1. Chapter 20 Acids and Bases

2. Describing Acids and Bases 1.Properties of Acids and Bases Acids Bases Contains H + Contains OH - Turns blue litmus red Turns red litmus blue Taste sour Taste bitter Can be electrolytes Can be electrolytes Reacts with bases to Reacts with acids to form water form water

3. 2. Quick review of naming acids  Hydrogen ions and acidity 1. Hydrogen Ions from water a. When water molecules lose a hydrogen ion it becomes OH - Anion endingExampleAcid nameExample -ideCl - chlorideHydro- (stem) –ic acid -IteSO 3 -2 (stem)-ous acid -ateSO 4 -2 (stem)-ic acid

4. b. When water molecules gain a hydrogen ion it becomes H 3 O + (called the hydronium ion) 2. Self-ionization a. When two water molecules produce ions b. H 2 O (l) H + (aq) + OH - (aq)

5. c. [H + ] = 1.0 x 10 -7 M d. [OH - ] = 1.0 x 10 -7 M e. When [H + ] and [OH - ] are equal it is a neutral solution f. When they are independent (not equal) [H + ] increases, [OH - ] decreases [H + ] decreases, [OH - ] increases

6. 3. Ion-product constant a. k w : product of concentration of H + and OH - in water b. K w = [H + ] [OH - ] = 1.0 x 10 -14 M 2 c. Acidic solution: one where [H + ] is greater than [OH - ] [H + ] > 1.0 x 10 -7 M

7. d. Basic solution: one where [H + ] is less than [OH - ] [H + ] < 1.0 x 10 -7 M e. Basic solution also known as Alkaline solution 4.The pH concept a. Better expressed using the pH scale

8. b. pH + pOH = 14 pH = -log[H + ] pOH = -log[OH - ] c. In a neutral solution [H + ] = 1.0 x 10 -7 M pH = -log (1 x 10 -7 ) = -(log 1 + log 10 -7 ) = -(0.0 + (-7.0)) = 7.0

9. d.

10. e. 5.Example problems: a. What is the pH of a solution with a hydrogen-ion concentration of 1.0 x 10 -10 M?

11. b. The pH of an unknown solution is 6.00. What is its hydrogen-ion concentration? c. What is the pOH of a solution if [OH - ] = 4.0 x 10 -11 M? d. What is [H + ] of a solution if the pH = 3.70?

12. 6. Measuring pH a. Acid-base Indicator 1. Indicator (In) is an acid or base that undergoes dissociation in a known pH range 2. Reaction form: HIn (aq) H + (aq) + I n - (aq) Acid form Base form

13. 3. Types: pH color Thymol blue 1.2-3.0 red  yellow 8.0-9.5 yellow  blue Bromphenol blue 3.0-4.6 yellow  blue Bromcresol green 3.7-5.3 yellow  blue methyl red 4.2-6.2 red  yellow Alizarin 4.5-6.0 yellow  red Bromthymol blue 6.0-7.5 yellow  blue Phenol red 6.9-8.2 yellow  orange Phenolphthalein 8.0-10.0 colorless  pink

14. alizarin yellow R 8.0 – 12.2 yellow  red 4. Useful at room temperature (25 °C) b. pH meter 1. Useful to make rapid, accurate pH measurements 2. more practical than liquid indicators

15.  Acid-Base Theories 1.Arrhenius Acids and Bases a. Acids are hydrogen containing compounds that ionize to yield H + in aq solutions b. Bases are compounds that ionize to yield OH - in aq solutions Acids c. Monoprotic acids have one hydrogen HCl

16. d. Diprotic acids : have two hydrogens H 2 SO 4 e. Triprotic acids: have three hydrogens H 3 PO 4 f. Only very polar bonds will dissociate H δ+ --Cl δ-  H + (aq) + Cl - (aq)

17. g. C-H bonds weakly polar will not dissociate ex. Ethanoic acid (CH 3 COOH):

18. Bases h. NaOH (s)  Na + (aq) + OH - (aq) i. Common bases: KOH, NaOH, Ca(OH) 2, Mg(OH) 2

19. 2.Bronsted-Lowery Acids and Bases a. Acid is a hydrogen-ion donor b. Base is a hydrogen-ion acceptor c. Conjugate acid – particle formed when a base gains a hydrogen ion d. Conjugate base- particle that remains when an acid has donated a hydrogen ion

20. e. Conjugate acid-base pair: two substances related by the loss or gain of a single hydrogen bond f. Examples: 1. NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) acceptor donor (base) (acid) (CA) (CB) 2. HCl (g) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) (acid) (base) (CA) (CB)

21. g. Amphoteric: a substance that can act like both an acid and base

22. Sample problems 1. Classify the following as Brønsted acids, bases or both. a) H 2 O b) OH - c) NH 3 d) NH 4 + 2. What is the conjugate base of the following acids? a) HClO 4 b) NH 4 + c) H 2 O d) HCO 3 - 3. What is the conjugate acid of the following bases? a) CN - b) SO 4 2- c) H 2 O d) HCO 3 -

23. 3.Lewis Acids and Bases a. Acid: a substance that accepts a pair of electrons to form a covalent bond b. Base: a substance that donate a pair of electrons to form a covalent bond c. Examples: 1. H + + acid base 2.

24.  Strengths of Acids and Bases 1. Strong acids and bases a. Strong acids: completely ionize (dissociate) HCl, HNO 3, H 2 SO 4, HBr, H I, HClO 4 b. Dissociation constant (K a ): the ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form

25. ***See page 600 Table 20.7 3. Equilibrium-constant expression K = [products] [reactants] ** Remember to raise the concentrations to the coefficient number.

26. 4. K a = [H + ][A - ] (Gives the ratio of ions [HA] vs molecules)  Weak acid has K a <1  Leads to small [H + ] and pH of 2-7 5. K b = [BH + ][OH - ] [B]  Weak bases has K b < 1  Leads to small [OH - ] and pH of 12-7 **Do not use water in the [ ]

27. 6. Examples: a. Calculate the [OH - ] of a 0.500 M solution of aqueous ammonia. The K b is 1.74 x 10 -5. NH 3 + H 2 O NH 4 + + OH - K b = [NH 4 + ][OH - ] [NH 3 ]

29. Step 2: Write the K a expression HC 2 H 3 O + H 2 0 H + + C 2 H 3 O - (HOAc) (OAc - ) K a = [H + ][OAc - ] [HoAc] 1.8 x 10 -5 = (x)(x) = x 2 (1.00 –x) (1.00 – x) This is a quadratic. Solve using the quadratic formula. OR you can make an approximation if x is very small. (Rule of thumb: 10 -5 or smaller is OK)

30. 1.8 x 10 -5 = x 2 1.00 x = [H + ] = [OAc -] = 4.2 x 10 -3 M pH = -log[4.2 x 10 -3 ] = 2.37

31. c. You have 0.010 M NH 3. Calculate the pH if the K b = 1.8 x 10 -5. NH 3 + H 2 O  NH 4 + + OH - [NH 3 ] [NH 4 + ] [OH - ] Initial 0.010 0 0 Change -x x x Equilibrium 0.010-x x x

32. K b = [NH 4 + ] [ OH - ] [NH 3 ] 1.8 x 10 -5 = (x)(x) 0.010 – x x = 4.2 x 10 -4 M At equilibrium: 0.010 -4.2 x 10 -4 = 0.00958≈0.01

33. Once you find [OH - ], you find the pOH pH + pOH = 14

34.  pH indicators 1. indicator (In) is an acid or base that dissociates in a known pH range HIn (aq) acid form OH -   H + H + (aq) + In - (aq) base form 2.

35. 3. Types of indicators a. Methyl red: dye that turns red in acids 0-4.4 : red 4.5-6.1: orange 6.2-above: yellow b. Phenolphthalein: colorless in acids, pink in bases below pH 8.2: colorless above pH 10: pink

36. c. Bromothymol blue: used for weak acids/bases below pH of 6.0 = yellow pH of 7.0 = green above pH of 7.6 = blue d. Universal indicator: used for acids and bases 0-3 3-6 7 8-11 11-14 red orange/ green blue purple yellow

37. Problems with indicators 1. Only work at room temperature (will change colors at different temp) 2. Salts in the solution may change the dissociation process  pH meter: equipment used to measure pH (best pH measurement)