Chapter 16: Chemical Equilibrium- General Concepts

1. Chapter 16: Chemical Equilibrium- General Concepts • WHAT IS EQUILIBRIUM?

2. The decomposition of N2O4(g) into NO2(g). The concentrations of N2O4 and NO2 change relatively quickly at first, but eventually stop changing with time when equilibrium is reached.

3. The equilibrium between N2O4 and NO2. • The equilibrium mixture is independent of whether we start on the “reactant side” or the “product side”

4. The same equilibrium composition is reached from either the forward or reverse direction, provided the overall system composition is the same.

5. There is a simple relationship among the concentrations of the reactants and products for any chemical system at equilibrium • Consider the equilibrium:

6. Four experiments to study the equilibrium among H2, I2, and HI gases. Different amounts of the reactants and products are placed in a 10.0 L reaction vessel at 440oC where the gases establish equilibrium. When equilibrium is reached, different amounts of reactants and products remain.

7. This average value is called the reaction quotient, Q

8. The reaction can be evaluated at any concentrations • At equilibrium (and 440oC) for this reaction the reaction quotient has the value 49.5 (a unitless number) • This relationship is called the equilibrium law for the system

9. The value 49.5 is called the equilibrium constant, Kc, and characterizes the system • For chemical equilibrium to exist, the reaction quotient Q must be equal to the equilibrium constant Kc • Consider the general chemical equation

10. The exponents in the mass action expression are the same as the stoichiometric coefficients • At equilibrium • The form is always “products over reactants” raised to the appropriate powers

11. Various operations can be performed on equilibrium expressions • Changing the direction of equilibrium – when the direction of an equilibrium is reversed, the new equilibrium constant is the reciprocal of the original

12. Multiplying the coefficients by a factor – when the coefficients in an equation are multiplied by a factor, the equilibrium constant is raised to a power equal to that factor

13. Adding chemical equilibria – when chemical equilibria are added, their equilibrium constants are multiplied

14. The gas law can be used to write the equilibrium constant in terms of partial pressures • Equilibrium constants written in terms of partial pressures are given the symbol Kp

15. The size of the equilibrium constant gives a measure of how the reaction proceeds • General statements can be made about the equilibrium constant (either Kc or KP)

16. The magnitude of K and the position of equilibrium. A large amount of product and very little reactant at equilibrium gives K>>1 (large K). When , approximately equal amounts of reactant and product are present at equilibrium. When K<<1, mostly reactant and very little product are present at equilibrium.

17. The two different forms of the equilibrium constants can be related

18. In a homogeneous reactions, all the reactants and products are in the same phase • Heterogeneous reactions involve more than one phase • For example the thermal decomposition of sodium bicarbonate (baking soda) • Heterogeneous reactions can come to equilibrium just like homogeneous systems

19. If NaHCO3 is placed in a sealed container, homogeneous equilibrium is established • The equilibrium law involving pure liquids and pure solids can be simplified

20. For a pure liquid or solid, the ratio of amount of substance to volume of substance is constant The concentration of a substance in a solid is constant. Doubling the number of moles doubles the volume, but the ratio of moles to volume remains the same.

21. The equilibrium law for a heterogeneous reaction is written without concentrations terms for pure solids or pure liquids. • The equilibrium constants found in tables represent all the constants combined

22. According to Le Châtelier’s principle: If an outside influence upsets an equilibrium, the system undergoes a change in the direction that counteracts the disturbing influence and, if possible, returns the system to equilibrium • We can consider some common “stresses” • Adding or removing a product or reactant • The equilibrium shifts to remove reactants or products that have been added • The equilibrium shifts to replace reactants or products that have been removed

23. Changing the volume • Reducing the volume of a gaseous reaction causes the reaction to decreases the number of molecules of gas, if it can • Moderate pressure changes have a negligible effect on reactions involving only liquids or solids • Changing the temperature • Increasing the temperature shifts a reaction in a direction that produces an endothermic (heat-absorbing) change • Decreasing the temperature shifts a reaction in a direction that produces an exothermic (heat-releasing) change

24. Catalysts have no effect on the position of equilibrium • Catalysts change how fast a system achieves equilibrium, not the relative distribution of reactants and products • Adding an inert gas at constant volume • If the added gas cannot react with any reactants or products it is inert towards the substances in the equilibrium • No concentration changes occur, so Q still equals K and no shift in equilibrium occurs

25. Equilibrium calculations can be divided into two main categories: • Calculating equilibrium constants from known equilibrium concentrations or partial pressures • Calculating one or more equilibrium concentrations or partial pressures using the known value of Kc or KP • Consider the decomposition of N2O4

26. Calculating the equilibrium constant this way is easy

27. More commonly, you will have a set of initial conditions and an equilibrium constant • If a KP describes the system, equilibrium will usually be described in terms of partial pressures • If a Kc describes the system, equilibrium will usually be describe in terms of concentration (molarity, mol/L) • The Initial, Change, Equilibrium or “ICE” table is a useful way to summarize the problem

28. Example: Ethyl acetate, CH3CO2C2H5, is produced from acetic acid and ethanol by the reaction At 25oC, Kc=4.10 for this reaction. Suppose 0.100 mol of ethyl acetate and 0.150 mol of water are placed in a 1.00 L reaction vessel. What are the concentrations of all species at equilibrium? ANALYSIS: Use an ICE table and the equilibrium constant to find the concentrations.

30. This can be solved by putting it in quadraticform:

31. Negative concentrations are not allowed, so • A similar procedure can be used to calculate partial pressures using KP

32. Sometime simplifications can be made Example: Nitrogen and oxygen react to form nitrogen monoxide with Kc=4.8x10-31. In air at 25oC and 1 atm, the N2 concentrations and O2 are initially 0.033 M and 0.00810 M. What are the equilibrium concentrations? ANALYSIS: The equilibrium constant is very small, very little of the reactants will be converted into products

34. Substituting: [N2]=0.033-x=0.033 M [O2]=0.00810-x=0.030810 M [NO]=2x=1.60x10-17 M