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Acid Base Equilibria and Solubility Equilibria

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1. Recognize the effect common ions exert on acid ionization and pH, and mechanisms by which solutions are buffered by ions.

2. Examine acid base titration and the Henderson-Hasselbalch equation and calculate the pH of a system at any stage of titration involving both strong and weak acid bases.

3. Apply Le Chatelier's Principle to explain the effects of common ions and pH on solubility.

4. Describe the effect of complex ion formation on solubility.

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1) Buffer solution: A solution which resists change in pH upon addition of small amount of acid or base to it. They are of two types' acidic buffer and basic buffers.

Acidic buffer is a mixture of weak acid and its salt
Eg : acetic acid buffer

It is a mixture of acetic acid and sodium acetate

If a small amount of acid is added to this system according to Le Chatlier principle
the system will adjust itself to nullify the effect made by us to the equilibrium
If a small amount of HCl is added to this system

HCl dissociate as H + and Cl- , H+ ions are added to this equilibrium (definition of acid = H+ ion donor)
The concentration of H+ ions will increase in the system, the equilibrium ( Ist equation) will shift to the backward direction ie the formation of acetic acid will be preferred there by maintaining the pH constant

Similarly if a small amount of base say NaOH is added (Base is OH- ion donor)

It combines with acetic acid and eliminates a molecule of water and forms acetate ion which is again in the equilibrium (if the any ion in the equilibrium is formed, the system will adjust itself and keep the pH constant)

2) Consider first a buffer solution containing a weak acid HA and its highly dissociated salt NaA. The hydrogen ion concentration of ...

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Aqueous Ionic Equilibrium

1- Consider the following titration curve of a weak acid titrated with strong base. At which point(s) on the graph is the solution a buffer? (see attachment for image)

2- Which of the following solutions has the highest buffer capacity? Choose from:
0.100 M NaCl / 0.100 M NaOH
pure H2O
0.050 M NaCN / 0.050 M HCN
0.100 M HCl / 0.100 M NaOH
0.025 M NaCN / 0.025 M HCN

3- A student adds solid silver chloride (AgCl) to each of two beakers: one containing 1.0 L of pure water, and one containing 1.0 L of 0.500 M NaCl. In which will AgCl be more soluble, and why? Choose from:
0.500 M NaCl: the sodium ions in the solution will complex with the chloride, allowing more AgCl to dissolve.
0.500 M NaCl: in the 0.500 M NaCl solution, the chloride ions that are already present will increase the amount of AgCl that dissolves by decreasing the value of
Neither: AgCl is completely insoluble, and will not dissolve in either.
Pure H2O: in the 0.500 M NaCl solution the chloride ions that are already present will inhibit the AgCl from dissolving by increasing the value of Q.
Both: AgCl is freely soluble in both.

4- Determine the pH of each solution.
0.16M KCHO 2
0.16M KCHO 2
0.25M KI

5- Ammonia, NH 3 , is a weak base with a K b value of 1.8×10 −5 .
What is the percent ionization of ammonia at this concentration?

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