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Chapter 15: ACIDS AND BASES

Vanessa N. Prasad-Permaul CHM 1046 Valencia Community College. Chapter 15: ACIDS AND BASES. Acid-Base Concepts: The Arrhenius Acid-Base Theory. Generalized Arrhenius Acid HA (aq)  H + (aq) + A - (aq)

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Chapter 15: ACIDS AND BASES

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  1. Vanessa N. Prasad-Permaul CHM 1046 Valencia Community College Chapter 15: ACIDS AND BASES

  2. Acid-Base Concepts: The Arrhenius Acid-Base Theory Generalized Arrhenius Acid HA(aq)  H+(aq) + A-(aq) Acids are substances that dissociate in water to produce hydrogen ions (HYRONIUM IONS, H3O+) Generalized Arrhenius Base MOH(aq) M+(aq) + OH-(aq) Bases are substances that dissociate in water to yield hydroxide ions Arrhenius theory has limitations 1.      Restricted to aqueous solutions 2.      Doesn’t account for the basicity of substances like ammonia (NH3) that don’t contain OH groups

  3. Acid-Base Concepts: The Bronsted-Lowry Theory Bronsted-Lowry Theory 1.  Acid is any substance that can transfer a proton (H+ ion) to another substance a) proton donors 2.   Base is any substance that can accept a proton b)  proton acceptor 3.   Acid-base reactions are proton-transfer reactions 4.   The products of a Bronsted-Lowry acid-base reaction are themselves acids and bases

  4. Acid-Base Concepts: The Bronsted-Lowry Theory 5.   The BH+ produced when base B accepts a proton from HA can itself donate a proton back to A-, BH+ is a Bronsted-Lowry acid 6.   A- produced when HA loses a proton can itself accept a proton back from BH+, A- is a Bronsted-Lowry base 7. Chemical species whose formulas differ only by one proton are said to be conjugate acid-base pairs

  5. Example 1: Which of the following is a Bronsted-Lowry base but not an Arrhenius base? • NaOH • HCl • NH3 • Mg(OH)2

  6. EXAMPLE 15.1: IDENTIFYING ACIDS AND BASES EXERCISE 15.1: IDENTIFYING ACIDS AND BASES LABEL EACH AS ACID OR BASE AND IT’S CONJUGATE PAIR H2CO3(aq) + CN-(aq) HCN(aq) + HCO3-(aq) HSO4- + NH3 SO42- + NH4+

  7. Example 2: A) Write a balanced equation for the dissociation of each of the following Bronsted-Lowry acids in water • What is the conjugate base of the acid in each case? 1. H2SO4(aq) 2. HSO4-(aq) 3. H3O+(aq) 4. NH4+(aq)

  8. Example 3: What is the conjugate acid of each of the following Bronsted-Lowry Bases? • HCO3-(aq) • CO32-(aq) • OH-(aq) • H2PO4-(aq)

  9. Lewis Acids and Bases Lewis acid – An electron pair acceptor Lewis base- An electron pair donor 1. All lewis bases are Bronsted-Lowry bases 2. All Bronsted-Lowry bases are lewis bases 3. Lewis acid is more general than the Bronsted- Lowry (all lewis acids are not also Bronsted- Lowry acids)

  10. Lewis Acids and Bases

  11. Example 12: For each of the following reactions, identify the lewis acid and the lewis base • AlCl3 + Cl- AlCl4- • 2 NH3 + Ag+ Ag(NH)2+ • SO2 + OH- HSO3- • 6 H2O + Cr3+ Cr(OH2)63+

  12. EXAMPLE 15.2: IDENTIFYING LEWIS ACIDS AND BASES EXERCISE 15.2: IDENTIFYING LEWIS ACIDS AND BASES IDENTIFY THE LEWIS ACID AND BASE H BF3 + CH3OH F3B:OCH3 O2- + CO2 CO32- ..

  13. Acid Strength and Base Strength View of an acid-dissociation equilibrium HA(aq) + H2O(l) H3O+(aq) + A-(aq) Acid Base Acid Base 1. Two bases, H2O and A- are competing for protons 2.  If H2O is a stronger base than A- the H2O molecules will get the protons and the solution will be mainly H3O+ and A- 3. If A- is a stronger base than H2O, the A- ions will get the protons and the solution will be mainly HA and H2O 4. With equal concentrations of reactants and products, the proton is always transferred to the stronger base

  14. HA(aq) + H2O(l) H3O+(aq) + A-(aq)Acid Base Acid Base • The direction of the reaction to reach equilibrium is proton transfer from the stronger acid to the stronger base to give the weaker acid and the weaker base 6. Strong acid is one that is almost completely dissociated in water a) Acid-dissociation equilibrium of a strong acid lies nearly 100% to the right (products) b)  Solution contains almost entirely H3O+ and A- ions c)  HClO4, HCl, HBr, HI, HNO3, H2SO4 d) Strong acids have very weak conjugate bases (weaker than H2O)

  15. HA(aq) + H2O(l) H3O+(aq) + A-(aq)Acid Base Acid Base 7. Weak acid is one that is only partially dissociated in water a) Solution contains mainly HA b) HNO2, HF, CH3CO2H c) Acid-dissociation equilibrium lies essentially 100% to the left (reactants) d) Very weak acids have strong conjugate bases

  16. RELATIVE STRENGTHS OF ACIDS AND BASES

  17. EXAMPLE 15.3: REACTANTS OR PRODUCTS IN AN ACID-BASE REACTION EXERCISE 15.3: REACTANTS OR PRODUCTS IN AN ACID-BASE REACTION DETERMINE THE DIRECTION OF THE FOLLOWING REACTION FROM THE RELATIVE STRENGTHS OF ACIDS AND BASES H2S(aq) + C2H302-(aq) HC2H3O2(aq) + HS-(aq)

  18. Example 4: If you mix equal concentrations of reactants and products, which of the following reactions proceed to the right and which proceed to the left? • HF(aq) + NO3-(aq) HNO3(aq) + F-(aq) • NH4+(aq) + CO32-(aq) HCO3-(aq) + NH3(aq)

  19. Dissociation of Water Most important property of water is its ability to act as an acid and a base 1. In the presence of a base water acts as an acid 2. In the presence of an acid water acts like a base Acid Base Acid Base

  20. Dissociation of Water • Dissociation of water a)      2 H2O(l) H3O+(aq) +OH-(aq) b)      Kw = [H3O+][OH-] c)      Kw = ion-product for water d)      Kw = 1 x10-14 @ 25C equilibrium constants are affected by temp e)      [H3O+] > [OH-] acidic f)       [H3O+] = [OH-] neutral g)      [H3O+] < [OH-] basic

  21. Example 5: • The concentration of OH- in a sample of seawater is 5 x 10-6 M. Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, or basic.

  22. EXAMPLE 15.4: CALCULATING [H30+] AND [OH-] EXERCISE 15.5: CALCULATING [H30+] AND [OH-] A SOLUTION OF BARIUM HYDROXIDE @ 25oC IS 0.125M. WHAT IS THE CONCENTRATION OF [H30+] AND [OH-]?

  23. The pH Scale Introduction The term pH is derived from the French puissance d’hydrogene (power of hydrogen) and refers to the power of 10 (the exponent) used to express the molar H3O+ concentration Only works for Strong Acids and Strong Bases pH = -log [H3O+] [H3O+] = 10-pH pOH = -log [OH-] pH + pOH = 14 kw = [H3O+] [ OH-] = 1.0 x10-14 Acidic solution pH < 7 Neutral solution pH = 7 Basic solution pH > 7

  24. Example 6: Calculate the pH of the following solutions • A sample of seawater that has OH- of 1.58 x 10-6 M b) A sample of seawater that has H3O+ 6.0 x 10 –5 M

  25. THE pH SCALE

  26. Example 7: Calculate the concentrations of H3O+ and OH- in each of the following • Human blood pH = 7.40 • A cola beverage pH = 2.8

  27. Example 8: If you have an acidic solution you would expect? • [H3O+] < 1 x 10-7 • [OH-] > 1 x 10-7 • [H3O+] > 1 x 10-7 • [OH-] = 1 x 10-7

  28. Measuring pH Acid-base indicator 1.    The approximate pH of a solution can be determined by using an acid-base indicator a. A substance that changes color in a specific pH range b. HIn abbreviated c.  Exhibit pH-dependent color change 2. The indicators are weak acids and have different colors in their acid (HIn) and conjugate base (In-) forms

  29. Measuring pH 2.   More accurate pH measurements can be made by using a pH meter a) A device that measures the pH-dependent electrical potential of the test solution

  30. The pH in Solutions of Strong Acids and Strong Bases Strong acids are nearly 100% dissociated in aqueous solution a) H3O+ and A- concentrations are equal to the initial concentration of the acid b) Concentration of undissociated HA molecules is essentially zero c) The pH of a solution of a strong monoprotic acid is calculated from the H3O+ concentration pH = -log[H3O+]

  31. The pH in Solutions of Strong Acids and Strong Bases • Calculation of the pH of a diprotic (H2SO4) solution is more complicated because 100% of the H2SO4 molecules dissociate to give H3O+ and HSO4- ions 1. Less than 100% of the resulting HSO4- ions dissociate to give H3O+ and SO42- ions 2. Majority of H3O+ comes from first dissociation

  32. The pH in Solutions of Strong Acids and Strong Bases The most familiar Strong bases are alkali metal hydroxides, MOH, (NaOH, KOH) a) Strong bases are nearly 100% dissociated in aqueous solution b) The pH of a strong base solution can be found 2 ways 1.       Example first way Find the pH of a 0.01 M NaOH

  33. The pH in Solutions of Strong Acids and Strong Bases • Example 2nd way Find the pH of a 0.01 M NaOH**

  34. EXERCISE 15.8: CALCULATING pH FROM [OH-] A SATURATED SOLUTION OF CALCIUM HYDROXIDE HAS AN [OH-] OF 0.025M. WHAT IS THE pH OF THE SOLUTION?

  35. EXAMPLE 15.5: CALCULATING pH FROM [H3O+] EXERCISE 15.7: CALCULATING pH FROM [H3O+] WHAT IS THE pH OF A SAMPLE OF GASTRIC JUICE WHOSE HYDRONIUM ION CONCENTRATION IS 0.045M?

  36. EXAMPLE 15.6: CALCULATING [H3O+] FROM pH EXERCISE 15.9: CALCULATING [H3O+] FROM pH A BRAND OF CARBONATED BEVERAGE HAS A pH OF 3.16. WHAT IS THE HYDRONIUM-ION CONCENTRATION OF THE BEVERAGE?

  37. Equilibria in Solutions of Weak Acids Introduction 1. A weak acid is not the same as a dilute solution of a strong acid 2. A strong acid is 100% dissociated in aqueous solution 3. A weak acid is only partially dissociated

  38. Equilibria in Solutions of Weak Acids The dissociation of a weak acid in water is characterized by an equilibrium equation HA(aq) + H2O(l) H3O+(aq) + A-(aq) Ka = [H3O+] [A-] / [HA] Ka is the acid-dissociation constant Equilibrium equation, pure water is left out!! pKa = -log Ka Larger the Ka the stronger the acid Smaller the pKa the stronger the acid

  39. Calculating Equilibrium Concentrations in Solutions of Weak Acids Once the Ka value for a weak acid has been measured it can be used to calculate equilibrium concentrations and the pH in a solution of the acid Example 9: Calculate the pH and the concentrations of all species present in a 0.10 M HCN (Ka = 4.9 x 10-10) General equation: HCN + H2O  H3O+ + CN-

  40. Example 10: Calculate the pH and the concentrations of all species present (H3O+, F-, HF) in 0.050 M HF (Ka= 6.8 x10-4) ax2 + bx + C = 0

  41. Percent Dissociation in Solutions of Weak Acids % dissociation = ([HA] dissociated / [HA] initial) x 100% 1. The higher the % dissociation the stronger the acid 2. % dissociation increases as Ka increases 3. % dissociation in weak acids increases as dilution increases

  42. Polyprotic Acids • Acids that contain more than one dissociable proton • Dissociate in a stepwise manner • Each dissociation has it’s own Ka labeled Ka1 Ka2 etc. 4. Ka1 > Ka2 > Ka3 it is harder to take a hydrogen from a compound that is already electron rich, proton deficient 5. The principal reaction is the first dissociation 6. Essentially all the H3O+ in the solution comes from the first dissociation

  43. Equilibria in Solutions of Weak Bases Weak bases accept a proton from water to give the conjugate acid of the base and OH- ions NH3(aq) + H2O(l) NH4+(l) + OH-(aq) Equilibrium Expression Kb = [NH4+] [OH-] / [NH3] H2O omitted because it is a pure liquid! Kb = base-dissociation Constant: Stronger base higher the Kb, lower the pKb

  44. Equilibria in Solutions of Weak Bases • Equilibria in solutions of weak bases are treated the same as weak acids. We will use the same procedure Example 11 • Calculate the pH and the concentrations of all species present in 0.40 M NH3 (Kb = 1.8 x 10-5)

  45. Relation Between Ka and Kb NH4+(aq) + H2O(l) H3O+(aq) + NH3(aq) Ka = [H3O+][ NH3] NH4+ NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Kb = [NH4+][OH-] NH3 Net: 2 H2O(l) H3O+(aq) + OH-(aq) Kw = [H3O+][OH-] = 1 x 10-14 Ka x Kb = Kw for conjugate acid-base pairs only!!!

  46. Relation Between Ka and Kb • The equilibrium constant for the net reaction equals the product of the equilibrium constants for the reactions added • For any conjugate acid-base pair, the product of the acid-dissociation constant for the acid and the base-dissociation constant for the base always equals the ion-product constant for water Kw = Ka x Kb • As the strength of an acid increases (larger Ka) the strength of its conjugate base decreases (smaller Kb)

  47. Acid-base Properties of Salts Introduction: • When an acid neutralizes a base, an ionic compound called a salt is formed. • Salt solutions can be neutral, acidic, or basic 3. The pH depends on the acid-base properties of the cations and anions that result from the reaction a) Strong acid + strong base  Neutral sol’n b) Strong acid + weak base  Acidic sol’n c) Weak acid + strong base  Basic sol’n

  48. Acid-base Properties of Salts a) Salts that yield neutral solutions The following ions do not react appreciably with water to produce either H3O+ or OH- ions: Cations from strong bases 1. Alkali metal cations of group 1A (Li+, Na+, K+) 2. Alkaline earth cations of group 2A (Ca2+, Sr2+, Ba2+) except for Be2+ Anions from strong monoprotic acids: 1. Cl-, Br-, I-, NO3-, CLO4- Salts that contain only these ions give neutral solutions in pure water

  49. Acid-base Properties of Salts b) Salts that yield Acidic solutions The anion is neither an acid nor a base but the cation is a weak acid 1. The acidity of hydrated main-group cations increases from left to right in the periodic table, as the metal ion charge increases and the metal ion size decreases (Li+ < Be2+;Na+ < Mg2+ < Al3+) 2. Transition metal cations , (Zn2+, Cr3+, Fe3+)

  50. Acid-base Properties of Salts • salts that yield basic solutions The cation is neither an acid nor a base, but the anion is a weak base 1. NO2-, F-, CH3CO2-, CO2-

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