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Chapter 8: Acids & Bases

Chapter 8: Acids & Bases. Acids are all around us. Foods like lemons, limes, and oranges contain acidic compounds. When wine is oxidized, it becomes acetic acid. Our muscles produce lactic acid. Our stomach uses hydrochloric acid to break down food. Acids.

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Chapter 8: Acids & Bases

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  1. Chapter 8: Acids & Bases • Acids are all around us. • Foods like lemons, limes, and oranges contain acidic compounds. • When wine is oxidized, it becomes acetic acid. • Our muscles produce lactic acid. • Our stomach uses hydrochloric acid to break down food.

  2. Acids • Arrhenius definition – an acid is a substance that produces H+ ion when added to water. HCl(aq) H+(aq) + Cl-(aq) • Acids taste sour, are electrolytes, and neutralize bases.

  3. Bases • Arrhenius definition – a base is a substance that produces OH- when added to water. NaOH(aq) Na+(aq) + OH-(aq) • Bases taste bitter, have a slippery, soapy feel, and neutralize acids.

  4. Bronsted-Lowery Theory • Expanded definition of acids & bases. • Acid = a proton (H+) donor • Base = a proton acceptor • Important note about the use of H+ in equations. • For weak acids and bases, the reactions are reversible, equilibria.

  5. Bronsted-Lowery Theory • HC2H3O2 + H2O  H3O+ + C2H3O2- • NH3 + H2O  NH4+ + OH-

  6. Conjugates • Write the conjugate base of an acid. • Remove an H+ 1. HBr – H+ = Br− 2. H2S – H+ = HS− • Write the conjugate acid of a base. • Add an H+ 1. NO2− + H+ = HNO2 2. NH3 + H+ =NH4+

  7. Learning Check 1. The conjugate base of HCO3− is a. CO32− b. HCO3− c. H2CO3 2. The conjugate acid of HCO3− is a. CO32− b. HCO3−c. H2CO3 3. The conjugate base of H2O is a. OH−b. H2O c. H3O + 4. The conjugate acid of H2O is a. OH−b. H2O c. H3O+

  8. Strong Acids • Strong acids completely ionize (dissociate) in water. • Strong acids are also strong electrolytes. • There are six: HClO4, H2SO4, HI, HBr, HCl, and HNO3.

  9. Weak Acids • Weak acids only partially dissociate to produce ions in solution. • Weak acids are weak electrolytes. • Too many to list, but some common ones are: H3PO4, HC2H3O2, HF, and HC6H7O6.

  10. Comparison

  11. Strong Bases • Strong bases completely ionize in water and are also strong electrolytes. • Only group 1A and 2A hydroxides are strong bases. • All other metal hydroxides are insoluble in water.

  12. Weak Bases • Weak bases only partially react with water (accepting a proton) and are, thus, weak electrolytes. • Ammonia, NH3, is the most common weak base. • Lone pair on N group will accept a proton. • Other organic weak bases.

  13. Ionization of Water • Water can act as both an acid and a base. • Any substance that does this is called amphoteric. • Pure water – two water molecules will occasionally react with each other where one is an acid and one is a base.

  14. Ionization of Water • H2O + H2O  H3O+ + OH- • [ X ] = symbol for molarity. • In pure water, [H3O+] = [OH-]. • Kw = [H3O+] x [OH-]; where Kw = 1 E-14. • Thus, [H3O+] = [OH-] = ____________

  15. Acidic, Basic, and Neutral Solutions

  16. Calculating [H3O+] and [OH-] • If we know [H3O+], then [OH-] = • If we know [OH-], then [H3O+] = • Ex) [H3O+] = 2.5 E-5 • Ex) [OH-] = 4.8 E-3

  17. pH Scale • One convenient method for measuring the acidity or basicity of a solution is to use the pH scale. • pH is a logarithmic (log) scale and equal to: pH = -log[H3O+]. • pOH = -log[OH-]. • pH + pOH = 14.

  18. pH Scale • A word about significant figures. • 2.4 x 10-3M pH = 2.62 • Red numbers are the significant digits. • Blue numbers are exact numbers.

  19. Basic Calculators Enter concentration Press “log” key Change the sign Record answer to proper s.f.’s TI-83 or TI-89 Press negative sign Press “log” key (select in catalog) Enter concentration, close parenthesis Enter key and round to proper s.f.’s Guide to Using Your Calculator

  20. pH to Concentration • To convert a pH back to a concentration, you will use the “antilog” key = 10x. • [H3O+] = 10-pH • [OH-] = 10-pOH • Can also use the universal power (^) key.

  21. Comparison of Values

  22. Measuring pH • Can be done with a meter, pH paper, or an indicator.

  23. Fill in the Chart

  24. Reactions of Acids • An acid will react with most metals. • Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)

  25. Reactions of Acids • Acids react with any carbonate (CO3-2) and bicarbonate (HCO3-) to generate CO2.

  26. Environmental Note: Acid Rain • Normally, rain is slightly acidic – pH of 5.5 to 6.2 – due to dissolved CO2. • Burning fossil fuels, which contain small amounts of Sulfur, produces SO3. • This is converted to H2SO4. • H2SO4 + CaCO3 CaSO4 + H2O + CO2

  27. Neutralization • Acids neutralize bases and vice versa. • Acid + Base  Water + Salt • HCl + NaOH  H2O + NaCl • Balancing more complex acid-base neutralization. • Each H+ needs one OH-. • Each H+ and OH- makes one H2O.

  28. Solution Stoichiometry • Acid-base titration – can use a known acid or base solution to determine an unknown counterpart. • Endpoint – when all of the unknown acid or base has reacted. • Indicator – a substance that changes color when it changes pH.

  29. Solution Stoichiometry • Requires precise glassware to deliver the known solution = buret. • Stopcock allows for delivery drop-by-drop. • Volumes can be read to nearest 0.05mL.

  30. Buffers • When a small amount of acid or base is added to pure water, the pH swings drastically. • Some solutions, though, will resist these wild swings in pH and are called buffer solutions. • Made from a ___________ and a salt containing the ______________.

  31. Buffers • 1.0L of pure water + 0.0100moles (0.365g) of HCl. • [H3O+] = 0.010 mol / 1.0L = 0.010M • pH = • 1.0L of pure water + 0.0100moles (0.400g) of NaOH. • [OH-] = 0.010 mol / 1.0L = 0.010M • pH =

  32. Buffers • Buffer of HF and NaF (note: Na+ is a spectator ion). • Ideal buffer would contain a 50 / 50 mixture of each. • 1.0L of 0.10 moles HF (2.0g) and 0.10 moles of NaF (4.2g) will have a pH of 3.17.

  33. Buffers • HF + H2O  H3O+ + F- 50% 50% (from NaF) • Addition of strong acid • reacts with F- ion to generate more HF • Addition of strong base • reacts with HF to produce water plus more F-

  34. Buffers • Buffer plus 0.010 moles of HCl. • pH = 3.08 (from 3.17). • Buffer plus 0.010 moles of NaOH. • pH = 3.26 (from 3.17).

  35. Buffers in Blood • Normal blood pH is 7.35 to 7.45. • Outside this range, cells cannot function properly. • Two buffer systems are present to maintain this pH. • H2CO3 / HCO3- • H2PO4- / HPO4-2

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