1. Chapter 14 Acids, Bases, and pH

2. Objectives 14.1 Distinguish acids from bases by their properties 14.1 Relate acids and bases to their reactions in water 14.1 Evaluate the central role of water in the chemistry of acids and bases

3. Objectives 14.2 Relate different electrical conductivities of acidic and basic solutions to their degree of dissociation or ionization 14.2 Distinguish strong and weak acids or bases by their degree of ionization 14.2 Compare and contrast the composition of strong and weak solutions of acids or bases 14.2 Relate pH to the strengths of acids and bases

4. General Info Bases Taste Bitter to the taste React with our skin to form soap (so it feels soapy) React with oils and greases (so often used in cleaning products (ammonia) Can cause dyes to change color. Electrolytes Corrosive

5. General Info about Acids Taste sour Corrosive Electrolytes Attack skin by dissolving fatty acids.

6. Some Common Acids and Bases Can you name a few?

7. Indicators Indicators change color at different pH’s and are useful for knowing what the pH of the solution is (they won’t give exact, but they will let you know more or less acidic than some number)

8. During a titration, you place indicators into the solution to let you know when the pH has crossed a certain point. Different titrations have different equivalent points, so choosing an appropriate one is important.

9. For a SA and SB reaction, indicator isn’t important. For WA and WB reactions, it is more important

11. What is an Acid and Base? 3 different definitions that describe what Acids and Bases are – Arrhenius Acids and Bases – Bronsted-Lowry Acids and bases – Lewis Acids and Bases

12. Arrhenius Background First definition of an acid came from the Swedish chemist, Svante Arrhenius in 1890. He was the first to recognize a solution became more acidic when hydronium ions were present in greater number.

13. Arrhenius Background His theory defined an acid as any substance that when added to water increases the hydronium ion concentration Bases as any substance when added to water that increase the hydroxide ion concentration.

14. Bronsted Acid and Bases The Bronsted definition of an acid comes from a man from Denmark who made his proposal in 1923. His theory helped to overcome the shortcomings of the Arrhenius definition by allowing us to describe solutions which were not aqueous.

15. Bronsted Acids and Bases His definition has acids as proton donors and bases as proton acceptors. – Proton = Hydrogen (has 1 Proton)

16. Lewis Acid and Bases The Lewis definition of an acid comes from the same guy who came up with Lewis electron diagrams (Lewis Dot Diagrams, which are very useful for drawing molecules showing electrons, lone pairs, and bonds).

17. Lewis Acids and Bases He was an American chemist who did most of his work with bonding and thermodynamics, but made important contributions to our understanding of acids and bases. He expanded the definition of acid to be a Lone pair acceptor and the base to be a lone pair donor (which forms a bond).

18. In Summary (Acids) An Arrhenius acid generates hydronium ions in water. A Bronsted acid donates protons A Lewis acid is a lone pair acceptor

19. In Summary (Bases) An Arrhenius base generates hydroxide ions in water A Bronsted base accepts protons A Lewis base is a lone pair donator

20. Bronsted Most effective for identifying acids/bases HA + B  A- + HB+ HA = AcidB = Base A- = Reacted Acid (Gave away hydrogen) HB+ = Reacted base (Accepted hydrogen)

21. Practice Identifying On the following slides, identify the acid and base (forward reaction) and then whether each acid/base definition works

22. Question 1 CH 3 COOH (aq) + H 2 O (L)  CH 3 COO - (aq) + H 3 O + (aq)

23. Question 2 HCl (aq) + NH 3 (aq)  NH 4 + (aq) + Cl-

24. Question 3 NH 3 (aq) + H 2 O (L)  NH 4 + (aq) + OH- (aq)

25. Amphiprotic/Amphoteric A substance that can act as an acid or a base

26. Amphiprotic/Amphoteric Some chemicals can act as a base and an acid, these substances get the above special names. Some notable chemicals that act as both are: Ammonia (NH 3 ) and Water (H 2 O)

27. Nomenclature Throw as many hydrogen’s onto the anion/polyatomic ion and change the ending from ate to ic and ite to ous. With the exception the halogens (Add hydrochloric acid = HCl) Perchlorate (ClO 4 -1 ) becomes Perchloric Acid (HClO 4 )

28. Nomenclature Nitrate (NO 3 -1 ) becomes Nitric Acid (HNO 3 ) Sulfite (SO 3 -2 ) becomes Sulfurous Acid (H 2 SO 3 ) Also: When writing the formula, if it is an acid, the H is placed at the beginning to denote that the chemical is an acid.

29. Major Ideas Strong and Weak Acids and Bases Conjugate Acids and Bases Bond Strength Acid/Base Equilibrium

30. Warm Up What is the name of the following acid: H 2 CO 3 (CO 3 -2 = Carbonate) How do acids and bases taste different from one another? What are the three definitions of acids and bases? Are any of the definitions better than the other ones?

31. What are Strong Acids/Bases? Strong acids and bases fully ionize when placed in water. The Unionized from of the acid is not present. There is no equilibrium, Strong acids and bases are completion reactions in water. – HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)

32. Know These Acids There are 6 strong acids you will need to know (in no particular order): – Nitric Acid (HNO 3 ) – Sulfuric Acid (H 2 SO 4 ) – PerChloric Acid (HClO 4 ) – HydroBromic Acid (HBr) – HydroChloric Acid (HCl) – HydroIodic Acid (HI)

33. Strong Acids Every one of those, when placed in water, will ionize and all you will have is Hydronium and the Anion. – HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) – Acid + Water  Hydronium + Anion

34. Strong Bases A strong base is any Alkali Metal with a Hydroxide. Such as: Sodium Hydroxide (NaOH), Potassium Hydroxide (KOH) Strong bases also fully Ionize in water

35. Molarity If 1.0 Mols of HCl are placed into 1.0 L of water, the Concentration of Hydronium is equal to the concentration of the Strong Acid placed into the solution. The concentration of Hydronium is 1.0 Molar in this example.

36. Conjugate Acids/Bases When an Acid gives up its proton, the molecule left (minus a hydrogen) is called a conjugate base. The conjugate base has a negative charge and, being negative, has the ability to attract a nearby hydrogen (which is positive) to bond to it.

37. Conjugate Acids/Bases When a Base receives a proton, the molecule (plus a hydrogen) is called a conjugate acid. The conjugate acid has a positive charge and is looking to give up its positive charge to another molecule.

38. Conjugate Acids/Bases Summarized Conjugate Acid  An acid that forms when a base gains a proton Conjugate Base  A base that forms when an acid loses a proton

39. Example Acid Base – HC 2 H 3 O 2 (aq) + H 2 O  – C 2 H 3 O 2 - (aq) + H 3 O + (aq) – Conj. Base Conj. Acid

40. Example Base Acid NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) Conj. Acid Conj. Base

41. Practice Questions What is the conjugate Base of the following Acids? HClH 2 SO 4 HydroniumAmmonium (NH 4 + )

42. Practice Label the following as Acid/Base/CB/CA HF + H 2 O  F - + H 3 O + CH 3 O - +H 2 O  CH 3 OH +OH -

43. What are Weak Acids/Bases? Weak acids and bases partially ionize when placed into water. An equilibrium is established where the Hydrogens are fought over (who gets to have the hydrogen?). Acetic acid (HC 2 H 3 O 2 ) when placed into water partially ionizes.

44. Weak Acid/Base Example Weak Acid – HC 2 H 3 O 2 (aq) + H 2 O  – C 2 H 3 O 2 - (aq) + H 3 O + (aq) Weak Base – NH 3 (aq) + H 2 O (l)  – NH 4 + (aq) + OH - (aq)

45. How weak is weak? If 100 weak acid molecules were put into a solution of water, only about 5 would react. Most weak acids are found with their hydrogen

46. Calculating the Concentration The concentration of Hydronium and Hydroxide are more difficult to find for weak acids and bases and requires a bit of math using our equilibrium expressions.

47. Acid/Base Equilibrium Where H + = Hydronium HA = Unionized acid, A - = Conjugate Base Where OH - = Hydroxide B = Unionized base HB + = Conj. Acid

48. Underlying Factor Bond Strength Ultimately, what determines how strong an acid or base is the bond strength between the hydrogen and the molecule. The strong acids are strong because they have weak bonds with hydrogen.

49. Bond Strength The hydrogen makes a much stronger bond with a fellow water than with its anion. The anion left is so weak it can’t attract a hydrogen to bond with it. The stronger the acid, the less attraction its conjugate base has for a positive hydrogen.

50. Strong Acids vs Weak Acids Strong acids don’t have values we talk about (way too high) Most weak acids have values below 0.01

51. Ka = Weak Acid Constant The general formula for Weak Acids is: – HF (aq) + H 2 O (l)  H 3 O + (aq) + F - (aq)

52. What it all stands for Where [H + ] refers to the hydronium concentration (H 3 O + ) Where [A - ] refers to the conjugate base concentration (F - in example) Where [HA] refers to the unionized acid concentration (HF in example) There is no Water in the equation because it is a pure liquid.

53. The Meaning of Ka The value of Ka is generally smaller than 1. Weak acids tend to ionize to a very small extent and their Ka values reflect this. As the value of Ka increases, the strength of the acid increases (strength being how well it loses its hydrogen) Write the equilibrium expression for the weak acid HC 5 H 7 O 5

54. Kb = Base Equilibrium Constant The general formula for Weak Bases is – NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq)

55. What it stands for Where [OH - ] refers to the hydroxide concentration (OH - in example) Where [HB + ] refers to the base plus an extra hydrogen (NH 4 + in example) Where [B] refers to the base before it grabs the hydrogen (NH 3 in example)

56. The Meaning of Kb The value of Kb is generally less than 1. Weak bases tend to grab hydrogen’s rarely compared to strong bases. As the value of Kb increases, the stronger the weak base and the more likely it is to grab a hydrogen from a water. Write the equilibrium expression for the weak base C 2 H 5 OH

57. pH Scale: Before we start… Quick Math: The Logarithm Scale On a logarithmic scale, a change in 1 represents a change in 10, a change in 2 represents a change of 100. – The Richter Scale (Earthquakes) An earthquake that registers a 5.0 is 1000x stronger than an earthquake that registers 2.0

58. The pH Scale Is a measure of how many Hydroniums are in the water. The pH means: Powers of Hydrogen – On the Log Scale A pH of 7 means the concentration of Hydronium is 0.0000001 or 1x10 -7

59. pH Scale pHpOH[H 3 O + ][OH - ] 01411 x10 -14 1130.1 1x10 -13 2120.011x10 -12 3111x10 -3 1x10 -11 4101 x10 -4 1 x10 -10 591 x10 -5 1 x10 -9

60. pH Scale pHpOH[H 3 O + ][OH - ] 681 x10 -6 1 x10 -8 771 x10 -7 1 x10 -7 861 x10 -8 1 x10 -6 951 x10 -9 1 x10 -5 1041 x10 -10 1 x10 -4 1131 x10 -11 1 x10 -3

62. Hydronium in Distilled Water pH: Why it adds up to 14 Water self ionizes (to a very small extent) 2 H 2 O (l) H 3 O + (aq) + OH - (aq) Kw = [H 3 O + ][ OH - ] Kw = 1.0 x10 -14 (This is a constant)

63. The pH scale When acids are added to water, they add to the hydronium concentration (decreasing the hydroxide concentration) When bases are added to water, they add to the hydroxide concentration (decreasing the hydronium concentration).

64. The pH Scale at neutral Kw = [OH][H] Kw = (1E -7 )(1E -7 ) = 1E -14 pH = 7 A pH of 7 is considered Neutral (it contains just as much base as acid)

67. Question Can you have a pH less than 0 and greater than 14? Yes, a 10.0 Molar HCl solution has a pH of -1

68. What is the pH? A solution has a hydronium concentration of 0.001? A solution has a hydroxide concentration of 0.00001? A solution has equal concentrations of hydronium and hydroxide?