Chapter 9: Acids, Bases, and Salts

1. Chapter 9:Acids, Bases, and Salts

2. Acids Arrhenius acids Produce H+ ions in water. H2O HCl(g) H+(aq) + Cl-(aq) Are electrolytes. Have a sour taste. Turn litmus red. Neutralize bases.

3. Naming Some Common Acids

4. Bases Arrhenius bases Produce OH− ions in water. Taste bitter or chalky. Are electrolytes. Feel soapy and slippery. Neutralize acids.

5. Some Common Bases Bases with OH− ions are named as the hydroxide of the metal in the formula. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide Al(OH)3 aluminum hydroxide Fe(OH)3 iron(III) hydroxide

6. BrØnsted-Lowry Acids and Bases According to the BrØnsted-Lowry theory, Acids donate a proton (H+). Bases accept a proton (H+).

7. NH3, A BrØnsted-Lowry Base In the reaction of ammonia and water, NH3 is the base that accept H+. H2O is the acid that donates H+.

8. Comparing Acids and Bases

9. Conjugate Acid-Base Pairs In any acid-base reaction, there are two conjugate acid-base pairs Each related by the loss and gain of H+ . One occurs in the forward direction. One occurs in the reverse direction. conjugate acid-base pair 1 HA + B A− + BH+ conjugate acid-base pair 2

10. Conjugate Acids and Bases In this acid-base reaction, An acid HF donates H+ toform its conjugate base F−. A base H2O accepts H+ to form its conjugate acid H3O+. There are two conjugate acid-base pairs.

11. Conjugate Acid-Base Pairs In the reaction of HF and H2O One conjugate acid-base pair is HF/F−. The other conjugate acid-base pair is H2O/H3O+. Each pair is related by a loss and gain of H+.

12. Conjugate Acid-Base Pairs In the reaction of NH3 and H2O, One conjugate acid-base pair is NH3/NH4+ The other conjugate acid-base is H2O/OH¯.

13. Lewis Acids and Bases Lewis Acids accept electron pairs Lewis Bases donate electron pairs Broadens definition of an acid Example: H+ + H2O  H3O+ Water donates a lone pair The Proton accepts the lone pair

14. Marble, a naturally occurring form of CaCO3, reacts with hydrochloric acid, HCl. Eggshells are also made of CaCO3.

15. Acids can react with and dissolve certain metals to yield hydrogen gas in a redox reaction. The activity series is a tabular representation of the tendencies of metals to react with H+.

16. The reaction of zinc metal with hydrochloric acid can be written as follows: The chloride ion (Cl-) is a spectator ion. The hydrogen ion gains an electron to be reduced, and therefore, the HCl is the oxidizing agent. The zinc metal loses electrons to be oxidized, and therefore, the zinc metal is the reducing agent. This reaction occurs because zinc is above the reactivity line that divides lead (reactive) from copper (unreactive) in the activity series.

17. Metals vary in their ability to reduce hydrogen ions (H+) to hydrogen gas (H2). The difference is apparent when iron, zinc, and magnesium (left to right) are put into hydrochloric acid (HCl) of the same molarity.

18. Strong acids completely ionizes (100%) in aqueous solutions. HCl(g) + H2O(l) H3O+ (aq) + Cl− (aq) Weak acids dissociate only slightly in water to form a solution of mostly molecules and a few ions. H2CO3(aq) + H2O(l) H3O+(aq) + HCO3− (aq) Strengths of Acids

19. Strong Acids A strong acid dissolved in water Dissociates into ions. Gives H3O+ and the anion (A-).

20. Weak Acids Weak acids dissolved in water, Dissociate only a few molecules. Remain mostly as the undissociated (molecular) form. Have low concentrations of H3O+ and anion (A-). HA(aq) + H2O(l) H3O+(aq) + A−(aq)

21. Strong vs Weak Acids In solution, HCl, a strong acid, dissociates 100%. acetic acid, CH3COOH a weak acid, is mostly molecules and only a few ions.

22. Strong acids Only a handful of acids are strong. The common ones are listed below. Have weak conj bases. Other acids are generally weak! Strong Acids

23. Strong Bases Strong bases Are Group 1A and 2A Metal Oxides and Hydroxides. Include LiOH, NaOH, KOH, and Na2O. Dissociate completely in water. KOH(s) K+(aq) + OH−(aq)

24. Weak Bases Weak bases Are most other bases. Dissociate only slightly in water. Form only a few ions in water. NH3(g) + H2O(l) NH4+(aq) + OH−(aq)

25. Acid Dissociation Constant In a weak acid, the rate of the dissociation of the acid is equal to the rate of the association. HA + H2O H3O+ + A- The equilibrium expression is Ka = [H3O+][A-] [HA]

26. Some Acid Dissociation Constants Ka values for some acids TABLE 10.4

27. Writing Ka for a Weak Acid Write the Ka for H2S. 1. Write the equation for the dissociation of H2S. H2S(aq) + H2O(l) H3O+(aq) + HS−(aq) 2. Set up the Ka expression Ka = [H3O+][HS-] [H2S]

28. In the ionization of water, H+ is transferred from one H2O molecule to another. One water molecule acts as an acid, while another acts as a base. H2O + H2O H3O+ + OH− .. .. .. .. H:O: + H:O: H:O:H+ + :O:H− .. .. .. .. HH H water water hydronium hydroxide ion (+)ion (-) Ionization of Water

29. Pure Water is Neutral In pure water, The ionization of water molecules produces small, but equal quantities of H3O+ and OH−ions. Molar concentrations are indicated in brackets as [H3O+] and [OH−]. [H3O+] = 1.0 x 10−7 M [OH−]=1.0 x 10−7 M

30. Acidic Solutions Adding an acid to pure water Increases the [H3O+]. Increases the [H3O+] to more than 1.0 x 10-7 M. Decreases the [OH−].

31. Basic Solutions Adding a base to pure water, Increases the [OH−]. Causes the [OH−] to exceed 1.0 x 10− 7M. Decreases the [H3O+].

32. The ion product constant, Kw, for water Is the product of the concentrations of the hydronium and hydroxide ions. Kw = [ H3O+] [ OH−] Can be obtained from the concentrations in pure water. Kw = [ H3O+] X [ OH−] Kw = [1.0 x 10− 7 M] X [ 1.0 x 10− 7 M] = 1.0 x 10− 14 Ion Product of Water, Kw

33. [H3O+] and [OH−] in Solutions In neutral, acidic, or basic solutions, the Kw is always 1.0 x 10−14.

34. Calculating [H3O+] What is the [H3O+] of a solution if [OH−] is 5.0 x 10-8M? STEP 1Write the Kw for water. Kw = [H3O+ ][OH−] = 1.0 x 10−14 STEP 2Rearrange the Kw expression. [H3O+] = 1.0 x 10-14 [OH−] STEP 3Substitute [OH−]. [H3O+] = 1.0 x 10-14 = 2.0 x 10-7 M 5.0 x 10- 8

35. pH Scale The pH of a solution Is used to indicate the acidity of a solution. Has values that usually range from 0 to 14. Is acidic when the values are less than 7. Is neutral with a pH of 7. Is basic when the values are greater than 7.

36. pH of Everyday Substances

37. Testing the pH of Solutions The pH of solutions can be determined using a) pH meter, b) pH paper, or c) indicators that have specific colors at different pH values.

38. Mathematically pH Is the negative log of the hydronium ion concentration, pH = - log [H3O+] For a solution with [H3O+] = 1.0 x 10−4 pH =−log [1.0 x 10−4 ] pH = - [-4.0] pH = 4.0 Calculating pH

39. Calculating pH pH = -log [H3O+] Do your values match the ones listed here? [H3O+] = 1.0x 10-4 pH = 4.0 [H3O+] = 8.00x 10-6 pH = 5.10 [H3O+] = 2.40x 10-8 pH = 7.62

40. [H3O+], [OH-] and pH Values

41. Calculating [H3O+] from pH Calculate the [H3O+] for pH = 3.8 is calculated as follows: STEP 1 Enter the pH value, change sign -3.8 STEP 2 Use 2ndand 10x keys or inverse and log keys 1.584893 -04 STEP 3Adjust the significant figures 1.6 x 10-4

42. Acid/Base Reactions Typically involve exchange of protons Reaction of acid with 4 common bases is shown below. Remember these common reactions! H+ (aq) + OH- (aq)  H2O (l) H+ (aq)+ HCO3- (aq)  H2O (l) + CO2 (g) 2 H+ (aq)+ CO32- (aq)  H2O (l) + CO2 (g) H+ (aq) + NH3 (aq)  NH4+ (aq)

43. Write the balanced equation for the neutralization of magnesium hydroxide and nitric acid. STEP 1 Write the base and acid formulas. Mg(OH)2(aq) + HNO3(aq) STEP 2 Balance OH- and H+. Mg(OH)2(aq) + 2HNO3(aq) STEP 3 Balance with H2O Mg(OH)2(aq) + 2HNO3(aq) salt + 2H2O (l) STEP 4Write the salt from remaining ions. Mg(OH)2(aq) + 2HNO3(aq) Mg(NO3)2(aq) + 2H2O(l) Balancing Neutralization Reactions

44. Antacids Antacids are used to neutralize stomach acid (HCl).

45. Acid-Base Titration Titration Is a laboratory procedure used to determine the molarity of an acid. Uses a base such as NaOH to neutralize a measured volume of an acid. Base (NaOH) Acid solution

46. Indicator An indicator Is added to the acid in the flask. Causes the acid solution to change color when the acid is neutralized.

47. End Point of Titration At the end point, The indicator has a permanent color. The volume of the base used to reach the end point is measured. The molarity of the acid is calculated using the neutralization equation for the reaction.

48. Calculating Molarity from A Titration with A Base What is the molarity of an HCl solution if 18.5 mL of a 0.225 M NaOH are required to neutralize 10.0 mL HCl? HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) STEP 1 Calculate the moles of base. 18.5 mL NaOH x 1 L x 0.225 mole NaOH 1000 mL 1 L = 0.00416 mole NaOH

49. Calculating Molarity (continued) STEP 2 Calculate the moles of HCl. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) 0.00416 mole NaOH x 1 mole HCl 1 mole NaOH = 0.00416 mole HCl STEP 3Calculate the molarity of HCl. 10.0 mL HCl = 0.010 L HCl 0.00416 mole HCl = 0.416 M HCl 0.0100 L HCl

50. Salts that Form Neutral Solutions A salt solution containing the anion of a strong acid and the cation of a strong base Does not produce or attract H+ from water. Is a neutral solution. Of KNO3, for example, is neutral because it contains a cation from a strong base (KOH) and an anion from a strong acid (HNO3).