HCl (aq) hydrochloric acid HBr (aq) hydrobromic acid HI (aq) hydroiodic acid

1. Binary acids – use the prefix “hydro” + the name of the negative (second ion), changing the ending to “-ic acid”.Oxy-acids – Name the polyatomic ion, change the ending to either “-ic acid” or “-ous acid” according to the “ate = ic, ite = ous” rule. • HCl (aq) hydrochloric acid • HBr (aq) hydrobromic acid • HI (aq) hydroiodic acid • HF (aq) hydrofluoric acid • H2S (aq) hydrosulfuric acid • H2SO4 sulfuric acid • H2SO3 sulfurous acid • H3PO4 phosphoric acid • H3PO3 phosphorous acid • HNO3 nitric acid

2. Acids and Bases

3. What is an Acid? Base? • What do you think and acid is? • What do you think a base is?

4. Properties of Acids and Bases! • Acids have specific properties • Sour taste • pH beow 7.00 S.U. • Turn litmus paper red • Dissolve metals • Bases have similar properties • Bitter taste • pH above 7.00 S.U. • Feel slippery • Turn litmus paper blue

5. What else can they do? • Acids and bases neutralize each other • In other words, they raise/lower the pH to a non-acid/non-base • Acids and Bases can both destroy human tissue

6. How do they do that? • Acids and bases are electrolytes • The ions that they produce give them their properties

7. Strength of Acids • A strong acid is one that dissolves completely in solution. • A weak acid is one that does not dissolve completely. • Why would an acid not dissolve completely?

8. Do you know any acids or bases? • List some common acids and bases • Vinegar is Acetic Acid • Ammonia is a Base • Sodium Bicarbonate (Baking soda) is a Base • Batteries contain Sulfuric Acid • Your stomach contains Hydrochloric Acid!

9. There are more than one • There are 3 different types of acids and bases • Arrhenius • Brönsted-Lowry • Lewis

10. Arrhenius Acids and Bases • An Arrhenius Acid is a substance that releases H3O+ ions in an aqueous solution • HCl H2O H3O+ + Cl- • An Arrhenius Base is a substance that releases –OH in an aqueous solution • NaOH H2O Na+ + OH-

11. Brönsted-Lowry Acids and Bases • A Brönsted-Lowry acid is any species that can donate a proton (H+ ion) to another species; a proton donor • A Brönsted-Lowry base is any speices that can accept a proton (H+ ion) from another species; a proton acceptor

12. HCl + H2O  H3O+ + Cl- • Which is the Brönsted-Lowry acid/base? • Acid is the proton donator • Base is the proton acceptor • HCl is the acid • H2O is the base

13. Water: Acid or Base? • Take a look at these two reactions 1. NH3 + H2O  NH4++ OH- 2. HCl + H2O  H3O++ Cl- • In reaction #1, which is the B-L acid? Base? • Water is the acid and ammonia is the base • In reaction #2? • HCl is the acid and water is the base

14. Amphiprotic • A compound that can be both a proton donor OR a proton acceptor in separate reactions is called amphiprotic or amphoteric

15. Lewis Acids and Bases • A Lewis Acid is any species that can accept a pair of electrons from another species; electron pair acceptor • A Lewis Base is any species that can donate a pair of electrons from another species; electron pair donor

16. Remember Lewis Structures! • Let’s look at a few examples H+ H+ is an electron pair acceptor :Ö-H- OH- is an electron pair donor :Ö-H Lewis acid-base product .. H ..

17. Compare the Definitions

18. Acid-Base Equilibria • In a chemical reaction involving both acids and bases, dynamic equilibrium can still be achieved • A species that is an acid in the forward direction, can become the base in the reverse direction

19. HNO2 + H2O  H3O+ + NO2- acid base acid base • Let’s take a look at this reaction • HNO2 donates a proton in the forward reaction and is therefore a B-L acid • H2O accepts a proton in the forward reaction and is the B-L base • In the reverse direction, H3O+ donates a proton; B-L acid • And NO2- accepts a proton; B-L base

20. Conjugate Pairs • When an acid donates a proton, it becomes a conjugate base • When a base accepts a proton, it becomes a conjugate acid • Why? • The acid-base reaction can be a reversible reaction and so when you move in the opposite direction, the compounds are redefined as acids and/or bases

21. Conjugates • For example, look at the following reaction HCl + H2O  H3O+ + Cl- • Which reactant is the acid? The base? • Now, in the reverse direction, which product is the acid? The base? • Now, what are the conjugate pairs?

22. Conjugates, some more • Look at this reaction: NH3 + H2O  NH4+ + OH- • Which reactant is the acid? The base? • Now, in the reverse direction, which product is the acid? The base? • Now, what are the conjugate pairs?

23. Conjugate Practice • What are the conjugate pairs for the following reactions? • HNO3 + OH-  H2O + NO3- • HNO3 is the acid and NO3- is the conjugate base • OH- is the base and H2O is the conjugate acid • H2SO4(aq) + SO32-(aq)  HSO4-(aq) + HSO3-(aq) • Sulfuric acid is the acid and HSO4- is the conjugate base • Sulfite ion is the base and HSO3- is the conjugate acid

24. Strength • If a acid is strong, then the conjugate base is weak and vice-versa

25. Strong/Weak Electrolytes • An electrolyteis a solution that will conduct an electrical current. • Any ionic compound that dissociates in solution will conduct electricity. • If a substance dissociates 100%, it is called a strongelectrolyte. • A weak electrolyte dissociates < 100% (usually < 5%)

26. Strong Acids • Binary acids: • HCl, HBr, and HI are strong acids. All other binary acids are weak. • Oxy acids: • If the number of oxygen atoms exceed the number of hydrogen atoms by 2 or more, the acid is strong.

27. Strong Bases • Hydroxides made from metals in group 1 or group 2 are strong bases. • All other hydroxides are weak bases.

28. Molecular Structure and the Strength of Acids • The strength of an acid depends on a number of factors, such as the properties of the solvent, the temperature, and the molecular structure of the acid. • We compare the strengths of two acids, in the same solvent and at the same temperature. That way we can focus on the structure of the acid.

29. Binary Hydrides • HA ---> H+ + A- • Two factors influence the extent to which the acid undergoes ionization. • One is the strength of the H-A bond • the stronger the bond, the more difficult it is for the HA molecule to break up and hence the weaker the acid.

30. Polarity versus Strength • The other factor is the polarity of the H-A bond. • The difference in electronegativities between H and A results in a polar bond. • If the bond is highly polarized, it is a stronger acid.

31. Hydrohalic Acids • The halogens form a series of binary acids called the hydrohalic acids. • The strengths of the hydrohalic acids increase in the following order: • HF<<HCl<HBr<HI

32. Other Non-metallic Acid Strengths • In any vertical column (Group) of nonmetallic elements, there is a tendency toward increasing acidity of the hydride with increasing atomic number (as you go down the group). • For example, among the group VIA elements the acid strength increases in the order: • H2O< H2S<H2Se<H2Te

33. Oxyacids • Many common acids contain one or more O-H bonds.

34. Practice Problems • HF • weak • H2S • weak • H2SO4 • strong • H2CO3 • weak • Fe(OH)3 • weak 6. Barium hydroxide strong 7. Chloric acid strong 8. Sulfurous acid weak 9. Hypochlorous acid weak 10. Tin(IV) hydroxide weak Ba(OH)2 HClO3 H2SO3 HClO Sn(OH)4

35. Review • In the following equations, name the acid and base. What was your justification? • NH3 + H2O  NH4+ + OH- • 2NaBr + H2SO4  Br2 + SO2 + 2NaOH

36. Review • What is the difference between a strong acid and a weak acid? • A buffer is a solution made from a weak acid and it’s conjugate base that neutralizes a small amount of acids or bases added to it

37. Calculating pH

38. pH • pH is a measure of how acidic, neutral or basic a solution is • The scale for pH is 0-14 • 7 is neutral • 0-7 is acidic • 7-14 is basic

39. The Numbers mean . . . • pH + pOH = 14 S.U. • pH = -log10[H3O+] • pOH = -log10[OH-]

40. What is pH? • The term pH refers to a scale that describes how strongly acidic or basic a solution is • pH + pOH = 14 • Acidic = 0-7 • Neutral = 7 • Basic = 7-14

41. The Kw Constant • Kw = 1.0 x 10-14 • at standard thermodynamic temperature (25oC or 298 K)

42. Ionization Constants of Acids and Bases • HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) • Keq= [H3O+][A-] / [HA][H2O] • In a dilute solution, the concentration of water is constant, so • Keq= [H3O+][A-] / [HA]

43. Ionization Constants • The Ionization Constant for water; Kw is derived as follows: • First: for the acids HA(aq) + H2O(l)  H3O+(aq) + A-(aq) Keq = Ka = [H3O+][A-] / [HA]

44. Now the bases . . . B(aq) + H2O(l)  BH+(aq) + OH-(aq) Keq = Kb = [BH+][OH-] / [B]

45. Combine the two HA(aq) + H2O(l)  H3O+(aq) + A-(aq) Keq = Ka = [H3O+][A-] / [HA] & B(aq) + H2O(l)  BH+(aq) + OH-(aq) Keq = Kb = [BH+][OH-] / [B] = Kw= [H3O+][OH-]

46. Calculations • Problem: What are the hydronium and hydroxide ion concentrations in a 0.10M HCl? HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) Kw = [H3O+][OH-] 1.0 x 10-14 = (0.1M) x [OH-] [OH-] = 1.0 x 10-13

47. Calculating pH • A pH is equal to the -log10 of the Hydronium ion concentration pH = -log10[H3O+] • If an acid has a [H3O+] of 1.0 x 10-2M; then the pH is 2 S.U. • S.U. = standard unit

48. Calculating [H3O+] • The concentration of Hydronium ions can be determined by the reverse calculation, sort of • If the pH of a solution is 5.05 S.U., then the concentration of [H3O+] is 10-pH. • 5.05 S.U. = 10-5.05 = 8.9 x 10-6 M

49. Tricks • If your concentration is 1.0 x 10-10 M, then your pH will be 10 S.U.! • The negative exponent becomes the pH • If your concentration is greater than 1.0, then the pH will be less than the negative exponent. • 2.5 x 10-10 M = >10 S.U. = 9.6 • This comes in handy when you are answering multiple choice questions about the pH of a solution.

50. Practice Problems • What is the pH of a solution if the [H3O+] is 5.0 x 10-3M? • What is the pH of a solution if the [OH-] is 2.0 x 10-3M? • What is the [H3O+] of a solution whose pH is 3.3? • What is the [OH-] of a solution whose pH is 8.1?