EXPERIMENT #10 Acid-Base Properties of Salt Solutions: Hydrolysis

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1 OBJECTIVES: EXPERIMENT #10 Acid-Base Properties of Salt Solutions: Hydrolysis Use ph indicators to determine the ph range of various solutions Use a ph meter to determine the ph of various solutions Explain why various salt solutions are acid, basic, or neutral BACKGROUND: When a strong acid or a strong base dissolves in water, it dissociates completely into ions. In a solution of a strong monoprotic * acid the concentration of hydronium ion, H3O +, is equal to the concentration of the strong acid. In a solution of a strong monobasic base, the concentration of hydroxide ion, OH -, is equal to the concentration of the strong base; in a dibasic strong base the concentration of hydroxide ion is twice the concentration of the strong base. (In each case the autoionization of water contributes a negligible quantity of hydronium ions or hydroxide ions.) An acid is a strong acid if it is one of the following: hydrochloric acid, HCl; hydrobromic acid, HBr; hydriodic acid, HI; nitric acid, HNO3; perchloric acid, HClO4; and the first proton of sulfuric acid, H2SO4. Any other material capable of donating a proton is a weak acid. A base is a strong base if it is one of the following: lithium hydroxide, LiOH; sodium hydroxide, NaOH; potassium hydroxide, KOH; or any other alkali metal hydroxide; and magnesium hydroxide, Mg(OH)2; calcium hydroxide, Ca(OH)2; or barium hydroxide; Ba(OH)2. Any other material capable of accepting a proton is a weak base. A weak acid only partially dissociates upon dissolving in water. The extent of dissociation is governed by the equilibrium constant Ka for the following process, HA + H2O H3O + + A - (1) where H O A 3 Ka = (2) HA The larger the Ka, the larger the extent of dissociation for the weak acid. Thus the stronger weak acids have larger Ka values, and the weaker have smaller Ka values. Often, both Ka and [H3O + ] are expressed in their logarithmic forms, pka and ph, where pka = -logka and ph = -log[h3o + ]. A weak base only partially dissociates upon dissolving in water. The extent of dissociation is governed by the equilibrium constant Kb for the following process, B + H2O BH + +OH - (3) where BH OH Kb = (4) B The larger the Kb, the larger the extent of dissociation for the weak base. Thus the stronger weak bases have larger Kb values, and the weaker have smaller Kb values. Often, both Kb and [OH - ] are expressed in their logarithmic forms, pkband poh, where pkb = -logkb and poh = -log = [OH - ]. * a monoprotic acid has one proton to donate a monobasic base accepts one proton 111 P a g e

2 EXPERIMENT #10 ACIDS/BASES/HYDROLYSIS For all aqueous solutions, the concentration of H3O + is related to the concentration of OH - by the ion product (auto-ionization) of water, Kw, which at 25 C has a value of 1.00 x H2O + H2O H3O + + OH - (5) Kw = [H3O + ][OH - ] (6) In logarithmic form, equation (6) becomes pkw = ph + poh (7) where pkw = -log10kw = Unlike strong acids and strong bases, weak acids and weak bases do not dissociate completely in solution. Because the reaction of a weak acid or base with water is less than complete, measurable quantities of both the weak acid and its conjugate base (or weak base and its conjugate acid) are present at equilibrium. While the Ka for a strong acid is very a very large number, e.g., 10 8, the Ka for a weak acid is small, e.g., 10-3 or 10-5, or less. Likewise, the Kb for a strong base is a very large number, but the Kb for a weak base is small. Values of Ka and Kb have been tabulated for a large number of acids and bases, and representative tables can be found in your textbook. A salt is the non-aqueous product of a reaction between an acid and a base. Most soluble salts dissociate completely into ions when they dissolve in water. The ions may show no interaction with water; such a solution will remain neutral. On the other hand, the ions may interact with water by donating a proton to make the solution acidic or accepting a proton from water to make the solution basic. It is important for the chemist to recognize how the ions of soluble salts affect the relative concentrations of H3O + and OH -. Anions that are conjugate bases of strong acids show no interaction with water and are termed pathetically weak bases; the ph of the solution is unaffected. The following anions belong to this class: chloride, Cl - ; bromide, Br - ; iodide, I - ; nitrate, NO3 - ; perchlorate, ClO4 - ; and hydrogen sulfate, HSO4 -. However, anions that are the conjugate bases of weak acids are themselves weak bases; they abstract a proton from water to produce hydroxide ions and the solution becomes basic. For example, the conjugate base of the weak acid HA would be A-, which is itself a weak base that can react to a small, but measurable, extent with water to produce hydroxide: A - + H2O HA + OH - (8) The value for the Kb of the conjugate base, A-, can be obtained from Kw and the Ka for the weak acid HA using the relationship Ka Kb = Kw. Metal cations of the Group IA and Group IIA hydroxide bases show no interaction with water and are termed pathetically weak acids; the ph of the solution is unaffected and the solution remains neutral. The following cations belong to this class: lithium, Li + ; sodium, Na + ; potassium, K + ; or any other alkali metal cation; and magnesium, Mg 2+ ; calcium, Ca 2+ ; or barium; Ba 2+. However, cations that are the conjugate acids of weak bases are themselves weak acids. For example, the conjugate acid of the weak base B would be BH +, which is itself a weak acid that can react to a small but significant extent in water to produce hydronium ion: BH + + H2O B + H3O + (9) The value for the Ka of the conjugate acid, BH +, can be calculated from the Kw and the Kb for the weak base using the relationship Kw = Ka Kb. 112 P a g e

3 EXPERIMENT #10 ACIDS/BASES/HYDROLYSIS There is another class of cations that can act as weak acids. To explain this class we need to consider the Lewis theory of acids and bases. Solutions of cations of several transition metals and Al 3+ are acidic. Most cations in aqueous solution are surrounded by several molecules of water, usually four or six. These water molecules are known as the solvation shell. This occurs because the partially negative charge on the oxygen of a water molecule is attracted to the positively charged cation. However, the charge to size ratio for many transition metal cations and Al 3+ is unusuallylarge. Additionally, the transition metal cations are electron pair acceptors (Lewis acids). The result is an extreme distortion of the electron cloud making up the hydrogen oxygen bond in the water molecules solvating the cation. The hydrogen of this water molecule becomes an acidic hydrogen. M(H2O)x n+ + H2O M(H2O)x-1(OH) (n-1) + + H3O + (10) In this experiment you will use two methods of differing sophistication to determine the ph of solutions of various substances. One method will use acid-base indicators, which are themselves weak acids. The other will use a ph meter and combination electrode. The method of lower less sophistication uses acid-base indicators. Indicators are weak acids which change color over a narrow range ( 2 ph units) of ph. Below the ph of this range the solution of the indicator (in its acid form) is one color; above the ph of this range (in its conjugate base form) another color; within the ph range a combination of the two colors is observed. For example, bromothymol blue changes color from yellow to blue over the ph range of 6.0 to 7.6. If bromothymol blue is placed in a solution and the color is yellow, the ph of the solution is 6.0 or less. On the other hand, if bromothymol blue is blue in the solution being tested, the ph of the solution is 7.6 or more. If the solution being tested is green, the ph is between 6.0 and 7.6. By using several acid-base indicators which change color over various sections of the ph scale, a solution s ph can be approximated. In this part of the experiment we are only deciding whether the solution is acidic, basic, or neutral. Once you know the ph of the solution, you can determine whether the cations and/or anions are pathetic, weak, or strong acids or bases. The acid-base indicators used in this experiment are listed in TABLE I. 113 P a g e

4 EXPERIMENT #10 ACIDS/BASES/HYDROLYSIS NOTE: ph paper is filter paper which has been impregnated with a mixture of acid-base indicators. When ph paper is wetted with a solution it will display a certain color dependent upon the ph of the solution used to wet it. ph paper is available in a wide-range form, which covers the ph range 0-14, and a short-range form, which covers a narrower ph range. The ph paper available in the laboratory has a color chart provided to guide in interpreting what ph is indicated by the various colors. TABLE I Selected Acid-Base Indicators Bromocresol Green Methyl Red Bromothymol Blue Phenol Red Phenolphthlein ph Range color change yellow - blue red yellow yellow - blue yellow - red colorless pink The method of greater sophistication is the potentiometric method which uses a voltmeter with special electrodes to measure the ph of the solution directly. The ph is converted to the hydronium ion concentration by the relationship [H3O + ] = 10 -ph PROCEDURE: I. Determination of ph using acid-base indicators 1. Place 1 ml of the first solution to be tested from TABLE II into each of five small test tubes. Add 1-2 drops of the first indicator to the first test tube and mix thoroughly. Record your observations in the table on the data sheet. Continue with the remaining indicators in sequence. 2. Repeat the process with each of the other solutions listed in TABLE II. TABLE II Solutions to Test with Indicator Solutions 0.10 M NaCl 0.10 M Na2SO M Na2HSO M Na2SO M Na2CO M NH4Cl 0.10 M NaC2H3O M ZnCl M Na2HPO M Cu(NO3)2 The following WEB Site contains useful information on the use of ph paper: P a g e

5 EXPERIMENT #10 ACIDS/BASES/HYDROLYSIS II. Determination of ph Using a ph Meter The most accurate and precise method of determining the ph of a solution requires the use of a ph-sensing electrode and a digital ph meter. The business end of a ph meter is a special glass membrane electrode which produces a tiny voltage dependent on the H3O + concentration in the solution into which it is immersed. The ph sensing electrode is normally paired with a reference electrode, which completes the electrical circuit and produces a constant, steady voltage to which the ph electrode voltage can be compared. The voltage generated by this electrode is measured by a high priced digital voltmeter which converts a voltage signal into a precise (±0.01 unit) ph reading. NOTES 1. The ph electrode should never be left to dry out. When the electrode is not being used, keep it immersed in a beaker of deionized water or the ph 4 or 7 buffer solution. When you are finished using the ph electrode, rinse it thoroughly with deionized water, close the filling hole, and place the business end into the appropriate buffer solution. 2. Do not leave the business end of the electrode immersed in the carbonate solutions (or any other basic solution)! Place the electrode in the solution only long enough to determine the ph, then remove the electrode and thoroughly rinse it with deionized water. Basic solutions attack ph electrodes, and lead to their premature failure and demise. 1. For each of the solutions in TABLE II measure the ph potentiometrically The following WEB Site contains useful information on the use of a ph meter: P a g e

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7 NAME Section Date Data And Observations: Acids/Bases/Hydrolysis I. Determination of ph using acid-base indicators NaCl Bromocresol Green Methyl Red Bromothymol Blue Phenol Red Phenolphthlein Acidic/Basic/ Neutral Na 2SO 4 Na 2CO 3 NaC 2H 3O 2 NaHSO 4 Na 2HPO 4 Na 2SO 3 NH 4Cl ZnCl 2 Cu(NO 3) 2 For each of the above salt solutions record the following beginning on the next page: (1) Write the formulas of the ions in solution. (2) Classify these ions as strong/weak/pathetic acids or bases by circling the proper abbreviation. (3) Predict whether solution is acidic, basic, or neutral by circling the proper letter. (4) Determine the ph of the solutions colorimetrically and circle the proper letter. (5) Determine the ph of the solutions potentiometrically. Divide this task among the students participating in lab. NaCl solution Formula Cation Formula Anion 117 P a g e

8 Na2SO4 solution Formula Cation Formula Anion Na2CO3 solution Formula Cation Formula Anion NaC2H3O2 solution Formula Cation Formula Anion 118 P a g e

9 NAME Section Date NaHSO4 solution Formula Cation Formula Anion Na2HPO4 solution Formula Cation Formula Anion Na2SO3 solution Formula Cation Formula Anion 119 P a g e

10 NH4Cl solution Formula Cation Formula Anion ZnCl2 solution Formula Cation Formula Anion Cu(NO3)2 solution Formula Cation Formula Anion 120 P a g e

11 NAME Section Date ADDITIONAL ASSIGNMENT I: Acids/Bases/Hydrolysis 1. What is the conjugate base of H2PO4 -? What is the conjugate acid of H2PO4 -? 2. Which of the following salts will undergo hydrolysis: KF NaNO3 NH4NO2 MgSO4 KCN C6H5COONa RbI CaCl2 Na2C2O4 CaCl2 HCOOK 3. Circle the following ions which undergo hydrolysis? Ca 2+ Fe 3+ Al 3+ F - PO4 3- Br - OBr - 4. Predict the ph (>7, <7, or 7) of aqueous solutions containing the following salts: (a) KBr (b) Al(NO3)3 (c) BaCl2 (d) Bi(NO3)3 121 P a g e

12 5. Predict whether the following solutions are acidic (A), basic (B), or nearly neutral (N). NaBr LiClO4 K2SO3 Cr(NO3)3 NaNO2 FeCl3 Na3PO4 CuSO4 ZnCl2 NaCN xc. Predict semi-quantitatively whether the following solutions are acidic, basic, or nearly neutral. (a) NH4NO2 (b) NH4F (c) NH4CN (d) NaHCO3 (e) K2HPO4 (f) (NH4)2SO4 The first part is solved below. Use similar reasoning to make predictions for the other solutions. (a) Ka for NH4 = 5.6 x and Kb for NO2 - is 2.2 x Since the Ka is about 25 times larger than Kb, the solution should be acidic. 122 P a g e