1. Acids and Bases: Properties of acids: Taste sour React with metals to form hydrogen gas React with carbonates to form CO 2 gas Form electrolyte solutions Properties of bases: Feel slippery Taste bitter Will react with some metals Form electrolyte solutions

2. pH scale 0714acidicbasic If pH = 7.0, the solution is neutral As the hydronium concentration ([H 3 O +1 ]) increases, the solution becomes more acidic and the pH drops.

3. Acids Examples: HC 2 H 3 O 2 vinegar H 2 C 6 H 6 O 6 ascorbic acid (Vitamin C) H 2 CO 3 carbonic acid (in sodas) HClhydrochloric acid (stomach acid) H 3 PO 4 phosphoric acid (in colas) H 2 SO 4 sulfuric acid (battery acid) The hydrogens that appear first in the formula are called ACIDIC PROTONS, or just PROTONS for short.

4. When acids dissolve in water they lose their acidic proton to a water molecule HCl + H 2 OH 3 O +1 + Cl -1 The H 3 O +1 ion that’s formed is called the HYDRONIUM ION Strong acids: an acid that dissociates completely into hydronium ions. None of the original acid is left over. Weak acids: an acid that only partially dissociates into hydronium ions. Some of the original acid is left over.

5. Bases Examples: Al(OH) 3 Aluminum hydroxide (antacids) Ca(OH) 2 Calcium hydroxide (soil additive) Mg(OH) 2 Magnesium hydroxide (antacid) NaOHSodium hydroxide (lye) NH 3 Ammonia (cleaning solution) Bases are usually ionic compounds that contain the hydroxide (OH-) group. Ammonia is an exception.

6. When bases dissolve in water, they increase the concentration of hydroxide ions (OH -1 ) in the solution. They can do this in one of two ways: NaOH (s)Na +1 + OH -1 NH 3 + H 2 ONH 4 +1 + OH -1 While ammonia does not have a hydroxide group, it increases the [OH -1 ] by reacting with water.

7. Strong bases: a base that ionizes completely when dissolved in water. None of the original compound remains. Example: NaOH Weak bases: A base that only partially ionizes when dissolved in water. Some of the original compound is left over. Example: NH 3

8. Acid-Base reactions The H +1 from the acid reacts with the OH -1 from the base to form water. H + + OH - H2OH2O This is called areaction

9. Examples: What would the products be if HCl and NaOH reacted? HCl + NaOHH 2 O +NaCl If HCl and KOH reacted? HCl + KOHH 2 O +KCl If HCl and Mg(OH) 2 reacted? HCl + Mg(OH) 2 H 2 O +MgCl 2 2 2

10. In general... *A salt is any ionic compound that is NOT an acid or a base.

11. Titrations Titrations allow the concentration of an acid or base to be determined using an acid-base reaction and an indicator. 1. Measure out a volume of the acid or base that has the unknown concentration*. 2. Add small volumes of the other reactant until the indicator changes color. 3. Use the ‘magic equation’ to calculate the unknown concentration. *The concentration unit we will be using is molarity (M).

12. Indicators: organic dyes whose color depends upon the pH of the solution. The point at which the indicator changes color is the endpoint of the titration.

13. C 1 V 1 = C 2 V 2 C 1 = unknown concentration V 1 = volume used for unknown C 2 = known concentration V 2 = total volume added of known concentration

14. Examples: 25 mL of HCl are titrated with 12.5 mL of 1.0 M NaOH. What is the concentration of the HCl? C 1 (25 mL) = (1.0 M)(12.5 mL) C 1 V 1 = C 2 V 2 C 1 = 12.5/25 C 1 = 0.5 M The concentration of the HCl is 0.5 M

15. Buffers: solutions that resist changes in pH There are two main buffering equilibria… H 2 CO 3 + H 2 O ⇋ HCO 3 -1 + H 3 O + H 2 PO 4 -1 + H 2 O ⇋ HPO 4 -2 + H 3 O + An important buffer example: The pH of blood must be maintained at 7.4 ± 0.2 or death may occur. Proteins also participate in the buffering effect