1. Basic Concepts One of the fundamental ideas of chemical equilibrium is that equilibrium can be established from either the forward or reverse direction. The rates of the forward and reverse reactions can be represented as: When system is at equilibrium: Rate f = Rate r Equilibrium constants are dimensionless because they actually involve a thermodynamic quantity called activity. Activities are directly related to molarity

2. The Equilibrium Constant K c is the equilibrium constant. K c is defined for a reversible reaction at a given temperature as the product of the equilibrium concentrations (in M) of the products, each raised to a power equal to its stoichiometric coefficient in the balanced equation, divided by the product of the equilibrium concentrations (in M) of the reactants, each raised to a power equal to its stoichiometric coefficient in the balanced equation.

3. Variation of K c with the Form of the Balanced Equation The value of K c depends upon how the balanced equation is written. This reaction has a K c =[PCl 3 ][Cl 2 ]/[PCl 5 ]=0.53 This reaction has a K c =[PCl 5 ]/=[PCl 3 ][Cl 2 ]=1.88

4. The Reaction Quotient The mass action expression or reaction quotient has the symbol Q. –Q has the same form as Kc The major difference between Q and Kc is that the concentrations used in Q are not necessarily equilibrium values. Why do we need another “equilibrium constant” that does not use equilibrium concentrations? Q will help us predict how the equilibrium will respond to an applied stress. To make this prediction we compare Q with K c. Q<K products favored Q>K reactants favored favored Q=K equilibrium

5. Disturbing a System at Equlibrium: Predictions LeChatelier’s Principle - If a change of conditions (stress) is applied to a system in equilibrium, the system responds in the way that best tends to reduce the stress in reaching a new state of equilibrium. –We first encountered LeChatelier’s Principle in Chapter 14. Some possible stresses to a system at equilibrium are: 1.Changes in concentration of reactants or products. 2.Changes in pressure or volume (for gaseous reactions) 3.Changes in temperature.

6. Relationship Between K p and K c The relationship between K p and K c is: Heterogeneous equilibria have more than one phase present. –For example, a gas and a solid or a liquid and a gas. How does the equilibrium constant differ for heterogeneous equilibria? –Pure solids and liquids have activities of unity. –Solvents in very dilute solutions have activities that are essentially unity. –The Kc and Kp for the reaction shown above are:

7. Relationship Between  G o rxn and the Equilibrium Constant  G (notice no o indicating standard state) is the free energy change at nonstandard conditions For example, concentrations other than 1 M or pressures other than 1 atm.  G is related to  G o by the following relationship.

8. Relationship Between  G o rxn and the Equilibrium Constant The relationships among  G o rxn, K, and the spontaneity of a reaction are:  G o rxn KSpontaneity at unit concentration < 0> 1Forward reaction spontaneous = 0= 1System at equilibrium > 0< 1Reverse reaction spontaneous

9. There are three classes of strong electrolytes. 1Strong Water Soluble Acids Remember the list of strong acids from Chapter 4. 2Strong Water Soluble Bases The entire list of these bases was also introduced in Chapter 4. 3Most Water Soluble Salts The solubility guidelines from Chapter 4 will help you remember these salts. Acid HCl Base NaOH ArrheniusProduces H + Produces OH - Brönsted-LoweryDonates H + Accepts H + LewisAccepts e - pairDonates e - pair

10. Ionization Constants for Weak Monoprotic Acids and Bases We can define a new equilibrium constant for weak acid equilibria that uses the previous definition. –This equilibrium constant is called the acid ionization constant. –The symbol for the ionization constant is K a.

11. Polyprotic Acids Many weak acids contain two or more acidic hydrogens. –Examples include H 3 PO 4 and H 3 AsO 4. The calculation of equilibria for polyprotic acids is done in a stepwise fashion. –There is an ionization constant for each step. Consider arsenic acid, H 3 AsO 4, which has three ionization constants. 1K a1 = 2.5 x 10 -4 2K a2 = 5.6 x 10 -8 3K a3 = 3.0 x 10 -13 This is a general relationship. –For weak polyprotic acids the K a1 is always > K a2, etc.

12. Polyprotic Acids Calculate the concentration of all species in 0.100 M arsenic acid, H 3 AsO 4, solution. 1Write the first ionization step and represent the concentrations. Approach this problem exactly as previously done. 2Substitute the algebraic quantities into the expression for K a1. 3.Use the quadratic equation to solve for x, and obtain both values of x. 4Next, write the equation for the second step ionization and represent the concentrations. 5Substitute the algebraic expressions into the second step ionization expression. 6Finally, repeat the entire procedure for the third ionization step. 7.Substitute the algebraic representations into the third ionization expression.

13. The Common Ion Effect and Buffer Solutions There are two common kinds of buffer solutions: 1Solutions made from a weak acid plus a soluble ionic salt of the weak acid. 2Solutions made from a weak base plus a soluble ionic salt of the weak base 1.Solutions made of weak acids plus a soluble ionic salt of the weak acid One example of this type of buffer system is: –The weak acid - acetic acid CH 3 COOH –The soluble ionic salt - sodium acetate NaCH 3 COO

14. The Common Ion Effect and Buffer Solutions The Henderson-Hasselbach equation is one method to calculate the pH of a buffer given the concentrations of the salt and acid. The Henderson-Hasselbach Equation can be used for bases by substituting OH - for H + and base for acid. Henderson-Hasselbach Equation

15. Buffering Action 1Calculate the pH of the original buffer solution. 2Next, calculate the concentration of all species after the addition of the gaseous strong acid or strong base. –This is another limiting reactant problem. 3Using the concentrations of the salt and base and the Henderson- Hassselbach equation, the pH can be calculated. 4Finally, calculate the change in pH.

16. Strong Acid/Strong Base Titration Curves We have calculated only a few points on the titration curve. Similar calculations for remainder of titration show clearly the shape of the titration curve.

17. Weak Acid/Strong Base Titration Curves We have calculated only a few points on the titration curve. Similar calculations for remainder of titration show clearly the shape of the titration curve.

18. Strong Acid/Weak Base Titration Curves Titration curves for Strong Acid/Weak Base Titration Curves look similar to Strong Base/Weak Acid Titration Curves but they are inverted. Weak Acid/Weak Base Titration curves have very short vertical sections. The solution is buffered both before and after the equivalence point. Visual indicators cannot be used.