Show that the pH at the halfway point of a titration of a weak acid with a strong base

The following boxes represent aqueous solutions containing a weak acid, HX, and its conjugate base, x-. Water molecules and cations are not shown. Which solution has the highest pH? Explain. [Section 17.1]

The beaker on the right contains 0.1 M acetic acid solution with methyl orange as an indicator. The beaker on the left contains a mixture of 0.1 M acetic acid and 0.1 M sodium acetate with methyl orange. (a) Using Figure 16.7, estimate the pH of each solution, and explain the difference. (b) Which solution is better able to maintain its pH when small amounts of NaOH are added? Explain. [Sections 17.1 and 17.2]

A buffer contains a weak acid, HX, and its conjugate base. The weak acid has a pKa of 4.5, and the buffer solution has a pH of 4.3. Without doing a calculation, predict whether [HX] = [X-], [HX] > [X-], or [HX] < [X-]. Explain. [Section 17.2]

The drawing on the left represents a buffer composed of equal concentrations of a weak acid, HX, and its conjugate base, x-. The heights of the columns are proportional to the concentrations of the components of the buffer. (a) Which of the three drawings, (1), (2), or (3), represents the buffer after the addition of a strong acid? (b) Which of the three represents the buffer after the addition of a strong base? (c) Which of the three represents a situation that cannot arise from the addition of either an acid or a base? [Section 17.2]

The following drawings represent solutions at various stages of the titration of a weak acid, HA, with NaOH. (The Na+ ions and water molecules have been omitted for clarity.) To which of the following regions of the titration curve does each drawing correspond: (a) before addition of NaOH, (b) after addition of NaOH but before equivalence point, (c) at equivalence point, (d) after equivalence point? [Section 17.3]

Match the following descriptions of titration curves with the diagrams: (a) strong acid added to strong base, (b) strong base added to weak acid, (c) strong base added to strong acid, (d) strong base added to polyprotic acid. [Section 17.3]

Equal volumes of two acids are titrated with 0.10 M NaOH resulting in the two titration curves shown in the following figure. (a) Which curve corresponds to the more concentrated acid solution? (b) Which corresponds to the acid with the largest Ka? Explain. [Section 17.3]

The following drawings represent saturated solutions of three ionic compounds of silver-AgX, AgY, and AgZ. (Na + cations, which might also be present for charge balance, are not shown.) Which compound has the smallest K,p? [Section 17.4]

The figures below represent the ions in a saturated aqueous solution of the slightly soluble ionic compound MX: MX(s) :;====' M2+(aq) + X2-(aq). (Only the M2+ and X2- ions are shown.) (a) Which figure represents a solution prepared by dissolving MX in water? (b) Which figure represents a solution prepared by dissolving MX in a solution containing Na2 X? (c) If xz- is a basic anion, which figure represents a saturated solution with the lowest pH? (d) If you were to calculate the Ksp for MX, would you get the same value in each of the three scenarios? Why or why not? [Sections 17.4 and 17.5]

What is the name given to the kind of behavior demonstrated by a metal hydroxide in this graph? [Section 17.5]

Three cations, Ni2 +, Cu2 +, and Ag+, are separated using two different precipitating agents. Based on Figure 17.22, what two precipitating agents could be used? Using these agents, indicate which of the cations is A, which is B, and which is C. [Section 17.7]

(a) What is the common-ion effect? (b) Give an example of a salt that can decrease the ionization of HN02 in solution.

(a) Consider the equilibrium B(aq) + H20(/) = HB+(aq) + OH-(aq). Using Le Chatelier's principle, explain the effect of the presence of a salt of HB+ on the ionization of B. (b) Give an example of a salt that can decrease the ionization of NH3 in solution.

Use information from Appendix D to calculate the pH of (a) a solution that is 0.060 M in potassium propionate (C2H5COOK or KC3H502) and 0.085 Min propionic acid (C2H5COOH or HC3H502); (b) a solution that is 0.075 M in trimethylamine, (CH3hN, and 0.10 M in trimethylammonium chloride, (CH3hNHCl; (c) a solution that is made by mixing 50.0 mL of 0.15 M acetic acid and 50.0 mL of 0.20 M sodium acetate.

Use information from Appendix D to calculate the pH of (a) a solution that is 0.150 M in sodium formate (HCOONa) and 0.200 M in formic acid (HCOOH); (b) a solution that is 0.210 M in pyridine (C5H5N) and 0.350 M in pyridinium chloride (C5H5NHC1); (c) a solution that is made by combining 125 mL of 0.050 M hydrofluoric acid with 50.0 mL of 0.10 M sodium fluoride.

(a) Calculate the percent ionization of 0.0075 M butanoic acid (Ka = 1.5 X 10-5 ). (b) Calculate the percent ionization of 0.0075 M butanoic acid in a solution containing 0.085 M sodium butanoate.

(a) Calculate th ercent ionization of 0.085 M lactic acid (Ka = 1.4 X 0 ) . (b) Calculate the percent ionization of 0.095 M lactic acid in a solution containing 0.0075 M sodium lactate.

Explain why a mixture of CH3COOH and CH3C00Na can act as a buffer while a mixture of HCl and NaCl cannot.

(a) Calculate the pH of a buffer that is 0.12 M in lactic acid and 0.11 M in sodium lactate. (b) Calculate the pH of a buffer formed by mixing 85 mL of 0.13 M lactic acid with 95 mL of 0.15 M sodium lactate.

(a) Calculate the pH of a buffer that is 0.105 M in NaHC03 and 0.125 M in Na2C03. (b) Calculate the pH of a solution formed by mixing 65 mL of 0.20 M NaHC03 with 75 mL of 0.15 M Na2C03.

A buffer is prepared by adding 20.0 g of acetic acid (CH3COOH) and 20.0 g of sodium acetate (CH3C00Na) to enough water to form 2.00 L of solution. (a) Determine the pH of the buffer. (b) Write the complete ionic equation for the reaction that occurs when a few drops of hydrochloric acid are added to the buffer. (c) Write the complete ionic equation for the reaction that occurs when a few drops of sodium hydroxide solution are added to the buffer.

A buffer is prepared by adding 7.00 g of ammonia (NH3) and 20.0 g of ammonium chloride (NH4CI) to enough water to form 2.50 L of solution. (a) What is the pH of this buffer? (b) Write the complete ionic equation for the reaction that occurs when a few drops of nitric acid are added to the buffer. (c) Write the complete ionic equation for the reaction that occurs when a few drops of potassium hydroxide solution are added to the buffer.

How many moles of sodium hypobromite (NaBrO) should be added to 1.00 L of 0.050 M hypobromous acid (HBrO) to form a buffer solution of pH 9.15? Assume that no volume change occurs when the NaBrO is added.

How many grams of sodium lactate [CH3CH(OH)COONa or NaC503] should be added to 1.00 L of 0.150 M lactic acid [CH3CH(OH)COOH or HC3H503] to form a buffer solution with pH 4.00? Assume that no volume change occurs when the sodium lactate is added.

A buffer solution contains 0.10 mol of acetic acid and 0.13 mol of sodium acetate in 1 .00 L. (a) What is the pH of this buffer? (b) What is the pH of the buffer after the addition of 0.02 mol of KOH? (c) What is the pH of the buffer after the addition of 0.02 mol of HN03?

A buffer solution contains 0.10 mol of propionic acid (C2H5COOH) and 0.13 mol of sodium propionate (C2H5COONa) in 1.50 L. (a) What is the pH of this buffer? (b) What is the pH of the buffer after the addition of 0.01 mol of NaOH? (c) What is the pH of the buffer after the addition of O.Dl mol of HI?

(a) What is the ratio of HC03- to H2C03 in blood of pH 7.4? (b) What is the ratio of HC03- to H2C03 in an exhausted marathon runner whose blood pH is 7.1?

You have to prepare a pH 3.50 buffer, and you have the following 0.10 M solutions available: HCOOH, CH3COOH, H04, HCOONa, CH3COONa, and NaH2P04. Which solutions would you use? How many milliliters of each solution would you use to make approximately a liter of the buffer?

You have to prepare a pH 4.80 buffer, and you have the following 0.10 M solutions available: formic acid, sodium formate, propionic acid, sodium propionate, phosphoric acid, and sodium dihydrogen phosphate. Which solutions would you use? How many milliliters of each solution would you use to make approximately a liter of the buffer?

The accompanying graph shows the titration curves for two monoprotic acids. (a) Which curve is that of a strong acid? (b) What is the approximate pH at the equivalence point of each titration? (c) How do the original concentrations of the two acids compare if 40.0 mL of each is titrated to the equivalence point with 0.100 M base?

How does titration of a strong, monoprotic acid with a strong base differ from titration of a weak, monoprotic acid with a strong base with respect to the following: (a) quantity of base required to reach the equivalence point, (b) pH at the beginning of the titration, (c) pH at the equivalence point, (d) pH after addition of a slight excess of base, (e) choice of indicator for determining the equivalence point?

Predict whether the equivalence point of each of the following titrations is below, above, or at pH 7: (a) NaHC03 titrated with NaOH, (b) NH3 titrated with HCI, (c) KOH titrated with HBr.

Predict whether the equivalence point of each of the following titrations is below, above, or at pH 7: (a) formic acid titrated with NaOH, (b) calcium hydroxide titrated with perchloric acid, (c) pyridine titrated with nitric acid.

Two monoprotic acids, both 0.100 M in concentration, are titrated with 0.100 M NaOH. The pH at the equivalence point for HX is 8.8, and that for HY is 7.9. (a) Which is the weaker acid? (b) Which indicators in Figure 16.7 could be used to titrate each of these acids?

Assume that 30.0 mL of a 0.10 M solution of a weak base B that accepts one proton is titrated with a 0.10 M solution of the monoprotic strong acid HX. (a) How many moles of HX have been added at the equivalence point? (b) What is the predominant form of B at the equivalence point? (c) What factor determines the pH at the equivalence point? (d) Which indicator, phenolphthalein or methyl red, is likely to be the better choice for this titration?

How many milliliters of 0.0850 M NaOH are required to titrate each of the following solutions to the equivalence point: (a) 40.0 mL of 0.0900 M HN03, (b) 35.0 mL of 0.0850 M CH3COOH, (c) 50.0 mL of a solution that contains 1.85 g of HCl per liter?

A 20.0-mL sample of 0.200 M HBr solution is titrated with 0.200 M NaOH solution. Calculate the pH of the solution after the following volumes of base have been added: (a) 15.0 mL, (b) 19.9 mL, (c) 20.0 mL, (d) 20.1 mL, (e) 35.0 mL.

A 30.0-mL sample of 0.150 M KOH is titrated with 0.125 M HCI04 solution. Calculate the pH after the following volumes of acid have been added: (a) 30.0 mL, (b) 35.0 mL, (c) 36.0 mL, (d) 37.0 mL, (e) 40.0 mL.

A 35.0-mL sample of 0.150 M acetic acid (CH3COOH) is titrated with 0.150 M NaOH solution. Calculate the pH after the following volumes of base have been added: (a) 0 mL, (b) 17.5 mL, (c) 34.5 mL, (d) 35.0 mL, (e) 35.5 mL, (f) 50.0 mL.

Consider the titration of 30.0 mL of 0.030 M NH3 with 0.025 M HCI. Calculate the pH after the following volumes of titrant have been added: (a) 0 mL, (b) 10.0 mL, (c) 20.0 mL, (d) 35.0 mL, (e) 36.0 mL, (f) 37.0 mL.

Calculate the pH at the equivalence point for titrating 0.200 M solutions of each of the following bases with 0.200 M HBr: (a) sodium hydroxide (NaOH), (b) hydroxylamine (NH20H), (c) aniline (C6HsNH2l

Calculate the pH at the equivalence point in titrating 0.100 M solutions of each of the following with 0.080 M NaOH: (a) hydrobromic acid (HBr), (b) lactic acid [CH3CH(OH)COOH], (c) sodium hydrogen chromate (NaHCr04).

(a) Why is the concentration of undissolved solid not explicitly included in the expression for the solubilityproduct constant? (b) Write the expression for the solubility-product constant for each of the following strong electrolytes: Agl, SrS04, Fe(OH)z, and Hg2Br2.

(a) Explain the difference between solubility and solubility-product constant. (b) Write the expression for the solubility-product constant for each of the following ionic compounds: MnCO:v Hg(OH)z, and Cu3(P04)z.

(a) If the molar solubility of CaF2 at 35 oc is 1.24 X 10-3 moljL, what is Ksp at this temperature? (b) It is found that 1.1 x 10-2 g of SrF2 dissolves per 100 mL of aqueous solution at 25 oc. Calculate the solubility product for SrF2. (c) The Ksp of Ba(I03)z at 25 oc is 6.0 X 10-IO What is the molar solubility of Ba(I03)z?

A 1.00-L solution saturated at 25 oc with calcium oxalate (CaC204) contains 0.0061 g of CaC204. Calculate the solubility-product constant for this salt at 25 oe,

A 1.00-L solution saturated at 25 oc with lead(II) iodide contains 0.54 g of Pbl2. Calculate the solubility-product constant for this salt at 25 oc.

Using Appendix 0, calculate the molar solubility of AgBr in (a) pure water, (b) 3.0 X 10-2 M AgN03 solution, (c) 0.10 M NaBr solution.

Calculate the solubility of LaF3 in grams per liter in (a) pure water, (b) 0.010 M KF solution, (c) 0.050 M LaC13 solution.

Calculate the solubility of Mn(OHh in grams per liter when buffered at pH (a) 7.0, (b) 9.5, (c) 11.8.

Calculate the molar solubility of Fe(OH)z when buffered at pH (a) 8.0, (b) 10.0, (c) 12.0.

Which of the following salts will be substantially more soluble in acidic solution than in pure water: (a) ZnC03, (b) ZnS, (c) Bil3, (d) AgCN, (e) Ba3(P04)z?

For each of the following slightly soluble salts, write the net ionic equation, if any, for reaction with acid: (a) MnS, (b) PbF2, (c) AuCI3, (d) Hg2C204, (e) CuBr.

From the value of Kr listed in Table 17.1, calculate the concentration of Cu2 + in 1.0 L of a solution that contains a total of 1 X 10-3 mol of copper(II) ion and that is 0.10 M in NH3.

By using the values of K,p for Agl and Kt for Ag(CNh-, calculate the equilibrium constant for the reaction Agl(s) + 2 CW(aq) Ag(CNh -(aq) + r-(aq)

Using the value of Ksp for Ag2S, Ka1 and Ka2 for H2S, and K1 = 1.1 x 105 for AgC12 -, calculate the equilibrium constant for the following reaction: Ag2S(s) + 4 Cr-(aq) + 2 H+(aq) 2 AgCI2 -(aq) + H2S(aq)

(a) Will Ca(OHh precipitate from solution if the pH of a 0.050 M solution of CaCI2 is adjusted to 8.0? (b) Will Ag2S04 precipitate when 100 mL of 0.050 M AgN03 is mixed with 10 mL of 5.0 x 10-2 M Na2S04 solution?

(a) Will Co(OHh precipitate from solution if the pH of a 0.020 M solution of Co(N03h is adjusted to 8.5? (b) Will Ag!03 precipitate when 20 mL of 0.010 M AgN03 is mixed with 10 mL of O.D15 M Na!03? (Ksp of Ag!03 is 3.1 X 10-8.)

Calculate the minimum pH needed to precipitate Mn(OHh so completely that the concentration of Mn2 + is less than 1 ,.,_g per liter [1 part per billion (ppb)].

Suppose that a 10-mLsample of a solution is to be tested for Cl- ion by addition of 1 drop (0.2 mL) of 0.10 M AgN03. What is the minimum number of grams of Cr that must be present for AgCl(s) to form?

A solution contains 2.0 X 10-4 M Ag+ and 1.5 x 10-3 M Pb2 + If Nal is added, will Agl (Ksp = 8.3 X 10-17) or Pbl2 (Ksp = 7.9 X 10-9) precipitate first? Specify the concentration of r- needed to begin precipitation.

A solution of Na2S04 is added dropwise to a solution that is 0.010 M in Ba2 + and 0.010 M in Sr2 + (a) What concentration of so/- is necessary to begin precipitation? (Neglect volume changes. BaS04: K,p = 1.1 X 10-10; SrS04: K,p = 3.2 x 10-7.) (b) Which cation r.recipitates first? (c) What is the concentration of S04 - when the second cation begins to precipitate?

A solution containing an unknown number of metal ions is treated with dilute HCl; no precipitate forms. The pH is adjusted to about 1, and H2S is bubbled through. Again, no precipitate forms. The pH of the solution is then adjusted to about 8. Again, H2S is bubbled through. This time a precipitate forms. The filtrate from this solution is treated with (NH4hHP04. No precipitate forms. Which metal ions discussed in Section 17.7 are possibly present? Which are definitely absent within the limits of these tests?

In the course of various qualitative analysis procedures, the following mixtures are encountered: (a) zn2 + and Cd2 +, (b) Cr(OH)J and Fe(OH)J, (c) Mg2 + and K+, (d) Ag + and Mn2 + Suggest how each mixture might be separated

Suggest how the cations in each of the following solution mixtures can be separated: (a) Na + and Cd2 +, (b) Cu2 + and Mg2 +, (c) Pb2 + and Al3 +, (d) Ag+ and Hg2 +

(a) Precipitation of the group 4 cations (Figure 17.22) requires a basic medium. Why is this so? (b) What is the most significant difference between the sulfides precipitated in group 2 and those precipitated in group 3? (c) Suggest a procedure that would serve to redissolve the group 3 cations following their precipitation.

A student who is in a great hurry to finish his laboratory work decides that his qualitative analysis unknown contains a metal ion from the insoluble phosphate group, group 4 (Figure 17.22). He therefore tests his sample directly with (NH4hHP04, skipping earlier tests for the metal ions in groups 1, 2, and 3. He observes a precipitate and concludes that a metal ion from group 4 is indeed present. Why is this possibly an erroneous conclusion?

Derive an equation similar to the HendersonHasselbalch equation relating the pOH of a buffer to the pK0 of its base component.

Benzenesulfonic acid is a monoprotic acid with pK, = 2.25. Calculate the pH of a buffer composed of 0.150 M benzenesulfonic acid and 0.125 M sodium benzensulfonate.

Furoic acid (HCsH303) has a Ka value of 6.76 X 10- 4 at 25 oc. Calculate the pH at 25 oc of (a) a solution formed by adding 25.0 g of furoic acid and 30.0 g of sodium furoate (NaC5H303) to enough water to form 0.250 L of solution; (b) a solution formed by mixing 30.0 mL of 0.250 M HC5H303 and 20.0 mL of 0.22 M NaC5H303 and diluting the total volume to 125 mL; (c) a solution prepared by adding 50.0 mL of 1.65 M NaOH solution to 0.500 L of 0.0850 M HC5H303.

The acid-base indicator bromcresol green is a weak acid. The yellow acid and blue base forms of the indicator are present in equal concentrations in a solution when the pH is 4.68. What is the pKa for bromcresol green?

Equal quantities of 0.010 M solutions of an acid HA and a base B are mixed. The pH of the resulting solution is 9.2. (a) Write the equilibrium equation and equilibriumconstant expression for the reaction between HA and B. (b) If Ka for HA is 8.0 x 10-5, what is the value of the equilibrium constant for the reaction between HA and B? (c) What is the value of K0 for B?

A biochemist needs 750 mL of an acetic acid-sodium acetate buffer with pH 4.50. Solid sodium acetate (CH3COONa) and glacial acetic acid (CH3COOH) are available. Glacial acetic acid is 99% CH3COOH by mass and has a density of 1.05 gfmL. If the buffer is to be 0.15 M in CH3COOH, how many grams of CH3C00Na and how many milliliters of glacial acetic acid must be used?

A sample of 0.2140 g of an unknown monoprotic acid was dissolved in 25.0 mL of water and titrated with 0.0950 M NaOH. The acid required 27.4 mL of base to reach the equivalence point. (a) What is the molar mass of the acid? (b) After 15.0 mL of base had been added in the titration, the pH was found to be 6.50. What is the Ka for the unknown acid?

Show that the pH at the halfway point of a titration of a weak acid with a strong base (where the volume of added base is half of that needed to reach the equivalence point) is equal to pK. for the acid.

Potassium hydrogen phthalate, often abbreviated KHP, can be obtained in high purity and is used to determine the concentrations of solutions of strong bases. Strong bases react with the hydrogen phthalate ion as follows: HP-(aq) + OH-(aq) ----> H20(1) + P 2 -(aq) The molar mass of KHP is 204.2 g/mol and Ka for the HP- ion is 3.1 x 10-6. (a) If a titration experiment begins with 0.4885 g of KHP and has a final volume of about 100 mL, which indicator from Figure 16.7 would be most appropriate? (b) If the titration required 38.55 mL of NaOH solution to reach the end point, what is the concentration of the NaOH solution?

If 40.00 mL of 0.100 M Na2C03 is titrated with 0.100 M HCI, calculate (a) the pH at the start of the titration; (b) the volume of HCl required to reach the first equivalence point and the predominant species present at this point; (c) the volume of HCl required to reach the second equivalence point and the predominant species present at this point; (d) the pH at the second equivalence point.

A hypothetical weak acid, HA, was combined with NaOH in the following proportions: 0.20 mol of HA, 0.080 mol of NaOH. The mixture was diluted to a total volume of 1.0 L, and the pH measured. (a) If pH = 4.80, what is the pKa of the acid? (b) How many additional moles of NaOH should be added to the solution to increase the pH to 5.00?

What is the pH of a solution made by mixing 0.30 mol NaOH, 0.25 mol Na2HP04, and 0.20 mol H3P04 with water and diluting to 1.00 L?

Suppose you want to do a physiological experiment that calls for a pH 6.5 buffer. You find that the organism with which you are working is not sensitive to the weak acid H2X (Kal = 2 X 10-2 ; Ka2 = 5.0 X 10-7 ) or its sodium salts. You have available a 1.0 M solution of this acid and a 1.0 M solution of NaOH. How much of the NaOH solution should be added to 1.0 L of the acid to give a buffer at pH 6.50? (Ignore any volume change.)

How many microliters of 1.000 M NaOH solution must be added to 25.00 mL of a 0.1000 M solution of lactic acid [CH3CH(OH)COOH or HC3H503] to produce a buffer with pH = 3.75?

For each pair of compounds, use Ksp values to determine which has the greater molar solubility: (a) CdS or CuS, (b) PbC03 or BaCr04, (c) Ni(OHh or NiC03, (d) Agi or Ag2S04.

Describe the solubility of CaC03 in each of the following solutions compared to its solubility in water: (a) in 0.10 M NaCI solution; (b) in 0.10 M Ca(N03h solution; (c) 0.10 M Na2C03; (d) 0. 10 M HCl solution. (Answer same, less soluble, or more soluble.)

Tooth enamel is composed of hydroxyapatite, whose simplest formula is Ca5 (P04h0H, and whose corresponding Ksp = 6.8 X 10- 2 7. As discussed in the "Chemistry and Life" box in Section 17.5, fluoride in fluorinated water or in toothpaste reacts with hydroxyapatite to form fluoroapatite, Ca5(P04hF, whose Ksp = 1.0 X 10__,;o. (a) Write the expression for the solubility-constant for hydroxyapatite and for fluoroapatite. (b) Calculate the molar solubility of each of these compounds.

Calculate the solubility of Mg(OHh in 0.50 M NH4Cl.

Seawater contains 0.13% magnesium by mass, and has a density of 1.025 g/mL. What fraction of the magnesium can be removed by adding a stoichiometric quantity of CaO (that is, one mole of CaO for each mole of Mg2 +)?

The solubility-product constant for barium permanganate, Ba(Mn04)u is 2.5 x 10-10 Assume that solid Ba(Mn04h is in equilibrium with a solution of KMn04. What concentration of KMn04 is required to establish a concentration of 2.0 x 10-8 M for the Ba2 + ion in solution?

Calculate the ratio of [Ca2 +] to [Fe2 +] in a lake in which the water is in equilibrium with deposits of both CaC03 and FeC03. Assume that the water is slightly basic and that the hydrolysis of the carbonate ion can therefore be ignored.

The solubility products of PbS04 and SrS04 are 6.3 X 10-7 and 3.2 x 10-7, respectively. What are the values of [5042-], [Pb2 +], and [Sr2 +] in a solution at equilibrium with both substances?

What pH buffer solution is needed to give a Mg2 + concentration of 3.0 x 10-2 M in equilibrium with solid magnesium oxalate?

The solubility product for Zn(OHh is 3.0 x 10-16 The formation constant for the hydroxo complex, Zn(OH)42-, is 4.6 X 10-17. What concentration of OHis required to dissolve O.Q15 mol of Zn(OHh in a liter of solution?

(a) Write the net ionic equation for the reaction that occurs when a solution of hydrochloric acid (HCl) is mixed with a solution of sodium formate (NaCH02). (b) Calculate the equilibrium constant for this reaction. (c) Calculate the equilibrium concentrations of Na+, CC H+, CH02 -,and HCH02 when 50.0 mL of 0.15 M HCl is mixed with 50.0 mL of 0.15 M NaCH02.

(a) A 0.1044-g sample of an unknown monoprotic acid requires 22.10 mL of 0.0500 M NaOH to reach the end point. What is the molecular weight of the unknown? (b) As the acid is titrated, the pH of the solution after the addition of 11.05 mL of the base is 4.89. What is the Ka for the acid? (c) Using Appendix D, suggest the identity of the acid. Do both the molecular weight and Ka value agree with your choice?

A sample of 7.5 L of NH3 gas at 22 oc and 735 torr is bubbled into a 0.50-L solution of 0.40 M HCI. Assuming that all the NH3 dissolves and that the volume of the solution remains 0.50 L, calculate the pH of the resulting solution.

Aspirin has the structural formula At bod temperature (37 C), Ka for aspirin equals 3 X 10- . If two aspirin tablets, each having a mass of 325 mg, are dissolved in a full stomach whose volume is 1 L and whose pH is 2, what percent of the aspirin is in the form of neutral molecules?

What is the pH at 25 oc of water saturated with C02 at a partial pressure of 1.10 atm? The Henry's law constant for C02 at 25 C is 3.1 X 10- 2 mol/L-atrn. The C02 is an acidic oxide, reacting with H20 to form H2C03.

Excess Ca(OH)z is shaken with water to produce a saturated solution. The solution is filtered, and a 50.00-mL sample titrated with HCl requires 11.23 mL of 0.0983 M HCl to reach the end point. Calculate K,p for Ca(OH)z. Compare your result with that in Appendix D. Do you think the solution was kept at 25 C?

The osmotic pressure of a saturated solution of strontium sulfate at 25 oc is 21 torr. What is the solubility product of this salt at 25 C?

A concentration of 10-100 parts per billion (by mass) of Ag + is an effective disinfectant in swimming pools. However, if the concentration exceeds this range, the Ag+ can cause adverse health effects. One way to maintain an appropriate concentration of Ag + is to add a slightly soluble salt to the pool. Using Ksp values from Aprendix D, calculate the equilibrium concentration of Ag in parts per billion that would exist in equilibrium with (a) AgCl, (b) AgBr, (c) Agl.