AqueousEquilibria

Chapter 17Additional Aspects of Aqueous

Equilibria

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know there is a little love for the central science!

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The Common ion effect How do they effect dissociation?

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Compare:

• Calculate the pH of a 0.25 M propionic acid solution (Ka=1.3 x 10-5)

• Calculate the pH of a 0.25 M propionic acid solution that also has 0.10 M sodium propionate add.

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Common Ion Effect

• How does LeChatelier support the previous calculations?

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Common Ion Effect

• Summarize the Common Ion Effect:

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Buffers:

• Solutions of a weak conjugate acid-base pair.

• They are particularly resistant to pH changes, even when strong acid or base is added.

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How do Buffers Resist pH Changes?

• Consider a buffer composed of equal concentrations of nitrous acid and nitrite ion.

• 1) Does this meet the criteria for a buffer? • 2) What would happen if a volume of HCl was added to

the buffer?

• 3) What would happen if a volume of NaOH was added to the buffer

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How do Buffers Resist pH Changes?

• The pH of a buffer will change somewhat, but not significantly. The Balance between the conjugate acid/base pair is

disrupted Either the conjugate acid or the conjugate base will be

present in a higher concentration after the addition This will cause a minor change to the pH

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Calculating the pH of a Buffer Solution

• In order to calculate the pH of a buffer solution you will need to use the Henderson-Hasselbalch Equation:

pH = pKa + log[base][acid]

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Henderson–Hasselbalch Equation

What is the pH of a buffer that is 0.12 M in benzoic acid and 0.20 M in sodium benzoate? Ka for benzoic acid is 6.3 10−5.

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Another Application of the Henderson-Hasselbalch Equation

• How many moles of NH4Cl must be added to 2.0 L of 0.10 M NH3 to form a buffer whose pH is 9.00? (Assume that the addition of NH4Cl does not change the volume of the solution.) Kb for NH3 is 1.8 x 10-5.

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pH Range and Buffer Capacity

• The pH range is the range of pH values over which a buffer system works effectively.

• It is best to choose an acid with a pKa close to the desired pH.

• Buffer Capacity is the amount of acid or base a buffer can neutralize before significant pH changes The higher the molarity or volume of the conjugate

pairs, the greater the capacity of the buffer.

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Calculating pH Changes in Buffers

A buffer is made by adding 0.300 mol HC2H3O2 and 0.300 mol NaC2H3O2 to enough water to make 1.00 L of solution. Calculate the original pH of the buffer and the pH after 0.020 mol of NaOH is added. Ka of acetic acid is 1.8 x 10-5

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Homework

• Ch. 17: 15, 17, 21, 23, 25

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Solubility Product Constant

Using Equilibrium to Determine the dissociation of a solid in solution.

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Solubility Products

Consider the equilibrium that exists in a saturated solution of BaSO4 in water:

Write the Equilibrium Expression for this reaction.

BaSO4(s) Ba2+(aq) + SO42−(aq)

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Solubility Products

The equilibrium constant expression for this equilibrium is

Ksp = [Ba2+] [SO42−]

where the equilibrium constant, Ksp, is called the solubility product.

The solubility product defines the dissociation of the solid in solution

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Practice Problem

• Write separate expressions for the solubility product constant for CaF2 and Silver Sulfate

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Practice Problem

• Determine the concentration of each ion in a saturated solution of zinc hydroxide. Zinc hydroxide has a Ksp=3.0 x 10-16

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Solubility Products

• Ksp is not the same as solubility.

• Solubility is generally expressed as the mass of solute dissolved in 1 L (g/L) or 100 mL (g/mL) of solution, or in mol/L (M).

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Ksp from Solubility Data

• Solid silver chromate is added to pure water at 25°C. Some of the solid remains undissolved at the bottom of the flask. Analysis of the equilibrated solution shows that its silver ion concentration is 1.3 10–4 M. Assuming that Ag2CrO4 dissociates completely in water and that there are no other important equilibria involving the Ag+ or CrO4

2– ions in the solution, calculate Ksp for this compound.

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Solubility from Ksp

• The Ksp for CaF2 is 3.9 10–11 at 25°C. Assuming that CaF2 dissociates completely upon dissolving and that there are no other important equilibria affecting its solubility, calculate the solubility of CaF2 in grams per liter.

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Refresher Problems

Getting our mind back into Chemistry Mode.

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Equilibrium Expressions

• Write equilibrium expressions for the following reactions:

• 1) Ni(OH)2(s) Ni2+(aq) + 2 OH-(aq)

• 2) 2NOBr(g) 2 NO(g) + Br2(g)

• 3) HClO3(ag) H+(aq) + ClO3-(aq)

• 4) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

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pH

• What is the pH of a solution that consists of 250 ml of 0.75 M hydrofluoric acid and 2.1 grams of sodium fluoride? (Ka=6.8 x 10-4)

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Solubility

• What’s the molar concentration of each ion in an equilibrated solution of lead (II) fluoride. Ksp of lead (II) fluoride= 3.6 x 10-8

• Can we figure out the pH of this solution?

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Buffer Solutions

• What is the pH of a 500.0 mL buffer solution that consists of 1.25 M acetic acid and 1.00 M sodium acetate if 15.0 mL of 0.750 M nitric acid is added to it? Ka of acetic acid is 1.80 x 10-5

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Homework

Read and Take Notes on section 17.3 Acid-Base Titrations

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Factors that Affect Solubility

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Factors that Affect Solubility

• From your understanding of equilibrium and how it relates to acids/bases, what would be some factors that either increase or decrease the dissociation of a solid in solution?

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Factors Affecting Solubility

1) The Common-Ion Effect If one of the ions in a solution equilibrium

is already dissolved in the solution, the equilibrium will shift to the left and the solubility of the salt will decrease.

BaSO4(s) Ba2+(aq) + SO42−(aq)

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Factors Affecting Solubility

2) pH If a substance has a

basic anion, it will be more soluble in an acidic solution.

Substances with acidic cations are more soluble in basic solutions.

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Factors Affecting Solubility

• pH Explain how the

solubility of Mg(OH)2 would be affected by the presence of an acid.

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Practice Problem

• Which of the following substances will be more soluble in acidic solution than in basic solution:

(a)Ni(OH)2(s)

(b) CaCO3(s)

(c) AgCl(s)

(d) BaF2(s)

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Evaluating a Solution

Will more solid dissolve or will a precipitate form?

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Will a Precipitate Form?

• In a solution, If Q = Ksp, the system is at equilibrium

and the solution is saturated. If Q < Ksp, more solid will dissolve until

Q = Ksp.

If Q > Ksp, the salt will precipitate until Q = Ksp.

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Practice Problem

• Will a precipitate form when 0.10 L of 8.0 10–3 M Pb(NO3)2 is added to 0.40 L of 5.0 10–3 M Na2SO4? (Ksp for PbSO4=6.3 x 10-9)