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Acids, Bases, and Salts

Acids, Bases, and Salts. Characteristics of Acids and Bases. Household Acids and Bases. pH of Some Household Materials. 1.0 battery acid (sulfuric acid) 1.8-2.0 limes 2.2-2.4 lemon juice 2.2 vinegar (acetic acid) 2.8-3.4 fruit jellies 2.9-3.3 apple juice, cola

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Acids, Bases, and Salts

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  1. Acids, Bases, and Salts

  2. Characteristics of Acids and Bases

  3. Household Acids and Bases

  4. pH of Some Household Materials 1.0 battery acid (sulfuric acid) 1.8-2.0 limes 2.2-2.4 lemon juice 2.2 vinegar (acetic acid) 2.8-3.4 fruit jellies 2.9-3.3 apple juice, cola 3.0-3.5 strawberries 3.7 orange juice 4.0-4.5 tomatoes 5.6 unpolluted rain 5.8-6.4 peas 6.0-6.5 corn 6.1-6.4 butter 6.4 cow's milk 6.5-7.5 human saliva 6.5-7.0 maple syrup 7.0 distilled water 7.3-7.5 human blood 7.6-8.0 egg whites 8.3 baking soda 8-9 laundry detergents 9.2 borax 10.5 milk of magnesia 11.0 ammonia 12.0 lime water 13.0 lye 14.0 CSSI 2010 PGCC Barbara A. Gage

  5. Indicators What color are most acids and bases? What do litmus, phenolphthalein, and universal indicator paper have in common? These substances change color depending on whether a solution is acidic, basic or neutral. They are called indicators.

  6. Are Acids and Bases Ionic? How can you tell if a substance is ionic? A solution of the substance will conduct electricity Many bases are ionic compounds that contain the polyatomic ions OH-, HCO3- or CO32-

  7. The electrical conductivity of ionic solutions.

  8. Are Acids and Bases Ionic? Acids are held together by covalent bonding (H form covalent bonds only) so should they conduct electricity? But acids do conduct so what’s happening? The hydrogen in acids is attached to an electronegative element so the bond is very polar. When you put the acid in water the hydrogen will “ionize” and leave its shared electrons with the other atom.

  9. Acids HCl + H2O  H3O+ + Cl- hydronium ion

  10. Strong and Weak Acids and Bases If an acid or base solution is a strong conductor it is called a strong acid or strong base. If the solution is a poor conductor it is called a weak acid or base. Strength of acids or bases IS NOT THE SAME AS concentration.

  11. Strong acid: HA(g or l) + H2O(l) H3O+(aq) + A-(aq) The extent of ionization for strong acids. H+ and H2O  H3O+ (hydronium ion)

  12. Weak acid: HA(aq) + H2O(l) H3O+(aq) + A-(aq) The extent of ionization for weak acids and bases. Weak acid are only partly ionized. Many intact molecules remain.

  13. Acids Strong ionizes completely in water hydrochloric acid, HCl stomach acid hydrobromic acid, HBr nitric acid, HNO3 used to make explosives sulfuric acid, H2SO4 battery acid perchloric acid, HClO4 used to digest plant matter Weak ionizes partially in water hydrofluoric acid, HF used to etch glass phosphoric acid, H3PO4 naval jelly, to remove rust acetic acid, CH3COOH (or HC2H3O2) vinegar carbonic acid, H2CO3 in carbonated drinks

  14. Bases (or alkalis) Strong Moderate Dissociates completely Dissociates completely but is not very soluble sodium hydroxide, NaOH magnesium hydroxide, Mg(OH)2 potassium hydroxide, KOH aluminum hydroxide, Al(OH)3 calcium hydroxide, Ca(OH)2 Weak strontium hydroxide, Sr(OH)2 barium hydroxide, Ba(OH)2 Dissociates partially ammonia, NH3 (NH4OH) carbonates, CO32- bicarbonates, HCO31-

  15. Acid and Base Definitions Arrhenius Acid = compound that forms hydrogen (H+) ions in water Base = compound that forms hydroxide (OH-) ions in water

  16. Acid and Base Definitions Bronsted-Lowry Acid = proton donor (H+ is a proton) Base = proton acceptor (picks up an H+)

  17. Acid Anhydrides Non-metal oxides react with water to form acidic solutions CO2 (g) + H2O (l)  H2CO3 (aq) N2O5 (s) + H2O (l)  2 HNO3 (aq) SO3 (g) + H2O (l)  H2SO4 (aq) Dissolved non-metal oxides cause acid rain. This is why tap/distilled water has a pH lower than 7!

  18. Basic Anhydrides Metal oxides react with water to form alkaline solutions Na2O (s) + H2O (l)  2 NaOH(aq) CaO (s) + H2O (l)  Ca(OH)2 (aq) Al2O3 (s) + 3 H2O (l)  2 Al(OH)3 (aq) Lime (CaO) is used on lawns and is converted to Ca(OH)2 when it rains. CaO is less hazardous to handle.

  19. What Happened When Acids and Bases are Mixed? What did you uncover when these materials are mixed? How could you tell? Acid + Base  Water and a “Salt” A “salt” is an electrolyte that is generally not an acid or base.

  20. An aqueous strong acid-strong base reaction on the atomic scale. MX is a “salt” – an electrolyte that is not an acid or base

  21. Acid-Base Reactions HCl + NaOH  H2SO4 + Mg(OH)2  HBr + Al(OH)3  HCl + NaHCO3 

  22. Acid-Base Reactions HCl + NaOH  NaCl + H2O H2SO4 + Mg(OH)2  MgSO4 + 2H2O 3 HBr + Al(OH)3  AlBr3 + 3 H2O HCl + NaHCO3  NaCl + H2O + CO2

  23. Acid-Base Reactions Acid-base reactions are also called neutralization reactions. Why? If you mix the right proportions of acid and base the solution should end up neutral.

  24. Antacids • What happens when you produce too much stomach acid? Acid may reflux into the unprotected esophagus and cause a burning sensation • Antacids (anti-acids) neutralize excess stomach acid.

  25. Antacids Tums, Rolaids: 2HCl + CaCO3 CO2 + CaCl2 + H20 Milk of Magnesia, Mylanta, Maaolox 2HCl + Mg(OH)2  MgCl2 + 2H2O (and Al(OH)3 in Maaolox) Alka-Seltzer: HCl + NaHCO3  NaCl + CO2 + H2O (and KHCO3)

  26. An acid-base titration. Start of titration Excess of acid Point of neutralization Slight excess of base

  27. What about pH? Why is the pH of a neutral solution = 7? First we need to talk about what happens in pure water.

  28. H2O(l) + H2O(l) H3O+(aq) + OH-(aq) Kw = [H3O+][OH-] = 1.0 x 10-14 at 250C A change in [H3O+] causes an inverse change in [OH-]. In an acidic solution, [H3O+] > [OH-] In a basic solution, [H3O+] < [OH-] In a neutral solution, [H3O+] = [OH-]

  29. The pH values of some familiar aqueous solutions. pH = -log [H3O+] pOH = -log [OH-] pH + pOH = 14

  30. Figure 18.6 The relations among [H3O+], pH, [OH-], and pOH.

  31. Out of Range pH’s • Can a pH be lower than 0 or higher than 14? Yes! • When you have solutions that are more concentrated than 1 M the value of pH will be out of the normal range • 6M HCl pH = -log[2.0] = -0.8 • 1.5 M NaOH pH = 14.2

  32. Buffers Solutions that resist change in pH Can maintain any pH value between 0 and 14 (not just neutral pH 7) Composed of a weak acid and a salt made from the weak acid or weak base and salt made from the weak base Examples: HC2H3O2 and NaC2H3O2 NH4OH and NH4Cl

  33. Buffers Reaction with acid: HC2H3O2 + C2H3O2- + H+ HC2H3O2 + HC2H3O2 Reaction with base: HC2H3O2 + C2H3O2- + OH- C2H3O2- + C2H3O2- + HOH A buffer regenerates it’s own components. The pH it maintains depends on the ratio of salt to acid (or base) and the nature of the acid (or base).

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