Chapter 15 Applications of Equilibrium

Common Ions An ion that is present in both –An acid and its conjugate base HNO 2 and NaNO 2 –A base and its conjugate acid NH 3 and NH 4 Cl This common ion effects the position of the equilibrium and the pH

How will the pH of a solution containing 0.10M HNO 2 compare to a solution containing 0.10M HNO 2 and 0.10M NaNO 2 ? Why?

Points To Remember In common ion situations you MUST consider the initial concentrations of any common ions in solution before you begin the problem The larger Ka or Kb will govern the equilibrium

Calculate the pH of a solution containing 0.10M HNO 2 and 0.10M NaNO 2 ? (Ka = 4.0x10 -4 )

Homework P. 774 #’s 21,23a&d, 34,36

Buffered Solutions Solutions that resists changes in pH when an acid or base is added Made of –Weak acid and its conjugate base –Weak base and its conjugate acid Made by adding both chemicals or reacting

Adding.1M HF NaF

Reacting.1M HF.1M NaOH

Determine the pH of a solution made by dissolving NaHCO 3 to 0.25M in M H 2 CO 3 (Ka = 4.3x10 -7 )

Henderson-Hasselbalch Equ. A way of calculating buffer problems ONLY USED 4 BUFFERS!!!!!!

HA  H + + A -

Determine the pH of a solution made by dissolving NaHCO 3 to 0.25M in M H 2 CO 3 (Ka = 4.3x10 -7 ) Use H-H Equ.

How Do Buffers Work Strong acid or base is not allowed to build up in solution Strong acid is replaced by weak acid Strong base is replaced by weak base –The strong acid or base reacts completely!!! –Then equilibrium kicks in

Acidic Buffers HF  H + + F - Add NaOH Add HCl

Basic Buffer NH 3 + H 2 O  NH OH - Add NaOH Add HCl

Determine the pH after 0.20 grams of NaOH is added to 500.mL a solution made by dissolving NaHCO 3 to 0.25M in M H 2 CO 3 (Ka = 4.3x10 -7 )

What Makes a Good Buffer? The best buffers have acids and bases at the same, large concentration –Added acid or base has little effect on pH –When the concentrations are the same pH=pKa

Homework P. 774 #’s 23d,31,37ab, 40, 42,46,48,49

Titrations The controlled addition of a chemical of known concentration to a chemical of unknown concentration Known Chemical – Standard or titrant Unknown Chemical - Analyte

Titrations are carried out to the equivalence point –Point when mole ratio of chemicals is reached –Often an invisible point Marked by and indicator Indicator is a chemical that changes color in different pH’s Indicator color change marks the endpoint –Hopefully endpoint and equivalence point are the same

Titration Curves Plotting of the pH of the titration solution as solution is added Used as a way of determining equivalence and Ka or Kb values Specific ranges for specific acid/base combinations –Strong/Strong or Strong/Weak

Titration Curve HCl vs NaOH

Titration Curve Problems Remember the process. Do NOT memorize! Treat each step as a separate problem –For Strong/Strong systems just stoichiometry For Strong/Weak First stoichiometry then equilibrium

Strong Acid / Strong Base Beginning – Just the pH of the chemical you have Before Equivalence – Stoichiometry Equivalence – pH is 7.00 After Equivalence - Stoichiometry

Strong Acid / Strong Base Titration #54 page mL of 0.100M Ba(OH) 2 is titrated with 0.400M HCl 0.00 mL 20.0 mL 30.0 mL 40.0 mL 80.0 mL

Weak / Strong Titration Beginning – Just the pH of the chemical you have (Might need equilibrium) Before Equivalence – Stoichiometry then equilibrium Equivalence – pH is NOT 7 – Stoichiometry then equilibrium After Equivalence - Stoichiometry

Weak / Strong Combos Weak Acid / Strong Base Equivalence Point is ABOVE seven –Conjugate base of a weak acid controls pH Weak Base / Strong Acid Equivalence Point is BELOW seven –Conjugate acid of a weak base controls pH At halfway pH = pKa or pOH = pKb!!!!!!!!

Weak Base / Strong Acid Titration

Weak Acid / Strong Base Titration

Strong Base/ Weak Acid Titration #55 page mL of 0.200M Acetic Acid is titrated with 0.100M KOH 0.00 mL 50.0 mL mL mL mL

Homework P. 775 #’s 52,53,56

Which Indicator to Use? Choosing the appropriate indicator is important when doing titrations There are so many good ones to use So how does a chemist know???????

Which Indicator to Use? Indicators are weak acids and change color with varying pH Pick the indicator that changes a little above your calculated equivalence point. Pick indicator with pKa value close to equivalence point Change color +/- 1 from pKa

Example Choose an indicator for the titration of HCl with NH 3 pH=5.36

Polyprotic Acids These acids have multiple equivalence points Carbonic acid has two Phosphoric acid has three Multiple step points on a titration curve

Diprotic Acid Titration

Molar Mass of an Acid You can find the molar mass of an acid by titration too Titrate a acid of known mass with a standard to determine number of moles Divided grams by moles –Molar Mass

Solubility Equilibrium Dissolving is an equilibrium process Chemicals that we say are insoluble are actually very slightly soluble –Meaning only a small amount of the solute dissolves –Concentrations are very small M to M are common

Solubility Equilibrium Iron (II) Sulfide is “insoluble” by solubility rules However, it still dissolved to a small extent Consider the reaction FeS (s)  Fe +2 + S -2 Since dissolving is an equilibrium process we can write an equilibrium expression Keq = [Fe +2 ] [S -2 ] FeS is omitted because it is a solid

Solubility Product Constant The previous equilibriums expression is a “special” type of equilibrium expression Ksp – solubility product constant Ksp = [Fe +2 ] [S -2 ] –Only used for slightly soluble salts –Never includes the reactants –Ksp’s are very small –Check out the table 15.4 page 753

Ksp Values Lead (II) Chloride 2.4X10 -4 Strontium Carbonate 3.8X10 -9 Nickel (II) Hydroxide 5.5X Copper (II) Hydroxide 2.2X Cadmium Sulfide 8.0X Silver Sulfide 6.0X And my personal favorite Bismuth Sulfide 1.6x10 -72

Ksp Values Q) What does a small Ksp value mean? A) Low concentration of ions Small solubility Q) Is the compound with the smallest Ksp the least soluble? A) Not necessarily –There are different numbers of ions that changes the expression

Solubility of Salts We can easily compare solubility for salts that have the same # of ions When there are the same # of ions the salt with the smallest Ksp is least soluble Which salt is least soluble, most soluble AgCl, FeS, BaSO 4 – FeS 6.0x Least soluble –AgCl 1.6x Most soluble –BaSO 4 1.1x Middle

Calculating Solubility We can calculate solubility of salts and the concentration of the ions in solutions from Ksp Deal with saturated solutions –The salt has dissolved as much as it can Ksp has been reached In saturated solutions the concentrations of the ions are related to the mole ratio

Example Calculate the solubility of silver chloride Ksp=1.6x10 -10

Example Calculate the solubility of mercury (II) sulfide and the concentration of each ion in a saturated solution Ksp=1.6x10 -54

Example Calculate the solubility of calcium phosphate Ksp=1.3x10 -32

Example A saturated solution of Iron (III) Hydroxide has a concentration of Fe +3 of 1.8x What is the Ksp of Iron (III) Hydroxide?

Common Ion Revisited Anyone ever have a barium shake for a scan? Contains barium sulfate. –Barium ions are very toxic to the body –Treated as calcium How can you have fewer barium ions in solution? Add Sodium Sulfate

What is the concentration of barium ions in a saturated solution of barium sulfate? –Ksp = 1.1x What is the concentration of barium ions in a when 0.10M Sodium Chloride is added to a saturated solution of barium sulfate

Precipitation If you combine Silver Nitrate with Sodium Chloride will a precipitate always form? Only if Ksp is exceeded Will a ppt. form if 2x10 -5 M Silver Nitrate is mixed with 3x10 -3 M Sodium Chloride? Ksp of AgCl is 1.6x No