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Basic Principles of Modern Chemistry II Week 10 Resources

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Basic Principles of Modern Chemistry II Week 10 Resources CHE 1302 – Basic Principles of Modern Chemistry II Week 10 Resources By Patrick Olademehin (03-22-2021) Hello everyone! This week’s CHE 1302 resource will focus on a review of Chapter 15 (Solubility and Complex Ion Equilibria) from the approved textbook, Chemistry; An Atoms First Approach. The content of this review will be very helpful to refresh your memory on the concepts discussed in Chapter 15. If you need a quick review on Chapter 14, check out last week’s resources using this link: https://www.baylor.edu/support_programs/doc.php/372074.pdf Remember, we conduct a Group Tutoring session every Monday 5:30 – 6:30 PM. Reserve a spot early at https://baylor.edu/tutoring and endeavor to come with a specific question in mind. Keywords: Solubility product constant, Relative solubilities, Selective precipitation, Complex ions. Topics Treated in Week 10 These include: 1. Solubility Equilibria and the Solubility Product a. Relative Solubilities b. Common Ion Effect c. pH and Solubility 2. Precipitation and Qualitative Analysis a. Selective Precipitation b. Qualitative Analysis 3. Equilibria Involving Complex Ions a. Complex Ions and Solubility Revision topics 1. Chemical Equilibrium (Chapter 12) – The foundation for solving equilibrium problems was laid in Chapter 12. A review on how to express equilibrium constants, Le Châtelier’s principle and using the ICE method to solve problems would be helpful. 2. Stoichiometry and Solubility Rules (Chapter 6) – At this point, students should be conversant with using a balanced chemical equation to show relationships between reactants and products of a reaction. Refer to Table 6-1 in Chapter 6 of the textbook for simple rules for the solubility of salts in water. This link can also be helpful: http://sayrechem.weebly.com/solubility-rules.html Highlights from Chapter 15 Topics 1. Solubility Equilibria and the Solubility Product – The equilibrium constant expression for a slightly soluble solid in equilibrium with its ions is called the solubility product constant (Ksp). 2+ - For the reaction: CaF2(s) Ca (aq) + 2F (aq), 2+ - 2 Ksp = [Ca ][F ] All contents and figures, unless stated otherwise, are taken from the textbook, Chemistry, An Atoms First Approach by Zumdahl and Zumdahl. Solubility, on the other hand, is the maximum amount of solute dissolved in a solvent at equilibrium (commonly expressed in mol/L or g/L). The solubility product (Ksp) is a constant value at a given temperature, while solubility is an equilibrium position which is subject to variation (e.g. in the presence of a common ion). Here is a link to a ~8 min. video explaining Ksp and showing how to find solubility and Ksp: https://www.youtube.com/watch?v=WjiXbemBXkE 2. Relative Solubilities – The Ksp values for different salts can be compared to predict their relative solubilities only if the salts being compared produce the same number of ions in solution. 3. Common Ion Effect – The presence of a common ion will decrease the solubility of a solid in solution (Le Châtelier’s principle). Here, when using the ICE table to find the equilibrium concentration (E) or solubility, the initial concentration (I) will not be zero for the ion already present in solution. 4. pH and Solubility – A decrease in pH (or increase in acidity) can affect the solubility of salts whose anions is the conjugate base of a weak acid. Examples of such anions include: - 2- 2- 2- OH , S , CO3 , and CrO4 . The addition of a strong acid, in this case, will increase the solubility of the salt (Le Châtelier’s principle), since the ions produced are being consumed by the acid. Also, an addition of OH- (increase in pH) will decrease the solubility of a salt that produces OH- in solution based on the common ion effect (Common Ion Effect). 5. Precipitation and Qualitative Analysis – The reaction quotient, Q, and the Solubility Product, Ksp, can be compared to determine if precipitation of a given salt will occur in solution. Precipitation occurs if Q > Ksp (means too many ions in solution), but if Q < Ksp, no precipitation occurs. Remember, the reaction quotient, Q, is the equilibrium constant expression using initial concentration (not equilibrium concentration). As an example, the figure below shows how to determine if precipitation of Ce(IO3)3 will occur in a solution + - containing Ce3 and IO3 ions: All contents and figures, unless stated otherwise, are taken from the textbook, Chemistry, An Atoms First Approach by Zumdahl and Zumdahl. 6. Selective Precipitation – An ion can be selectively precipitated from a mixture by taking advantage of the differences in Ksp values. Also, a mixture of metal ions can be separated by adding a reagent containing an anion that forms a precipitate with only one or few metal ions in the mixture. 7. Qualitative Analysis – A scheme for qualitative analysis of a mixture containing all the common cations are shown in the figure below: 8. Equilibria Involving Complex Ions – Complex ions consist of a metal ion (Lewis Acid) surrounded by ligands (Lewis Bases). The equilibrium constant for the formation of complex ion is known as the formation constant, Kf. The formation of a complex ion can be used to selectively dissolve solids. Common Mistakes Students Make 1. Not knowing the difference between solubility and solubility product constant. 2. Comparing Ksp values for salts with different total number of ions to determine the relative solubilities of the salts. 3. Forgetting to include the concentration of common ion present as initial concentration when trying to get the solubility of a salt present in a solution with a common ion. 4. In calculating the equilibrium concentration after precipitation occurs, failure to only use the excess ion present after precipitation as the initial concentration in the ICE method. 5. Not including the effect of dilution when mixing solution to get precipitates. 6. To get the formation constant expression for a complex ion formation, students sometimes forget to put the complex as a product and leave it as a reactant. All contents and figures, unless stated otherwise, are taken from the textbook, Chemistry, An Atoms First Approach by Zumdahl and Zumdahl. .
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