1. Acids and Bases Part 2

2. Classifying Acids and Bases Arrhenius Acid ◦ Increases hydrogen ions (H + ) in water ◦ Creates H 3 O + (hydronium) Base ◦ Increases OH - in water

3. Classifying Acids and Bases Brønsted - Lowry Acid ◦ Proton donor (H + ) ◦ Must have hydrogen in formula Base ◦ Proton acceptor (H + ) Water can be an acid or a base

4. Classifying Acids and Bases Lewis Acid ◦ Electron pair acceptor ◦ Usually positive ions Base ◦ Electron pair donor ◦ Usually negative ions

5. Pairs Acid + Base Conjugate acid + Conjugate base This is an equilibrium. The acid becomes the conjugate base after it has donated the H + The base becomes the conjugate acid once it accepts the H +

6. Conjugate Acids and Bases Identify the acid, base, conjugate acid, and conjugate base in the below reaction: HF + H 2 O → F - + H 3 O + Acid Base conjugate conjugate base acid

7. Water Water conducts electricity ◦ Electrolyte Thus, water self-ionizes H 2 O → H + + OH - Water is also Amphoteric ◦ Can act as either an acid or a base Therefore: 2H 2 O → H 3 O + + OH -

8. Types of Acids Monoprotic ◦ Only 1 acidic hydrogen Polyprotic Acids ◦ More than 1 acidic hydrogen ◦ Diprotic – H 2 SO 4 ◦ Triprotic – H 3 PO 4 Oxyacids ◦ Proton is attached to the oxygen of an ion. Organic acids ◦ Contain the Carboxyl group –COOH ◦ H attached to O – ex: CH 3 COOH – acetic acid ◦ Generally very weak

9. Strong Acids They completely dissociate in water ◦ HCl ◦ HNO 3 ◦ HI ◦ H 2 SO 4 ◦ HClO ◦ HBr

10. pH of Strong Acid Calculate the pH and [OH]: ◦ For a 1 x 10 -3 M solution of HClO 4

11. Bases The OH - is a strong base. Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. Others are weak.

12. Bases without OH - Bases are proton acceptors. NH 3 + H 2 O NH 4 + + OH - It is the lone pair on nitrogen that accepts the proton.

13. Salts as acids and bases Salts are ionic compounds. Salts of the cation of strong bases and the anion of strong acids are neutral. Salts of a strong acid and weak base are acidic salts Salts of strong base and weak acids are basic salts

14. Lewis Acids and Bases Most general definition. Acids are electron pair acceptors. Bases are electron pair donors. BF F F :N:N H H H

15. Ion concentrations The concentrations of H 3 O + and OH - are based on water Kw – ion-product constant of water At 25°C, Kw = 1.0x10 -14 Kw = (K a )(K b ) [H 3 O + ][OH - ] [H 3 O + ] – concentration of hydronium ion [OH - ] – concentration of hydroxide ion Neutral: [H 3 O + ] = [OH - ]= 1.0 x10 -7 Acidic:[H 3 O + ] > [OH - ] Basic: [H 3 O + ] < [OH - ]

16. Acid dissociation constant Acid dissociation constant K a Base dissociation constant K b The larger the number the stronger the substance The smaller the number the weaker the substance

17. Calculating pH pH is a measure of hydronium [H 3 O + ] concentration in solution Lower pH = more acidic pH = -log [H 3 O + ]

18. pH Scale

19. pH Calculations What is the pH of a neutral solution ([H 3 O + ] = 1.0x10 -7 M)? ◦ pH = 7 What is the pH of a solution if: [H 3 O + ] = 1.0x10 -4 M? ◦ pH = 4 [H 3 O + ] = 0.0015M? ◦ pH = 2.8

20. Calculating pOH pOH is a measure of hydroxide [OH - ] concentration in solution Higher pOH = more acidic pOH = - log [OH - ] pH + pOH = 14

21. pOH Calculations What is the pOH of a neutral solution ([OH - ] = 1.0x10 -7 M)? What is the pOH of a solution if: [OH - ] = 8.22x10 -6 M? ◦ pOH = 5.09 [OH - ] = 0.0541M? ◦ pOH = 1.27

22. BasicAcidicNeutral 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 [H + ] 013579 111314 pH Basic10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 [OH - ] 013579111314 pOH