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Acids and Bases: Definitions, Reactions, and Strength

This text provides an overview of acids and bases, including their definitions, reactions, and strength. Learn about Arrhenius and Brønsted-Lowry concepts, conjugate acids and bases, acid-base strength, and the autoionization of water.

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Acids and Bases: Definitions, Reactions, and Strength

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  1. Updates • Assignment 04 is up on ACME and is due Mon., Feb. 26 (in class) • Midterms not marked yet (soon!)

  2. Acids and Bases Chapter 16

  3. Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases. 16.1

  4. Some Definitions • Arrhenius • Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions. • Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions. • Brønsted–Lowry • Acid: Proton donor • Base: Proton acceptor

  5. Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water 16.1

  6. HCl is also a Brønsted-Lowry acid; and NH3 is also a Brønsted-Lowry base A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor 16.1

  7. Brønsted-Lowry acid-base reaction • Arrhenius concept of acids and bases limited to aqueous solutions • Brønsted-Lowry concept includes acid-base reactions in nonaqueous solutions

  8. Conjugate Acids and Bases: • From the Latin word conjugare, meaning “to join together.” • Reactions between acids and bases always yield their conjugate bases and acids. • Conjugate acid-base pairs differ by a proton • In an equilibrium reaction, the conjugate base and conjugate acid react to form the acid and base as written on the reactant side

  9. A Brønsted–Lowry acid must have a removable (acidic) proton. A Brønsted–Lowry base must have a pair of nonbonding electrons.

  10. A Brønsted–Lowry acid must have a removable (acidic) proton. A Brønsted–Lowry base must have a pair of nonbonding electrons.

  11. Acid and Base Strength • Strong acids are completely dissociated in water. • Their conjugate bases are quite weak. • Weak acids only dissociate partially in water. • Their conjugate bases are weak bases.

  12. Acid and Base Strength • Substances with negligible acidity do not dissociate in water. • Their conjugate bases are exceedingly strong.

  13. Acid and Base Strength In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. base conj. base HCl(aq) + H2O(l) H3O+(aq) + Cl−(aq) H2O is a much stronger base than Cl−, so the equilibrium lies so far to the right K is not measured (K>>1).

  14. C2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2−(aq) Acid and Base Strength base conj. base Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1).

  15. H2O stronger base than Cl- • Acetate stronger base than H2O • Remember reaction equilibrium will favor the direction that puts the proton on the stronger base

  16. H2O(l) + H2O(l) H3O+(aq) + OH−(aq) Autoionization of Water • Water is amphoteric, meaning it can act as an acid or a base • In pure water, a few molecules act as bases and a few act as acids. • This is referred to as autoionization.

  17. + H H O O [ ] + + - H H H O H O H H H2O (l) H+(aq) + OH-(aq) H2O + H2O H3O+ + OH- Acid-Base Properties of Water autoionization of water conjugateacid base conjugatebase acid 16.2

  18. H2O(l) + H2O(l) H3O+(aq) + OH−(aq) Ion-Product Constant • The equilibrium expression for this process is Kc = [H3O+] [OH−] • This special equilibrium constant is referred to as the ion-product constant for water, Kw. • At 25°C, Kw = 1.0  10−14

  19. pH pH is defined as the negative base-10 logarithm of the hydronium ion concentration. pH = −log [H3O+]

  20. pH • In pure water, Kw = [H3O+] [OH−] = 1.0  10−14 • Because in pure water [H3O+] = [OH−], [H3O+] = (1.0  10−14)1/2 = 1.0  10−7

  21. pH • Therefore, in pure water, pH = −log (1.0  10−7) = 7.00 • An acid has a higher [H3O+] than pure water, so its pH is <7 • A base has a lower [H3O+] than pure water, so its pH is >7.

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