Understanding Chemical Equilibrium: Key Concepts & Applications

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: Ratef = Rater • Equilibrium constants are dimensionless because they actually involve a thermodynamic quantity called activity. • Activities are directly related to molarity

2. The Equilibrium Constant • Kc is the equilibrium constant . • Kc 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. The value of Kc depends upon how the balanced equation is written. This reaction has a Kc=[PCl3][Cl2]/[PCl5]=0.53 Variation of Kc with the Form of the Balanced Equation • This reaction has a Kc=[PCl5]/=[PCl3][Cl2]=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 Kc. • 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: • Changes in concentration of reactants or products. • Changes in pressure or volume (for gaseous reactions) • Changes in temperature.

6. Relationship Between Kp and Kc • The relationship between Kp and Kc 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 BetweenDGorxn and the Equilibrium Constant • DG (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. • DG is related to DGo by the following relationship.

8. Relationship BetweenDGorxn and the Equilibrium Constant • The relationships among DGorxn, K, and the spontaneity of a reaction are:

9. There are three classes of strong electrolytes. • Strong Water Soluble Acids Remember the list of strong acids from Chapter 4. • Strong Water Soluble Bases The entire list of these bases was also introduced in Chapter 4. • Most Water Soluble Salts The solubility guidelines from Chapter 4 will help you remember these salts.

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 Ka.

11. Polyprotic Acids • Many weak acids contain two or more acidic hydrogens. • Examples include H3PO4 and H3AsO4. • The calculation of equilibria for polyprotic acids is done in a stepwise fashion. • There is an ionization constant for each step. • Consider arsenic acid, H3AsO4, which has three ionization constants. • Ka1 = 2.5 x 10-4 • Ka2 = 5.6 x 10-8 • Ka3 = 3.0 x 10-13 • This is a general relationship. • For weak polyprotic acids the Ka1 is always > Ka2, etc.

12. Polyprotic Acids • Calculate the concentration of all species in 0.100 M arsenic acid, H3AsO4, solution. • Write the first ionization step and represent the concentrations. Approach this problem exactly as previously done. • Substitute the algebraic quantities into the expression for Ka1. • Use the quadratic equation to solve for x, and obtain both values of x. • Next, write the equation for the second step ionization and represent the concentrations. • Substitute the algebraic expressions into the second step ionization expression. • Finally, repeat the entire procedure for the third ionization step. • 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: • Solutions made from a weak acid plus a soluble ionic salt of the weak acid. • Solutions made from a weak base plus a soluble ionic salt of the weak base • 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 CH3COOH • The soluble ionic salt - sodium acetate NaCH3COO

14. The Common Ion Effect and Buffer Solutions • Henderson-Hasselbach Equation 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.

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