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Acids & Bases

Acids & Bases. Unit 13. Overview. Acid/base properties Definitions Arrhenius Bronsted-Lowry Lewis Acid-Base Reactions Neutralization Sulfides Carbonates pH/ pOH Scale Acid/base strength Factor affecting K a , K b , K w Percent ionization. Vocabulary Polyprotic Acids

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Acids & Bases

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  1. Acids & Bases Unit 13

  2. Overview • Acid/base properties • Definitions • Arrhenius • Bronsted-Lowry • Lewis • Acid-Base Reactions • Neutralization • Sulfides • Carbonates • pH/pOH Scale • Acid/base strength • Factor affecting • Ka, Kb, Kw • Percent ionization • Vocabulary • Polyprotic Acids • Amphoteric • Anhydrides • Acids/Bases & Salts • Determine acidity • Calculations • Common Ion Effect • Buffers • Henderson-Hasselbalch • Titration • Indicators • 4 types of curves

  3. Acids • Sour taste • React with active metals to produce hydrogen gas • Change the color of acid-base indicators • React with bases to produce salt and water • Conduct an electric current (electrolytes) • Turn litmus paper red

  4. Common Acids • Sulfuric Acid • Car batteries; production of metals, paints, dyes, detergents • Nitric Acid • Explosives, pharmaceuticals, rubber, plastics, dyes • Phosphoric acid • Soda, fertilizers, animal feed, detergents • Hydrochloric Acid • Stomach acid, cleaning metals, found in hardware stores (muriatic acid) • Acetic Acid • Vinegar, food supplements, fungicide • Citric Acid • Fruit juices

  5. Acids • Binary acids • Contain only two different elements • Name as “hydro - ic acid” • Example: HCl (hydrochloric acid) • Oxyacids • Acid consisting of hydrogen and a polyatomic anion that contains oxygen (oxyanion) • To name, drop ending of polyatomic ion and ad “- ic acid” • Example: HNO3 (nitric acid)

  6. Bases • Taste bitter • Feel slippery • Change the color of acid-base indicators • React with bases to produce salt and water • Conduct an electric current (electrolytes) • Turn litmus paper blue

  7. Common Bases • Ammonium hydroxide • Household cleaners, window cleaner • Ammonia • (Gas) inhalant to revive unconscious person • Sodium bicarbonate (baking soda) • Acid neutralizers in acid spills • Antacids for upset stomachs • Sodium hydroxide • Drain cleaner (drano), oven cleaner, production of soap • Magnesium hydroxide • Antacids, milk of magnesia, laxatives

  8. Definition: Arrhenius • Acid • Substance that ionizes in water and produces H+ ions • Example: HCl H+ + Cl- • Base • Substance that ionizes in water and produces OH- ions • Example: NaOH Na+ + OH-

  9. Definition: Bronsted-Lowry • Acid • Substance that is capable of donating a proton (H+ ion) • Base • Substance that is capable of accepting a proton (H+ ion)

  10. Examples: Bronsted-Lowry HC2H3O2 + H2O ↔ C2H3O2- + H3O+ Acids: HC2H3O2 and H3O+ Bases: H2O and C2H3O2- NH3 + H2O ↔ NH4+ + OH- Acids: H2O and NH4+ Bases: NH3 and OH- Notice that water can act as an acid or a base

  11. Bronsted-Lowry • Conjugate Pair – a BL acid/base pair(one with H+ and one without H+) • Examples: • HC2H3O2 and C2H3O2- • H3O+ and H2O • H2O and OH- • NH4+ and NH3

  12. Bronsted-Lowry • The more easily a substance gives up a proton, the less easily the conjugate base accepts a proton (and vice versa) • The stronger the acid, the weaker the conjugate base • The stronger the base, the weaker the conjugate acid

  13. Acid Base Reactions • Neutralization • Salt + water • Sulfides • Salt + sulfide gas • Carbonates • Salt + CO2 + H2O

  14. Neutralization • Solution of an acid and solution of a base are mixed • Products have no characteristics of either the acid or the base • Acid + Base (metal hydroxide)  salt + water • Salt comes from cation of base and anion of acid HY + XOH  XY + H2O HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

  15. Sulfides • Acid reacts with a sulfide • Gaseous product (H2S) has a foul odor (rotten eggs) • Acid + metal sulfide  salt + hydrogen sulfide • Salt comes from cation of sulfide and anion of acid HY + XS  XY + H2S HCl(aq) + Na2S(aq)  NaCl(aq) + H2S(g)

  16. Carbonates • Carbonates and bicarbonates react with acids HY + XHCO3 XY + H2CO3 H2CO3 is not stable so breaks into H2O and CO2 Then HY + XHCO3 XY + H2O + CO2 HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2CO3 (aq) HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g)

  17. The pH Scale • pH scale: measures concentration of hydrogen ions in solution pH = -log[H+]therefore [H+] = 10-pH • Example: What is the pH of a solution with a [H+] of 1.4×10-5? pH = -log[1.4×10-5] =4.9

  18. pH Scale Acids: pH < 7 Neutral: pH = 7 Bases: pH > 7 Increasing [H+] means decreasing pH Increasing pH means decreasing [H+]

  19. The pOH Scale • pOH scale: measures concentration of hydroxide ions in solution pOH = -log[OH-]therefore [OH-] = 10-pOH • Example: What is the [OH-] of a solution with a pOH of 6.2? [OH-] = 10-6.2= 6.3×10-7

  20. Comparing pH and pOH pH + pOH = 14 • An acid has a pH of 4, what is the pOH? 4 + pOH = 14 pOH = 10

  21. Measuring pH • pH meter • Electrodes measure [H+] • Acid-base indicators • Change color in presence of acid or base (or certain pH ranges) • Litmus paper, phenolphthalein, cabbage juice, methyl orange, thymol blue…

  22. Strong Acids and Bases • Completely ionize in solution (strong electrolytes) • If acid/base is not on this list, it is a weak acid/base

  23. Weak Acids and Bases • Do not completely ionize in water (weak electrolytes) • Common weak acids: • HF, acids with -COOH group • Common weak bases: • NH3

  24. Factors Affecting Acid Strength • Electronegativity of element bonded to H • Binary acids • More electronegative bond = stronger acid • Example: HCl stronger than HBr • Bond Strength • Stronger bonds do not allow hydrogen to dissociate as easily • Reason why HF is not a strong acid (F is most electronegative, but H-F bond is strongest bond) • Stability of Conjugate base • More stable the conjugate base, the stronger the acid

  25. Factors Affecting Acid Strength • For polyatomic ions, the more electronegative the nonmetal, the stronger the acid (when comparing acids with same number of O atoms) • Example: HClO3 is stronger than HBrO3 • For polyatomic ions, when nonmetal is the same, the more O atoms, the stronger the acid • Example: HClO3 is stronger than HClO2

  26. Percent Ionization • Tells us what percent of an acid (or base) is ionized in water • Helps determine the strength of an acid (or base) Percent Ionization = ×100 [H+] at equilibrium Initial Acid Concentration

  27. Percent Ionization (Example) • A 0.035 M solution of HNO2 contains 3.7×10-3 M H+(aq). Calculate the percent ionization. = = 11% This means that 11% of the acid will dissociate in water. 3.7×10-3 M 0.035 M

  28. Strong Acids HA(aq) + H2O(l)  H+(aq)+ A-(aq) acid water protonconjugate base • Strong acids dissociate completely • The dissociation is not reversible • The acid is the only significant source of H+ ions, so pH can be calculated directly from the [H+] • Example: A 0.20 M solution of HNO3 has an [H+] of 0.20 M • pH = -log[H+]

  29. Strong Bases • Strong bases dissociate completely • The dissociation is not reversible • The base is the only significant source of OH- ions, so pOH can be calculated directly from the [OH-] • Example: A 0.30 M solution of NaOH has a [OH-] of 0.30 M • pOH = -log[OH-]

  30. Equilibrium Time!

  31. Weak Acids and Bases • Do not dissociate completely • Reversible reactions • Need to use equilibrium to solve for [H+] K = [Products] [Reactants]

  32. Acid Dissociation HA(aq) + H2O(l)  H+(aq)+ A-(aq) acid water protonconjugate base • Write the equilibrium expression for the acid dissociation constant, Ka.

  33. Base Dissociation B(aq) + H2O(l) ↔BH+(aq)+ OH-(aq) base water conjugate hydroxide acid ion • Write the equilibrium expression for the base dissociation constant, Kb.

  34. Size of K • The greater the Ka, the stronger the acid • The smaller the Ka, the weaker the acid • The greater the Kb, the stronger the base • The smaller the Kb, the weaker the base

  35. Example: Weak Acid Equilibrium Problem • What is the pH of a 0.50 M solution of acetic acid, HC2H3O2, Ka = 1.8 x 10-5 ? • Step #1: Write the dissociation equation HC2H3O2 C2H3O2- + H+

  36. Example: Weak Acid Equilibrium Problem • What is the pH of a 0.50 M solution of acetic acid, HC2H3O2, Ka = 1.8 x 10-5 ? • Step #2: ICE HC2H3O2 C2H3O2- + H+ 0.50 0 0 +x +x - x x x 0.50 - x

  37. Example: Weak Acid Equilibrium Problem • What is the pH of a 0.50 M solution of acetic acid, HC2H3O2, Ka = 1.8 x 10-5 ? • Step #3: Set up the equilibrium expression If percent ionization is less than 5%, you can ignore using the quadratic.

  38. Example: Weak Acid Equilibrium Problem • What is the pH of a 0.50 M solution of acetic acid, HC2H3O2, Ka = 1.8 x 10-5 ? • Step #5: Solve for pH You can use the Kb expression to solve for pOH using the same method!

  39. Example 2: Weak Acid Equilibrium - Solving for Ka • A student prepared a 0.10 M solution of formic acid (HCOOH) and measured its pH. The pH at 25°C was found to be 2.38. Calculate the Ka for formic acid at this temperature. • Step #1: Solve for [H+] from pH [H+] = 10-2.38 = 4.2×10-3 M

  40. Example 2: Weak Acid Equilibrium - Solving for Ka • A student prepared a 0.10 M solution of formic acid (HCOOH) and measured its pH. The pH at 25°C was found to be 2.38. Calculate the Ka for formic acid at this temperature. • Step #2: Set up ICE table HCOOH(aq)  HCOO- + H+ 0.10 0 0 4.2×10-3

  41. Example 2: Weak Acid Equilibrium - Solving for Ka • A student prepared a 0.10 M solution of formic acid (HCOOH) and measured its pH. The pH at 25°C was found to be 2.38. Calculate the Ka for formic acid at this temperature. • Step #3: Use stoichiometry to complete table HCOOH(aq)  HCOO- + H+ 0.10 0 0 4.2×10-3 4.2×10-3 4.2×10-3 4.2×10-3 4.2×10-3 0.0096

  42. Example 2: Weak Acid Equilibrium - Solving for Ka • A student prepared a 0.10 M solution of formic acid (HCOOH) and measured its pH. The pH at 25°C was found to be 2.38. Calculate the Ka for formic acid at this temperature. • Step #4: Solve for Ka using equilibrium expression Ka = 1.8×10-4

  43. Rule of Thumb The larger the value of Ka, the stronger the acid

  44. (Self-) Auto-ionization of Water • According to Bronsted Lowry, H2O can act as either an acid or a base • Auto-ionization: One water molecule can donate a proton to another water molecule • Extremely rapid reaction and no molecule remains ionized for long • At room temperature 1 out of every 109 molecule are ionized at a given instant • Water is a nonelectrolyte and consists almost entirely of H2O molecules H2O(l) + H2O(l) ↔ H3O+ + OH-

  45. Auto-ionization of Water • H2O(l) + H2O(l) ↔ H3O+ + OH- • Auto-ionization of water is an equilibrium process (use Kw - ion product constant) Kw = [H3O+][OH-] Also written as Kw = [H+][OH-] • At 25°C, Kw =1.4×10-14

  46. Auto-ionization of Water 1.4×10-14 = [H+][OH-] • In basic solutions, [OH-] > [H+] • In acidic solutions, [H+] > [OH-] • In neutral solutions, [H+] = [OH-]

  47. Auto-ionization of Water 1.4×10-14 = [H+][OH-] • If the concentration of one ion is known, you can solve for the concentration of the other ion Example: Calculate the concentration of H+ in a solution in which the concentration of OH- is 0.010M. 1.4×10-14 = [H+][0.010] [H+] = 1.0×10-12 M

  48. Relating pKw to pKa and pKb • Acid or base dissociation constants are sometimes expressed as pKa and Kb. • pKa = –logKa • pKb = -logKb pKw = 14 = pKa+ pKb

  49. Polyprotic Acids • Acids with more than one ionizable H+ ion • The acid-dissociation constants are Ka1, Ka2, etc… • The first proton is most easily removed • As protons are removed, it becomes more and more difficult to remove protons • Ka1>Ka2>Ka3…. H2SO3(aq) ↔ H+(aq) + HSO3-(aq) Ka1 = 1.7×10-2 HSO3-(aq) ↔ H+(aq) + SO3-2(aq) Ka2 = 6.4×10-8

  50. Polyprotic Acids • To calculate the overall K for the reaction, treat it as a multi-step equilibrium Overall reaction… H2SO3(aq) ↔ H+(aq) + HSO3-(aq) Ka1 = 1.7×10-2 HSO3-(aq) ↔ H+(aq) + SO3-2(aq) Ka2 = 6.4×10-8 H2SO3(aq) ↔ 2H+(aq) + SO3-2(aq) Ka = Ka1 × Ka2 Ka = (1.7×10-2)(6.4×10-8) = 1.1×10-9

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