# Calculate the percent ionization of 0.125 M lactic acid

Use information from Appendix D to calculate the pH of (a) a solution that is 0.250 M in sodium formate (HCOONa) and 0.100 M in formic acid (HCOOH); (b) a solution that is 0.510 M in pyridine (C5H5N) and 0.450 M in pyridinium chloride (C5H5NHCl); (c) a solution that is made by combining 55 mL of 0.050 M hydrofluoric acid with 125 mL of 0.10 M sodium fluoride.

For Appendix D, see:
https://online.vitalsource.com/#/books/9781269447980/content/id/app04
pg 1062

(a) Calculate the percent ionization of 0.125 M lactic acid (Ka = 1.4 × 10−4). (b) Calculate the percent ionization of 0.125 M lactic acid in a solution containing 0.0075 M sodium lactate.

A buffer contains 0.15 mol of propionic acid (C2H5COOH) and 0.10 mol of sodium propionate (C2H5COONa) in 1.20 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 0.01 mol of HI?

Consider the titration of 30.0 mL of 0.050 M NH3 with 0.025 M HCl. Calculate the pH after the following volumes of titrant have been added:

(a) 0 mL

(b) 20.0 mL

(c) 59.0 mL

(d) 60.0 mL

(e) 61.0 mL

(f) 65.0 mL

Using the value of Ksp for Ag2S, Ka1 and Ka2 for H2S, and Kf = 1.1 × 105 for AgCl2−, calculate the equilibrium constant for the following reaction:
Ag2S(s) + 4 Cl−(aq) + 2 H+(aq) ⇌ 2 AgCl2−(aq) + H2S(aq)

A solution of Na2SO4 is added dropwise to a solution of 0.010 M in Ba2+ and 0.010 M in Sr2+. (a) What concentration of SO42− is necessary to begin precipitation? (Neglect volume changes. BaSO4: Ksp = 1.1 × 10−10; SrSO4: Ksp = 3.2 × 10−7.) (b) Which cation precipitates first? (c) What is the concentration of SO42− when the second cation begins to precipitate?

A sample of 0.1687 g of an unknown monoprotic acid was dissolved in 25.0 mL of water and titrated with 0.1150 M NaOH. The acid required 15.5 mL of a base to reach the equivalence point. (a) What is the molecular weight of the acid? (b) After 7.25 mL of the base had been added in the titration, the pH was 2.85. What is the Ka for the unknown acid?

The element gallium (Ga) freezes at 29.8°C, and its molar enthalpy of fusion is ΔHfus = 5.59 kJ/mol. (a) When molten gallium solidifies to Ga(s) at its normal melting point, is ΔS positive or negative? (b) Calculate the value of ΔS when 60.0 g of Ga(l) solidifies at 29.8°C.

Predict the sign of ΔSsys for each of the following processes: (a) Molten gold solidifies. (b) Gaseous Cl2 dissociates in the stratosphere to form gaseous Cl atoms. (c) Gaseous CO reacts with gaseous H2 to form liquid methanol, CH3OH. (d) Calcium phosphate precipitates upon mixing Ca(NO3)2(aq) and (NH4)3PO4(aq).

Calculate ΔS° values for the following reactions using tabulated S° values from Appendix C. In each case, explain the sign of ΔS°.
(a) HNO3(g) + NH3(g) → NH4NO3(s)
(b) 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g)
(c) CaCO3(s, calcite) + 2HCl(g) → CaCl2(s) + CO2(g) + H2O(l)
(d) 3 C2H6(g) → C6H6(l) + 6 H2(g)

From the values given for ΔH° and ΔS°, calculate ΔG° for each of the following reactions at 298 K. If the reaction is not spontaneous under standard conditions at 298 K, at what temperature (if any) would the reaction become spontaneous?

The Kb for methylamine (CH3NH2) at 25°C is given in Appendix D. (a) Write the chemical equation for the equilibrium corresponding to Kb. (b) Using Kb’s value, calculate ΔG° for the equilibrium in part (a). (c) What is the value of ΔG at equilibrium? (d) What is the value of ΔG when [H+] = 6.7 × 10-9 M, [CH3NH3+] = 2.4 × 10-3 M, and [CH3NH2] = 0.098 M?

Consider the following three reactions:
(i) Ti(s) + 2 Cl2(g) → TiCl4(g)
(ii) C2H6(g) + 7 Cl2(g) → 2 CCl4(g) + 6 HCl(g)
(iii) BaO(s) + CO2(g) → BaCO3(s)
(a) For each of the reactions, use data in Appendix C to calculate ΔH°, ΔG°, and ΔS° at 25°C. (b) Which of these reactions are spontaneous under standard conditions at 25°C? (c) For each reaction, predict how the change in free energy varies with an increase in temperature.

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