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Acid and Base Equilibria

Acid and Base Equilibria. Memorize – strong acids and bases. Definitions. Arrhenius : acids produce H + ions in water and bases produce OH - ions. Bronsted-Lowry : an acid is a proton donor and a base is a proton acceptor. HC 2 H 3 O 2 + H 2 O  C 2 H 3 O 2 - + H 3 O +

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Acid and Base Equilibria

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  1. Acid and Base Equilibria Memorize – strong acids and bases

  2. Definitions • Arrhenius: acids produce H+ ions in water and bases produce OH- ions. • Bronsted-Lowry: an acid is a proton donor and a base is a proton acceptor. HC2H3O2 + H2O  C2H3O2- + H3O+ acid base base acid NH3 + H2O  NH4+ + OH- base acidacidbase

  3. Species that differ by a proton, like H2O and H3O+, are called conjugate acid-base pairs The reaction of HCl and H2O. HCl is the acid because it donates a proton. Water is the base because it accepts a proton.

  4. (a) Formic acid transfers a proton to a water molecule. HCHO2 is the acid and H2O is the base. (b) When a hydronium ion transfers a proton to the CHO2- ion, H3O+ is the acid and formate ion is the base.

  5. Conjugate Acid/Base pairs Acid/Base HC2H3O2/C2H3O2- NH4+/NH3 H3O+/H2O H2O/OH- • H2O is amphoteric because it can act as either an acid or a base.

  6. An amphoteric substances can act as either an acid or base • For example, the hydrogen carbonate ion:

  7. Strong and Weak Acids and Bases • The strength of an acid is a measure of its ability to transfer a proton • Acids that dissociate completely with water (like HCl and HNO3) are classified as strong • Acids that are less than completely ionized are called weak acids • Bases can be classified in a similar fashion.

  8. Acetic acid (HC2H3O2) is a weak acid • It ionizes only slightly in water • The hydronium ion is a better proton donor than acetic acid (it is a stronger acid) • The acetate ion is a better proton acceptor than water (it is a stronger base) • The position of an acid-base equilibrium favors the weaker acid and base (the reactants are favored in this example)

  9. The strengths of the binary acids • increases from left to right within the same period • For example, PH3 < H2S < HCl • increase from top to bottom within a group • For example, HCl < HBr < HI

  10. Strength of All Acids 1. The POLARITY of the X-H bond • The greater the polarity of the X-H bond the greater the strength of the acid.  Polarity is measured by the difference in electronegativity between the bonded atoms. • When an acid dissociates in water, the X-H bond is broken.  The greater the polarity of the bond, the easier it is to break and produce H+ ions, and thus the stronger the acid.

  11. Explain the difference in the Ka Values • FormulaKa valueDEN • HF 7.2 x 10-4 1.8 • H2O 1.8 x 10-16 1.2 • NH3 1 x 10-33 0.8 • CH4 1 x 10-49 0.4

  12. 2.  The CHARGE on the acid or base • Compounds becomes less acidic and more basic as the negative charge increases. • It is easier to remove a positive ion (H+) from a neutral atom or molecule than a negatively charged one.   • Bases (H+ acceptors) become stronger as their negative charge increases because they have a stronger force of attraction for pulling in extra H+ ions.

  13. Compounds becomes less acidic and more basic as the negative charge increases. Formula pH • H3PO4 1.5 • H2PO4- 4.4 • HPO42- 9.3 • PO43- 12.0

  14. Oxyacids When the polarity, size, and charge of two compounds are all the same (e.g. oxyacids of the same element) we must find another way to measure the relative strengths of these acids. • Trends in oxoacids (acids of hydrogen, oxygen, and one other element)

  15. Oxyacid • Oxyacid - An acid in which the acid hydrogen atoms are attached to an oxygen atom • Examples of oxyacids of the same element: H2SO4 and H2SO3 HNO3 and HNO2

  16. 3. Oxidation State • As the oxidation state of an atom increases, its tendency to draw electrons increases.   • In an oxyacid, the central atom pulls electrons away from the oxygen, consequently making the oxygen more electronegative.   • The O-H bond, therefore becomes more polar, making it easier to form ions and thus increasing the strength of the acid.

  17. As oxidation state increases so does the acidity of the oxyacid. Oxyacid      Ka value     Oxidation # of Cl HClO 2.9 x 10-8 HClO2 1.1 x 10-2 HClO3 5.0 x 102 HClO4 1 x 103 +1 +3 +5 +7

  18. When the central atom holds the same number of oxygen atoms, the trend is the same as for binary acids across a period, but the reverse for down a column. • Acid strength: HClO4 > HBrO4 > HIO4 • Acid strength:HClO4 > H2SO4 > H3PO4 • For a given central atom, the acid strength of an oxoacid increases with the number of oxygens held by the central atom • Acid strength: H2SO4 > H2SO3

  19. There is a third definition for acid and bases • It is a further generalization, or broadening, of the species that can be classified as either an acid or base • The definitions are based on electron pairs and are called Lewis acids and bases

  20. Definitions • Lewis acid - accepts a pair of electrons to form a coordinate covalent bond. • Lewis base – donate a pair of electrons Cl H H Cl N: + B  H N B H Cl H Cl Cl Cl H

  21. NH3 (a Lewis base) forms a coordinate covalent bond with BF3 (a Lewis acid) during neutralization. NH3BF3 is called an addition compound because it was made by joining two smaller molecules.

  22. Carbon dioxide (Lewis acid) reacts with hydroxide ion (Lewis base) in solution to form the bicarbonate ion. The electrons in the coordinate covalent bond come from the oxygen atom in the hydroxide ion.

  23. Lewis acids: • Molecules or ions with incomplete valence shells (for example BF3 or H+) • Molecules or ions with complete valence shells, but with multiple bonds that can be shifted to make room for more electrons (for example CO2) • Molecules or ions that have central atoms capable of holding additional electrons (usually, atoms of elements in Period 3 and below, for example SO2)

  24. Lewis bases: • Molecules or ions that have unshared pairs of electrons and that have complete shells (for example O2- or NH3) • All Brønsted acids and bases are Lewis acids and bases, just like all Arrhenius acids and bases are Brønsted acids and bases

  25. Neutral solutions: [H+] = [OH-] • Acidic solutions: [H+] > [OH-] • Basic solutions: [H+] < [OH-] • To make the comparison of small values of [H+] easier, the pH was defined: • In terms of the pH: • Neutral solutions: pH = 7.00 • Acidic solutions: pH < 7.00 • Basic solutions: pH > 7.00

  26. The pH of some common solutions. [H+] decreases, while [OH-] increases, from top to bottom.

  27. The pH of a solution can be measured with a pH meter or estimated using a visual acid-base indicator • An acid-base indicator is a species that changes color based on the pH • Calculating the pH of a strong acid or base is “easy” because they are 100% dissociated in aqueous • For example, the pH of 0.10 M HCl is 1.00 and the pH of 0.10 M NaOH is 13.00

  28. In the last example it was assumed that the total concentration of [H+] was due to the strong acid (HCl) and [OH-] was due to the strong base (NaOH) • This assumption is valid because the autoionization of water is suppressed in strongly acidic or strongly basic solutions • This assumption fails for very dilute solutions of acids or bases (less than 10-6M)

  29. Equilibrium Constant Expression • The equilibrium or ionization constants for weak acids and bases and water are given the labels of Ka, Kb and Kw. • A weak acid: HA + H2O  H3O+ + A- Ka = [H+][A-] [HA]

  30. Equilibrium Constant Expressions • A weak base: B + H2O  BH+ + OH- Kb = [BH+][OH-] [B] • Water dissociation: H20  H+ + OH- Kw = [H+][OH-] (1 x 10-7)(1x 10-7)= 1 x 10-14

  31. Equilibrium Constant Relationships • The product of the Ka and Kb for an acid and its conjugate base is the Kw KaKb = Kw = 1 x 10-14 • The greater the Ka or Kb the greater the dissociation of the acid or base • Ka and Kb values are usually very small since they are weak acids and bases. • pKa and pKb are used to show the equilibrium concentrations.

  32. Ionization Constant Relationships • pKb = -log Kb • pKa = -log Ka • pKw = -log Kw = -log(1 x 10-14) pKw = 14 So, pKa + pKb = 14 And, pH + pOH = 14 • If you know ka for a weak acid you can always find kb for its conjugate base and visa-versa.

  33. Example: Morphine is very effective at relieving intense pain and is a weak base. What is the Kb, pKb, and percentage ionization of morphine if a 0.010 M solution has a pH of 10.10? At equilibrium, [OH-] = x = 10-pOH

  34. SOLUTION: Use pOH = 14.00 – pH, substituting:

  35. Ionization Constant Calculations • A monoprotic acid solution has a concentration of 0.100 M and the pH is 2.44 @ 25oC. Calculate the Ka and pKa. Ka = 1.36 x 10-4, pKa = 3.86 • Calculate the pH of a 0.010 M solution of HCl. pH = 2.0 • Hydrazine, N2H4 has a concentration of 0.025 M. Calculate the pH and % ionization. The Kb = 1.7 x 10-5 pH = 10.81, % ionization = 2.6 %

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