Presentation on theme: "Applications of Aqueous Equilibria Electrolyte Effect Chapter 8."— Presentation transcript:
Applications of Aqueous Equilibria Electrolyte Effect Chapter 8
Ideal vs. real ionic solutions Ideal solution: Cations and anions do not interact. Real solution: Electrostatic interactions Shaded region: Solvation cage ideal real
Ionic Strength The effect of added electrolyte on equilibria is independent of the chemical nature of electrolyte, but depends on a property of the solution called the ionic strength. [A], [B], [C]…: molar species concentration species Z A, Z B, Z C..: ionic charge
Molality Seawater (SW)Lake Water (LW) Na x Mg x 10 ‑ 3 Ca x K x Cl x SO x HCO x I SW = 1/2 (m Na × m Mg × m Ca × m K × m Cl × m SO4 × m HCO3 × 1 2 ) = 0.72 mol kg -1 I LW = = 1.5 × mol kg ‑ 1
Salt Effect The consequence of this effect is a decrease in overall attraction between barium and sulfate ions and an increase in solubility. Ba +2 SO 4 -2 Cl - Na + Small net negative chargeSmall net positive charge BaSO 4 dissolved in NaCl (aq)
The effect of added NaCl to increase the size of the K sp for BaSO 4 At 0 M NaCl, K sp has a value of 1.1 x At 1 x M NaCl, K sp has a value of approximately 1.8 x At 1 x M NaCl, K sp has a value of approximately 2.85 x
The effect of added NaCl to increase the size of the K a for acetic acid At 0 M NaCl, K a has a value of 1.75 x At 1 x M NaCl, K a has a value of approximately 2.1 x At 1 x M NaCl, K a has a value of approximately 2.7 x
Activity Coefficients The activity, or effective concentration, of species X depends on the ionic strength of the medium and is defined as: a X = x [X] a X : activity x : activity coefficient The a X & x vary with ionic strength. ex. X m Y n precipitate
Properties of Activity Coefficients In very dilute solutions, where the ionic strength is minimal, the effectiveness becomes constant. X →1 and a X →[X], then K’ sp →K sp As ionic strength increases, an ion loses some of its effectiveness and its activity coefficient decreases. At high ionic strength, > 0.1m, often increases and may even becomes greater than unity.
Properties of Activity Coefficients In solutions that are not too concentrated, the for a given species is independent of the nature of the electrolyte and dependent only on the ionic strength. The activity coefficient of an unchanged molecule is approximately unity, regardless of ionic strength.
Debye-H ü ckel Equation X =activity coefficient Z X =charge =ionic strength X =effective diameter of the hydrated ion X in nm
Ion α X, nm H3O+H3O+ 0.9 Li +, C 6 H 5 COO Na +, IO 3 -, HSO 3 -, HCO 3 -, H 2 PO 4 -, H 2 ASO 4 -, OA C OH -, F -, SCN -, HS -, ClO 3 -, ClO 4 -, BrO 3 -, IO 4 -, MnO K +, Cl -, Br -, I -, CN -, NO 2 -, NO 3 -, HCOO Rb +, Cs +, Tl +, Ag +, NH Mg 2+, Be Ca 2+, Cu 2+, Zn 2+, Sn 2+, Mn 2+, Fe 2+, Ni 2+, Co 2+, Phthalate Sr 2+, Ba 2+, Cd 2+, Hg 2+, S Pb 2+, CO 3 2-, SO 3 2-, C 2 O Hg 2 2+, SO 4 2-, S 2 O 3 2-, CrO 4 2-, HPO Al 3+, Fe 3+, Cr 3+, La 3+, Ce PO 4 3-, Fe(CN) Th 4+, Zr 4+, Ce 4+, Sn Fe(CN) Activity Coefficients for Ions at 25 ℃
General Rules for Activity Coefficients i 1 as 0 i.e. activity = concentration at infinite dilution i is decreased as increased i.e., the free ion activity coefficient decreases with ionic strength 2+ < +, the activity corrections decrease with increasing charge
Solutions of Acids or Bases Containing a Common Ion Common ion effect: The shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction.
Buffered Solutions A buffered solution is one that resists a change in pH when either hydroxide ions or protons are added. A buffered solution may contain a weak acid and its salt or a weak base and its salt. (HF+NaF, NH 3 +NH 4 Cl)
How Do the H + /OH - Ions Work in Buffered Solutions The equilibrium concentration of H + and the pH are determined by the ratio [HA]/[A - ].
The Effect of Added Bases When OH - are added, HA is converted to A -, causing the ratio [HA]/[A - ] to decrease. If the amount of HA and A - originally present are very large compared with the amount OH - added, the change in [HA]/[A - ] ratio is small.
The Effect of Added Acids When protons are added to a buffered solution, the added H + ions react with A - to form the weak acid. If [HA] and [A - ] are large compared with the [H + ] added, only a slight change in the pH occurs.
Exact Treatment of Buffered Solutions charge balance
[HCOOH]=0.4 M, [HCOONa]=1.0 M K a =1.8x10 -4
K b =1.75X10 -5 [NH 3 ]=0.2 M, [NH 4 Cl]=0.3 M
Unique Properties of Buffer Solution The Effect of Dilution － The pH of a buffer solution remains essentially independent of dilution. The Effect of Added Acids and Bases － A buffer solution resists pH change after addition of small amounts of strong acids or bases Buffer Capacity
The Effect of Dilution
Buffer Capacity The number of moles of a strong acid or a strong base that causes 1L of buffer to undergo a 1 unit change in pH. The capacity of a buffered solution is determined by the magnitudes of [HA] and [A - ].
Calculate Buffer Capacity 1. Calculate the buffer capacity (B) for a mixture of 0.01 moles of acetic acid and 0.03 moles of NaOAc in 100 mL of total solution. 2. Calculate the buffer capacity (B) for a mixture of 0.03 moles of acetic acid and 0.03 moles of NaOAc in 100 mL of total solution.
Solution 1. B=2.303 × (0.1×0.3)/( ) =0.172 mol/L per pH 2. B=2.303 ×(0.3×0.3)/( ) =0.345 mol/L per pH
Prepare an Optimum Buffered Solution The optimum buffering will occur when [HA] is equal to [A - ]. Reasons: (1) The pK a of the weak acid selected for the buffer should be as close as possible to the desired pH. (2) It can provide best buffer capacity.
Titrations and pH Curves The equivalence point is defined by stoichiometry, no by the pH The pH value at equivalence point is affected by the acid strength or base strength. The strength of a weak acid or weak base have significantly effect on the shape of pH curves.
Titrations of Weak Acid with Strong Bases Titration Curve Calculations A stoichiometry problem: The reaction of hydroxide ion with the weak acid is assumed to run completion. An equilibrium problem: The position of the weak acid equilibrium is determined, and the pH is calculated.
1. No NaOH has been added ml of 0.1M NaOH has been added ml of 0.1M NaOH has been added ml of 0.1M NaOH has been added ml of 0.1M NaOH has been added ml of 0.1M NaOH has been added ml of 0.1M NaOH has been added. The titration of 50ml of 0.1M acetic acid (K a =1.8 × ) with 0.1M NaOH
Titration of Polyprotic Acids
Titration of Triprotic Acids If the above reaction is the only important reaction involving these species, then [H 3 A]=[HA -2 ]
Titration of Polyprotic Acids Amphiprotic Salts
K b2 >>K a2, the solution is basic
Titration of Diprotic Acids Point A: Treat the system as if it contained a single monoprotic acid. (K a1 >>K a2 ) Region B: Treat the system as if a simple buffer solution consisting of the H 2 A and NaHA. Point C: Use equation Region D: The second buffer solution consisting of the HA - and Na 2 A. Point E: Use equation Region F: The excess NaOH dominated the pH.
Question A mL sample of the weak acid H 3 A (0.100 M) is titrated with M NaOH. What are the major species after 40.0 mL of M NaOH is added in the titration (water is always assumed to be a major species)? H 3 A H 2 A –, HA 2– H 3 A, H 2 A – HA 2– H 2 A –
Answer c)H 3 A, H 2 A – Section 8.7, Titration of Polyprotic Acids The concentrations are equal, so the first equivalence point is at mL of NaOH.
Question A mL sample of the weak acid H 3 A (0.100 M) is titrated with M NaOH. What are the major species after mL of M NaOH is added in the titration (water is always assumed to be a major species)? H 3 A H 2 A –, HA 2– H 3 A, H 2 A – HA 2– H 2 A –
Answer b) H 2 A –, HA 2– Section 8.7, Titration of Polyprotic Acids The concentrations are equal, so the second equivalence point is at mL of NaOH.
Determine the Equivalence Point of an Acid-Base Titration Use a pH meter to monitor the pH and then plot a titration curve. Use an acid-base indicator, which marks the endpoint of a titration by changing color.
Acid-Base Indicators Add a few drops of the phenolphthalein indicator to a acidic solution. (pH=1) The ratio shows that the predominant form of the indicator is HIn, resulting in a colorless solution.
As OH - is added to this solution, [H + ] decreases and the equilibrium shift to right, changing HIn to In -. For most indicators, about 1/10 of the initial form must be converted to the other form before a new color is apparent.
Application of Henderson- Hasselbalch Equation
Bromthymol Blue Ken O'Donoghue
pH Ranges of Common Indicators
Solubility Equilibria and the Solubility Product:Common Ion Effect x x Initial Equilibrium
Relative Solubilities Relative solubilities can be predicted by comparing K sp values only for salts that produce the same total number of ions. K sp : CaSO 4 >CuI>AgI Relative solubilities: CaSO 4 >CuI>AgI
Precipitation Use the ion product to predict whether a precipitation will form. For CaF 2 Q=[Ca +2 ][F - ] 2 Q>K sp, precipitation occurs Q
Qualitative Analysis Group I-insoluble chlorides When dilute HCl is added to a solution containing a mixture of the common cations, only Ag +, Pb +2 and Hg 2 +2 will precipitate as insoluble chlorides. All other chlorides are soluble and remain in solution.
Group II-sulfides insoluble in acid solution H 2 S is added to precipitate those ions Hg +2, Cd +2, Bi +3, Cu +2 and Sn +4. Low concentration of S -2 is in the weak acidic solution.
Group III-sulfides insoluble in basic solution The solution is made basic and more H 2 S is added. A basic solution produces a higher [S -2 ]. The cations precipitate as sulfides at this stage are Co +2, Zn +2, Mn +2, Ni +2 and Fe +2. If Cr +3 and Al +3 ions are present, they will also precipitate.
Group IV-insoluble carbonates At this point, all the cations have been precipitated except those from Groups 1A and 2A of the periodic table. The Group 2A cations form insoluble carbonates and can be precipated by the addition of CO 3 -2.(eg Ca +2 and Mg +2 )
Group V-alkali metal and ammonium ions The only ions remaining in solution at this point are the Group 1A cations and the NH 4 + ion, all of which form soluble salts with the common anions.
Complex Ion Equilibria A complex ion is a charged species consisting of a metal ion surrounded by ligands. Formation constants K 1 =2.1X10 3 K 2 =8.2X10 3