1. Chapter 14

2. 2 Arrhenius Concept: Acids produce H + in solution, bases produce OH  ion. In aqueous solutions. Brønsted-Lowry: Acids are H + donors, bases are proton acceptors. This is the more general definition and can include solids and gases. There are two ways to view Acids & Bases: The Two Views

3. Arrhenius Arrhenius - In aqueous solution… Arrhenius - In aqueous solution… HCl + H 2 O  H 3 O + + Cl – AcidsAcids form hydronium ions (H 3 O + ) H HHHH H Cl OO – + acid

4. Arrhenius Arrhenius - In aqueous solution… Arrhenius - In aqueous solution… BasesBases form hydroxide ions (OH - ) NH 3 + H 2 O  NH 4 + + OH - H H H H H H N NO O – + H H H H base

5. Brønsted-Lowry Brønsted-Lowry Brønsted-Lowry HCl + H 2 O  Cl – + H 3 O + AcidsAcids are proton (H + ) donors. BasesBases are proton (H + ) acceptors. conjugate acid conjugate base baseacid

6. Lewis Acid and Bases Lewis Lewis AcidsAcids are electron pair acceptors. BasesBases are electron pair donors. Lewis base Lewis acid

7. Properties electrolytes  electrolytes  turn litmus red  sour taste  react with metals to form H 2 gas  slippery feel  turn litmus blue  bitter taste ChemASAP  vinegar, milk, soda, apples, citrus fruits  ammonia, lye, antacid, baking soda

8. Strength Strong Acid/Base Strong Acid/Base 100% ionized in water strong electrolyte - + HCl HNO 3 H 2 SO 4 HBr HI HClO 4 NaOH KOH Ca(OH) 2 Ba(OH) 2

9. Strength  Weak Acid/Base does not ionize completely weak electrolyte HF CH 3 COOH H 3 PO 4 H 2 CO 3 HCN NH 3

10. Strength K a =[H 3 O + ][A - ] [HA] Acid Dissociation Constant (K a ): expression for the dissociation of an acid (larger value = stronger acid) K b = [HB + ][OH - ] [B] Base Dissociation Constant (K b ): expression for the dissociation of an base (larger value = stronger base)

11. Strength

12. Polyprotic Acids Some compounds have more than 1 ionizable hydrogen. HNO 3 nitric acid - monoprotic H 2 SO 4 sulfuric acid - diprotic - 2 H + H 3 PO 4 phosphoric acid - triprotic - 3 H + Having more than one ionizable hydrogen does not make the acid stronger!

13. Hydrogen Ions from Water Water naturally ionizes, or breaks into ions: H 2 O  H + + OH - This is the “self ionization” of water, and occurs to a very small extent: [H 1+ ] = [OH 1- ] = 1 x 10 -7 M Brackets [ ] mean “concentration of”

14. Hydrogen Ions from Water H 2 O  H + + OH - Since the [H + ] and [OH - ] equal, a neutral solution results from water [H + ] x [OH - ] = 1 x 10 -14 M 2 Molarity is squared This is true of ALL solutions. As [H + ] increases, [OH - ] decreases

15. pH & pOH pH = -log[H + ]pOH = -log[OH - ] pH + pOH = 14.00

16. Acid – Base Neutralization Rules for Acid-Base Reactions 1.List the species present in the combined solution and decide what reaction will take place (remember solubility!) 2.Write the balanced equation 3.Calculate the moles of reactants (for solutions use the volumes to convert to moles) 4.Determine a limiting reactant if necessary 5.Calculate the moles required of reactant or product (in this case H + or OH - ) 6.Convert to grams or volume as required

17. Titration Titration: delivery of a measured volume of solution of a known concentration (titrant) to a substance being analyzed (analyte). Equivalence Point: where enough titrant has been added to react completely with the analyte (marked by an indicator) Endpoint: point where the indicator changes color