1. Acids and Bases

2. Properties of Acids/Bases  Acids are substances which…  Bases are substances which…

3. Properties of Acids and Bases  Acids:  Sour taste  React with active metals to produce hydrogen gas  Neutralize bases  Change the color of indicators  Conduct electricity  Bases:  Bitter taste  Slippery feel  Neutralize acids  Change the color of indicators  Conduct electricity

4. Arrhenius Acid/Base Theory  Acids are substances that produce H + (H 3 O + ) ions in solution.  HCl  H + + Cl -  Bases produce OH - ions in solution.  NaOH  Na + + OH -

5. Acid Nomenclature  The presence of O in an acid molecule determines how the acid is named.  Acids without O use the prefix “hydro” and the suffix “ic.” HCl is hydrochloric acid.  Acids with O use a suffix ending of “ous” or “ic” depending on the amount of O present.  H 2 SO 4 = sulfuric acid; H 2 SO 3 = sulfurous acid.

6. Name the following acids: HBr H 2 CO 3 H 3 PO 4 H 2 Se HCN H 2 C 2 O 4

7. Strong Acids and Strong Bases  Strong acids ionize completely when placed in water.  There are six strong acids, all the rest are weak acids.  HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 4  Strong bases dissociate completely when placed in water.  Strong bases include all elements in group I and heavy group 2 elements.

8. Brønsted-Lowry Acids and Bases  Acids are substances capable of donating a proton.  Bases are capable of accepting a proton.  B/L theory is applicable to reactions that do not occur in water.  Can include gas phase reactions  NH 3 (g) + H 2 O(l)  NH 4 + (aq)+ OH - (aq)

9. HBr(aq) + H 2 O(l)  H 3 O + (aq) + Br - (aq) CO 3 2- (aq) + H 2 O(l)  HCO 3 - (aq) + OH - (aq) HSO 4 - (aq) + HCO 3 - (aq)  SO 4 2- (aq) + H 2 CO 3 (aq) HCHO 2 (aq)+ PO 4 3- (aq)  CHO 2 - (aq) + HPO 4 - (aq)

10. Identify the conjugate acid-base pairs in the following reactions: H 2 SO 3 (aq)+H 2 O(l)  HSO 3 - (aq) + H 3 O + (aq) HPO 4 2- (aq)+H 2 O(l)  H 2 PO 4 - (aq) + OH - (aq) HSeO 3 - (aq)+H 2 O(l)  H 3 O + (aq) + SeO 3 2- (aq)

12. Lewis Acids and Bases  According to Lewis, an acid is an electron pair acceptor.  Acids generally have incomplete octets.  A base is an electron pair donor.  Bases will have lone pairs.  NH 3 + BF 3  H 3 NBF 3

13. Strengths of Acids and Bases  Strong acids ionize completely.  HCl  H + + Cl -  Strong acids produce the maximum number of ions in solution so they are good conductors of electricity.  Strong bases dissociate completely.  NaOH  Na + + OH -  Strong bases produce the maximum number of hydroxide ions in solution.

14. Weak acids produce an equilibrium mixture of molecules and ions in solution. Because an equilibrium is established, an equilibrium expression or K a can be written for a weak acid. HCOOH  H + + HCOO - K a = [H + ][HCOO - ] [HCOOH]

15. Determine the equilibrium constant for a 0.15 M solution of acetic acid if the H + concentration is 1.45 x 10 -6 M.

16. The K a for hydrofluoric acid is 6.8 x 10 -4. Determine the [H + ] of a 0.500 M solution.

17. Table of Acid and Base Dissociation Constants Weak Acids NameChemical FormulaKAKA Acetic acidHC 2 H 3 O 2 1.80  10 -5 Acetylsalicylic acidHC 9 H 7 O 4 3.2  10 -4 Benzoic acidHC 7 H 5 O 2 6.0  10 -5 Chloroacetic acidClCH 2 COOH 1.40  10 -3 Formic acidHCOOH 1.77  10 -4 Hydrocyanic acidHCN 4.00  10 -10 Hypochlorous acidHOCl 4.00  10 -8 Hydrofluoric acidHF 7.20  10 -4 Hypoiodous acidHOI 2.30  10 -11 Lactic AcidHC 3 H 5 O 3 1.5×10 -4 Nicotinic acidHC 6 H 4 NO 2 1.40  10 -5 Nitrous AcidHNO 2 4.5  10 -4 Weak Bases NameChemical FormulaKBKB AmmoniaNH 3 1.80  10 -5 Diethylamine(CH 3 CH 2 ) 2 NH 8.6  10 -4 MethylamineCH 3 NH 2 4.2  10 -4 PyridineC5H5NC5H5N1.5×10 -9 QuinolineC9H7NC9H7N 6.3  10 -8 Trimethylamine(CH 3 ) 3 N7.4×10 -5 Carbonic = 4.3 x 10 -7 Sulfurous = 1.7 x 10 -2 Phosphoric = 7.5 x 10 -3 Phosphorous = 1.58 x 10 -2 Hydrosulfuric = 9.5 x 10 -8

18. Autoionization of Water  Protons transfer from one water molecule to another.  H 2 O(l)  H + (aq) + OH - (aq)  K w = [H + ] [OH - ]  K w = 1.0 x 10 -14  The product of [H + ] [OH - ] must always equal 1.0 x 10 -14

19. Determine the [OH - ] if the [H + ] = 2.5 x 10 -5 M. Calculate the [H + ] if the [OH - ] is 6.5 x 10 -8 M.

21. The K a for phosphoric acid is 7.08 x 10 -3. Write the equilibrium expression for phosphoric acid and determine the hydrogen ion concentration of a 0.145 M solution.

22. pH  The amount of H + ions in solution can be expressed quantitatively using a logarithmic scale.  pH = -log[H + ]; pOH = -log[OH - ]  At 298 K pH values range from 0-14.  pH + pOH = 14  A solution is considered acidic if its pH <7.  A solution is considered basic if its pH >7.  At a pH of 7 [H + ] = [OH - ] and the solution is considered to be neutral.

23. [H + ] [OH - ] pH pOH 6.5 x 10 -3 7.8 x 10 -8 10.56 Complete the following table:

24. Determine the pH of the following solutions: 1.0 M HI 0.050 M HNO 3 1.5 M KOH 0.0075 M Ba(OH) 2 2.5 M HCl 0.0045 M NaOH

25. pH + pOH =14 [H + ] [OH - ] = 1.0 x 10 -14

26. Determine the K a of a 0.10 M solution of formic acid with a pH of 2.37.

27. Determine the pH of a 0.50 M solution of acetic acid. The K a for acetic acid is 1.8 x 10 -5.

28. Neutralization Reactions  Occur between acids and bases.  Produce a salt and water.  The pH of the resulting salt depends on the nature of the acid and the base.  HNO 3 + CsOH  H 2 O + CsNO 3

29. Determine the pH of a 0.450 M solution of sulfurous acid. The K a of sulfurous acid is 1.7 x 10 -2.

30. Write equations for the following reactions. Predict the pH of the salt formed in each case. Hydrobromic acid + calcium hydroxide  Sulfuric acid + potassium hydroxide  Acetic acid + ammonium hydroxide 

31. Titration  Method for determining the concentration of a solution by neutralization.  A solution of known concentration is reacted with a set volume of another solution of unknown molarity.  Titration proceeds until the equivalence point is reached.  The equivalence point is where moles of H + ions = moles of OH - ions.

33. Titration Procedure: 1. A measured volume of acid or base is placed in a beaker. 2. A buret is filled with a solution of known concentration. 3. A pH meter or indicator is used to signal the endpoint of the procedure.

35. An indicator or pH meter is used to signal the end of the titration.

36. Titration problems can be solved by relating moles of acid to moles of base using the formula for Molarity: Molarity = moles liter of solution At the equivalence point: moles of acid =moles of base or M A V A = M B V B

37. What is the molarity of a solution of KOH if 30.0 mL of the solution is titrated with 26.5 mL of 0.250 M HBr solution?

38. What volume of 0.590 M HCl is needed to neutralize 55.00 mL of 0.125 M NaOH?

39. Determine the concentration of 35.0 mL of sulfuric acid solution that is titrated with 45.0 mL of 0.150 M KOH.