate was a milky orange colour due to the presence of precipitates. The solubility of barium dichromate was more soluble since the solution was still visible relative to barium chromate in which the solution was cloudy and non-transparent.
9. No precipitation was formed in K2Cr2O7 because there was a lower concentration and the Ba(NO3)2 was unable to produce precipitates. In step 7 and 8,