Acids, Bases, and Salts

1. Acids, Bases, and Salts - Acids taste sour, will change the color of an indicators (chemical dyes), and can be strong or weak electrolytes (aqueous solutions that conduct electricity). - Bases - taste bitter, feel slippery, will change the color of an indicator, and can be strong or weak electrolytes in an aqueous solution.

2. Arrhenius Acids and Bases • Arrhenius said that acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solutions. • He defined bases as compounds that ionize to yield hydroxide ions (OH-) in aqueous solutions.

3. Arrhenius says that acids can be: • 1) Monoprotic – acids that contain only one ionizable hydrogen (HNO3, HCl). • 2) Diprotic – acids that contain two ionizable hydrogens (H2SO4, H2S). • 3)Triprotic – acids that contain 3 ionizable hydrogens (H3PO4). • Not all compounds that contain H are acids.

4. Not all hydrogens in an acid may be released as hydrogen ions. Only the hydrogens in very polar bonds are ionizable. • HCl ----- H+ + Cl- • HCl + H2O --- H3O+ + Cl-

5. Bronsted-Lowry Acids and Bases • The Bronsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor. • NH3 (aq) + H2O (l) -- NH4+ (aq) + OH- (aq) • In this reaction ammonia is a base because it accepts a H+ and water is an acid because it donates a H+.

6. Conjugate Acids and Bases • A conjugate acid is the particle formed when a base gains a hydrogen ion. (NH4) • A conjugate base is the particle formed when an acid has donated a hydrogen ion. (OH-) • A conjugated acid-base pair consists of two substances that are related by the loss or gain of a single hydrogen ion. • A substance that can both accept a hydrogen ion or donate one is called amphoteric.

7. Lewis Acids and Bases • Lewis proposed that an acid accepts a pair of electrons during a reaction, while a base donates a pair of electrons. • H+ + -O—H --- H—O—H • OH- = base • H+ = acid

8. pH • pH - a scale that determines if a substance is acidic, basic, or neutral. • The scale is based from 0 to 14. • 0 = strongly acidic • 7 = neutral • 14 = strongly basis

9. Formulas for calculating pH • pH = -log[H3O+] • pOH = -log[OH-] • pH + pOH = 14 • Kw = [H+] [OH-] = 1.00 X 10-14

10. Strong and Weak Acids and Bases • Strong acids are completely ionized in water. • HCl (g) + H2O (l) - H3O+ (aq) + Cl- (aq) (100% ionized) • Weak acids only slightly ionize in aqueous solution.

11. Acid Dissociation Constant • An acid dissociation constant (Ka) is the ratio of the concentration of the dissociated (or ionized) form of the acid to the concentration of the undissociated (nonionized) form. • Weak acids have small Ka values. • Strong acids have large Ka values.

12. Acid-Base Reactions • Neutralization rxn. – a rxn between an acid and a base which produces a salt and water. • H2SO4 + 2KOH  K2SO4 + H2O • Salts are compounds that consist of an anion from an acid and a cation from a base.

13. Titration • Titration – the process used to determine the concentration of a solution (often an acid or a base) in which a solution of known concentration (the standard) is added to a measured amount of the solution of unknown concentration until an indicator signals the end point.

14. Steps in Titrating • 1) A measured volume of an acid solution of unknown concentration is added to a flask. • 2) Several drops of indicator are added to the solution and the flask is swirled gently. • 3) A base with known concentration is added slowly to the solution of acid until the indicator turns pink ever so slightly.

15. The point at which the indicator changes color is known as the end point. • The unknown concentration can then be determined using the following equation: MaVa = MbVb The end point is also know as the point of neutralization or the equivalence point (when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.