1. Acids and Bases

2. Properties of Acids  Sour taste  React w/ metals to form H 2  Most contain hydrogen  Are electrolytes  Change color in the presence of indicators (turns litmus red)  Has a pH lower than 7

3. Two Types of Acids  Strong acids –Any acid that dissociates completely in aqueous sol’n  Weak acids –Any acid that partially dissociates in aqueous sol’n

4. Properties of Bases  Bitter taste  Slippery feel  Are electrolytes  Change color in the presence of indicators (turns litmus blue)  Has a pH higher than 7

5. Types of Bases  Strong Base –Any base that dissociates completely in aqueous sol’n  Weak Base –Any base that partially dissociates in aqueous sol’n

6. Neutralization  Neutralization rxn: a rxn of an acid and a base in aqueous sol’n to produce a salt and water  Salt: compound formed from the positive ion of a base and a negative ion of an acid  Properties of the acid and base cancel each other

7. Arrhenius Model of Acids and Bases  Proposed the model in 1887  Acid: any compound that produces H + ions in aqueous (water) sol’n  Base: any compound that produces OH - (hydroxide) ion in aqueous sol’n  Offers an explanation of why acids and bases neutralize each other (H + + OH - = H 2 O)

8. Problems with Model  Restricts acids and bases to water sol’ns (similar reactions occur in the gas phase)  Does not include certain compounds that have characteristics of bases (e.g., ammonia)

9. Brønsted-Lowry Model of Acids and Bases  Brønsted acid: a hydrogen ion donor (H +, or proton)  Brønsted base: a hydrogen ion acceptor  Defines acids and bases independently of how they behave in water  Amphiprotic: having the property of behaving as an acid and a base –Also called amphoteric, e.g., water

10. Conjugate Acid-Base Pairs  The rxn between Brønsted-Lowry acids and bases can proceed in the reverse direction (reversible reactions) HX (aq) + H 2 O (l)  H 3 O + (aq) + X - (aq)  The water molecule becomes a hydronium ion (H 3 O + ), and is an acid because it has an extra H + to donate  The acid HX, after donating the H +, becomes a base X -

11. Conjugate Acids and Bases HX (aq) + H 2 O (l)  H 3 O + (aq) + X - (aq) Acid BaseConjugate Acid Conjugate Base Forward reaction: Acid and base Reverse reaction: Conjugate acid and conjugate base

12.  Conjugate Acid: species produced when a base accepts a hydrogen ion from an acid  Conjugate Base: species produced when an acid donates a hydrogen ion to a base  Conjugate Acid-Base Pair: two substances related to each other by the donating and accepting of a single hydrogen ion

15. Types of Acids  Monoprotic acids: acids that contain only 1 hydrogen; e.g., HCl  Diprotic acids: acids that contain 2 hydrogens; e.g. H 2 CO 3  Triprotic acids: acids that contain 3 hydrogens; e.g. H 3 PO 4

16. More Types of Acids  Binary acids: acids that contain only 2 elements; e.g. HF  Polyatomic acids: acids that contain more than 2 elements; e.g. H 2 SO 4 – These acids contain polyatomic ions – Also called ternary or oxy- acids

17. Naming Binary Acids  Start with the prefix hydro-  Put it in front of the root word of the anion (- charged ion)  Add –ic to the end  Examples – Hydrobromic (HBr) – Hydrofluoric (HF) – Hydroiodic (HI) – Hydrochloric (HCl)

18. Naming Polyatomic Acids  Start with the root word of the name of the polyatomic ion  Add –ous if name ends in –ite  Add -ic if name ends in –ate  Examples: – Chlorous (from chlorite, ClO 2 - ) – Nitric (from nitrate, NO 3 - ) – Sulfurous (from sulfite, SO 3 -2 )

19. pH and [H 3 O + ]  pH: number that is derived from the concentration of hydronium ions ([H 3 O + ]) in sol’n – pH = -log [H 3 O + ] – As pH increases, [H 3 O + ] decreases  Scale ranges from 0 – 14 – pH = 7 is neutral – pH < 7 is acidic – pH > 7 is basic

20. p[OH]  pOH = - log [OH - ]  pH + pOH = 14.00  Calculating ion concentrations from pH  [H + ] = antilog (-pH)  [OH - ] = antilog (-pOH)

21. Dissociation Constants  Acid dissociation constant: (K a ): the equilibrium constant for the rxn of an aqueous weak acid and water  Base dissociation constant: (K b ): the equilibrium constant for the rxn of an aqueous weak base w/ water  Both are derived from the ratio of the concentration of the products and reactants at equilibrium

22. Acid Dissociation Constant K a = [H 3 O + ] [A - ] [HA]  K a is a measure of the strength of an acid  K a values for weak acids are always less than one  Used mostly w/ weak acids because the K a values for strong acids approach infinity

23. Examples  HMnO 4 (aq) + H 2 O (l)   H 2 S (aq) + H 2 O (l) 

24. Base Dissociation Constant K b = [HB + ] [OH - ] [B]  K b is a measure of the strength of a base  K b values for weak bases are always less than 1  K b values for strong bases approach infinity

25. Examples  H 2 NOH (aq) + H 2 O (l)   NH 3 (aq) + H 2 O (l) 

26. Water  Water can dissociate into its component ions, H + and OH - – 2H 2 O (l)  H 3 O + (aq) + OH - (aq)  One water molecule acts as a weak acid, and the other acts as a weak base  The ions are present in such small amounts they can’t be detected by a conductivity apparatus  In pure water, [H 3 O + ] =1.0 x 10 –7 M and [OH - ] = 1.0 x 10 -7 M

27. Dissociation Constant for Water  It is defined as K w : the ion product constant for water  K w = [H 3 O + ] [OH - ]  K w = (1.0 x 10 -7 )(1.0 x 10 -7 )  K w = 1.0 x 10 -14  The value of K w can always be used to find the concentration of either H 3 O + or OH - given the concentration of the other

28. Examples What is the pH of a 0.001 M sol’n of HCl, a strong acid?

29. Examples What is the pH of a sol’n if [H 3 O + ] = 3.4 x 10 -5 M?

30. Examples The pH of a sol’n is measured with a pH meter and determined to be 9.00. What is the [H 3 O + ]?

31. Examples The pH o f a sol’n is measured with a pH meter and determined to be 7.52. What is [H 3 O + ]?

32. Calculating K a  In these problems, remember that the concentration of the [H 3 O + ] ions will equal the concentration of the conjugate base ions. –This is because for every molecule of weak acid that dissociates, there will be an equal number of H 3 O + ions and base ions

33. Example Assume that enough lactic acid is dissolved in sour milk to give a solution concentration of 0.100 M lactic acid. A pH meter shows that the pH of the sour milk is 2.43. Calculate K a for the lactic acid equilibrium system.

34. Titrations  An analytical procedure used to determine the concentration of a sample by reacting it with a standard sol’n  In a titration, an indicator is used to determine the end point  Standard sol’n: a sol’n of precisely known concentration  Indicator: any substance in sol’n that changes color as it reacts with either an acid or a base

35. Titrations  Each indicator changes its color over a particular range of pH values (transition interval)  An unknown acid sol’n will be titrated with a standard sol’n that is a strong base  An unknown base sol’n will be titrated with a standard sol’n that is a strong acid

36. Titrations  Equivalence point: point at which the concentration of H 3 O + ions is the same as the concentration of OH - ions; [H 3 O + ] = [OH - ]  Endpoint: the point at which the indicator changes color  Titration curve: graph that shows how pH changes in a titration

37. Titrations  The equivalence point is at the center of the steep, vertical region of the titration curve  At the equivalence point, pH increases greatly w/ only a few drops

39. Example Problem 1 What is the molarity of a CsOH solution if 20.0 mL of the solution is neutralized by 26.4 mL of 0.250M HBr solution? HBr + CsOH → H 2 O + CsBr

40. Example Problem 2 What is the molarity of a nitric acid solution if 43.33 mL 0.200M KOH solution is needed to neutralize 20.00 mL of unknown solution?

41. Example Problem 3 What is the concentration of a household ammonia cleaning solution (NH 4 OH) if 49.90 mL of 0.5900M H 2 SO 4 is required to neutralize 25.00 mL solution?