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Acids and Bases

Dive into the world of acids and bases using the Bronsted-Lowry model, learn about strong and weak acids, conjugate acids and bases, acid strength, and the structure affecting acid stability. Discover the fascinating chemistry behind these fundamental substances.

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Acids and Bases

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  1. Acids and Bases

  2. Acids & Bases The Bronsted-Lowry model defines an acid as a proton donor. A base is a proton acceptor. Note that this definition is based on the transfer of a proton from the acid to the base.

  3. Acids & Bases H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) In this reaction, water accepts a proton from HCl. Water is a base, and HCl is an acid. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor B-L base B-L acid

  4. Acids & Bases Since this reaction goes to completion (note the one-way arrow), we classify HCl(aq) as a strong acid. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor

  5. Acids & Bases Hydrochloric acid dissociates 100%, and exists as a solution of hydronium and chloride ions. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor Although a bottle may be labeled 1.0M HCl, it really contains 1.0 M H3O+(aq) and 1.0 M Cl1-(aq).

  6. Acids & Bases Hydrochloric acid dissociates 100%, and exists as a solution of hydronium and chloride ions. H2O(l) + HCl(g)  H3O+(aq) + Cl1-(aq) proton proton acceptor donor There is no reverse reaction because chloride has no tendency to accept a proton to form HCl.

  7. Acids & Bases There are only a few common strong acids. They are: HCl(aq), HBr(aq), HI(aq), HNO3(aq), HClO4(aq) and H2SO4(aq)* *for the first proton only

  8. Acids & Bases Most acids are weak acids in which only a small percentage of the molecules dissociate to protonate water. HA(aq) + H2O(aq) ↔ H3O+(aq) + A-(aq) HA is a generic monoprotic weak acid such as HF, HCN or CH3COOH.

  9. Weak Acids Weak acids in water form an equilibrium with hydronium ion and the deprotonated anion of the acid. The reaction does not go to completion. H-A + H2O ↔ H3O++ A-

  10. Acids The equilibrium of acids in water can be viewed as a competition between the forward reaction and the reverse reaction. H-A + H2O ↔ H3O++ A-

  11. Acids The forward reaction involves HA, a generic acid, and water. Water acts as a base by accepting a proton from the acid. H-A + H2O ↔ H3O++ A-

  12. Acids The forward reaction involves HA, a generic acid, and water. Water acts as a base by accepting a proton from the acid. Acid HA proton donor Base H2O proton acceptor H-A + H2O ↔ H3O+ + A-

  13. Acids The reverse reaction involves A- accepting a proton and acting as a base. H3O+ donates a proton, and is an acid. Base A- proton acceptor Acid H3O+ proton donor Acid HA proton donor Base H2O proton acceptor H-A + H2O ↔ H3O+ + A-

  14. Ka Values Acid strength is determined by measuring the equilibrium constant for the following reaction: HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) Ka = [H3O+][A-] [HA]

  15. Ka Values HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) Ka = [H3O+][A-] [HA] Water is left out of the equilibrium constant expression because it is a pure liquid (with constant concentration). The “a” subscript stands for acid.

  16. Conjugate Acids & Bases The deprotonated acid is a base, and the protonated base is an acid. These acids and bases are called conjugate acids and bases. Base A- proton acceptor Acid H3O+ proton donor Acid HA proton donor Base H2O proton acceptor H-A + H2O ↔ H3O+ + A-

  17. Acids For weak acids, the equilibrium lies to the left, indicating that A- is a stronger base than water. Base A- proton acceptor Acid H3O+ proton donor Acid HA proton donor Base H2O proton acceptor

  18. Acids Acid HA and base A- are related, and differ only by the addition or removal of H+. Acid H3O+ Acid HA Base H2O Base A- remove H+ add H+

  19. Acids HA and A- are called conjugate acid-base pairs. A- is the conjugate base of the acid HA. Acid H3O+ Acid HA Base H2O Base A- remove H+ add H+

  20. Acids Likewise, H2O and H3O+ are related, and differ only by the addition or removal of H+. Acid H3O+ Acid HA Base H2O Base A- add H+ remove H+

  21. Acids H2O and H3O+ are conjugate acid-base pairs. H2O is the conjugate base of H3O+. Acid H3O+ Acid HA Base H2O Base A- add H+ remove H+

  22. Conjugate Acid & Bases • Provide the formulas for the conjugate bases of: H2SO4 HCN H2O CH3COOH

  23. Conjugate Acid & Bases • Provide the formulas for the conjugate acids of: NH3 OH1- HCO31- H2O

  24. Structure and Acid Strength For binary acids of the general formula HnX (where X is a non-metal), the acidity reflects the H-X bond strength and polarity. If the bond is strong and non-polar, such as with carbon (in CH4), the compound is non-acidic, and the C-H bonds remain intact.

  25. Structure and Acids & Bases The effects of bond polarity can be seen in the hydrogen halides: HF, HCl, HBr and HI. HF has the most polar bond, yet it is the weakest acid of the group. The HF bond is the strongest of the group, and this very high bond strength results in a lower tendency for dissociation in water.

  26. Binary acid strength depends upon the electronegativity of the non-metal, and also the bond strength of the bond with hydrogen.

  27. The bond strength is greatest when there is a good match in energy between the 1s orbital on hydrogen, and the bonding orbital on the non-metal.

  28. Structure and Acids & Bases The acids HCl, HBr and HI are all strong acids. In all cases, the bonds are polar, but they are also weaker than the H-F bond, and are more easily broken in aqueous solution.

  29. Structure and Acids & Bases

  30. The Oxyacids The oxyacids typically have one or more oxygen atoms attached to a central non-metal. Attached to one or more of the oxygen atoms are acidic hydrogens. Examples include H2SO4, HNO3, HNO2, H3PO4 and HClO4. It is important to remember that in all cases, the acidic hydrogen is attached to oxygen (X-O-H).

  31. The Oxyacids

  32. Oxyacids For a given central atom, the greater the number of oxygen atoms attached, the more acidic the acid. The larger number of oxygen atoms polarizes and weakens the O-H bonds.

  33. Oxyacids The greater the number of oxygen atoms, the weaker the O-H bonds, and the stronger the acid.

  34. Oxyacids The nature of the central atom also affects the acidity of the acid. The more electronegative X is (in X-O-H), the weaker the O-H bond, and the stronger the acid.

  35. Oxyacids The nature of the central atom also affects the acidity of the acid. The more electronegative X is (in X-O-H), the weaker the O-H bond, and the stronger the acid.

  36. Strong and Weak Acids HX is a weak acid and forms only a small amount of H3O+ and X-

  37. Strong and Weak Acids HY is a strong acid and dissociates completely to form H3O+ and Y-

  38. Strong and Weak Acids

  39. Acid Strength The conjugate bases of strong acids have no tendency to pick up a proton. There is no reverse reaction. The conjugate bases of infinitely strong acids will not act as a base and accept a proton.

  40. Ka Values

  41. Bases Strong bases are very effective at accepting protons. The most common strong bases are the soluble group IA and IIA metal hydroxides. In general, the metal ion is non-reactive, and serves as a spectator ion. Another strong base is oxide ion, O2-. Oxide reacts with water to become fully protonated. O2-(aq) + H2O(l)  2 OH-(aq)

  42. Bases O2-(aq) + H2O(l)  2 OH-(aq) In this reaction, water is donating a proton, and hence acting as an acid. In previous reactions, water accepted a proton and served as a base. Substances that can behave as either an acid or a base are called amphoteric.

  43. Amphoteric Nature of Water Depending upon its environment, water may donate a proton, acting as an acid, or accept a proton, and act as a base. This behavior is characteristic of amphoteric substances. Pure water molecules can react with each other, to a very small extent, to form hydronium and hydroxide ions.

  44. Autoionization of Water This process is called autoionization or selfionization. One water molecule donates a proton to another. The result is the formation of equal amounts of hydronium and hydroxide ions. H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)

  45. Autoionization of Water H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq) Kw = [H3O+][OH-] = 1.0 x 10-14 at 25oC In pure water, the concentration of hydronium ion equals the concentration of hydroxide. Both ions have a concentration of 1.0 x 10-7M.

  46. [H3O+] and [OH-] in Aqueous Solution The product of the hydroxide and hydronium concentration in any aqueous solution must equal Kw. As a result, when a solution is acidic, the hydronium concentration increases, and the hydroxide concentration decreases. Kw = [H3O+][OH-] = 1.0 x 10-14 at 25oC

  47. [H3O+] and [OH-] in Aqueous Solution Likewise, in basic solutions, the hydroxide ion concentration is greater than the hydronium ion concentration. There is always some hydroxide ion and some hydronium ion present in any aqueous solution.

  48. [H3O+] and [OH-] in Aqueous Solution

  49. The pH Scale A scale of acidity, the pH scale, is used to indicate the degree of acidity of aqueous solutions. pH = -log[H3O+] or –log[H+] The scale generally runs from 0-14, though negative pH values are possible. A neutral solution will have a pH = 7.00

  50. The pH Scale A one unit change in pH is a ten-fold change in the concentration of hydronium ion. Acidic solutions have pH values less than 7.00, and basic solutions have pH values greater than 7.00

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