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Acid and Base Equilibria

Acid and Base Equilibria. The concept of acidic and basic solutions is perhaps one of the most important topics in chemistry. Acids and bases affect the properties of foods, biochemical reactions, pharmaceuticals, and industrial materials. Acid and Base Equilibria. Properties of Acids

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Acid and Base Equilibria

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  1. Acid and Base Equilibria The concept of acidic and basic solutions is perhaps one of the most important topics in chemistry. Acids and bases affect the properties of foods, biochemical reactions, pharmaceuticals, and industrial materials.

  2. Acid and Base Equilibria Properties of Acids Sour or tart taste. Corrosive (deteriorate). Electrolytes Electrolytes are able to conduct an electrical current because of the presence of ions in aqueous solutions.

  3. Acid and Base Equilibria Properties of Acids (cont.) Will react with most metals to form hydrogen gas. Some acids are ‘stronger’ than others All acids contain a hydrogen that they can give away. HA(aq) + H2O(l)  H3O+(aq) + A-(aq)

  4. Acid and Base Equilibria Properties of Bases Bitter tasting Slippery Caustic – They will degrade biological tissue. Chemical burns from strong bases are nasty. Bases form the hydroxide ion (OH-) in water.

  5. Acid and Base Equilibria Arrhenius Acids Svante Arrhenius (1900) defined an acid. Acids are hydrogen containing compounds that yield a hydrogen ion (H+) in water. An Arrhenius acid donate an H+ ion. HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)

  6. Acid and Base Equilibria Arrhenius Acids HNO3(aq) + H2O(l)  H3O+(aq) + NO3-(aq) H2SO4(aq) + H2O(l)  H2O(l)+ H2O(l) 

  7. Acid and Base Equilibria Arrhenius Acids Acids that have one hydrogen ion to donate are called monoprotic. Acids that have 2 hydrogen ions to donate are called diprotic. Acids that have 3 hydrogen ions to donate are called ____protic.

  8. Acid and Base Equilibria Arrhenius Bases Compounds that produce the hydroxide ion(s) (OH-) in water are called Arrhenius Bases. NaOH(aq)

  9. Acid and Base Equilibria Arrhenius Acids and Bases Problem with the Arrhenius definition; Some bases can form OH- ions in solution but not have an OH- ion in the chemical formula. NH3(g) + H2O(l) NH4+(aq) + OH-(aq)

  10. Acid and Base Equilibria Bronsted-Lowry Acids and Bases A better definition of an acid and a base Bronsted-Lowry Acid – A molecule that donates an H+to another molecule. Bronsted-Lowry Base – A molecule that accepts an H+ from the B-L acid.

  11. Acid and Base Equilibria Bronsted-Lowry Acids and Bases Identify the Bronsted-Lowry acid and base; NH3(g) + H2O(l) NH4+(aq) + OH-(aq)

  12. Acid and Base Equilibria Bronsted-Lowry Acids and Bases Conjugate Acid – becomes the H+ donor in the reverse reaction. Conjugate Base – becomes the H+ acceptor in the reverse reaction. Identify the BL acid, base, conjugate acid and conjugate base; HClO2(aq) + H2O(l)  H3O+1(aq) + ClO2-(aq)

  13. Acid and Base Equilibria The Ion-Product of Water H2O(l)  H3O+(aq) + OH-(aq) Keq = Kw = Kw = 1.0 x 10-14

  14. Acid and Base Equilibria The Ion-Product of Water H2O(l)  H3O+(aq) + OH-(aq) Keq = Kw = [H3O+] [OH-] = 1.0 x 10-14 What would be the concentration of [H3O+] [OH-] in pure water?

  15. Acid and Base Equilibria What would be the concentration of [OH-] in a solution where[H3O+] = 0.025 M?

  16. Acid and Base Equilibria The pH Scale – Measures the concentration of H+ ion in an aqueous solution.

  17. Acid and Base Equilibria The pH Scale Remember that a water molecules ionizes; H2O(l) H+(aq) + OH-(aq) In pure water, the concentration of H+ and OH- each is 1.0 x 10-7M

  18. Acid and Base Equilibria The pH Scale H2O(l) H+(aq) + OH-(aq) Therefore, the product of [H+] and [OH-] must be equal to 1.0 x 10-14M2. An aqueous solution will always have the concentration of H+ and OH- equal to 1.0x10-14M2.

  19. Acid and Base Equilibria The pH Scale H2O(l) H+(aq) + OH-(aq) Keq = [H+] x [OH-] Kw = [1.0x10-7] x [1.0x10-7] = 1.0x10-14 Kw is called the autoionization constant for water.

  20. Acid and Base Equilibria The pH Scale Keq = [H+] x [OH-] [1.0x10-7] x [1.0x10-7] = 1.0x10-14 If the addition of an acid makes the [H+] increase, then the [OH-] will decrease. If the addition of a base makes the [OH-] increase, the [H+] will decrease.

  21. Acid and Base Equilibria The pH Scale Calculate the [H+] if 0.05 moles of HCl is added to 1.0 L of water.

  22. Acid and Base Equilibria The pH Scale pH is a measure of the concentration of H+. pH = -log[H+]

  23. Acid and Base Equilibria The pH Scale What is the pH of an aqueous solution where [H+] = 1.0 x 10-7? What is the pH of an aqueous solution where [H+] = 1.0 x 10-2?

  24. Acid and Base Equilibria The pH Scale What is the pH of an aqueous solution where [H+] = 2.7 x 10-1? What is the pH of an aqueous solution where [H+] = 8.0 x 10-12?

  25. Acid and Base Equilibria The pH Scale Since Kw = 1 x 10-14 = [H+] x [OH-] 14 = pH + pOH

  26. Acid and Base Equilibria The pH Scale Calculate the pOH of an aqueous solution that has an [H+] = 1.0 x 109.

  27. Acid and Base Equilibria The pH Scale Calculate the [H+] of an acid with a pH of 4. Calculate the [H+] of a base with a pH of 10.8.

  28. Acid and Base Equilibria The pH Scale Calculate the [OH-] of an acid with a pH of 2. Calculate the [OH-] of a base with a pH of 12.9.

  29. Acid and Base Equilibria Strong Versus Weak Acids What makes some acids ‘strong’ and some ‘weak’?

  30. Acid and Base Equilibria Strong Versus Weak Acids We can quantify the relative strength of an acid by using it’s equilibrium expression (Ka). Ka = [products]x= [H+] x [conj. base] [reactants]y [acid]

  31. Acid and Base Equilibria Strong Versus Weak Acids Write the Kaexpression for HCl. Ka =

  32. Acid and Base Equilibria Strong Versus Weak Acids Write the Kaexpression for H3PO4. Ka =

  33. Acid and Base Equilibria Strong Versus Weak Acids As the [H+] increases, Ka increases. Therefore the greater the value of Ka, the more [H+] present, the stronger the acid.

  34. Acid and Base Equilibria Strong Versus Weak Acids Calculate the Ka of a 0.10 M solution of acetic acid that has a pH of 5.0.

  35. Acid and Base Equilibria Salt Hydrolysis Sometimes an ion from a salt can make an aqueous solution acidic or basic. What happens when sodium bicarbonate dissolves in water?

  36. Acid and Base Equilibria Salt Hydrolysis NaHCO3(aq)  Na+(aq) + HCO3-(aq) HCO3-(aq) + H2O(l)  H2CO3(aq) + OH-(aq) Now there will more OH- than H+ in the solution making it basic.

  37. Acid and Base Equilibria Salt Hydrolysis Will a solution of ammonium chloride, NH4Cl, be acidic or basic?

  38. Acid and Base Equilibria Buffers A buffer is a solution that resists changes in pH when either an acid or a base is added. Buffers consist of either a weak acids with one of its salts, or a weak base with one of its salts.

  39. Acid and Base Equilibria Buffers For example, if a solution is make by dissolving carbonic acid (weak acid) and sodium bicarbonate (salt of the acid) we get the following; H2CO3(aq) + H2O(l)  H3O+(aq) + HCO3-1(aq)

  40. Acid and Base Equilibria Buffers H2CO3(aq) + H2O(l)  H3O+(aq) + HCO3-1(aq) If we add a base to this buffered solution, the H3O+ will scoop it up. H3O+(aq) +OH(aq)-  2H2O(l)

  41. Acid and Base Equilibria Buffers H2CO3(aq) + H2O(l)  H3O+(aq) + HCO3-1(aq) If we add an acid to this buffered solution, the HCO3-1 will scoop it up. H+(aq) + HCO3-1 (aq) H2CO3(aq)

  42. Acid and Base Equilibria Buffers Write the chemical reaction for the phosphoric acid – dihydrogen phosphate buffer reaction.

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