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CH 104: ACID-BASE PROPERTIES OF AQUEOUS SOLUTIONS

CH 104: ACID-BASE PROPERTIES OF AQUEOUS SOLUTIONS. In the Arrhenius theory an acid produces H + in aqueous solution and a base produces OH – in aqueous solution. The more general Brønsted-Lowry theory defines an acid as a H + donor and a base as a H + accepter.

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CH 104: ACID-BASE PROPERTIES OF AQUEOUS SOLUTIONS

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  1. CH 104: ACID-BASE PROPERTIES OF AQUEOUS SOLUTIONS • In the Arrhenius theory an acid produces H+ in aqueous solution and a base produces OH– in aqueous solution. • The more general Brønsted-Lowry theory defines an acid as a H+ donor and a base as a H+ accepter. Svante Arrhenius Johannes Brønsted Thomas Lowry

  2. HYDROGEN, HYDRONIUM, AND HYDROXIDE IONS • The symbol H+(aq) is convenient to use; however, it is not accurate. Hydrogen ion (H+) is a proton without an electron. It is hydrated in water and exists as hydronium ion (H3O+(aq)). • The self-ionization of water: • Preferred: 2H2O(l) = H3O+(aq) + OH–(aq) • Accepted: H2O(l) = H+(aq) + OH–(aq) • H3O+ (or H+) is acid. • Hydroxide ion (OH–) is base.

  3. pH AND HYDRONIUMION CONCENTRATION • The pH scale measures acidity. It typically ranges from 0 to 14. The acidity is neutral at pH 7. Values less than pH 7 are increasingly acidic. Values greater than pH 7 are increasingly basic.

  4. pH AND HYDRONIUMION CONCENTRATION • If pH = 0.0, then [H3O+] = 1 M = 1x10–0 M • If pH = 1.0, then [H3O+] = 0.1 M = 1x10–1 M • If pH = 2.0, then [H3O+] = 0.01 M = 1x10–2 M • If pH = 3.0, then [H3O+] = 0.001 M = 1x10–3 M • If pH = 4.0, then [H3O+] = 0.0001 M = 1x10–4 M • If pH = 5.0, then [H3O+] = 0.00001 M = 1x10–5 M • If pH = 6.0, then [H3O+] = 0.000001 M = 1x10–6 M • If pH = 7.0, then [H3O+] = 0.0000001 M = 1x10–7 M • If pH = 8.0, then [H3O+] = 0.00000001 M = 1x10–8 M • If pH = 9.0, then [H3O+] = 0.000000001 M = 1x10–9 M • If pH = 10.0, then [H3O+] = 0.0000000001 M = 1x10–10 M • If pH = 11.0, then [H3O+] = 0.00000000001 M = 1x10–11 M • If pH = 12.0, then [H3O+] = 0.000000000001 M = 1x10–12 M • If pH = 13.0, then [H3O+] = 0.0000000000001 M = 1x10–13 M • If pH = 14.0, then [H3O+] = 0.00000000000001 M = 1x10–14 M • pH 2.0, 0.01 M, and 1x10–2 M each have 1 significant figure. For pH and other logarithms, the numbers to the right of the decimal are significant. The numbers to the left of the decimal are NOT significant. The 0 in pH 2.0 is significant. The 2 in pH 2.0 is NOT significant, it defines the 2 in 1x10–2 M.

  5. THE pH OF COMMON SUBSTANCES

  6. THE pH OF ACID PRECIPITATION • This is the distribution of precipitation pH for North America. • The combustion of sulfur-containing coal from Midwestern power plants is a major cause of acid precipitation. • 2S + 3O2 + 2H2O → 2H2SO4 • The prevailing winds carry this acid from the Midwest to the East.

  7. MATHEMATICS, ACIDS, AND BASES All concentrations are in moles per liter. (1) The Ion Product of Water = Kw = 1.0x10–14 = [H3O+][OH–] Rearranging Equation 1. (2) [H3O+] = (1.0x10–14) / [OH–] (3) [OH–] = (1.0x10–14) / [H3O+] “p” is the negative base 10 logarithm. (4) pH = –log10[H3O+] = log10(1 / [H3O+]) (5) pOH = –log10[OH–] = log10(1 / [OH–])

  8. MATHEMATICS, ACIDS, AND BASES Taking the “p” of Equation 1 and rearranging. Kw = 1.0x10–14 = [H3O+][OH–] pKw = 14.00 =pH + pOH (6) pH = 14.00 – pOH (7) pOH = 14.00 – pH Taking the antilogarithm of Equation 4. pH = –log10[H3O+] (8) [H3O+] = 10(–pH) Taking the antilogarithm of Equation 5 and inserting Equation 7. pOH = –log10[OH–] [OH–] = 10(–pOH) (9) [OH–] = 10(pH – 14.00)

  9. MATHEMATICS, ACIDS, AND BASES • In summary, • (1) The Ion Product of Water = Kw = 1.0x10–14 = [H3O+][OH–] • (2 and 8) [H3O+] = (1.0x10–14) / [OH–] = 10(–pH) = 10(pOH – 14.00) • (3 and 9) [OH–] = (1.0x10–14) / [H3O+] = 10(–pOH) = 10(pH – 14.00) • (4 and 6) pH = –log10[H3O+] = log10(1 / [H3O+]) = 14.00 – pOH • (5 and 7) pOH = –log10[OH–] = log10(1 / [OH–]) = 14.00 – pH • Complete this table. 2.9x10–11 M 3.46 10.54 2.0x10–5 M 4.70 9.30 4.2x10–6 M 2.4x10–9 M 8.62 3.8x10–7 M 2.6x10–8 M 6.42

  10. STRONG ACIDS AND STRONG BASES • A strong acid or a strong base in distilled water will almost completely ionize. • Strong acid: HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq) • Strong base: NaOH(s) + H2O(l) → Na+(aq) + OH–(aq) • Common strong acids and strong bases. • a H2SO4 ionizes in 2 steps. The first ionization goes to completion. The second ionization does not go to completion.

  11. WEAK ACIDS AND WEAK BASES • Most acids and most bases are weak. That is, most acids and most bases in distilled water do not completely ionize. • Weak acid: HC2H3O2(l) + H2O(l) = H3O+(aq) + C2H3O2–(aq) • A weak acid (HC2H3O2) is in equilibrium with its conjugate base (C2H3O2–). • Ionization Constant = Ka = [H3O+][C2H3O2–] / [HC2H3O2] = 1.74x10–5 at 25° C. • Weak base: NH3(aq) + H2O(l) = NH4+(aq) + OH–(aq) • A weak base (NH3) is in equilibrium with its conjugate acid (NH4+). • Ionization Constant = Kb = [NH4+][OH–] / [NH3] = 1.74x10–5 at 25° C.

  12. BUFFERS AND THE HENDERSON-HASSELBALCH EQUATION • A buffer is a solution that resists drastic changes in pH when an acid or base is added. • Furthermore, a buffer resists drastic changes in pH when it is diluted. • Buffers are used to control pH. For example, human blood is buffered at pH 7.4±0.1. The ability of blood to carry oxygen depends on the pH being within this range. • A buffer is a mixture of a weak acid and a salt of its conjugate base, or a weak base and a salt of its conjugate acid. • For example, a mixture of acetic acid (HC2H3O2) and sodium acetate (NaC2H3O2) is a common buffer. • What is the conjugate base of acetic acid? • Acetate (C2H3O2–).

  13. BUFFERS AND THE HENDERSON-HASSELBALCH EQUATION • This Henderson-Hasselbalch equation shows how a buffer of a weak acid and its salt resists drastic changes in pH. • (A similar Henderson-Hasselbalch equation would show how a buffer of a weak base and its salt resists drastic changes in pH.)

  14. BUFFERS AND THE HENDERSON-HASSELBALCH EQUATION • Therefore, the buffering of a weak acid and its salt depends on the relative concentrations of its conjugate base (A–) and its unionized acid (HA). • If a small amount of strong acid is added, it will combine with A– to make HA. If the change in [A–]/[HA] is small, the change in pH will be small. • Conversely, if a small amount of strong base is added, it will react with HA to make A–. If the change in [A–]/[HA] is small, the change in pH will be small. • What has a larger buffering capacity (a larger resistance to changes in pH)? A solution with [A–] = 0.001 M and [HA] = 0.001 M. Or a solution with [A–] = 1 M and [HA] = 1 M. • The solution with [A–] = 1 M and [HA] = 1 M has a larger buffering capacity.

  15. MEASURING pH BY GLASS ELECTRODE • The glass membrane of a pH electrode is made out of silicate groups with exchangeable hydrogen ions, Si-O-H+. These Si-O– groups are attached to the glass membrane. These H+ ions are in equilibrium with the surface of the glass membrane and the sample. • Glass membrane-Si-O-H+(s) = • Glass membrane-Si-O–(s) + H+(aq) • If the number of H+ ions in the sample is large, then the number of H+ ions on the glass membrane is large and the electrode voltage is small. Conversely, if the number of H+ ions in the sample is small, then the number of H+ ions on the glass membrane is small and the electrode voltage is large. This voltage is converted to a pH value.

  16. MEASURING pH BY GLASS ELECTRODE

  17. MEASURING pH BY PAPER • A wide variety of dyes are used to make pH paper. These dyes change color with pH.

  18. SAFETY • Give at least 1 safety concern for the following procedure. • Using acids and bases. • These are irritants. Wear your goggles at all times. Immediately clean all spills. If you do get either of these in your eye, immediately flush with water. • Your laboratory manual has an extensive list of safety procedures. Read and understand this section. • Ask your instructor if you ever have any questions about safety.

  19. SOURCES • Christian, G.D. 1986. Analytical Chemistry, 3rd ed. New York, NY: John Wiley & Sons, Inc. • Harris, D.C. 1999. Quantitative Chemical Analysis, 5th ed. New York, NY: W.H. Freeman Company. • Hill, J.W., D.K. Kolb. 2007. Chemistry for Changing Times, 11th ed. Upper Saddle River, NJ: Pearson Prentice Hall. • McMurry, J., R.C. Fay. 2004. Chemistry, 4th ed. Upper Saddle River, NJ: Prentice Hall. • Park, J.L. 2004. ChemTeam: Photo Gallery Menu. Available: http://dbhs.wvusd.k12.ca.us/webdocs/Gallery/GalleryMenu.html [accessed 9 October 2006]. • Petrucci, R.H. 1985. General Chemistry Principles and Modern Applications, 4th ed. New York, NY: Macmillan Publishing Company.

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