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Chapter 17

Chapter 17. Acids and Bases. Acids, Bases, and Matter. Classification of Matter. Arrhenius Definition of Acids and Bases. Acids Acidus – Latin for sour Substances that, when added to water, produce hydrogen ions, H+ (protons), or hydronium ions, H 3 O +. Bases

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Chapter 17

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  1. Chapter 17 Acids and Bases

  2. Acids, Bases, and Matter Classification of Matter

  3. Arrhenius Definition of Acids and Bases Acids • Acidus – Latin for sour • Substances that, when added to water, produce hydrogen ions, H+ (protons), or hydronium ions, H3O+.

  4. Bases • Alkali – Arabic for ashes. Product of burning plants is potash (KOH) which feels soapy and tastes bitter. • Substance that when added to water produce hydroxide ions, OH-.

  5. Salts • Substance formed in addition to water when an acid reacts with a base. The cation is from the base and the anion is form the acid.

  6. Limitations of Arrhenius Definition • Arrhenius's concept is limited to aqueous solutions because it refers to ions derived from water.

  7. The Hydronium Ion and Water Autoionization • There is an equilibrium between these two ions in water or any aqueous solution: H2O + H2O <---------> H3O+(aq) + OH-(aq)

  8. . The Bronsted Concept of Acids and Bases • Classifies an acid as a proton donor and a base as a proton acceptor. Another way to look at the model is an acid is a substance from which a proton can be removed and a base is a substance that can remove a proton from an acid.

  9. Acids • Recognizing acids is fairly simple - since an acid has to be able to "donate" a proton it must contain an ionizable hydrogen. Generally speaking this means that an acid is something that begins with a hydrogen. It can be hydrogen attached to a polyatomic ion or hydrogen attached to a nonmetal.

  10. Acids Continued • Acids that are capable of donating one proton are called monoprotic acids. Some acids are capable of donating more than one proton and are called polyprotic acids.

  11. Bases • Bases are not so simple- in most cases bases will be ionic compounds containing the OH- ion. When the compound is dissolved in water the OH- ion dissociates from the positive ion and is then able to "accept" the proton. In other cases regular anions act as bases. The only bases that don't fit into these categories are ammonia and derivatives of ammonia. The ammonia molecule - for reasons which aren't important to us - is able to "accept" a proton and become the ammonium ion, NH4+.

  12. Bases Continued • Bases can also be monoprotic or polyprotic depending on how many protons they accept.

  13. Amphiprotic substances • These are molecules or ions that can behave as a Bronsted acid or base. One of the best examples of this is water.

  14. Conjugate Acid-Base Pairs • A pair of compounds or ions that differ be the presence of one H+ unit is called a conjugate acid-base pair. When you have the proton you are the acid. If you do not you are the conjugate base.

  15. Conjugate Acid-Base Pairs • Every acid-base reaction involving H+ transfer has two conjugate acid-base pairs.

  16. Relative Strengths of Acids and Bases • The strength of an acid in solution has to do with the relative ability of an acid to donate protons, where an acid is considered strong only if it is such a good proton donor that each and every acid molecule will give up at least one hydrogen ion to a water molecule (100% acid dissociations). (The stronger the acid the weaker it holds the proton.)

  17. In the Bronsted model, an acid donates a proton and produces a conjugate base. this model also informs us that, in general, the stronger the acid, the weaker it conjugate base.

  18. Strong Acids • Completely ionize in water to form H3O+ ions: HX(aq) + H2O(l) ----------> H3O+(aq) + X-(aq) • Include: HCl, HBr, HI, HNO3, HClO4, H2SO4

  19. Strong Bases • Completely ionize to form OH- ions in solution. • Hydroxides of the Group I metals: • NaOH(s) ----------> Na+(aq) + OH-(aq) • Hydroxides of the heavier Group II metals • M(OH)2(s) ----------> M2+(aq) + 2OH-(aq) • M = Ca, Sr, Ba

  20. Weak Acids • Partially dissociate in water to form H3O+ ions: HX(aq) + H2O(l) <-------> H3O+(aq) + X-(aq) • Generally, concentration of HX molecules >> concentration of H3O+ions. Concentrations are governed by equilibrium constants - Ka.

  21. General Equation HB(aq) + H2O(l) <-----> H3O+(aq) + B-(aq) • Ka = [H3O+] [B-] / [HB] • The smaller the ionization constant, the weaker the acid. See table 17.4 pg 808

  22. Types of Weak Acids Molecular Weak Acids • (acid-base indicators are molecular weak acid) HIn(aq) + H2O(l) -----> H3O+(aq) + In-(aq) blue Yellow

  23. Types of Weak Acids Anions containing an ionizable H atom • HSO4-(aq), HCO3-, H2PO4-, etc

  24. Types of Weak Acids • All cations except those of Group I metals, Ca2+, Sr2+, Ba2+ • Ammonium, NH4+ Zn(H2O)42+ + H2O(l) <-------> H3O+(aq) + Zn(H2O)3OH+

  25. Weak Bases Concentrations are governed by equilibrium constants - Kb. • General Equation B-(aq) + H2O(l) <-------> HB(aq) + OH-(aq) • Kb = [HB] [OH-] / [B-] • The smaller the ionization constant, the weaker the base.

  26. Types of Weak Bases Molecular: NH3(aq) + H2O(l) <-----> NH4+(aq) + OH-(aq)

  27. Types of Weak Bases Anions derived from weak acids (conjugate bases): F-(aq) + H2O(l) <------> HF(aq) + OH-(aq)

  28. Predict the Predominate Direction of Acid-Base Reactions • The following chart can be used to predict whether the equilibrium in an acid-base reaction lies predominately to the left or the right. • Examples. • 1. CH3COOH and NaCN • 2. NH4Cl and Na2CO3

  29. The Water Ionization Constant, Kw • Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 C

  30. Water, Acids and Bases Pure Water [H3O+]= [OH-] = 1.00 x 10-7 Acidic [H3O+] > [OH-] [H3O+] > 1.00 x10-7 Basic [H3O+] < [OH-] [H3O+] < 1.00 x10-7

  31. Autoionization of Water • All acids/bases dissolved in water must obey equation for the ionization of water. • They either add H3O+ or OH to water. • Most of the acids in this chapter will be stronger than water and add significantly to the hydronium ion concentration.

  32. Examples The hydronium ion concentration of an acidic solution was 1.00x105 M. What was the [OH]? What is the hydronium ion concentration if the hydroxide concentration was 2.50x103 M?

  33. The pH Scale • pH = - log 10 [H+] = - log 10 [H3O+] • Relation to Aqueous Solution Neutral: pH = 7.00 Acidic: pH < 7.00 Basic: pH > 7.00

  34. Methods of Measuring pH pH paper is used that has compounds in it which are change to different colors for different pH ranges.

  35. Methods of Measuring pH • An colored indicator can be placed in the solution and its color correlated with pH. HIn(aq) + H2O(l)  H3O+(aq) + In(aq). E.g. phenolphthalein is colorless in acid form but pink in basic form. • The pH at which they change color depends on their equilibrium constant.

  36. Methods of Measuring pH • More accurate and precise measurements are made with a pH meter. A combination of voltmeter and electrodes

  37. The pOH • pOH= - log 10 [OH-] • Given: • pKw = - log 101.00 x 10-14 = 14 • pKw = 14 = pOH + pH

  38. Examples Determine the pH of a solution in which [H3O+] = 5.40x106 M Determine the pH of a solution in which the [OH] = 3.33x103 M

  39. Examples Determine the pOH of a solution in which the [OH] = 3.33x103 M Determine the [H3O+] if the pH of the solution is 7.35.

  40. The Ionization Constants for an Acid and Its Conjugate Base Ka x Kb = [H3O+] [B-] / [HB] x [HB] [OH-] / [B-] = [H3O+] [OH-] = 1.00x 10-14 = Kw • The strength of the base is inversely related to that of its conjugate weak acid.

  41. Equilibria Involving Weak Acids Calculating Ka or Kb from Initial Concentrations and Measured pH Calculation of Ka (in water) [H3O+] = [B-] = antilog –pH • [HB] = original concentration - [H3O+] Percent Ionization Percent Ionization HB = [H3O+] / original concentration HB

  42. Weak Acids and the ICE Table

  43. Equilibria Involving Weak Bases • Same procedure as with Ka (H+ and OH- exchange)

  44. Weak Bases and the ICE Table

  45. Calculating Equilibrium Concentrations and pH from Initial Conc. and Ka or Kb Dissolving a weak acid, HB, in water • [H3O+] = [B-] • [HB] = original concentration - [H3O+]

  46. Making it easier • In general, Ka is seldom known to better than +/- 5%, Hence, in the expression • Ka = [H3O+] [B-] / [HB] = x2/(a-x) if x/a < 0.05 we take a - x = a

  47. Calculation of [OH-] in Solution of Weak Base • Same procedure as the calculation of [H3O+] in solution of weak acid.

  48. Acid-Base Properties of Salts Salt • A salt is an ionic compound that COULD have been formed by the reaction of an acid with a base; a salt's positive ions come from the base , and its negative ions come from the acid.

  49. Hydrolysis reaction • A hydrolysis reaction is said to have occurred when a salt dissolves in water and leads to changes in the hydronium or hydroxide ion concentrations of the water. To determine whether a salt is acidic or basis you must consider the effect of the cation and the anion separately and then combine these effects to give the overall result for the salt.

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