Unit 9 Chemical Equilibrium & Acid-Base Chemistry

1. Unit 9Chemical Equilibrium & Acid-Base Chemistry

2. Reversible Reactions • A chemical reaction in which the products can react to re-form the reactants is called areversible reaction. • A reversible reaction is written using double arrows to show that the reaction is proceeding in both directions. Example:

3. Dynamic Equilibrium • A reversible reaction reaches dynamic equilibrium when therate of its forward reaction equalsthe rate of its reverse reaction and the concentrations of its products and reactants remain unchanged. • At equilibrium, both reactions continue, but there is no net change in the composition of the system. Visual Concept

4. Equilibrium  Equal • At equilibrium, the rates of the forward and reverse reactions are equal. But the concentrations aren’t necessarily equal. • Some reactions reach equilibrium only after almost all reactants are consumed (productsare favored.) • Others reach equilibrium when only a small percentage of reactants are consumed (reactants are favored.)

5. Le Châtelier’s Principle • Le Châtelier’sPrinciple: When a system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance. • A shift in equilibrium will result from a change to any of the following: • Concentration • Volume/Pressure • Temperature

6. Change in Concentration • If you increase the concentration byadding more of areactant or product, the system will shift to produce less of that substance. • If you lower the concentration by removing some of a reactant or product, the system will shift to produce more of that substance.

7. Change in Volume/Pressure • When you increase the pressure (usually by decreasing volume of the container), the system shifts so the least numberof gas molecules are formed (less collisions= lower pressure.) • When you decrease thepressure, the system willshift so the greatest numberof gas molecules are formed.

8. Change in Temperature • For every reversible reaction, one direction is endothermic and the other is exothermic. • If the temperature is increased, the endothermic reaction will be favored (because it takes in some of the excess heat.) • If the temperatureis decreased,the exothermicreaction will befavored (produces heat.)

9. Le Châtelier’sPrincipleSample Problem 2 SO2(g) + O2(g) Û 2 SO3(g) DH° = -198 kJ How will the reversible reaction above shift in response to each of the following stresses? • adding more O2 to the container • condensing and removing SO3 • compressing the gases • cooling the container • doubling the volume of the container • warming the mixture Shift right Shift right Shift right Shift right Shift left Shift left

10. The Law of Mass Action • The relationship between the chemical equation and the concentrations of reactants and products is called the Law of Mass Action. • for the general equation aA + bB cC + dD, the Law of Mass Action is: • Lowercase letters represent coefficients. • Always products over reactants. • Pure solids and pure liquids are not included. [C]c [D]d K = [A]a [B]b

11. The Equilibrium Constant • The equilibrium constant (K) reflects how the concentrationsof the reactants and products compare at equilibrium. • It can also be a ratio of pressures(in atmospheres)if a reaction involves gaseous reactants and/or products. • K is unitless.

12. The Value of K • K > 1 more product molecules present than reactant molecules (the position of equilibrium favors products.) • K < 1 more reactant molecules present than product molecules (the position of equilibrium favors reactants.) • K = 1 reactant and product particles are present in exact equal concentrations at equilibrium.

13. Equilibrium ConstantSample Problem Equilibrium concentrations of [H2] = 0.033 M, [I2] = 0.53 M and [HI] = 0.934 M were observed at 445oC for the reaction: H2(g) + I2(g) Û 2 HI(g) Write an equilibrium expression for the above reaction. 2. Calculate the value of Kc for this reaction at 445oC. [HI]2 K = [H2][I2] [0.934]2 [HI]2 = 49.9 = Kc = [0.033] [0.53] [H2][I2]

14. The Reaction Quotient • When a reaction is not at equilibrium, how do you know in which direction it will proceed? • the answer is to compare the equilibrium constant to a ratio of current concentrations called the reaction quotient (Q). • for the general equation aA + bB cC + dD: • The non-equilibrium concentrations (or pressures) are used. [C]c [D]d Q = [A]a [B]b

15. Q vs. K • We calculate Q in order to compare it with K. • Q < K means the reaction will proceed in the forward direction ([products] increase and [reactants] decrease.) • Q > K means the reaction will proceed in the reverse direction ([products] decrease and [reactants] increase.) • Q = K means the reaction is at equilibrium ([products] and [reactants] will not change.)

16. Q, K, and the Direction of Reaction

17. Reaction QuotientSample Problem For the reaction below, which direction will it proceed if PI2 = 0.114 atm, PCl2 = 0.102 atm & PICl = 0.355 atm? I2(g) + Cl2(g)  2 ICl(g) Kp = 81.9 First calculate Q: Then, compare it with K: (0.355)2 (ICl)2 = 10.8 = Q = (0.114) (0.102) (I2)(Cl2) Reaction will proceed to the right Q < K 10.8 81.9

18. Properties of Acids • Taste sour. • React with metals to release H2 gas. • React with bases to produce salts and water. • Change the color of acid-base indicators. • Conduct electric current.

19. Properties of Bases • Taste bitter. • Feel slippery. • React with acids to produce salts and water. • Change the color of acid-base indicators. • Conduct electric current.

20. Arrhenius Acids and Bases • An Arrhenius acidproduces hydrogen ions, H+, in aqueous solution. • An Arrhenius base produces hydroxide ions, OH−, in aqueous solution. • A strong acid (or base) ionizes completely. • A weak acid (or base) releases only a few ions.

21. HCl ionizes in water, producing H+ and Cl– ions NaOH dissociates in water, producing Na+ and OH– ions Arrhenius Theory

22. Hydronium Ion • The H+ ions (protons) produced by the acid are so reactive they cannot exist in water. • instead, they react with a water molecule to form a hydronium ion, H3O+. H+ + H2O  H3O+ • Chemists use H+ and H3O+ interchangeably.

23. Brønsted-Lowry Acids and Bases • In a Brønsted-Lowry Acid-Base reaction, an H+ ion (proton) is transferred. • Does not have to take place in aqueous solution. • Broader definition than Arrhenius. • ABrønsted-Lowry acid is a molecule or ion that is a proton donor. • A Brønsted-Lowry base is a molecule or ion that isa proton acceptor.

24. Conjugate Pairs • Brønsted-Lowry theory allows for reversible reactions. • The original base has an extra H+ after the reaction. It will act as an acid in the reverse process. • The original acid has a lone pair of electrons after the reaction. It will act as a base in the reverse process. • each reactant and the product it becomes is called a conjugate pair.

25. Brønsted-Lowry Acid-Base ReactionsSample Problem Identify the Brønsted-Lowry Acids and Bases and Their Conjugates in the Reactions below: a. b. H2SO4 + H2O  HSO4– + H3O+ conjugate base conjugate acid acid base HCO3– + H2O  H2CO3 + HO– conjugate acid conjugate base base acid

26. Amphoteric Compounds • An amphoteric substance is one that can react as either an acid or a base. Example:water • Water can act as an acid. • Water can act as a base. • base acid Visual Concept • acid base

27. Polyprotic Acids • Molecules with more than one ionizable H are called polyproticacids. • 1 H = monoprotic, 2 H = diprotic, 3 H = triprotic(Ex: HCl = monoprotic, H2SO4 = diprotic, H3PO4 = triprotic) • Polyproticacids ionize in steps (each ionizableH removed sequentially.) • Removing the first H automaticallymakes removing the second H harder.(Ex: H2SO4 is a stronger acid than HSO4)

28. Lewis Acids and Bases • A Lewis acidis an atom, ion, or molecule that accepts an electron pair to form a covalent bond. • A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent bond. • The Lewis definition is the broadest of the three definitions.

29. Comparing the Three Definitions Visual Concept

30. Strong and Weak Acids and Bases • Strong acid – fully dissociatesin water (almostevery molecule breaks up to form H+ions. • Weak acid – partially dissociates in water. • Strong base – fully dissociates in water (almostevery molecule breaks up to form OH-ions. • Weak base – partially dissociates in water. A Strong Acid A Weak Acid

31. Ionization Constants for Weak Acids and Bases • The acid ionization constant (Ka) is the equilibrium constant for the ionization reaction of a weak acid: • A base ionization constant (Kb) can also be created for the ionization reaction of a weak base: HA(aq) + H2O(l)  H3O+(aq) + A-(aq) [H3O+] [A-] Ka = [HA] B(aq) + H2O(l)  BH+(aq) + OH-(aq) [BH+] [OH-] Kb = [B]

32. Acid Ionization ConstantSample Problem Calculate the Ka of a 0.100 M solution of acetic acid with a measured [H3O+] of 1.34 x 10-3 M. HC2H3O2(aq) + H2O(l)  H3O+(aq) + C2H3O2-(aq) • For every H3O+ produced, there is also a C2H3O2- produced, so the concentrations must be the same. • The equilibrium concentration of original acid is the original concentration decreased by the amount ionized (0.100M – 1.34 x 10-3 M = 0.0987 M) [H3O+][C2H3O2-] [1.34 x 10-3M] [1.34 x 10-3M] 1.82 x 10-5 = Ka= = [0.0987 M] [HC2H3O2]

33. Ka & Kb and Strength of Acids/Bases • The strength of an acid or base is measured by the size of its equilibrium constant when it reacts with H2O.

34. Six Strong Acids and Bases • Because these acids and bases are known to dissociate (ionize) to essentially 100% completion, it is meaningless to connect them to equilibrium:

35. Strengths ofConjugate Pairs • the stronger an acid is at donating H, the weaker the conjugate base is at accepting H. (i.e. strong acidshave weak conjugatebases, and weak acidshave strong conjugatebases.) Increasing Basicity Increasing Acidity

36. Predicting Acid Strength • The strength of an acid dependson its tendency to ionize (let go of its hydrogen.) • For binary acids, the strength of the acid dependson two factors: • The stronger the bond, the weaker the acid. • The more polar the bond, the stronger the acid.

37. Periodic Trends are a BEAR All of these increase in the direction toward their letter: B=BasicityE=Electronegativity, ionization Energy, & Electron AffinityA=AcidityR=Radius B E A R

38. Predicting Acid Strength (continued) • The strength of Oxyacidsof the form H-O-Y, where Y is any atom (besides H) bonded to O, depends on two factors: • The more electronegative the element Y, the stronger the acid. • The greater the number of oxygen atomsbonded to Y, thestronger the acid.

39. Predicting Acid StrengthSample Problem Predict the relative strengths of the following acids: • HCl, HBr, and HI • HClO, HBrO, and HIO • HNO3 and HNO2 Predict the relative strengths of the following bases: • Cl-, Br- and I- • H2PO3-and H2PO4- • LiOH and Mg(OH)2 HCl < HBr < HI HIO < HBrO < HClO HNO2 < HNO3 I- < Br- < Cl- H2PO4- < H2PO3- LiOH < Mg(OH)2

40. Autoionization of Water • Water is actually an extremely weak electrolyte. • About 1 out of every 10 million water molecules form ions through a process called autoionization. H2O + H2O Û H3O+ + OH– • In pure water at 25°C, [H3O+] = [OH–] = 10-7M.

41. Ion Product of Water • The product of the H3O+ and OH– concentrations is always the same number, called the ion product of water (Kw). at 25°C • If you measure one of the concentrations, you can calculate the other. • As [H3O+] increases the [OH–] must decrease so the product stays constant. Kw = [H3O+] [OH-] = 1 x 10-14

42. Acidic and Basic Solutions • Neutral solutions have equal [H3O+] and [OH–] • [H3O+] = [OH–] = 1 x 10-7 • Acidic solutions have a larger [H3O+] than [OH–] • [H3O+] > 1 x 10-7; [OH–] < 1 x 10-7 • Basic solutions have a larger [OH–] than [H3O+] • [H3O+] < 1 x 10-7; [OH–] > 1 x 10-7

43. Acidic, Basic or Neutral?Sample Problem Calculate [OH] at 25°C when [H3O+] = 1.5 x 10-9 M, and determine if the solution is acidic, basic, or neutral. First calculate [OH-]: Then, compare [H3O+] with [OH-]: [H3O+][OH-] Kw = Kw 1.0 x 10-14 = 6.7 x 10-6 M [OH-] = = 1.5 x 10-9 [H3O+] < [H3O+] [OH-] The solution is basic 1.5 x 10-9 M 6.7 x 10-6 M

44. The pH Scale • The acidity/basicity of a solution is often expressed as pH. • pH is defined as the negative of the common logarithm of the hydronium ion concentration. pH = −log [H3O+] • pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral.

45. pH of Some Common Substances

46. The pH ScaleSample Problem Calculate the pH at 25°C when the [OH] = 1.3 x 10-2 M, and determine if the solution is acidic, basic, or neutral First calculate [H3O+]: Then, calculate pH: [H3O+][OH-] Kw = Kw 1.0 x 10-14 = 7.7 x 10-13 M [H3O+] = = 1.3 x 10-2 [OH-] pH = −log [H3O+] = -log(7.7 x 10-13 M) pH = 12.1 The solution is basic

47. Finding the Ionization Constant from pHSample Problem A 0.100 M weak acid (HA) solution has a pH of 4.25. Find Ka for the acid. First calculate [H3O+]: Then, calculate Ka: -log [H3O+] pH = 5.6 x 10-5 M 10-4.25 10-pH = = = [H3O+] HA(aq) + H2O(l)  H3O+(aq) + A-(aq) [H3O+] [A-] (5.6 x 10-5) (5.6 x 10-5) 3.1 x 10-8 = Ka = = [HA] (0.100 - 5.6 x 10-5)

48. The pOH Scale • Another way of expressing the acidity/basicity of a solution is pOH. • pOH is defined as the negative of the common logarithm of the hydroxide ion concentration. pOH = −log [OH-] • pOH < 7 is basic; pOH > 7 is acidic, pOH = 7 is neutral

49. Relationship Between pH and pOH • The sum of the pH and pOH of a solution is 14. pH + pOH = 14.0

50. Neutralization Reactions • A neutralization reaction is a double displacement reaction in which an acid and a base in an aqueous solution react to produce a salt and water. • A salt is an ionic compound made up of a cation from a base and an anion from an acid.