1. Acids, Bases, and Salts

2. AcidBase (Alkali) Litmus color Phenolphthalein color pH range (from universal indicator paper) Taste Formula component Other? Characteristics of Acids and Bases

3. SubstanceAcidBase (Alkali) Neutral Vinegar Epsom salts Lye Aspirin Baking soda Windex (window cleaner) Apple juice Bubble bath Washing soda Vitamin C Household Acids and Bases

4. CSSI 2010 PGCC Barbara A. Gage 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 8-9laundry detergents 9.2 borax 10.5 milk of magnesia 11.0 ammonia 12.0 lime water 13.0 lye 14.0 pH of Some Household Materials

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 -, HCO 3 - or CO 3 2-

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 + H 2 O  H 3 O + + 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 ) + H 2 O( l ) H 3 O + ( aq ) + A - ( aq ) The extent of ionization for strong acids. H + and H 2 O  H 3 O + (hydronium ion)

12. The extent of ionization for weak acids and bases. Weak acid: HA( aq ) + H 2 O( l ) H 3 O + ( aq ) + A - ( aq ) Weak acid are only partly ionized. Many intact molecules remain.

13. Acids Strong hydrochloric acid, HCl stomach acid hydrobromic acid, HBr nitric acid, HNO 3 used to make explosives sulfuric acid, H 2 SO 4 battery acid perchloric acid, HClO 4 used to digest plant matter Weak hydrofluoric acid, HF used to etch glass phosphoric acid, H 3 PO 4 naval jelly, to remove rust acetic acid, CH 3 COOH (or HC 2 H 3 O 2 ) vinegar ionizes completely in water ionizes partially in water carbonic acid, H 2 CO 3 in carbonated drinks

14. Bases (or alkalis) Strong Weak sodium hydroxide, NaOH calcium hydroxide, Ca(OH) 2 potassium hydroxide, KOH strontium hydroxide, Sr(OH) 2 barium hydroxide, Ba(OH) 2 ammonia, NH 3 (NH 4 OH) Moderate Dissociates completelyDissociates completely but is not very soluble aluminum hydroxide, Al(OH) 3 magnesium hydroxide, Mg(OH) 2 Dissociates partially carbonates, CO 3 2- bicarbonates, HCO 3 1-

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 CO 2 (g) + H 2 O (l)  H 2 CO 3 (aq) N 2 O 5 (s) + H 2 O (l)  2 HNO 3 (aq) SO 3 (g) + H 2 O (l)  H 2 SO 4 (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 Na 2 O (s) + H 2 O (l)  2 NaOH (aq) CaO (s) + H 2 O (l)  Ca(OH) 2 (aq) Al 2 O 3 (s) + 3 H 2 O (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  H 2 SO 4 + Mg(OH) 2  HBr + Al(OH) 3  HCl + NaHCO 3 

22. Acid-Base Reactions HCl + NaOH  NaCl + H 2 O H 2 SO 4 + Mg(OH) 2  MgSO 4 + 2H 2 O 3 HBr + Al(OH) 3  AlBr 3 + 3 H 2 O HCl + NaHCO 3  NaCl + H 2 O + CO 2

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 + CaCO 3  CO 2 + CaCl 2 + H 2 0 Milk of Magnesia, Mylanta, Maaolox 2HCl + Mg(OH) 2  MgCl 2 + 2H 2 O (and Al(OH) 3 in Maaolox) Alka-Seltzer: HCl + NaHCO 3  NaCl + CO 2 + H 2 O (and KHCO 3 )

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. [H 3 O + ][OH - ]K w = A change in [H 3 O + ] causes an inverse change in [OH - ]. = 1.0 x 10 -14 at 25 0 C H 2 O( l ) + H 2 O( l ) H 3 O + ( aq ) + OH - ( aq ) In an acidic solution, [H 3 O + ] > [OH - ] In a basic solution, [H 3 O + ] < [OH - ] In a neutral solution, [H 3 O + ] = [OH - ]

29. The pH values of some familiar aqueous solutions. pH = -log [H 3 O + ] pOH = -log [OH - ] pH + pOH = 14

30. Figure 18.6 The relations among [H 3 O + ], 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: HC 2 H 3 O 2 and NaC 2 H 3 O 2 NH 4 OH and NH 4 Cl

33. Buffers Reaction with acid: HC 2 H 3 O 2 + C 2 H 3 O 2 - + H +  HC 2 H 3 O 2 + HC 2 H 3 O 2 Reaction with base: HC 2 H 3 O 2 + C 2 H 3 O 2 - + OH -  C 2 H 3 O 2 - + C 2 H 3 O 2 - + 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).