chemical equilibrium - general and inorganic chemistry

chemical equilibrium - general and inorganic chemistry

1. CHEMICAL EQUILIBRIUM GENERAL CHEMISTRY 2

2. Equilibrium The reaction of the reactants to form the products is called the forward reaction. The reaction of the products to form the reactants is called the reverse reaction.

3. Suppose we put some hydrogen gas and iodine vapor in a sealed container. They will immediately begin to react with each other to form hydrogen iodide (the forward reaction). As they do, the concentration of hydrogen and iodine in the container will begin to drop (see graphs below - the brackets refer to concentrations). Since a decrease in concentration usually results in a decreased reaction rate, the forward reaction slows down. At some point the rate of the forward reaction will equal the rate of the reverse reaction. When this occurs the system is said to be in equilibrium.

4. Equilibrium State and the Equilibrium Constant

5. at equilibrium, reactant and product concentrations are constant because a change in one direction is balanced by a change in the other as the forward and reverse rates become equal: k fwd and k rev are the forward and reverse rate constants, respectively, and the subscript “eq” refers to concentrations at equilibrium.

6. By rearranging, we set the ratio of the rate constants equal to the ratio of the concentration terms: The ratio of constants creates a new constant called the equilibrium constant ( K ): The equilibrium constant K is a number equal to a particular ratio of equilibrium concentrations of product(s) to reactant(s) at a particular temperature.

7. Reaction Quotient and Equilibrium Constant law of chemical equilibrium, or the law of mass action. “at a given temperature, a chemical system reaches a state in which a particular ratio of product to reactant concentrations has a constant value.”

8. Changing Value of the Reaction Quotient The particular ratio of concentration terms that we write for a given reaction is called the reaction quotient (Q , also known as the mass-action expression). As the reaction proceeds toward equilibrium, the concentrations of reactants and products change continually, and so does their ratio, the value of Q. at a given temperature, at the beginning of the reaction, the concentrations have initial values, and Q has an initial value; a moment later, the concentrations have slightly different values, and so does Q

9. The changes continue, until the system reaches equilibrium . At that point, reactant and product concentrations have their equilibrium values and no longer change. Thus, the value of Q no longer changes and equals K at that temperature:

10. The ratio of initial concentrations (fourth column) varies widely but always gives the same ratio of equilibrium concentrations (rightmost column). The individual equilibrium concentrations are different in each case, but the ratio of these equilibrium concentrations is constant. Thus, monitoring Q tells whether the system has reached equilibrium or, if it hasn’t, how far away it is and in which direction it is changing.

11. The curves in Figure show experiment in the last slide . Note that [N2O4] and [NO2] change smoothly during the course of the reaction (as indicated by the changing brown color at the top), and so does the value of Q . Once the system reaches equilibrium (constant brown color), the concentrations no longer change and Q equals K . In other words, for any given chemical system, K is a special value of Q that occurs when the reactant and product concentrations have their equilibrium values.

12. Writing the Reaction Quotient (Q) Constructing the Reaction Quotient The most common form of the reaction quotient shows reactant and product terms as molar concentrations, which are designated by square brackets, [ ]. Thus, from now on, we designate the reaction quotient based on concentrations as Qc. (We also designate the equilibrium constant based on concentrations as Kc.) For the general equation:

13. Therefore: Q is a ratio of product concentration terms multiplied together and divided by reactant concentration terms multiplied together, with each term raised to the power of its balancing coefficient. Two steps in writing reaction quotient: 1. Start with the balanced equation. 2. Arrange the terms and exponents.

14. Example

15. Example: Evaluating a Reaction Quotient Gaseous nitrogen dioxide forms dinitrogen tetroxide according to this equation: When 0.10 mol NO2 is added to a 1.0-L flask at 25 °C, the concentration changes so that at equilibrium, [NO2] = 0.016M and [N2O4] = 0.042M. (a) What is the value of the reaction quotient before any reaction occurs? (b) What is the value of the equilibrium constant for the reaction?

16. Solution:

17. Practice: For the reaction 2SO 2 (g) + O 2 (g) ⇌ 2SO 3 (g), the concentrations at equilibrium are [SO2] = 0.90M, [O2] = 0.35M, and [SO3] = 1.1M. What is the value of the equilibrium constant, Kc?

18. By its definition, the magnitude of an equilibrium constant explicitly reflects the composition of a reaction mixture at equilibrium, and it may be interpreted with regard to the extent of the forward reaction. A reaction exhibiting a large K will reach equilibrium when most of the reactant has been converted to product, whereas a small K indicates the reaction achieves equilibrium after very little reactant has been converted.

19. Predicting the Direction of Reaction Determine in which direction the reaction proceeds as it goes to equilibrium in each of the three experiments shown. If Q < K Forward Reaction If Q = K Equilibrium If Q > K Reverse Reaction

20. Solution

21. Practice Calculate the reaction quotient and determine the direction in which each of the following reactions will proceed to reach equilibrium.

22. Why Q and K are unitless each term in Q represents the ratio of the quantity of the substance (molar concentration or pressure) to its thermodynamic standard-state quantity. the standard states are 1 M for a substance in solution, 1 atm for a gas, and the pure substance for a liquid or solid.

23. Writing Q for an Overall Reaction If an overall reaction is the sum of two or more reactions, the overall reaction quotient (or equilibrium constant) is the product of the reaction quotients (or equilibrium constants) for the steps:

24. Example

25. Writing Q for a Forward and a Reverse Reaction A given reaction quotient depends on the direction in which the balanced equation is written. Consider, for example, the oxidation of sulfur dioxide to sulfur trioxide. This reaction is a key step in acid rain formation and sulfuric acid production. The balanced equation is The reaction quotient for this equation as written is

26. If we had written the reverse reaction, the decomposition of sulfur trioxide, the reaction quotient would be the reciprocal of Qc( fwd ):

27. Thus, a reaction quotient (or equilibrium constant) for a forward reaction is the reciprocal of the reaction quotient (or equilibrium constant) for the reverse reaction:

28. Homogeneous Equilibira A homogeneous equilibrium is one in which all reactants and products (and any catalysts, if applicable) are present in the same phase. By this definition, homogeneous equilibria take place in solutions. These solutions are most commonly either liquid or gaseous phases,

29. For gas-phase solutions, the equilibrium constant may be expressed in terms of either the molar concentrations (Kc) or partial pressures ( Kp ) of the reactants and products. A relation between these two K values may be simply derived from the ideal gas equation and the definition of molarity: where P is partial pressure, V is volume, n is molar amount, R is the gas constant, T is temperature, and M is molar concentration.

30. For the gas-phase reaction

31. Example

32. Assigment :

33. Heterogeneous Equilibria A heterogeneous equilibrium involves reactants and products in two or more different phases, as illustrated by the following examples: N ote that concentration terms are only included for gaseous and solute species, as discussed previously.

34. Two of the above examples include terms for gaseous species only in their equilibrium constants, and so Kp expressions may also be written:

35. Coupled Equilibria Many systems involve two or more coupled equilibrium reactions, those which have in common one or more reactant or product species. Since the law of mass action allows for a straightforward derivation of equilibrium constant expressions from balanced chemical equations, the K value for a system involving coupled equilibria can be related to the K values of the individual reactions.

36. Three basic manipulations are involved in this approach, as described below. 1. Changing the direction of a chemical equation essentially swaps the identities of “reactants” and “products,” and so the equilibrium constant for the reversed equation is simply the reciprocal of that for the forward equation. 2. Changing the stoichiometric coefficients in an equation by some factor x results in an exponential change in the equilibrium constant by that same factor:

37. 3. Adding two or more equilibrium equations together yields an overall equation whose equilibrium constant is the mathematical product of the individual reaction’s K values: The net reaction for these coupled equilibria is obtained by summing the two equilibrium equations and canceling any redundancies: Comparing the equilibrium constant for the net reaction to those for the two coupled equilibrium reactions reveals the following relationship:

38. Example: Equilibrium Constants for Coupled Reactions A mixture containing nitrogen, hydrogen, and iodine established the following equilibrium at 400 °C: Use the information below to calculate Kc for this reaction.

39. Solution The equilibrium equation of interest and its K value may be derived from the equations for the two coupled reactions as follows. Reverse the first coupled reaction equation: Multiply the second coupled reaction by 3: Finally, add the two revised equations:

40. Practice