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580336_Equilibrium & LeChatlier's Principle.indd CHEMISTRY Equilibrium and Le Châtelier’s Principle Investigation Manual EQUILIBRIUM AND LE CHÂTELIER’S PRINCIPLE Table of Contents 2 Overview 2 Objectives 2 Time Requirements 3 Background 6 Materials 7 Safety 7 Preparation 8 Activity 1 8 Activity 2 8 Activity 3 9 Activity 4 9 Disposal and Cleanup 10 Data Tables Overview A copper(II) chloride aqueous solution is in equilibrium with copper hexahydrate. This lab investigates how this equilibrium can be shifted to the right or left by changing the temperature and concentration of the reactants or products. The shifts in equilib- rium can be monitored by changes in the color of the solution, as more products or more reactants are produced. Objectives • Understand what it means when a reaction is in equilibrium. • Explore Le Châtelier’s Principle. • Determine the influence of the common ion effect on an equilibrium reaction. • Determine the effect of temperature on an equilibrium reaction. • Determine the effect of concentration on an equilibrium reaction. Time Requirements Preparation .................................................................... 10 minutes Activity 1: Common Ion Effect ...................................... 15 minutes Activity 2: Addition of Silver Nitrate ................................. 5 minutes Activity 3: Addition of Water .......................................... 10 minutes Activity 4: Effect of Temperature on Equilibrium ........... 15 minutes 2 Carolina Distance Learning Key Personal protective equipment (PPE) goggles gloves apron warning corrosion flammable toxic environment health hazard follow photograph stopwatch link to results and required video submit 47 amber-colored NO2 after a brief period of time. How is the equilibrium for NO2 established? Initially, only the forward reaction occurs, as N2O4, which is colorless, changes into NO2, which is amber. This is demonstrated in the following equation: 57.2 kJ + N2O4(g) 2NO2(g) colorless amber However, in a short period of time, a reverse reaction takes place, and some of the NO2 reacts to form N2O4 until a dynamic equilibrium exists between the two compounds. After that, the concentrations remain at fixed values because the rate of the forward reaction (Rforward) equals the rate of the reverse reaction (Rreverse). Note that heat is a reactant in the above forward reaction. The enthalpy value (heat energy; ΔH°) for this reaction under standard conditions is 57.2 kJ. This value is positive for endothermic reactions because energy is consumed in the reaction. If this enthalpy had been negative (–57.2 kJ), the forward reaction would have been exothermic, with the heat value appearing as a product on the right side of the equation, as shown here: N2O4(g) 2NO2(g) + 57.2 kJ. Figure 1. Background Equilibrium Reactions A chemical reaction occurs when reactants react with each other to form products. If the reaction is unidirectional, reactions in which the reac- tants create products and the products cannot convert back to the reactant, then the reaction is called an irreversible reaction. Another type of reaction is a reversible reaction, where the forward reaction forms the products, and a reverse reaction reforms the original reactants. In a reversible reaction, the reactants and products are never fully consumed – they are constantly reacting and being produced. At the beginning of the reversible reaction (time 0), there are no products. When the reactants collide, the forward reaction starts forming products. As the product concentration increases, a reverse reaction occurs in which some of the products are converted back into reactants. Eventually, the rate of the reverse reaction equals the rate of the forward reaction, and a state of dynamic equilibrium is reached. When a dynamic equilibrium is reached, the concentrations of the reactants and products remain fixed, but this does not mean that the forward and reverse reactions have stopped. Reactants are still being converted into prod- ucts, and products are still being converted into reactants. However, the rates of conversion for the forward and reverse reactions are the same after a given period of time, thus resulting in no net change in concentrations. A good analogy is the act of walking down an “up” escalator at the same speed as the escalator moves upward. There will be no change in your position due to the balance of the opposing motions. Figure 1 depicts a graph showing the equilibrium that is established for colorless N2O4 and www.carolina.com/distancelearning 3 ©2015, Carolina Biological Supply Company 4 Background Equilibrium Reactions A chemical reaction occurs when reactants react with each other to form products. If the reaction is unidirectional, reactions in which the reactants create products and the products cannot convert back to the reactant, then the reaction is called an irreversible reaction. Another type of reaction is a reversible reaction, where the forward reaction forms the products, and a reverse reaction reforms the original reactants. In a reversible reaction, the reactants and products are never fully consumed – they are constantly reacting and being produced. At the beginning of the reversible reaction (time 0), there are no products. When the reactants collide, the forward reaction starts forming products. As the product concentration increases, a reverse reaction occurs in which some of the products are converted back into reactants. Eventually, the rate of the reverse reaction equals the rate of the forward reaction, and a state of dynamic equilibrium is reached. When a dynamic equilibrium is reached, the concentrations of the reactants and products remain fixed, but this does not mean that the forward and reverse reactions have stopped. Reactants are still being converted into products, and products are still being converted into reactants. However, the rates of conversion for the forward and reverse reactions are the same after a given period of time, thus resulting in no net change in concentrations. A good analogy is the act of walking down an “up” escalator at the same speed as the escalator moves upward. There will be no change in your position due to the balance of the opposing motions. Figure 1 depicts a graph showing the equilibrium that is established for colorless N2O4 and amber-colored NO2 after a brief period of time. How is the equilibrium for NO2 established? Initially, only the forward reaction occurs, as N2O4, which is colorless, changes into NO2, which is amber. This is demonstrated in the following equation. 57.2 kJ + N2O4(g) 2NO2(g) colorless amber continued on next page EQUILIBRIUM AND LE CHÂTELIER’S PRINCIPLE 4 Carolina Distance Learning Copper Chloride Equilibrium When copper(II) chloride is dissolved in water, the following reaction steps take place until an equilibrium is established: CuCl2(s) → Cu2+(aq) + 2Cl ¯ (aq) Cu2+(aq) + 6H2O(l) → Cu(H2O)62+(aq) Cu(H2O)62+(aq) + 4Cl ¯(aq) CuCl42 ¯(aq) + 6H2O(l) blue green If Le Châtelier’s principle holds, the equilibrium should be shifted by adding or removing either chloride ion (Cl¯) or water (H2O). Adding more chloride ion is called the “common ion” effect. In this investigation, the chloride ion is added and removed by using a sodium chloride solution and a silver nitrate solution, respectively, as shown in the reactions below. NaCl(s) → Na+(aq) + Cl ¯(aq) AgNO3(s) → Ag+(aq) + NO3¯(aq) Ag+(aq) + Cl ¯(aq) → AgCl(s) reaction to the right to replace the removed heat. The pressure is increased, which also shifts the reaction to the right to relieve the excess pressure. In this case, two moles of NH3 exert less pressure than four moles of reactants. The simultaneous application of all three stresses (concentration, heat, and pressure) brings about an “optimum yield” of NH3 for the reaction. In industry, the challenge with reversible reac- tions is to keep the products from decomposing back into reactants. Chemical engineers prefer a 100% product yield from a reaction. Such is the case for ammonia production. Fritz Haber, a German chemist, won the 1918 Nobel Prize in chemistry for developing the Haber process. This process involves taking nitrogen directly from the atmosphere and combining it with hydrogen to produce ammonia (NH3), which is a vital component of inorganic fertilizers and explosives. N2(g) + 3H2(g) 2NH3(g) + heat When the temperature, pressure, and reactant concentrations are normal, the percent yield of ammonia is quite low, due to its decomposition. However, as Henri Le Châtelier discovered in the late 1800s, placing a stress on a reaction at equilibrium can force it to make more prod- ucts or reactants. A stress can be a change in temperature or concentration. If the reactants or products are gases, the stress can be a change in pressure for the whole reaction system, like the inside of an enclosed tank. Le Châtelier’s principle states that if a stress is placed upon a reaction at equilibrium, the reaction will shift to relieve that stress. To efficiently produce NH3, chemical engineers need to shift the equilibrium to the right (the product side). This is achieved by concurrently increasing the concentrations of N2 and H2 and removing the ammonia gas (NH3) as soon as it is formed. This forces the reaction to the right to lower the concentration of the additional N2 and H2 and to replace the NH3. At the same time, heat is removed, thus further shifting the Background continued ©2015, Carolina Biological Supply Company 5 However, in a short period of time, a reverse reaction takes place, and some of the NO2 reacts to form N2O4 until a dynamic equilibrium exists between the two compounds. After that, the concentrations remain at fixed values, because the rate of the forward reaction (Rforward) equals the rate of the reverse reaction (Rreverse). Note