1. Acids and Bases Dr. Ron Rusay Summer 2004 © Copyright 2004 R.J. Rusay

2. Introduction to Aqueous Acids  Acids: taste sour and cause certain dyes to change color.

3. Introduction to Aqueous Bases  Bases: taste bitter, feel soapy and cause certain dyes to turn color.

4. Electrolytes  Aqueous solutions can be categorized into 3 types: non-electrolytes, strong electrolytes or weak electrolytes based on their ability to conduct electricity.  A solution must have ions to conduct.  Pure Water does not conduct.  Aqueous solutions can be tested for conductivity which will determine the degree of ionization of the solutes.  It is possible to have full or partial ionization. © Copyright 1995-2004 R.J. Rusay

5. Solution Test Apparatus for Electrolytes

6. Conductivity

7. Electrolytes / Ionization

8. Electrolytes  Almost all ionic compounds and a few molecular compounds are strong electrolytes.  Several molecular compounds are weak conductors, most are non-conductors.  Conductivity is directly related to the amount of ionization, i.e. ions in solution. Table salt, sodium chloride, is completely ionized: NaCl (s) + H 2 O (l) ---> NaCl (aq) ---> 0.10molNa + (aq) + Cl - (aq) © Copyright 1995-2004 R.J. Rusay 0.00mol 0.10mol 0.10mol

9. Strong vs. Weak Electrolytes

10. Models of Acids and Bases  Arrhenius Concept: Acids produce H + in solution, bases produce OH  ion.  Brønsted-Lowry: Acids are H + donors, bases are proton acceptors. HCl + H 2 O  Cl  + H 3 O + acid base acid base

11. Lewis Acids and Bases  Lewis Acid: electron pair acceptor  Lewis Base: electron pair donor

12. Conjugate Acid/Base Pairs HA(aq) + H 2 O(l)  H 3 O + (aq) + A  (aq) conj conj conj conj acid 1 base 1 acid 2 base 2  conjugate base: everything that remains of the acid molecule after a proton is lost.  conjugate acid: formed when the proton is transferred to the base.

13. Acid Strength  100% of the acid is ionized. For example nitric acid, HNO 3. Other common strong acids are sulfuric and hydrochloric.  100% of the acid is ionized. For example nitric acid, HNO 3, produces 100% H + (aq). Other common strong acids are sulfuric and hydrochloric.  Stong acids produce very weak conjugate bases, eg. (NO 3  ) Strong Acids:

14. Structure and Acid Strength  Bond polarity & bond strength affects acidity. In binary compounds:  Bond Polarity (The higher the bond polarity, the stronger the bond, the weaker the acid) eg. HF  Bond Strength (The lower the bond strength, the higher the resulting H + ionization and the stronger the acid. ) eg. HCl

15. Oxides  Acidic Oxides (Acid Anhydrides):  O  X bond is strong and covalent. SO 2, NO 2, CrO 3  Basic Oxides (Basic Anhydrides):  O  X bond is ionic. K 2 O, CaO

16. Dissociation of Strong and Weak Acids

17. Acid Strength (continued)  A weak acid is not 100% ionized. For example acetic acid, CH 3 COOH,. Most acids, particularly organic acids, are weak acids.  A weak acid is not 100% ionized. For example acetic acid, CH 3 COOH, produces <100% H + (aq). Most acids, particularly organic acids, are weak acids.  Weak acids produce a much stronger conjugate base than water, eg. The acetate ion: (CH 3 COO  ) Weak Acids :

19. Multiprotic Acids  Monoprotic acids have 1 acidic H, diprotic have 2, eg. Sulfuric acid H 2 SO 4 etc.  In a strong multiprotic acid, like H 2 SO 4, only the first H is strong; transferring the second H is usually weak H 2 SO 4 + H 2 O  H 3 O +1 + HSO 4 -1 HSO 4 -1 + H 2 O  H 3 O +1 + SO 4 -2

20. Aqueous Bases  Any compound that accepts a proton is a base.  The common bases are group IA & IIA metal hydroxide compounds. “Strong” and “weak” are used in the same sense for bases as for acids.  Strong = complete dissociation (100% hydroxide ion is supplied to the solution) An example of a weak base is ammonia. NH 3 (g) + H 2 O (l) NH 3 (aq) NH 4 + (aq) + OH - (aq) © Copyright 1995-2004 R.J. Rusay

21. Bases (continued)  Weak bases have very little dissociation (or reaction with water), eg. methyl amine like ammonia has <100% hydroxide ion in aqueous solution.  H 3 CNH 2 (aq) + H 2 O(l)  H 3 CNH 3 + (aq) + OH  (aq)  Organic bases are weak bases; for example, dopamine (neurotransmitter), cadaverine (product of cellular decomposition), morphine (narcotic pain killer) and cocaine are weak bases.

22. Natural Indicators

24. Reactions of Acids & Bases A use for natural indicators: “Neutralization Reactions” Titrations Examples

25. Water as an Acid and a Base  Amphoteric substances can act as either an acid or a base Water acting as an acid:Water acting as an acid: NH 3 + H 2 O  NH 4 +1 + OH -1 NH 3 + H 2 O  NH 4 +1 + OH -1 Water acting as a base:Water acting as a base: HCl + H 2 O  H 3 O +1 + Cl -1 HCl + H 2 O  H 3 O +1 + Cl -1 Water reacting with itself as both:Water reacting with itself as both: H 2 O + H 2 O  H 3 O +1 + OH -1

26. Water as an Acid and a Base  Water is amphoteric. It can behave either as an acid or a base. H 2 O (l) + H 2 O (l)  H 3 O + (aq) + OH  (aq) conj conj conj conj acid 1 base 1 acid 2 base 2 acid 1 base 1 acid 2 base 2  The equilibrium expression for pure water is: K w = [H 3 O + (aq) ] [OH  (aq) ]

27. Water: Self-ionization

28. Autoionization of Water  Water is an extremely weak electrolyte therefore are only a few ions present: K w = [H 3 O +1 ] [OH -1 ] = 1 x 10 -14 @ 25°C NOTE: the concentration of H 3 O +1 and OH -1 are equalNOTE: the concentration of H 3 O +1 and OH -1 are equal [H 3 O +1 ] = [OH -1 ] = 10 -7 M @ 25°C[H 3 O +1 ] = [OH -1 ] = 10 -7 M @ 25°C K w is called the ion product constant for water: as [H 3 O +1 ] increases, [OH - ] decreases and vice versa.K w is called the ion product constant for water: as [H 3 O +1 ] increases, [OH - ] decreases and vice versa.

29. Acidic and Basic Solutions  Acidic solutions have: a larger [H +1 ] than [OH -1 ]  Basic solutions have: a larger [OH -1 ] than [H +1 ]  Neutral solutions have [H +1 ] = [OH -1 ] = 1 x 10 -7 M [H +1 ] = 1 x 10 -14 [OH -1 ] [OH -1 ] = 1 x 10 -14 [H +1 ]

30. The pH Scale  pH   log[H + ]  log[H 3 O + ]  pH in water ranges from 0 to 14. K w = 1.00  10  14 = [H + ] [OH  ] pK w = 14.00 = pH + pOH  As pH rises, pOH falls (sum = 14.00).

31. pH & pOH  pH = -log[H 3 O +1 ]pOH = -log[OH -1 ] pH water = -log[10 -7 ] = 7 = pOH waterpH water = -log[10 -7 ] = 7 = pOH water  [H +1 ] = 10 -pH [OH -1 ] = 10 -pOH  pH 7 is basic, pH = 7 is neutral  The lower the pH, the more acidic the solution; The higher the pH, the more basic the solution  1 pH unit corresponds to a factor of 10  pOH = 14 - pH

32. There are no theoretical limits on the values of pH or pOH. (e.g. pH of 2.0 M HCl is -0.301, the pH at Iron Mountain, California is ~ -2 to -3)

33. What’s in these household products? Acids or bases? Strong or weak? Should you be concerned about safety?

34. The pH of Some Familiar Aqueous Solutions [H 3 O + ] [OH - ] [OH - ] = KWKW [H 3 O + ] neutral solution acidic solution basic solution [H 3 O + ]> [OH - ] [H 3 O + ]< [OH - ] [H 3 O + ] = [OH - ] What’s your diet? Your urine will tell!

35. Example #1  Determine the given information and the information you need to find Given [H +1 ] = 10.0 MFind [OH -1 ]  Solve the Equation for the Unknown Amount Determine the [H +1 ] and [OH -1 ] in a 10.0 M H +1 solution

36.  Convert all the information to Scientific Notation and Plug the given information into the equation. Given [H +1 ] = 10.0 M= 1.00 x 10 1 M K w = 1.0 x 10 -14 Example #1 (continued) Determine the [H +1 ] and [OH -1 ] in a 10.0 M H +1 solution

37. Example #2  Find the concentration of [H +1 ] Calculate the pH of a solution with a [OH -1 ] = 1.0 x 10 -6 M

38.  Enter the [H +1 ] concentration into your calculator and press the log key log(1.0 x 10 -8 ) = -8.0  Change the sign to get the pH pH = -(-8.0) = 8.0 Example #2 continued Calculate the pH of a solution with a [OH -1 ] = 1.0 x 10 -6 M

39.  Enter the [H +1 ] or [OH -1 ]concentration into your calculator and press the log key log(1.0 x 10 -3 ) = -3.0  Change the sign to get the pH or pH pOH = -(-3) = 3.0  Subtract the calculated pH or pOH from 14.00 to get the other value pH = 14.00 – 3.0 = 11.0 Calculate the pH and pOH of a solution with a [OH -1 ] = 1.0 x 10 -3 M Example #3

40.  If you want to calculate [OH -1 ] use pOH, if you want [H +1 ] use pH. It may be necessary to convert one to the other using 14 = [H +1 ] + [OH -1 ] pOH = 14.00 – 7.41 = 6.59  Enter the pH or pOH concentration into your calculator  Change the sign of the pH or pOH -pOH = -(6.59)  Press the button(s) on you calculator to take the inverse log or 10 x [OH -1 ] = 10 -6.59 = 2.6 x 10 -7 Example #4 Calculate the [OH -1 ] of a solution with a pH of 7.41

41. Calculating the pH of a Strong, Monoprotic Acid  A strong acid will dissociate 100% HA  H +1 + A -1  Therefore the molarity of H +1 ions will be the same as the molarity of the acid  Once the H +1 molarity is determined, the pH can be determined pH = -log[H +1 ]

42. Example #5  Determine the [H +1 ] from the acid concentration HNO 3  H +1 + NO 3 -1 0.10 M HNO 3 = 0.10 M H +1  Enter the [H +1 ] concentration into your calculator and press the log key log(0.10) = -1.00  Change the sign to get the pH pH = -(-1.00) = 1.00 Calculate the pH of a 0.10 M HNO 3 solution

43. The pH Scale What is the pH of 6M hydrochloric acid?

44. Neutralization Reactions How would indicator be used?

45. Aqueous Reactions: Neutralization Net Ionic Equations HCl (aq) + NaOH (aq) ---> NaCl (aq) + H 2 O (l) ___________________________________________________   HCl (aq) ---> H + (aq) + Cl - (aq)   NaOH (aq) ---> Na + (aq) + OH - (aq)   NaCl (aq) ---> Na + (aq) + Cl - (aq) ________________________________________________ Na + (aq) + OH - (aq) + H + (aq) + Cl - (aq) ---> Na + (aq) + Cl - (aq) + H 2 O (l) _______________________________________________________ © Copyright 1995-2000 R.J. Rusay H + (aq) + OH - (aq) ---> H 2 O (l)

46. Acid-Base Titration Acids have pH 7, and 7 is neutralAcids have pH 7, and 7 is neutral Without a pH meter how can the progress of reaction be monitored?Without a pH meter how can the progress of reaction be monitored?

47. Acid-Base Titration http://www.dartmouth.edu/~chemlab/techniques/titration.html http://chemistry.fullerton.edu/~chemdev/director/titrate.html pH & Water

49. Stomach Chemistry

50. Buffered Solutions  Buffered Solutions resist change in pH when an acid or base is added to it.  Used when need to maintain a certain pH in the system, eg. Blood.  A buffer solution contains a weak acid and its conjugate base  Buffers work by reacting with added H +1 or OH -1 ions so they do not accumulate and change the pH.  Buffers will only work as long as there is sufficient weak acid and conjugate base molecules present.

52. Buffers