1. Introduction to Acids and Bases Chapter 19

2. What is and Acid? Arrhenius Acid Defined as any chemical that increases the concentration of hydrogen ions (H + )in solution. (usually by a dissociation reaction Examples Hydrochloric acid HCl  H + + Cl - Sulfuric acid H 2 SO 4  2H + + SO 4 2- Phosphoric acid H 3 PO 4  3H + + PO 4 3-

3. Not all Hydrogens are acidic The hydrogen must be part of a polar bond in order to dissociate. For example: HF is acidic, but CH 4 is not. In CH 4, the hydrogen is part of a non-polar covalent bond and does not dissociate in solution! HF is a polar bond and HF  H + + F -

4. Bronsted – Lowry Acid Defined as a molecule or ion that is a hydrogen ion donor. Also known as a proton donor because H + is a proton. The acid will donate its H + ion to a base in an acid base reaction. H + + OH -  H 2 O Acid + Base

5. What is a Base? Arrhenius Base Defined as any chemical that increases the hydroxide ions (OH-) concentration in solution. Examples- NaOH  Na + + OH - KOH  K + + OH - Ca(OH) 2  Ca 2+ + 2OH -

6. Bronsted-Lowry Base Defined as a hydrogen ion acceptor. In an acid-base reaction the base “accepts” the hydrogen ion from the acid. NH 3 + H +  NH 4 + NH 3 accepts the H+ from the acid.

7. Physical Properties of Acids Taste Sour Feel Sticky React with metals to produce hydrogen gas Conduct electricity – because they produce ions in solution Are corrosive (they can burn skin and other materials) Are found in many naturally colored solutions Turn Litmus Indicator Red

8. Physical Properties of Bases Taste Bitter Feel slippery or soapy Also conduct electricity Rarely found in colored solutions Most do not have a smell (except ammonia) Turn Litmus indicator Blue Turn Phenolphthalein Indicator Pink

9. Physical Properties of Neutral Compounds Do not taste bitter or sour, but rather have a sweet or oily taste, or no taste at all. Many feel like water, or feel greasy/ oily May or may not conduct electricity They have varied reactivities. Many have a strong chemical smell or no smell at all.

10. Using pH to identify Acids and Bases The pH Scale

11. What is pH? pH = -log [H + ] pH= -log[H 3 O + ] pH stands for the “power of the hydrogen ion.” It is based on a logarithmic scale which has a base power of 10. –A pH of 1 differs from a pH of 2 by a factor of 10.

12. pH Scale - Continued Created to express acidity as a more simple number than molarity. pH = - log [H+] If... [H+] = 1 x 10 -7 then pH = 7 (neutral) [H+] = 1 x 10 -3 then pH = 3 (acid) [H+] = 1 x 10 -10 then pH = 10 (base)

13. Why is the pH scale 0 – 14? The Self-Ionization of Water H 2 O (l) ↔ H + (aq) + OH - (aq) The equilibrium expression for this reaction is Kw = [H + ][OH - ] = 1 x 10 -14 This low equilibrium constant means very few water molecules ionize.

14. Water is considered neutral because in water [H+] = [OH-] = 1 x 10 -7 M pH = -log (1 x 10 -7 ) = 7 *All aqueous solutions have H+ and OH- ions Acids HCl  H+ + Cl- Acidic solutions increases the H+ concentration and [H+] > [OH-] Bases NaOH  Na+ + OH- Basic solutions increases the OH- concentration and [H+] < [OH-]

15. pH scale ranges from 0-14 pH<7 indicates an acidic solution pH=7 indicates a neutral solution pH>7 indicates a basic solution Strong Acids have pH of 0 or 1 Strong Bases have a pH of 13 or 14

16. Strength of Acids and Bases “ Strength” refers to how much an acid or base ionizes in a solution. STRONGWEAK Ionize completely (~100%) Example: HCl  H + + Cl - NaOH  Na + + OH - Ionize partially (usually <10%) An equilibrium reaction! Example: HF ↔ H + + F - NH 3 + H 2 O ↔ NH 4 + + OH -

17. Strengths of Acids and Bases