Chapter 17

Additional Aspects of Solubility Equilibria

17.1 The Common Ion Effect

17.2 Buffer Solutions

17.2 How Buffers Work

17.2 Calculating pH of Buffer Solutions

17.2 Buffer Example Problem

17.2 Choosing the Proper Buffer Solution

17.2 Buffer Example Problem #2

17.3 Acid-Base Titrations

17.3 Weak Acid-Strong Base Titration pH Before Base Added

17.3 Weak Acid-Strong Base Titration pH After Base is Added

17.3 Weak Acid-Strong Base Titration pH After More Base is Added

17.3 Weak Acid-Strong Base Titration pH at Endpoint

Weak Acid-Strong Base Titration pH Beyond Endpoint

17.3 Weak Acid-Strong Base Titration Curve

In chapter 4 we defined solubility as the amount of a substance that dissolves in a given quantity of solvent at a given temperature to form a saturated solution. We approached solubility in a very simple manner by classifying salts as soluble or insoluble and solubility guidelines were established indicating if a precipitate would form when two solution are combined.

17.4 Overview of Solubility

Experimentally, we can consider solubility equilibria to make predictions about the amount of a given salt that will dissolve. First, we need to define what we mean by solubility and we need to consider what we already know about solubility and compare that to what a chemist thinks about when they hear the term solubility. When chemists discuss solubility they do so in terms of the solubility product constant. Keep in mind that the solubility product constant and molar solubility are two completely different terms. As chemists, we are able to calculate one from the other, but directly comparing the two without a calculation can lead to some improper assumptions. It is very important to recognize the similarities/differences between solubility in terms of the amount (in grams) that dissolve per liter of solution, the molar solubility, and the Ksp.

17.4 The Solubility Product Constant

Example Problem: 1.4 x 10-6 grams of ZnCO3 dissolve in 1.000 mL of solution and 2.8 x 10-6 grams of BaCrO4 dissolve in 1.000 mL of solution, which one has the largest Ksp? Solution: They are the same. Click below to see the explanation.

One common mistake students make is to simply look at the Ksp and think that constant tells them everything they need to know about the solubility when two slightly soluble salts are compared. If you are thinking this way, then you certainly aren’t thinking like a chemist. A chemist would think through the process shown in the example below.

17.4 Ranking the solubility of slightly soluble salts given the Ksp Part I

17.4 Ranking the solubility of slightly soluble salts given the Ksp Part II

Once you have a handle on saturated aqueous solutions, we can start to analyze and manipulate the solubility of these solutions and determine the criteria for precipitation.

17.6 Criteria For Precipitation

With a background on precipitate formation, we can now predict if a precipitate will form when two solutions are combined.

17.6 If Two Solutions are Mixed Will a Precipitate Form?

In certain cases, it is beneficial to separate a mixture of ions in solution. By utilizing the Ksp values and performing calculations, we can calculate the concentration needed to give the best separation of ions in solution.

17.6 Order of Precipitation, Minimum Concentration Needed to Facilitate Precipitation, and Best Separation

Chemists love to manipulate things. Ksp expressions are equilibrium expressions and as Le Chatlier showed us, equilibria can be manipulated if you apply a stress. There are many factors that influence solubility and the factors we will investigate in this class are: Common Ion, pH, Complex Ion Formation, and Amphoterism. In all of these effects, it is very important to WRITE OUT THE EQUILIBRIUM EXPRESSION so you can properly analyze how the solubility will be influenced.

The Common Ion effect is simply a direct application of Le Chatlier’s principle.

17.5 Common Ion Effect

The pH of a solution has a dramatic influence of the solubility. When investigating pH effects, it is important for you to know the strong acids, which will allow you to identify all the neutral anions in solution.

17.5 pH Effects

17.5 How does adding acid/base influence solubility?

Le Chatlier’s principle does a great job of explaining the solubility in terms of the common ion effect and pH effects, but in some instances some puzzling results are obtained. For instance, when concentrated NH3 is added to a saturated solution of Zn(OH)2 its solubility increases. Now we need to determine why this is, and if Le Chatlier’s principle is still valid.

17.5 Does zinc hydroxide follow the rules we’ve discussed so far?

It turns out Le Chatlier’s principle still applies, but there is something else going on in solution responsible for these results. The concept responsible for the difference in solubility is complex ion formation.

17.5 Complex ion formation and coordination complexes

This allows us to investigate zinc hydroxide in a more complex way and it allows us to perform several calculations based on complex ion formation.

17.5 Re-analyzing zinc hydroxide

17.5 Solubility of zinc hydroxide in 15 M NH3

17.5 Determining the concentration of free metal cations in solution

Amphoterism is a special case where a slightly soluble salt has its solubility increased due to pH effects in acidic solution and has its solubility increased due to complex ion formation under basic conditions.

17.5 Amphoterism

17.5 Solubility of Al(OH)3 in 15 M NH3

17.5 Molar Solubility of Al(OH)3 in 15 M NH3 continued

17.5 Amphoteric Effects on Solubility

A common laboratory experiment constructed to illustrate the principles of solubility is Qualitative Analysis. This experiment is designed to answer what is present. Unfortunately, there is no spot test for each individual cation, so a scheme was developed, which groups cations based on their solubility characteristics. Each ion is then isolated and identified.

17.7 Applied Qualitative Analysis Scheme

The Group I cations precipitate as chlorides under acidic conditions.

17.7 Group I Separations

17.7 Why does our HCl need to be cold and dilute?

17.7 Group I Analysis

The theory behind the separations of the Group II and Group III encompass all these effects and utilizes the solubility of the sulfide ion in solution.

17.7 Group II and Group III Sulfide Solubility

17.7 Will FeS precipitate?

17.7 At what pH will FeS begin to precipitate?

One Response to Chapter 17

  1. Robert Burkhart says:

    I am in another Chem 1220 course at OSU and found your lectures extremely helpful. Thanks!

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