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Strong and Weak Acids and Bases

Strong and Weak Acids and Bases. Strong and Weak Acids and Bases. You have learned that some substances, such as water and the hydrogen carbonate ion, are amphoteric. How could you predict which role the ion plays in each reaction?

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Strong and Weak Acids and Bases

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  1. Strong and Weak Acids and Bases

  2. Strong and Weak Acids and Bases • You have learned that some substances, such as water and the hydrogen carbonate ion, are amphoteric. • How could you predict which role the ion plays in each reaction? • This will depend on the strength of the ion in relation to the other substances present in the reaction.

  3. Strong Acids • In terms of acid-base reactions, strength refers to the extent to which a substance dissociates in its solvent. • An acid that dissociates completely is called a strong acid. • Hydrochloric acid is a strong acid. • All the molecules of HCl in an aqueous solution dissociate into H+ and Cl- ions.

  4. Strong Acids

  5. Weak Acids • A weak acid is an acid that dissociates only slightly in a water solution. • Thus, only a small percentage of the acid molecules dissociate. Most of the acid molecules remain intact. • For example, acetic acid is a weak acid. The percent dissociation of acetic acid is only about 1% in a 0.1 M solution.

  6. Conductivity • Strong acids almost completely dissociate in water, while weak acids dissociate slightly. This can be shown by measuring the conductivity of solutions.

  7. Monoprotic Acids • A few acids only contain a single hydrogen ion that can dissociate. • These acids are called monoprotic acid • Hydrochloric acid, hydrobromic acid and hydroiodic acid are strong monoprotic acids. • Hydrofluoric acid is a weak monoprotic acid.

  8. Diprotic Acids • Many acids contain two or more hydrogen ions that can dissociate. Sulfuric acid is a strong acid, but only for its first dissociation. H2SO4(aq)H+(aq) + HSO4-(aq) • The resulting aqueous hydrogen sulfate ion, HSO4-, is a weak acid. It dissociates to form the sulfate ion in the following equilibrium dissociation. HSO4-(aq)H+(aq) + SO42-(aq) • Acids that contain two hydrogen ions dissociate to form two anions. These acids are sometimes called diprotic acids.

  9. Triprotic Acids • Acids that contain three hydrogen ions capable of dissociatingare called triprotic acids. • Phosphoric acid, H3PO4(aq), is a triprotic acid. It gives rise to three anions.

  10. Strong Bases • Like a strong acid, a strong base dissociates completely into ions in water. • All oxides and hydroxides of the alkali metals—Group 1 (IA)—are strong bases. • The oxides and hydroxides of the alkaline earth metals—Group 2 (IIA)—below beryllium are also strong bases.

  11. Strong Bases

  12. Weak Bases • Most bases are weak. • A weak base dissociates very slightly in a water solution. • The most common weak base is aqueous ammonia.

  13. Relative Strengths of Acids and Bases • Over the centuries, chemists have performed countless experiments involving acids and bases. • These experiments have allowed chemists to rank acids and bases according to their strengths in relation to one another.

  14. Relative Strengths of Acids and Bases • Being able to compare the relative strengths of acids and bases can allow you to predict the direction an acid-base reaction will proceed. • The direction of an acid-base reaction usually proceeds from a stronger acid and a stronger base to a weaker acid and a weaker base.

  15. Describing Acids and Bases Quantitatively • As you know, all aqueous solutions contain ions. Even pure water contains a few ions that are produced by the dissociation of water molecules. H2O(l) + H2O(l) H3O+(aq) + OH-(aq) • It has been determined that the concentration of hydronomium ions in pure water at 25oC is 1.0x10-7 mol/L. • The dissociation of water also produces the same very small number of hydroxide ions: 1.0x10-7 mol/L.

  16. Describing Acids and Bases Quantitatively • Because the dissociation of water is an equilibrium, you can write an equilibrium constant expression for it. • The resulting constant is called the ion product constant for water, Kw. Kw = [H3O+][OH-] • The equilibrium value of Kw at 25oC is as follows: Kw = (1.0x10-7 mol/L)(1.0x10-7mol/L) = 1.0x10-14

  17. Describing Acids and Bases Quantitatively • The concentration of H3O+ in the solution of a strong acid is equal to the concentration of the dissolved acid. • Consider [H3O+] in a solution of 0.1 mol/L of HCl. All the molecules of HCl dissociate in water, forming a hydronium ion concentration that equals 0.1 mol/L.

  18. Describing Acids and Bases Quantitatively • When either [H3O+] or [OH-] is known, you can use the ion product constant for water, Kw, to determine the concentration of the other ion.

  19. Example • What is the [H3O+] and [OH-] in a 2.5 mol/L solution of nitric acid? • Nitric acid is a strong acid that dissociates completely in aqueous solution so you can use its concentration to determine [H3O+]. You can find the concentration of the other ion using Kw: [HNO3] = 2.5 mol/L, so [H3O+] = 2.5 mol/L Kw = 1.0 x 10-14 = [H3O+][OH-] [OH-] = 1.0x10-14 mol/L 2.5 = 4.0x10-15 mol/L

  20. Practice Problems • Try problems 12-15 on p. 566

  21. The pH Scale • The concentration of hydronium ions ranges from about 10 mol/L for a concentrated strong acid to about 10-15 mol/L for a concentrated strong base. • This wide range of concentrations are not convenient to work with. So, a method for converting the concentrations to positive numbers was suggested. pH = -log[H3O+] • pH stands for the power of hydrogen ions

  22. The pH Scale • For example, recall that [H3O+] of neutral water at 25oC is 1.0x10-7 mol/L. pH = -log[H3O+] = -log(1.0x10-7) = -(-7.00) = 7.00

  23. The pH Scale • Since the pH scale is logarithmic, a difference of one pH unit represents a 10-fold change in the H+ ion concentration. • Example • Describe the difference in [H+] for the following pair of solutions: orange juice with a pH of 3.5 and lemon juice with a pH of 2.5 • The pH of the lemon juice (2.5) is one unit lower than the pH of the orange juice (3.5); therefore, the [H+] of the lemon juice is 10 times greater than that of the orange juice.

  24. pH Values for Some Commonly Used Substances

  25. Significant Figures • How do you determine the number of significant figures in a pH? • You only count the digits to the right of the decimal point in the pH value. • Suppose that the concentration of hydronium ions in a sample of orange juice is 2.5x10-4 mol/L. • This number has two significant figues. • The pH of the sample is –log(2.5x10-4) = 3.602059 • The digit to the left of the decimal (the 3) is derived from the power of 10. • Only the two digits to the right of the decimal are significant. • Thus, the pH value is rounded off to 3.60

  26. pOH • Just as pH refers to the exponential power of the hydronium ion concentration in a solution, pOH refers to the power of hydroxide ion concentration. • You can calculate the pOH of a solution from the [OH-]. pOH = -log[OH-]

  27. pH and pOH • Kw = [H3O+][OH-] = 1.0x10-14 at 25oC • pH + pOH = 14 Because: -log(1.0 x 10-7) + -log(1.0 x 10-7) = 7 + 7 = 14

  28. Another Way to Find [H3O+] and [OH-] • You can calculate [H3O+] or [OH-] by finding the antilog of the pH or pOH. [H3O+] = 10-pH [OH-] = 10-pOH

  29. Practice Problems • Try problems 20-25 on p. 572

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