1. Chapter 18 Acids and Bases

2. 18.1 Acids Arrhenius Acid – a compound containing hydrogen that ionizes to produce hydrogen ions (H + ) in water Names: Hydrochloric Acid (HCl), Sulfuric Acid (H 2 SO 4 ), Nitric Acid (HNO 3 )

3. Properties of Acids Acid have a sour or tart taste.  Example: vinegar (acetic acid), lemons (citric acid), tomatoes (ascorbic acid) Acids carry charge in aqueous solutions. React easily with metals to produce hydrogen gas

4. Properties of Acids acids that contain one ionizable hydrogen are called monoprotic (HNO3); acids that contain two hydrogens are called diprotic (H2SO4); acids that contain three hydrogens are called triprotic (H3PO4)

5. Properties of Acids only the hydrogens in very polar bonds are ionizable (CH 4 is NOT an acid)

6. Bases Arrhenius Base – a compound that contains a hydroxide group (OH - ) and dissociates to produce a hydroxide ion in water

7. Properties of Bases Bases also can carry charges in aqueous solutions Bitter taste; slippery feel

8. Neutralization Bases combine with acids to neutralize an aqueous solution.  Example: “Tums” (Milk of Magnesium) is a base used to treat excess stomach acid. Acid + Base  Water & “Salts” HCl + NaOH → H2O + NaCl

9. Bronsted-Lowry Definitions more broad (includes bases such as ammonia and sodium carbonate that do not have hydroxide ions) An acid is a proton (H + ) donor and a base is a proton acceptor.

10. Bronsted-Lowry Acids and bases always come in pairs. HCl(g) + H2O(l)  H3O+ + Cl- water acts as a base to make hydronium ion (the conjugate acid) HCl acts as an acid to form chloride ion (the conjugate base) a substance that can act as either an acid or a base (like water) is called amphoteric

11. 18.2 Strength of Acids and Bases Strong and Weak Acids and Bases strong acids and bases are completely ionized in aqueous solution (HCl and NaOH) weak acids and bases ionize only slightly in aqueous solution (acetic acid and ammonia)

12. Acid and Base Ionization Constants if the value of the equilibrium constant is small, then the degree of ionization is small (weak) if the value of the equilibrium constant is large, then the degree of ionization is high (strong)

13. Acid and Base Ionization Constants For the reaction HCN + H 2 O  H 3 O + + CN - K a = For the reaction NH 3 + H 2 O  NH 4 + + OH - K b =

14. 18.3 pH the pH of a solution is the negative log of the hydrogen-ion concentration pH = -log [H+] ranges from 0 (acidic) to 14 (basic); neutral solutions have a pH of 7

15. pH and pOH the pOH equals the negative log of the hydroxide-ion concentration  pOH = -log [OH-] pH + pOH = 14

16. Ion Product Constant for H 2 O The equilibrium constant expression for water is called the ion product constant and gives a relationship between hydroxide and hydronium concentrations. K w = [H + ][OH - ] = 1.0 x 10 -14 In pure water, both ion concentrations are 1.0 x 10 -7 M.

17. 18.4 Acid/Base Indicators acids or bases that undergo dissociation at a known pH range can be used as indicators usually best accuracy at a given temperature color can be distorted unless solution is colorless; often indicator strips used to eliminate these problems

18. pH pH meters can be used to make precise measurements; shows a continuous recording of pH changes; typically more accurate (hospitals use them to determine the pH of blood/body fluids; sewage is also monitored using pH meters)

19. Titrations Acid-base titrations can be used to determine the concentration of a solution by reacting a known volume of the solution with a solution of known concentration. The point at which the indicator changes color is called the end point of the titration.

20. Buffers Control of pH is important in certain instances (such as your body). Buffers are solutions that resist changes in pH when limited amounts of acid or base are added.  usually a conjugate acid/base pair