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Chapter 19

Chapter 19. Acids and Bases. 19.1 Properties of acids and Bases. Acids Tastes sour Burns Reacts w/metal to make H 2 Electrolyte Formula usually starts with “H” Indicators Litmus: blue to red Phenolphthalein: red to clear. Bases Tastes bitter Slippery

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Chapter 19

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  1. Chapter 19 Acids and Bases

  2. 19.1 Properties of acids and Bases Acids • Tastes sour • Burns • Reacts w/metal to make H2 • Electrolyte • Formula usually starts with “H” • Indicators • Litmus: blue to red • Phenolphthalein: red to clear Bases • Tastes bitter • Slippery • Does not react with metals • Electrolyte • Formula usually ends with “OH” • Indicators • Litmus: red to blue • Phenolphthalein: clear to red

  3. H+ OH- H+ OH- H+ OH- 19.1 Acidic or Basic Ionization of Water H2O(l)qe H+(aq) + OH-(aq) • Normally, [H+] = [OH-] • Acidic solution : [H+] is greater than [OH-] • Basic solution : [OH-] is greater than [H+]

  4. 19.1 Arrhenius Model (1887) • Acid - compound that ionizes to make H+ (proton). • monoprotic acid - 1 ionizable hydrogen (a proton), HCl. • diprotic acid - 2 protons, H2SO4. • triprotic acid - 3 protons, H3PO4. • Base - compound that ionizes to make OH-.

  5. Arrhenius definition is easy to use, but is too limited. Some things act like acids and bases without H+ or OH-.

  6. 19.1 Bronsted- Lowery Model (1923) • Acid - proton or hydrogen-ion donor. HX + H2O qe H3O+ + X- • Base - proton or hydrogen-ion acceptor. Base Acid

  7. 19.1 Bronsted- Lowery Model (1923) NH3 + H2O qe NH4+ + OH- In the forward reaction, a hydrogen-ion or proton, leaves water and combines with NH3 to make NH4+. H2O is the proton donor, acid, and NH3 is the proton acceptor, base. Base Acid

  8. 19.1 Bronsted- Lowery Model (1923) NH3 + H2O qe NH4+ + OH- In the reverse reaction, NH4+ donates (acid) the proton to OH-(base). Because they are formed from NH3 and H2O, they are called conjugates (differ by a proton). Conjugate Acid Conjugate Base Base Acid

  9. 19.1 Bronsted- Lowery Model (1923) • Conjugate acid-base pair – two substances that differ from each other by a proton. NH3 + H2O qe NH4+ + OH- Conjugate Acid Conjugate Base Base Acid

  10. 19.1 Bronsted- Lowery Model (1923) • Amphoteric - a substance that can act as both an acid or a base. HCO3- + H2O qe H2CO3 + OH- HCl + H2O qe H3O+ + Cl-

  11. 19.2 Strong Verses Weak Acids and Bases • Strong Acid/Base • Acids and bases that ionize completely • This means that they are converted into 100% products and are represented in the reaction as a single arrow pointing to the products.

  12. 19.2 Strong Verses Weak Acids and Bases Strong Acid HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Strong Base NaOH(s) Na+(aq) + OH-(aq)

  13. 19.2 Strong Verses Weak Acids and Bases • Weak Acids and Bases • Weak acids and bases only partially ionize and are therefore equilibrium reactions. Weak Acid HC2H3O2(aq) + H2O(l)qe H3O+(aq) + C2H3O2-(aq) Weak Base CH3NH2(aq) + H2O(l)qe CH3NH3+(aq) + OH-(aq)

  14. 19.2 Strong Verses Weak Acids and Bases • Because they are equilibrium reactions, we can determine the Keq HCl(aq) + H2O(l)qe H3O+(aq) + Cl-(aq) Kashould be small for a weak acid The larger the Ka the stronger the acid

  15. 19.2 Strong Verses Weak Acids and Bases • The same applies for weak bases • Keq = Kb • Kb is small for a weak base • The larger the Kb the stronger the base

  16. Concentrated/dilute vs. Strong/weak Concentrated Weak Acid Dilute Strong Acid Concentrated Strong Acid Dilute Weak Acid

  17. 19.3 What is pH? • Remember: H2O(l) H+(aq) + OH-(aq) • Normally, [H+] = [OH-] • Acidic solution : [H+] > [OH-] • Basic solution : [OH-] > [H+]

  18. 19.3 What is pH? H2O(l) H+(aq) + OH-(aq)

  19. 19.3 What is pH? At 298 K for pure water [H+] and [OH-] both equal 110-7M. [H+][OH-] = 110-14

  20. 19.3 What is pH? What is the [OH-] if the [H+]=2.310-4? [H+][OH-] = 110-14 2.310-4[OH-] = 110-14 [OH-] = 4.310-11 M

  21. 19.3 What is pH? Because [H+] and [OH-] are such small numbers, scientists invented pH to simplify these odd numbers! pH = 7 neutral pH < 7 acidic pH > 7 basic

  22. 19.3 What is pH? • Where does pH come from? • p (in pH) means the negative log of the H+ concentration What is the pH if [H+]=2.310-4? pH = 3.6 pH = -log[H+] pOH = -log[OH-]

  23. 19.3 What is pH? • How do you find the original concentration from pH and pOH? • Just do the opposite! • To go to pH we used the -log[H+] • [H+] = antilog(-pH) • [OH-] = antilog(-pOH)

  24. 19.3 What is pH? • How are pH and pOH related? pH + pOH = 14 pOH 14 13 12 11 10 9 8 7 6 5 4 3 2 1 0

  25. Finding pH from solution concentrations Strong acids • Monoprotic acids fully ionize HCl  H+ + Cl- • [H+] = solution concentration • 0.002 M HCl  [H+] = 0.002 M Strong bases • All strong bases fully ionize NaOH  Na+ + OH- Ca(OH)2 Ca2+ + 2OH- • [OH-] = (solution concentration)(# hydroxide ions) • 0.002 M NaOH  [OH-] = 0.002 M • 0.002 M Ca(OH)2  0.004 M

  26. Types of Acid-Base Reactions • The reaction of an acid and a base is called a neutralization reaction. • Hydrochloric acid, HCl, is a common household and laboratory acid. • Sodium hydroxide, NaOH, is a common strong base found in drain cleaning

  27. Neutralization Reactions • A solution of hydrochloric acid, HCl, is added to exactly the amount of a solution of basic sodium hydroxide, NaOH, that will react with it. Click box to view movie clip.

  28. Neutralization Reactions • Litmus papers show that the resulting salt solution is neither acidic nor basic.

  29. Strong Acid + Strong Base • A typical type of acid-base reaction is one in which both the acid and base are strong.

  30. Strong Acid + Strong Base • Instead of an overall equation, an ionicequation, in which substances that primarily exist as ions in solution are shown as ions, can be written.

  31. Spectator Ions and the Net Ionic Reaction • When ions common to both sides of the equation are removed from the equation, the result is called the net ionic equation for the reaction of HCl with NaOH.

  32. Acid-Base Titrations • The general process of determining the molarity of an acid or a base through the use of an acid-base reaction is called an acid-base titration. • The known reactant molarity is used to find the unknown molarity of the other solution. • Solutions of known molarity that are used in this fashion are called standard solutions. • In a titration, the molarity of one of the reactants, acid or base, is known, but the other is unknown.

  33. Acid-Base Titrations • You know that NaOH and HCl react completely. • You know the concentration of the NaOH solution, so it is your standard solution.

  34. Acid-Base Titrations • You can use the reaction, the volumes of acid and base used, plus the molarity of the base to determine the molarity of the unlabeled HCl.

  35. Acid-Base Titrations

  36. Titration • At neutralization, moles H+ = moles OH- • How do you find moles?

  37. Titration At neutralization (equivalence point), moles H+ = moles OH- MaVa = MbVb

  38. Titration What is the molarity of a HCl solution if it takes 68 mL of 1.2 M NaOH to neutralize 25 mL of HCl? HCl + NaOH  NaCl + HOH If [NaOH] = 1.2 M, what does [OH-] = MaVa = MbVb X(25 mL)=(1.2 M)(68 mL) X = 3.3 M 1.2 M

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