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Chapter 16: Problem 2

Devise as many ways as you can to experimentally determine the \(K_{\mathrm{sp}}\) value of a solid. Explain why each of these would work.

Short Answer

Expert verified
There are multiple experimental methods to determine the \(K_{sp}\) value of a solid, such as: 1. Direct measurement of ion concentrations using a spectrophotometer. 2. Conductometric measurements of ion concentrations. 3. Ion-selective electrode measurements. 4. Titration. Each method involves measuring the concentrations of the constituent ions of the solid in equilibrium, thus allowing for the calculation of the \(K_{sp}\) value using the equation \(K_{sp} = [\textrm{cation}]^m [\textrm{anion}]^n\). The methods differ in the techniques used for measuring ion concentrations, providing flexibility for experimenters.

Step by step solution

01

Method 1: Direct measurement of ion concentrations using a spectrophotometer

In this method, a known amount of solid is dissolved in a known volume of water, and the resultant solution is allowed to equilibrate. After reaching equilibrium, the concentrations of the dissolved ions can then be determined using a spectrophotometer, which measures the light absorbed by the ions present in the solution. By knowing the concentrations of the constituent ions, we can subsequently calculate the \(K_{sp}\) value using the equation: \[K_{sp} = [\textrm{cation}]^m [\textrm{anion}]^n\] Where m and n are the stoichiometric coefficients of the cation and anion in the balanced dissolution equation of the compound.
02

Method 2: Conductometric measurements of ion concentrations

This method involves dissolving a known quantity of the solid in a solution and measuring the increase in its conductivity as the ionic compound dissociates. We can then convert these conductivity values into concentrations using a calibration curve (determined by measuring the conductivities of standard solutions with known ion concentrations). By knowing the concentrations of the constituent ions at equilibrium, the \(K_{sp}\) can be calculated in the same manner as described in Method 1.
03

Method 3: Ion-selective electrode measurements

Ion-selective electrodes (ISEs) allow for the direct measurement of the concentration of a specific ion in a solution. In this method, a known quantity of the solid is dissolved in a solution, and the concentrations of the constituent ions at equilibrium can be measured using appropriate ISEs. Once the concentrations of the ions are known, we can calculate \(K_{sp}\) using the equation mentioned in Method 1.
04

Method 4: Titration

A known mass of the solid can be dissolved in a known volume of water to form an equilibrated solution. We can then titrate this solution with a strong titrant, such as a strong acid or a strong base, to determine the concentration of the anion or the cation, respectively, in the solution. Once we have this information, we can calculate the \(K_{sp}\) value using the equation provided in Method 1. Each of these methods works by measuring the concentrations of the ions in equilibrium with the solid, allowing us to calculate the \(K_{sp}\) value. They differ in the techniques used to measure these ion concentrations, providing flexibility to experimenters in choosing the most appropriate method for a particular solid.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Spectrophotometry
Spectrophotometry is a powerful analytical method that helps determine the concentration of solute in a solution by measuring the intensity of light absorbed by the solution. When a solute, such as an ion, is present in a solution, it can interact with light at specific wavelengths. This interaction is measured by a spectrophotometer.

This technique is particularly useful when dealing with colored ions because the absorption is very noticeable in the visible spectrum. To find the solubility product constant (\(K_{sp}\)), the ions in equilibrium are quantified. The spectrophotometer provides a reading based on how much light the solution absorbs.

The concentration of ions can then be deduced using the Beer-Lambert law, which relates the absorbance to the concentration of the absorbing species. By knowing the stoichiometry of the dissolution reaction, you can plug these concentration values into the equation for the solubility product constant, \(K_{sp} = [\textrm{cation}]^m [\textrm{anion}]^n\). This allows us to precisely determine the \(K_{sp}\) value by calculating it from the measured ion concentrations.
Conductivity Measurements
Conductivity measurements provide another way to determine the solubility product constant. This technique is based on the principle that ions contribute to the electrical conductivity of a solution. When a solid dissolves, it dissociates into ions, increasing the solution's conductivity.

By dissolving a known quantity of solid in a solution and measuring the resulting increase in conductivity, we can infer the concentration of ions present. A calibration curve, created by measuring the conductivity of standards with known ion concentrations, is used to convert conductivity to concentration.

Once the concentrations of ions are known, the solubility product constant can be calculated using the same formula: \(K_{sp} = [\textrm{cation}]^m [\textrm{anion}]^n\). Conductivity measurements are effective for salts that fully dissociate in water, offering an accurate measure for these types of solute systems.
Ion-Selective Electrodes
Ion-selective electrodes (ISEs) offer a versatile and highly specific means of measuring ionic concentrations. ISEs are designed to respond to specific ions in a solution, making them particularly useful when measuring ions in complex mixtures.

The process involves immersing an electrode specific to one of the ions in question into the solution. The electrode generates a measurable electric potential that is related to the concentration of the target ion. By using ISEs for each ion type in the dissolution reaction, their concentrations in the equilibrated solution can be measured precisely.

The \(K_{sp}\) of the solid can then be calculated using the concentrations measured by the electrodes, through the standard solubility product equation. This direct method eliminates much of the interference that can occur with other measurement techniques. It is especially beneficial for systems where one ion might be more mobile or reactive than others.
Titration Methods
Titration methods are a classical technique used to determine the concentration of an unknown solution. In the context of solubility product constant determinations, titration involves using a standard solution with a known concentration to react with the ions present in the saturated solution.

For instance, if you're determining the concentration of an anion, you'd titrate with a standard solution containing a cation that reacts with the anion to form a precipitate, or vice versa. The point at which the reaction is complete (the endpoint) is often signalled by a color change or monitored using a pH meter or conductivity sensor.

By calculating the amount of titrant needed to reach the endpoint, the concentration of the ion in solution can be deduced. Once you know the ion concentrations, you can use the \(K_{sp}\) equation, \(K_{sp} = [\textrm{cation}]^m [\textrm{anion}]^n\), to find the solubility product constant. Titrations are especially effective for solutes that are only slightly soluble or when dealing with large volumes where other methods might not be feasible.

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Most popular questions from this chapter

\(K_{f}\) for the complex ion \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}{ }^{+}\) is \(1.7 \times 10^{7} . K_{\mathrm{sp}}\) for \(\mathrm{AgCl}\) is \(1.6 \times 10^{-10}\). Calculate the molar solubility of \(\mathrm{AgCl}\) in \(1.0 \mathrm{M}\) \(\mathrm{NH}_{3} .\)

Calculate the concentration of \(\mathrm{Pb}^{2+}\) in each of the following. a. a saturated solution of \(\mathrm{Pb}(\mathrm{OH})_{2}, K_{\mathrm{sp}}=1.2 \times 10^{-15}\) b. a saturated solution of \(\mathrm{Pb}(\mathrm{OH})_{2}\) buffered at \(\mathrm{pH}=13.00\) c. Ethylenediaminetetraacetate \(\left(\mathrm{EDTA}^{4-}\right)\) is used as a complexing agent in chemical analysis and has the following structure: Solutions of EDTA \(^{4-}\) are used to treat heavy metal poisoning by removing the heavy metal in the form of a soluble complex ion. The reaction of EDTA \(^{4-}\) with \(\mathrm{Pb}^{2+}\) is \(\begin{aligned} \mathrm{Pb}^{2+}(a q)+\mathrm{EDTA}^{4-}(a q) \rightleftharpoons \mathrm{PbEDTA}^{2-}(a q) & \\ K &=1.1 \times 10^{18} \end{aligned}\) Consider a solution with \(0.010\) mole of \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) added to 1.0 L of an aqueous solution buffered at \(\mathrm{pH}=13.00\) and containing \(0.050 M\) NasEDTA. Does \(\mathrm{Pb}(\mathrm{OH})_{2}\) precipitate from this solution?

a. Using the \(K_{\mathrm{sp}}\) value for \(\mathrm{Cu}(\mathrm{OH})_{2}\left(1.6 \times 10^{-19}\right)\) and the overall formation constant for \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}\left(1.0 \times 10^{13}\right)\), calculate the value for the equilibrium constant for the following reaction: \(\mathrm{Cu}(\mathrm{OH})_{2}(s)+4 \mathrm{NH}_{3}(a q) \rightleftharpoons \mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}(a q)+2 \mathrm{OH}^{-}(a q)\) b. Use the value of the equilibrium constant you calculated in part a to calculate the solubility (in \(\mathrm{mol} / \mathrm{L}\) ) of \(\mathrm{Cu}(\mathrm{OH})_{2}\) in \(5.0 \mathrm{M} \mathrm{NH}_{3} .\) In \(5.0 \mathrm{M} \mathrm{NH}_{3}\) the concentration of \(\mathrm{OH}^{-} \mathrm{is}\) \(0.0095 M .\)

In the presence of \(\mathrm{NH}_{3}, \mathrm{Cu}^{2+}\) forms the complex ion \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}\). If the equilibrium concentrations of \(\mathrm{Cu}^{2+}\) and \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}\) are \(1.8 \times 10^{-17} M\) and \(1.0 \times 10^{-3} \mathrm{M}\), respec- tively, in a \(1.5-M \mathrm{NH}_{3}\) solution, calculate the value for the overall formation constant of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}\). \(\mathrm{Cu}^{2+}(a q)+4 \mathrm{NH}_{3}(a q) \rightleftharpoons \mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}(a q) \quad K_{\text {overall }}=?\)

Assuming that the solubility of \(\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}(s)\) is \(1.6 \times 10^{-7}\) \(\mathrm{mol} / \mathrm{L}\) at \(25^{\circ} \mathrm{C}\), calculate the \(K_{\text {op }}\) for this salt. Ignore any potential reactions of the ions with water.
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