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# Acidandbase chm141 thursday[1]goodday

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### Acidandbase chm141 thursday[1]goodday

1. 1. Lewis acid Lewis Basesby: Dr. Robert D.Craig, Ph.D.
2. 2. Will finish chapter 3…..
3. 3. Acids bases and buffer solutions
4. 4. Strong Acids and Bases• Give the names and formulas of some strong acids and bases.• Explain the pH scale, and convert pH to concentration (will do later)• Evaluate solution pH and pOH of strong acids or bases.
5. 5. In a glass of water
6. 6. Autoionization of Water• Autoionization of WaterThe equilibrium product Kw = [H+] [OH-]is a constant at a definite temperature due to the autoionization of water, H2O = H+ + OH-.• At 298 K, Kw = 10-14 and the following relationship in any aqueous solution is obvious,• ***pOH + pH = 14 at 298 K.
7. 7. Will need this soon• ***pOH + pH = 14 at 298 K.pH = -log[H+]
8. 8. We say• The pH scale is defined as the negative log of the concentration of H+: pH = -log[H+]• The pOH scale is defined as the negative log of the concentration of OH-, [OH-]:• pOH = -log[OH-] With this scale, calculating the pOH can be done in the same manner as the pH scale.
9. 9. .
10. 10. Adapted from• http://www.science.uwaterloo.ca/~cchieh/cact/c 123/stacids.html• http://www.chem1.com/acad/webtext/chembon d/cb03.html• http://www.epa.gov/acidrain/education/site_stu dents/phscale.html
11. 11. Arrhenius Acids and Bases• Arrhenius Acids and Bases• The Arrhenius definition of acids and bases is one of the oldest. An Arrhenius acid is a substance that when added to water increases the concentration of H1+ ions present.
12. 12. Arrhenius Acids and Bases• The chemical formulas of Arrhenius acids are written with the acidic hydrogens first. An Arrhenius base is a substance that when added to water increases the concentration of OH1- ions present. HCl is an example of an Arrhenius acid and NaOH is an example of an Arrhenius base.
13. 13. Arrhenius Acids and Bases• HCl is an example of an Arrhenius acid and NaOH is an example of an Arrhenius base.••
14. 14. Arrhenius Acids and Bases• The H1+ ion produced by an Arrhenius acid is always associated with a water molecule to form the hydronium ion, H3O1+(aq).
15. 15. Arrhenius acids• Arrhenius acids are frequently referred to as proton donors, hydrogen ion donors, or hydronium ion donors,
16. 16. Arrhenius Acids and Bases• To represent the transfer of the H1+ ion to water to form the hydronium ion, we must include H2O in the chemical equation for acid ionization.
17. 17. Arrhenius Acids and Bases
18. 18. Brønsted–Lowry concept• It follows that, if a compound is to behave as an acid, donating a proton, there must be a base to accept the proton. So the Brønsted–Lowry concept can be defined by the reaction:• acid + base <-> conjugate base + conjugate acid.
19. 19. Brønsted-Lowry Style• Sample Equations written in the Brønsted- Lowry Style• A. Reactions that proceed to a large extent:• HCl + H2O <==> H3O+ + Cl¯• HCl - this is an acid, because it has a proton available to be transfered.• H2O - this is a base, since it gets the proton that the acid lost.
20. 20. Brønsted–Lowry concept• The conjugate base is the ion or molecule remaining after the acid has lost a proton, and the conjugate acid is the species created when the base accepts the proton.
21. 21. Now, here comes an interesting idea:• H3O+ - this is an acid, because it can give a proton.• Cl¯ - this is a base, since it has the capacity to receive a proton.• Notice that each pair (HCl and Cl¯ as well as H2O and H3O+ differ by one proton (symbol = H+). These pairs are called conjugate pairs.
22. 22. .• The reaction can proceed in either forward or backward direction; in each case, the acid donates a proton to the base.
23. 23. Acid base Pairs2 examples:CH3COOH + H2O <->CH3COO− + H3O+ H2O + NH3 <-> OH− + NH4+
24. 24. An example• Which of the following is usually referred to as strong acid in water solution?• HF, HNO2, H2CO3, H2S, HSO4-, Cl-, HNO3, HCN• Answer HNO3 All others are weak acids
25. 25. Water is amphoteric• Water is amphoteric and can act as an acid or as a base. In the reaction between acetic acid, CH3CO2H, and water, H2O, water acts as a base.CH3COOH + H2O <->CH3COO− + H3O+
26. 26. conjugate base of acetic acid• The acetate ion, CH3CO2-, is the conjugate base of acetic acid and the hydronium ion, H3O+, is the conjugate acid of the base, water
27. 27. act as an acidWater can also act as an acid, for instance when it reacts with ammonia.The equation given for this reaction is:H2O + NH3 <-> OH− + NH4 +• in which H2O donates a proton to NH3
28. 28. reaction as a base!!!H2O + NH3 <-> OH− + NH4+
29. 29. Strong acid weak acid• A strong acid, such as hydrochloric acid, dissociates completely.• A weak acid, such as acetic acid, may be partially dissociated; the acid dissociation constant, pKa, is a quantitative measure of the strength of the acid.
30. 30. the acid dissociation constant, pKa• A weak acid, such as acetic acid, may be partially dissociated; the acid dissociation constant, pKa, is a quantitative measure of the strength of the acid.
31. 31. Brønsted–Lowry framework• A wide range of compounds can be classified in the Brønsted–Lowry framework: mineral acids and derivatives such as sulfonates, phosphonates, etc., carboxylic acids, amines, carbon acids, and many more
32. 32. Solvent –Not Water????• Brønsted–Lowry base as the pair of electrons can be donated to a proton.• This means that the Brønsted–Lowry concept is not limited to aqueous solutions.• Any donor solvent S can act as a proton acceptor.•AH + S: <-> A − + SH +
33. 33. donor solvents• Typical donor solvents used in acid-base chemistry, such as dimethyl sulfoxide or liquid ammonia have an oxygen or nitrogen atom with a lone pair of electrons that can be used to form a bond with a proton.
34. 34. Do it here Rob• 87. Most naturally occurring acids are weak acids. Lactic acid is one example.• CH3COOH + H2O <->CH3COO− + H3O+
35. 35. Buffer solutions• If you place some lactic acid in water, it will ionize to a small extent, and an equilibrium will be established.• Suggest an experiment???
36. 36. Who was G.N lewis???
37. 37. G.N. Lewis (1875-1946)• G.N. Lewis (1875-1946) created the College of Chemistry at the University of California, Berkeley, and made it into one of the world’s most productive centers of chemistry research.• His other notable work included acid-base theory, the thermodynamics of solutions, the first isolation of heavy water (D2O), and the phosphorescence and magnetic properties of molecules.
38. 38. Lewis definition
39. 39. Heavy water –nuclear reactors
40. 40. .• At the time Lewis began developing his ideas in 1902, it was widely believed that chemical bonding involved electrostatic attraction between ion-like entities.
41. 41. Could not explain• This seemed satisfactory for compounds such as NaCl that were known to dissociate into ions when dissolved in water, but it failed to explain the bonding in non-electrolytes such as CH4
42. 42. .• .
43. 43. The ammonium ion-an acid!!!• The ammonium ion is mildly acidic, reacting with Brønsted bases to return to the uncharged ammonia molecule:• NH4 + + :B- <→ HB + NH 3• Thus, treatment of concentrated solutions of ammonium salts with strong base gives ammonia.
44. 44. .• When ammonia is dissolved in water, a tiny amount of it converts to ammonium ions• H3O+ + NH3 <-> H2O + NH4+
45. 45. depends on the pH• The degree to which ammonia forms the ammonium ion depends on the pH of the solution. If the pH is low, the equilibrium shifts to the right: more ammonia molecules are converted into ammonium ions.
46. 46. Just look –last time
47. 47. Strong Acids and Bases• Acids and bases that are completely ionized when dissolved in water are called strong acids and strong bases There are only a few strong acids and bases, and everyone should know their names and properties. These acids are often used in industry and everyday life
48. 48. concentrations of acids and bases• The concentrations of acids and bases are often expressed in terms of pH, and as an educated person, you should have the skill to convert concentrations into pH and pOH. The pH is an indication of the hydrogen ion concentration, [H+].
49. 49. The term Lewis acid• The term Lewis acid refers to a definition of acid published by Gilbert N. Lewis in 1923, specifically: An acid substance is one which can employ a lone pair from another molecule in completing the stable group of one of its own atoms.[1] Thus, H+ is a Lewis acid, since it can accept a lone pair, completing its stable form, which requires two electrons
50. 50. Just look – do not write . . .
51. 51. A Lewis base,• A Lewis base, then, is any species that donates a pair electrons to a Lewis acid to form a Lewis adduct. For example, OH− and NH3 are Lewis bases, because they can donate a lone pair of electrons
52. 52. Please just look !!!!!!!
53. 53. ammonium ion• The ammonium ion is mildly acidic, reacting with Brønsted bases to return to the uncharged ammonia molecule:• H3O+ + NH3 <-> H2O + NH4
54. 54. NH4+ + :B- → HB + NH3• Thus, treatment of concentrated solutions of ammonium salts with strong base gives ammonia. When ammonia is dissolved in water, a tiny amount of it converts to ammonium ions: (a buffer)
55. 55. How do buffer solutions work?• A buffer solution has to contain things which will remove any hydrogen ions or hydroxide ions that you might add to it - otherwise the pH will change. Acidic and alkaline buffer solutions achieve this in different ways.
56. 56. .• The degree to which ammonia forms the ammonium ion depends on the pH of the solution. If the pH is low, the equilibrium shifts to the right: more ammonia molecules are converted into ammonium ions. H3O+ + NH3 <-> H2O + NH4+ NH4 + + :B- → HB + NH 3
57. 57. Alpha curve
58. 58. Buffer solutions• Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A-.• HA <-> H+ + A-• When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in accordance with Le Chateliers principle
59. 59. These equations are complex
60. 60. Haber process• Example of l
61. 61. Le Chateliers Principle• In 1884 the French chemist and engineer Henry-Louis Le Chatelier proposed one of the central concepts of chemical equilibria. Le Chateliers principle can be stated as follows: A change in one of the variables that describe a system at equilibrium produces a shift in the position of the equilibrium that counteracts the effect of this change.
62. 62. Le Chateliers principle• Le Chateliers principle describes what happens to a system when something momentarily takes it away from equilibrium.
63. 63. • (1) changing the concentration of one of the components of the reaction• (2) changing the pressure on the system• (3) changing the temperature at which the reaction is run