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Chapter 16 Acids and Bases

Chemistry, The Central Science , 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten. Chapter 16 Acids and Bases. John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc. HW. CHAPTER 16 –ACID-BASE EQUILIBRIUM

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Chapter 16 Acids and Bases

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  1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 16Acids and Bases John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

  2. HW • CHAPTER 16 –ACID-BASE EQUILIBRIUM • Bronsted Lowry 15,16,17,19,23,25 • Kw 29 • pH scale 35, 37 • Strong acids and bases 41, 43 (a to c) • 45, 47 • WEAK ACIDS 53, 55, 57, 61-a , 65 • 73, 75, 81, 85, 87 • 91 to 110 red problems ONLY

  3. Electrolytes • Substances that dissociate into ions when dissolved in water. • Anonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

  4. Electrolytes and Nonelectrolytes Soluble ionic compounds tend to be electrolytes.

  5. Electrolytes and Nonelectrolytes Molecular compounds tend to be nonelectrolytes, except for acids and bases.

  6. Electrolytes • A strong electrolyte dissociates completely when dissolved in water. • A weak electrolyte only dissociates partially when dissolved in water.

  7. Strong Electrolytes Are… • Strong acids

  8. Strong Electrolytes Are… • Strong acids • Strong bases

  9. Some Definitions • Arrhenius • Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions. • Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions.

  10. Some Definitions • Brønsted–Lowry • Acid: Proton donor • Base: Proton acceptor

  11. A Brønsted–Lowry acid… …must have a removable (acidic) proton. A Brønsted–Lowry base… …must have a pair of nonbonding electrons.

  12. If it can be either… ...it is amphiprotic. HCO3− HSO4− H2O

  13. What Happens When an Acid Dissolves in Water? • Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid. • As a result, the conjugate base of the acid and a hydronium ion are formed.

  14. Conjugate Acids and Bases: • From the Latin word conjugare, meaning “to join together.” • Reactions between acids and bases always yield their conjugate bases and acids.

  15. Conjugate Acid-Base Pairs • Whatever is left of the acid after the proton is donated is called its conjugate base. • Similarly, whatever remains of the base after it accepts a proton is called a conjugate acid. • Consider • After HA (acid) loses its proton it is converted into A- (base). Therefore HA and A- are a conjugate acid-base pair. • After H2O (base) gains a proton it is converted into H3O+ (acid). Therefore, H2O and H3O+ are a conjugate acid-base pair. • Conjugate acid-base pairs differ by only one proton.

  16. Example – Identify the acid and the base in each equation, and identify each acid-base pair. • HNO3 + NH3 NO3- + NH4+ • CH3COOH + OH-  H2O + CH3COO- Identify the acid and base for the reverse reaction in each example.

  17. Acid and Base Strength • Strong acids are completely dissociated in water. • Their conjugate bases are quite weak. • Weak acids only dissociate partially in water. • Their conjugate bases are weak bases.The weaker the acid the stronger its conjugate base.

  18. Acid and Base Strength • Substances with negligible acidity do not dissociate in water. • Their conjugate bases are exceedingly strong.

  19. Relative Strengths of Acids and Bases • The stronger the acid, the weaker the conjugate base. • H+ is the strongest acid that can exist in equilibrium in aqueous solution. • OH- is the strongest base that can exist in equilibrium in aqueous solution.

  20. Acid and Base Strength In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. HCl(aq) + H2O(l) H3O+(aq) + Cl−(aq) H2O is a much stronger base than Cl−, so the equilibrium lies so far to the right K is not measured (K>>1).

  21. HC2H3O2(aq) + H2O H3O+(aq) + C2H3O2−(aq) Acid and Base Strength Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1).

  22. H2O(l) + H2O(l) H3O+(aq) + OH−(aq) Autoionization of Water • As we have seen, water is amphoteric. • In pure water, a few molecules act as bases and a few act as acids. • This is referred to as autoionization.

  23. Ion-Product Constant • The equilibrium expression for this process is Kc = [H3O+] [OH−] • This special equilibrium constant is referred to as the ion-product constant for water, Kw. • At 25°C, Kw = 1.0  10−14

  24. The pH Scale • In most solutions [H+(aq)] is quite small. • We define • In neutral water at 25 C, pH = pOH = 7.00. • In acidic solutions, [H+] > 1.0  10-7, so pH < 7.00. • In basic solutions, [H+] < 1.0  10-7, so pH > 7.00. • The higher the pH, the lower the pOH, the more basic the solution.

  25. pH • Therefore, in pure water, pH = −log (1.0  10−7) = 7.00 • An acid has a higher [H3O+] than pure water, so its pH is <7 • A base has a lower [H3O+] than pure water, so its pH is >7.

  26. pH These are the pH values for several common substances.

  27. Other “p” Scales • The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions). • Some similar examples are • pOH −log [OH−] • pKw−log Kw

  28. Watch This! Because [H3O+] [OH−] = Kw = 1.0  10−14, we know that −log [H3O+] + −log [OH−] = −log Kw = 14.00 or, in other words, pH + pOH = pKw = 14.00

  29. February 28 • Measuring pH: a/b indicators • pH meter • Section 16.5 Strong a/b – key points • Section 16.6 Weak acids • Ka • Problems a) Calculating Ka and % Ionization from measured pH and initial concentration

  30. HW • WEAK ACIDS 53, 55, 57, 61-a , 65

  31. Daily Quiz • 1. Write the formation of nitrous acid from its anhydride • 2. Calculate the pH of an aqueous solution of LiOH that has a pH of 12.5

  32. How Do We Measure pH? For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

  33. Most pH and pOH values fall between 0 and 14. • There are no theoretical limits on the values of pH or pOH. (e.g. pH of 2.0 M HCl is -0.301.) • Examples: • Consider a solution with [H+] = 6.2 x 10-3. • Calculate the pH, pOH, and [OH-] • Is this solution acidic or basic? • Consider a solution with pOH = 13.65 • Calculate the pH, [H+], and [OH-] • Is this solution acidic or basic?

  34. Measuring pH • Most accurate method to measure pH is to use a pH meter. • However, certain dyes change color as pH changes. These are indicators. • Indicators are less precise than pH meters. • Many indicators do not have a sharp color change as a function of pH.

  35. How Do We Measure pH? • For less accurate measurements, one can use • Litmus paper • “Red” paper turns blue above ~pH = 8 • “Blue” paper turns red below ~pH = 5 • An indicator

  36. Strong Acids • You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4. • These are, by definition, strong electrolytes and exist totally as ions in aqueous solution. • For the monoprotic strong acids, [H3O+] = [acid].

  37. Strong Bases • Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).

  38. Strong bases are strong electrolytes and dissociate completely in solution. • The pOH of a strong base is given by the initial molarity of the hydroxide ion. Be careful of stoichiometry. • In order for a hydroxide to be a base, it must be soluble. • Bases do not have to contain the OH- ion: • O2-(aq) + H2O(l)  2OH-(aq) • H-(aq) + H2O(l)  H2(g) + OH-(aq) • N3-(aq) + 3H2O(l)  NH3(aq) + 3OH-(aq)

  39. Strong basic solutions • Ionic metal oxides Na2O and CaO are used in industry to produce strong basic solutions. • Find the pH of a solution formed by dissolving 0.01 mol of Na2O in enough water to produce a liter of solution. • 12.3

  40. Examples: Calculate the pH and pOH of each: • 0.25 M NaOH • 12.00 M HCl • 6.00 M KOH • 0.0050 M HNO3 • 6.5 x 10-12 M RbOH • 0.10 M HClO4

  41. Examples: Calculate the pH and pOH of each: pH pOH • 0.25 M NaOH 13.40 0.60 • 12.00 M HCl -1.08 15.08 • 6.00 M KOH 14.78 -0.78 • 0.0050 M HNO32.30 11.70 • 6.5 x 10-12 M RbOH 7.00 7.00 • 0.10 M HClO41.00 13.00

  42. Weak Acids • Weak acids are only partially ionized (dissociated) in solution. • There is a mixture of ions and unionized acid in solution. • Therefore, weak acids are in equilibrium:

  43. Kc = [H3O+] [A−] [HA] HA(aq) + H2O(l) A−(aq) + H3O+(aq) Dissociation Constants • For a generalized acid dissociation, the equilibrium expression would be • This equilibrium constant is called the acid-dissociation constant, Ka.

  44. Dissociation Constants • The larger the Ka the stronger the acid (i.e. the more ions are present at equilibrium relative to unionized molecules). • If Ka >> 1, then the acid is completely ionized and the acid is a strong acid

  45. Percent Ionization • Percent ionization is another method to assess acid strength. • For the reaction

  46. Calculating Ka from pH • Weak acids are simply equilibrium calculations. • The pH is used to calculate the equilibrium concentration of H+. • Write the balanced chemical equation clearly showing the equilibrium. • Write the equilibrium expression. • Use an ICE Chart / find the [H+] from the pH • Find the value for Ka.

  47. [H3O+] [COO−] [HCOOH] Ka = Calculating Ka from the pH • The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature. • We know that

  48. Calculating Ka from the pH • The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature. • To calculate Ka, we need the equilibrium concentrations of all three things. • We can find [H3O+], which is the same as [HCOO−], from the pH.

  49. Calculating Ka from the pH pH = −log [H3O+] 2.38 = −log [H3O+] −2.38 = log [H3O+] 10−2.38 = 10log [H3O+] = [H3O+] 4.2  10−3 = [H3O+] = [HCOO−]

  50. Calculating Ka from pH Now we can set up a table…

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