Le Chateliers free essay sample

Henry-Louis Le Chatelier, (born Oct. 8, 1850, Paris, France—died Sept. 17, 1936, Miribel-les-Echelles), French chemist who is best known for Le Chatelier’s principle, which makes it possible to predict the effect a change of conditions (such as temperature, pressure, or concentration of reaction components) will have on a chemical reaction. His principle proved invaluable in the chemical industry for developing the most-efficient chemical processes. The most remarkable feature of a system at equilibrium is its ability to return to equilibrium after a change in conditions moves it away from the state. this drive to reattain equilibrium is state in Le Chatelier’s principle: when a chemical system at equilibrium is disturbed, it reattain equilibrium by undergoing a net reaction that reduce the effect of the disturbance. Two phrases in this statement need further explanation. First, what does it means to ‘disturb’ a system? At equilibrium, Q equals K. When a change in conditions forces the system temporarily out of equilibrium (Q, we say that the system has been disturbed or stresses. Three common disturbances are a change in concentration of a component (that appears in Q), a change in pressure (caused by a change in volume), or a change in temperature. The other phrase, â€Å"net reaction,† is often referred to as a shift in the equilibrium position of the system to the right or left. The equilibrium position is just the specific equilibrium concentrations (or pressures). A shift in the equilibrium position to the right means that there is the net reaction to the right (reactant to product) until equilibrium is reaatained; a shift to the left means that there is a net reaction to the left (product to reactant). Thus, when the disturbance occurs, we say that the equilibrium position shifts, which means that concentrations (or pressures) change in a way that reduces the disturbance, and the system attains a new equilibrium position (Q = K again). Le Chatelier’s principle allows us to predict the direction of the shift in equilibrium position. Most importantly it helps research and industrial chemists create conditions that maximize yields. The basis of Le Chatelier’s principle holds for any system at equilibrium, whether in the natural or social sciences. 2. 1Concentration effect If at equilibrium state, the concentration of one or more substances present in equilibrium mixture is changed, the equilibrium system will no longer remain in equilibrium state. The system will undergo changes in concentrations of various substances so as to minimize the effect or restore equilibrium state. Consider a reversible reaction: aA + bB cC + dD Kc for the reaction is Kc = [C]c[D]d/[A]a[B]b If we add some more amount of A or B at equilibrium, the equilibrium system will be disturbed . According to Le-chateliers principle, to restore equilibrium, the reaction will shift in the forward direction to cancel the effect of change in concentration. Actually an increase in concentration of reactants decreases the ratio [C]c[D]d/[A]a[B]b than Kc. To regain equilibrium, according to Le-chateliers principle, at equilibrium the concentration of A and B will decrease and the concentrations of C and D will increase i. e the reaction will shift in the forward direction. However if at equilibrium the concentration of C or D is increased, the reaction will shift in the backward direction. In this experiment, we will determine the effect of these reactions: Equilibrium of Fe (III)-SCN ion complexes Fe3+ (aq) + (SCN)-(aq) Fe(SCN)2+ (aq) (yellow) (dark red) 2. 2Temperature effect According to Le-chateliers principle a change in temperature is a stress on an equilibrium system. If at equilibrium the temperature of system is changed the system will no longer at remain at equilibrium. To restore equilibrium, the reaction will in either forward or backward direction. By applying Le-chateliers principle, we can predict the direction of reaction when temperature of an equilibrium system is changed. Effect of change in temperature is related to the nature of reaction whether it is an endothermic reaction or an exothermic reaction. For endothermic reaction, Kc increases with the increase in temperature. For exothermic reaction, Kc decreases with the increase in temperature. Consider a reversible reaction: aA + bB cC + dD For endothermic reaction Increase in temperature: In an endothermic reaction, an increase in temperature favours the reaction to occur in the forward direction. At equilibrium, the concentration of A and B will decrease and the concentration of C and D will increase. Decrease in temperature: In an endothermic reaction, a decrease in temperature favours the reaction to occur in the backward direction. At equilibrium, the concentration of C and D will decrease and the concentration of A and B will increase. For exothermic reaction Increase in temperature: In an exothermic reaction, an increase in temperature favours the reaction to occur in the backward direction. At equilibrium, the concentration of C and D will decrease and the concentration of A and B will increase. Decrease in temperature: In an exothermic reaction, a decrease in temperature favours the reaction to occur in the forward direction. At equilibrium, the concentration of A and B will decrease and the concentration of C and D will increase. 2. 2. 1Equilibrium of Co (II) ion complexes Co(H2O)62+ (aq) + 4 Cl- (aq) CoCl42- (aq) + H2O (l) (red) (blue) 2. 2. 2Equilibrium of Cu (II) ion complexes CuCl42- (aq) + H2O (l) Cu(H2O)62+ (aq) + 4Cl- (aq) (yellow) (blue) 4. PROCEDURE 4. 1 Concentration effect on Fe (III)-SCN ion complexes 4. 1. 1. 1 ml Fe (NO3)3 and 1 ml of KSCN solution is mixed in 100ml beaker. 4. 1. 2. 25ml distilled water is added into the mixture and the solution is stirred. 4. 1. 3. 4 clean test tubes are cleaned and prepared as A1, A2, A3 and A4. 4. 1. 4. The solution is divided evenly in all the test tubes. 4. 1. 5. An additional 1ml Fe(NO3)3 is added into test tube A1, additional 1ml of KSCN solution into test tube A2 and 8 drops of NaOH solution into test tube A3. 4. 1. 6. All the solutions are stirred gently and the solutions colours are compared with the solution in test tube A4. 4. 2 Temperature effect on CO (II) and Cu (II) ion complexes 4. 2. 1. 5ml of CoCl2 and 5ml of CuCl2 is prepared in two separate conical flasks. 4. 2. 2. 3ml HCL is added into both flasks (The solution is prepared in fume cupboard). 4. 2. 3. Both solutions are stirred until cobalt solution turn purple (red + blue) and cuprum solution turn to green (blue + yellow). 4. 2. 4. 6 clean test tubes are cleaned and prepared as B1, B2, B3 and C1, C2, C3. 4. 2. 5. The cobalt solution is divided evenly to test tubes B1, B2, B3 and cuprum solution to C1, C2, C3. 4. 2. 6. B1 and C1 test tubes are soaked in the ice.