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Acids and Bases

Acids and Bases. Ch. 16. Properties of Acids and Bases. Acids: taste sour, react with metals to produce hydrogen gas, conduct electricity. Bases: taste bitter, feel slippery, conduct electricity. Definitions of Acids and Bases. Arrhenius definition

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Acids and Bases

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  1. Acids and Bases Ch. 16

  2. Properties of Acids and Bases • Acids: taste sour, react with metals to produce hydrogen gas, conduct electricity. • Bases: taste bitter, feel slippery, conduct electricity.

  3. Definitions of Acids and Bases • Arrhenius definition Acid- produces hydrogen ions in aqueous solution Base- produces hydroxide ions in aqueous solution • This definition is limited only to aqueous solutions and allows for only one kind of base.

  4. Definitions of Acids and Bases • Bronsted-Lowrey definition Acid- proton (H+) donor Base-proton acceptor • This is a more general definition • Under this definition, water can act as both an acid and a base and is said to be amphoteric. • When water acts as a base and accepts a proton, the hydronium ion (H3O+) is formed.

  5. Definitions of Acids and Bases • Lewis Acid-Base model acid-electron pair acceptor (has an empty orbital that can be used to share an electron pair base-electron pair donor (has a lone pair of electrons) • Covers many reactions not covered by the other definitions. • Example: BF3 + NH3 F3BNH3 Lewis Lewis acid base

  6. Formation of Acids/Bases • General reaction when an acid is dissolved in water: HA + H2O  H3O+ + A- acid base conjugate conjugate acid base • Conjugate base- remains of the acid particle after the proton has been donated. • Conjugate acid- is formed when the proton is transferred to the base.

  7. Application of Equilibrium • The equation representing the formation of conjugate acids/bases really represents a competition for the proton between the two bases, H2O and A-. • If H2O is a much stronger base (has a stronger affinity for the proton) than A-, the equilibrium position lies far to the right. (Most of the acid will be in the ionized form). • If A- is a much stronger base than H2O, the equilibrium position will lie far to the left. (Most of the acid will be present as HA).

  8. Acid Dissociation Constant • The equilibrium constant for the dissociation of an acid is called the acid dissociation constant (Ka). • Because H2O does not appear in the equilibrium expression (pure liquid), it is often not included in the equation. • If water is not included in the equation, H3O+ will be represented simply as H+.

  9. Acid Strength • Strong acid- almost completely ionized; equilibrium lies far to the right/ a strong acid produces a weak conjugate base (much weaker than water). • Weak acid- dissociates only to a small extent; equilibrium lies far to the left/ a weak acid produces a strong conjugate base (much stronger than water).

  10. Autoionization of Water • Because water is amphoteric, water can transfer a proton from one water molecule to another to produce a hydronium ion and a hydroxide ion. • 2 H2O (l) <---> H3O+ (aq) + OH- (aq) • Kw = [H3O+][OH-] • Kw is the dissociation constant for water

  11. Dissociation Constant for Water • Experiments show that at 25oC, [H3O+] = [OH-] = 1x10-7 M • Therefore, at 25oC Kw = [H3O+][OH-] = (1x10-7)(1x10-7) = 1x10-14

  12. The Significance of Kw • In any aqueous solution at 25oC, no matter what it contains, the product of [H3O+] and [OH-] must always equal 1x10-14. There are three possible situations: 1. A neutral solution, where [H3O+] = [OH-] 2. An acidic solution, where [H3O+] > [OH-] 3. A basic solution, where [H3O+] < [OH-]

  13. Sample Problems • Calculate the [H3O+] and [OH-] for each of the following and state whether the solution is acidic, basic, or neutral. a) 1.0 x 10-5 M OH- b) 1.0 x 10-7 M OH- c) 10.0 M H30+

  14. Sample Problem • At 60oC, the value of Kw is 1x10-13. a) Using LeChatelier’s Principle, predict whether the reaction: 2H20 < -- > H3O+ + OH- is exothermic or endothermic. b) Calculate [H3O+] and [OH-] in a neutral solution at 600C.

  15. The pH Scale • Because [H3O+] is typically very small, the pH scale provides a convenient way to represent solution acidity. • pH = -log[H3O+] • For a solution where [H3O+] = 1.0 x 10-7 M, pH = 7.00 • Rules for significant digits: the number of decimal places in the log is equal to the number of significant digits in the original number.

  16. Similar log scales • Similar log scales are used for representing other quantities: pOH = -log[OH-] pK = -logK **** pH + pOH = 14.00

  17. Polyprotic Acids • Polyprotic acids can supply more than one proton. • Examples include: H2SO4, H3PO4, etc. • Polyprotic acids always dissociate in a step wise manner, one proton at a time. • Ka1 > Ka2> Ka3 and this is generally true for all inorganic acids. • Successive ionization constants often decrease by a factor of 104 to 106. • This means that each step in the ionization of a polyprotic acid occurs to a much lesser extent than the previous step.

  18. Acid-Base Properties of Salts • Salt is another name for an ionic compound. • When a salt is dissolved in water, it breaks up into ions. • Under certain conditions, these ions can behave as acids or bases.

  19. Salts that Produce Neutral Solutions • Salts that consist of the cations of strong bases and the anions of strong acids have no effect on the [H+] when dissolved in water. • This means that aqueous solutions of salts such as KCl, NaCl, NaNO3, and KNO3 are neutral (pH=7).

  20. 10 Which salt will be neutral (pH =7)? • NaC2H3O2 • KCN • FeCl3 • NaNO3

  21. 10 Which salt will be neutral? • KCl • NaF • KNO2 • Li3PO4

  22. 10 Which salt will be neutral? • KOCl • NH4OCl2 • KOCl3 • NH4OCl4

  23. Salts that Produce Basic Solutions • For any salt whose cation has neutral properties (such as Na+ or K+) and whose anion is the conjugate base of a weak acid, the aqueous solution will be basic. • The Kb value for the ion can be obtained from the relationship Kb = Kw/Ka.

  24. Example of a Basic Salt • NaC2H3O2 completely ionizes in water to form Na+ and C2H3O2- ions. • Na+ ions have no affinity for H+ ions and cannot produce H+ so they have no effect on the pH • C2H3O2- ions are the conjugate base of a weak acid and have a significant affinity for protons, therefore acting as a base. • C2H3O2- + H2O < -- > HC2H3O2 + OH- • A base reacts with water to produce hydroxide ions and a conjugate acid.

  25. 10 Which of the following salts is basic (pH>7)? • KNO3 • KCN • K2SO4 • KCl

  26. 10 Which salt is basic? • FeCl3 • NaNO2 • NH4ClO3 • Al2(SO4)3

  27. 10 Which salt is basic? • NaOCl • KNO3 • FeSO4 • LiBr

  28. Salts that Produce Acidic Solutions • Salts in which the anion is not a base (from a strong acid) and the cation is the conjugate acid of a weak base produce acidic solutions. • The Ka value for the ion can be obtained from the relationship Ka = Kw / Kb • A second type of salt that produces an acidic solution is one that contains a highly charged metal ion (ex. Al(H2O)63+) • Typically, the higher charge on the metal ion, the stronger the acidity of the hydrated ion.

  29. Example of an Acidic Salt • NH4Cl completely ionizes in water to form NH4+ and Cl- ions. • Cl- ions have no affinity for protons and does not affect the pH. • NH4+ ions are the conjugate acid of a weak base and therefore act as an acid. • NH4+ + H2O < -- > NH3 + H3O+ • An acid reacts with water to produce hydronium ions and a conjugate base.

  30. 10 Which salt is acidic (pH<7)? • NaNO3 • NaC2H3O2 • NH4Cl • KCN

  31. 10 Which salt is acidic? • KNO2 • KNO3 • NaCN • NH4OCl3

  32. 10 Which salt is acidic? • CH3NH3Cl • KF • RbBr • NaSO3

  33. Salts in Which Both Ions Have Acidic or Basic Properties • Salts, such as NH4C2H3O2, consist of two ions that can both affect the pH of the solution (both are ions of weak acids/bases). • We will only consider the qualitative nature of these situations. • A prediction can be made about whether the solution will be acidic, basic, or neutral by comparing the Ka value of the acidic ion with the Kb value of the basic ion. • If Kb > Ka, the solution will be basic. • If Ka > Kb, the solution will be acidic. • If Ka = Kb, the solution will be neutral.

  34. 10 KNO3 is • Acidic • Basic • Neutral

  35. 10 NaNO2 is • Acidic • Basic • Neutral

  36. 10 C5H5NHClO4 is • Acidic • Basic • Neutral

  37. 10 NH4NO2 is • Acidic • Basic • Neutral

  38. 10 KOCl is • Acidic • Basic • Neutral

  39. NH4OCl is • Acidic • Basic • Neutral

  40. 10 KCl is • Acidic • Basic • Neutral

  41. 10 NH4C2H3O2 is • Acidic • Basic • Neutral

  42. 10 CH3NH3Cl is • Acidic • Basic • Neutral

  43. 10 KF is • Acidic • Basic • Neutral

  44. 10 NH4F is • Acidic • Basic • Neutral

  45. 10 CH3NH3CN is • Acidic • Basic • Neutral

  46. 10 An unknown salt is either NaCN, NH4Cl, NH4C2H3O2, or NaOCl3. When 0.100 mole of the salt is dissolved in 1.0 L of solution, the pH is > 7. What is the identity of the salt? • NaCN • NH4Cl • NH4C2H3O2 • NaOCl3

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