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Prentice Hall 2003Chapter 17 Chapter 17 Additional Aspects of Aqueous Equilibria Slide 2 Prentice Hall 2003Chapter 17 The solubility of a partially soluble acid is decreased when a common ion is added HC 2 H 3 O 2(aq) + H 2 O (l) H 3 O + (aq) + C 2 H 3 O 2 - (aq) Consider the addition of C 2 H 3 O 2 - This is a common ion From a salt such as NaC 2 H 3 O 2 Therefore, [C 2 H 3 O 2 - ] increases and the system is no longer at equilibrium So, [H + ] must decrease (shift leftLeChtelier!) 17.1: The Common Ion Effect Slide 3 Prentice Hall 2003Chapter 17 Slide 4 Prentice Hall 2003Chapter 17 Composition and Action of Buffered Solutions A buffer consists of a mixture of a weak acid (HX) and its conjugate base (X - ): The K a expression is 17.2: Buffered Solutions Slide 5 Prentice Hall 2003Chapter 17 A buffer resists a change in pH when a small amount of OH - or H + is added When OH - is added to the buffer, the OH - reacts with HX to produce X - and water The [HX]/[X - ] ratio remains more or less constant, so the pH is not significantly changed When H + is added to the buffer, X - is consumed to produce HX the pH does not change significantly Slide 6 Text, P. 665 Slide 7 Prentice Hall 2003Chapter 17 Buffer Capacity and pH Buffer capacity is the amount of acid or base neutralized by the buffer before there is a significant change in pH It depends on the composition of the buffer The greater the amounts of conjugate acid-base pair (molar concentration), the greater the buffer capacity The pH of the buffer depends on K a Slide 8 Prentice Hall 2003Chapter 17 Recall: If K a is small (the equilibrium concentration of the undissociated acid is close to the initial concentration), then. the Henderson- Hasselbalch Equation! Slide 9 Prentice Hall 2003Chapter 17 Addition of Strong Acids or Bases to Buffers The amount of strong acid or base added results in a neutralization reaction: X - + H 3 O + HX + H 2 O HX + OH - X - + H 2 O Text, P. 668 Slide 10 Prentice Hall 2003Chapter 17 Problems 3, 5, 9, 15, 17, 19 Slide 11 Prentice Hall 2003Chapter 17 Strong Acid-Strong Base Titrations A plot of pH versus volume of acid (or base) added is called a titration curve Consider adding a strong base (NaOH) to a solution of a strong acid (HCl): 17.3: Acid-Base Titrations Slide 12 Text, P. 672 pH is determined by ? Appropriate indicator: dramatic color change in the desired range Slide 13 Prentice Hall 2003Chapter 17 The equivalence point in a titration is the point at which the acid and base are present in stoichiometric quantities The end point in a titration is the observed point The difference between equivalence point and end point is called the titration error Slide 14 Strong Base-Strong Acid Titrations Add HCl to NaOH Text, P. 674 Slide 15 Prentice Hall 2003Chapter 17 Weak Acid-Strong Base Titrations Consider the titration of acetic acid, HC 2 H 3 O 2 and NaOH Before any base is added, the solution contains only weak acid As strong base is added, the strong base consumes a stoichiometric quantity of weak acid: HC 2 H 3 O 2 (aq) + NaOH(aq) C 2 H 3 O 2 - (aq) + H 2 O(l) Slide 16 Prentice Hall 2003Chapter 17 Text, P. 674 There is an excess of acid before the equivalence point so there is a mixture of weak acid and its conjugate base The pH is given by the buffer calculation First the amount of C 2 H 3 O 2 - generated is calculated, as well as the amount of HC 2 H 3 O 2 consumed (Stoichiometry) Then the pH is calculated using equilibrium conditions (H-H) pH is determined by ? pH is determined by Slide 17 Prentice Hall 2003Chapter 17 Text, P. 675 Slide 18 Prentice Hall 2003Chapter 17 pH is determined by ? Note that pH is above 7 the acetate ion is a weak base Slide 19 Prentice Hall 2003Chapter 17 Compare initial pH values Compare pH values at eq. points Compare pH change near eq. points Weak Acid/Strong Base CurveStrong Acid/Strong Base Curve Slide 20 Prentice Hall 2003Chapter 17 The influence of acid strength on the shape of the curve for the titration with NaOH Text, P. 676 Slide 21 Text, P. 677 The titration of a weak base with a strong acid Slide 22 Prentice Hall 2003Chapter 17 Titrations of Polyprotic Acids In polyprotic acids, each ionizable proton dissociates in steps Therefore, in a titration there are n equivalence points corresponding to each ionizable proton In the titration of H 3 PO 3 with NaOH, The first proton dissociates to form H 2 PO 3 - Then the second proton dissociates to form HPO 3 2- Slide 23 Text, P. 677 Slide 24 Prentice Hall 2003Chapter 17 Problems 25, 27, 29, 31, 33 Slide 25 Prentice Hall 2003Chapter 17 The Solubility-Product Constant, K sp Consider equilibria that are heterogeneous Some common applications: Tooth enamel and soda, salts and kidney stones, stalactites and stalagmites Example: for which K sp is the solubility product constant 17.4: Solubility Equilibria Slide 26 Prentice Hall 2003Chapter 17 In general: the solubility product is the molar concentration of ions raised to their stoichiometric powers Solubility is the amount (grams) of substance that dissolves to form a saturated solution Affected by pH concentrations of other ions in solution Molar solubility is the number of moles of solute dissolving to form a liter of saturated solution Slide 27 Prentice Hall 2003Chapter 17 Solubility and K sp To convert solubility to K sp : Solubility needs to be converted into molar solubility (via molar mass) Molar solubility is converted into the molar concentration of ions at equilibrium (equilibrium calculation) K sp is the product of equilibrium concentration of ions Text, P. 697 Slide 28 Prentice Hall 2003Chapter 17 Sample Problems # 37 & 39 Slide 29 Prentice Hall 2003Chapter 17 The Common Ion Effect Solubility is decreased when a common ion is added Le Chteliers principle: as F - is added (from NaF), the equilibrium shifts away from the increase CaF 2 (s) is formed and precipitation occurs 17.5: Factors that Affect Solubility Slide 30 Prentice Hall 2003Chapter 17 Solubility and pH If the F - is removed, then the equilibrium shifts right and CaF 2 dissolves F - can be removed by adding a strong acid: As pH decreases, [H + ] increases and solubility increases The effect of pH on solubility is dramatic The more basic the anion, the more solubility is influenced by pH Slide 31 Prentice Hall 2003Chapter 17 Formation of Complex Ions The formation of Ag(NH 3 ) 2 + : The Ag(NH 3 ) 2 + is called a complex ion NH 3 (the attached Lewis base) is called a ligand Lewis bases share their nonbonded electron pairs with vacant orbitals on the metal atom The equilibrium constant for the reaction is called the formation constant, K f : Slide 32 Prentice Hall 2003Chapter 17 Text, P. 687 Slide 33 Prentice Hall 2003Chapter 17 Amphoterism Amphoteric oxides will dissolve in either a strong acid or a strong base Examples: hydroxides and oxides of Al 3+, Cr 3+, Zn 2+, and Sn 2+ The hydroxides generally form complex ions with four hydroxide ligands attached to the metal: Slide 34 Prentice Hall 2003Chapter 17 Hydrated metal ions act as weak acids Thus, the amphoterism is interrupted: Slide 35 Prentice Hall 2003Chapter 17 Problems 41, 43, 49 Slide 36 Prentice Hall 2003Chapter 17 At any instant in time, Q = [Ba 2+ ][SO 4 2- ] If Q > K sp, precipitation occurs until Q = K sp If Q = K sp, equilibrium exists If Q < K sp, solid dissolves until Q = K sp Based on solubilities, ions can be selectively removed from solutions 17.6: Precipitation and Separation of Ions Slide 37 Prentice Hall 2003Chapter 17 Selective Precipitation of Ions Ions can be separated from each other based on their salt solubilities Example: if HCl is added to a solution containing Ag + and Cu 2+ the silver precipitates (K sp for AgCl is 1.8 10 -10 ) while the Cu 2+ remains in solution Removal of one metal ion from a solution is called selective precipitation Slide 38 Prentice Hall 2003Chapter 17 Problems 51, 53, 55 Slide 39 Qualitative analysis is designed to detect the presence of metal ions Quantitative analysis is designed to determine how much metal ion is present See Text, P. 692-695 17.7: Qualitative Analysis for Metallic Elements Slide 40 Prentice Hall 2003Chapter 17 End of Chapter 17 Additional Aspects of Aqueous Equilibria