A bottled carbonated mineral water (i.e. closed system) has a pH of 4.9 and an alkalinity of 4×10-³ Considering the species H₂O, H*, OH, H₂CO³, HCO³ and CO² and the reactions: H₂O = OH + H* Kw = 10-14 H₂CO3 = HCO3 + H* K₁ = 10-6.35 HCO3 = CO3+H* K₂ = 10-10.33 Calculate the total C₁ = [H₂CO3 ]+[HCO3]+[CO3²] (Hint: you can use a Tableau). Verify any assumptions that you may have made. Now the bottle is opened to the atmosphere and CO₂ is allowed to degas to reach a new equilibrium at pH = 8.1 (Hint: consider what is the dominant carbonate specie and conservative entitie/s that de not change). a. b. C. Calculate how much CO₂ is degassed (or total carbon lost).

A bottled carbonated mineral water (i.e. closed system) has a pH of 4.9 and an alkalinity of 4×10-³ Considering the species H₂O, H, OH, H₂CO³, HCO³ and CO² and the reactions: H₂O = OH + H Kw = 10-14 H₂CO3 = HCO3 + H* K₁ = 10-6.35 HCO3 = CO3+H* K₂ = 10-10.33 Calculate the total C₁ = [H₂CO3 ]+[HCO3]+[CO3²] (Hint: you can use a Tableau). Verify any assumptions that you may have made. Now the bottle is opened to the atmosphere and CO₂ is allowed to degas to reach a new equilibrium at pH = 8.1 (Hint: consider what is the dominant carbonate specie and conservative entitie/s that de not change). a. b. C. Calculate how much CO₂ is degassed (or total carbon lost).

Chemistry: An Atoms First Approach

2nd Edition

ISBN:9781305079243

Author:Steven S. Zumdahl, Susan A. Zumdahl

Publisher:Steven S. Zumdahl, Susan A. Zumdahl

Chapter13: Acids And Bases

Section: Chapter Questions

Problem 3RQ

See similar textbooks

Similar questions

At 50 °C, the ion-product constant for H2O has the valueKw = 5.48 x 10-14. (a) What is the pH of pure water at50 °C? (b) Based on the change in Kw with temperature,predict whether ΔH is positive, negative, or zero for theautoionization reaction of water:2 H2O(l) ⇌ H3O+(aq) + OH-(aq)
A diprotic acid, H, A, has acid dissociation constants of K,j = 1.45 x 10-4 and K2 = 5.12 x 10-". Calculate the pH and molar concentrations of H,A, HA, and A²- at equilibrium for each of the solutions. A 0.184 M solution of H,A. pH = 2.2865 [H,A] = 0.1788 [HA"] = 5.17 xI10-3 [A?-] = M 5.12 x10-11 M A 0.184 M solution of NaHA. pH= [H,A] = M. [HA-] = [A2-] = M M
A solution is prepared at 25 °C that is initially 0.36M in chlorous acid (HClO₂), a weak acid with K=1.1 × 10¯², and 0.29M in potassium chlorite (KC1O₂). -2 Calculate the pH of the solution. Round your answer to 2 decimal places. pH = 0 X Ś
COHSOH(ag) + H2On + CeHsO (aq) + H3O*(a9) Ka= 1.12 x 10-10 (a) Phenol is a weak acid that partially dissociates in water according to the equation above. Write the equilibrium-constant expression for the dissociation of the acid in water. (b) What is the pH of a 0.75 M CaHsOH(ag) solution? (C) For a certain reaction involving CaHsOH(ag) to proceed at a significant rate, the phenol must be primarily in its deprotonated form, C3H5O (eg). In order to ensure that the CsHsOH(aq) is deprotonated, the reaction must be conducted in a buffered solution. On the number scale below, circle each pH for which more than 50 percent of the phenol molecules are in the deprotonated form (CoHsO (aq). Justify your answer. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Justification: (d) CeHsOH(ag) reacts with NaOH(ag). Write a net ionic equation representing this reaction (aka: invasion equation). (e) What is the pH of the resulting solution when 30 mL of 0.40 M CSH5OH(aq) is added to 25 mL of 0.60 M NAOH. Show all work…
Propionic acid, HC3H5O2, has Ka= 1.34 x 10–5. (a) What is the molar concentration of H3O+ in 0.15 M HC3H5O2 and the pH of the solution? (b) What is the Kb value for the propionate ion, C3H5O2–? (c) Calculate the pH of 0.15 M solution of sodium propionate, NaC3H5O2. (d) Calculate the pH of solution that contains 0.12 M HC3H5O2 and 0.25 M NaC3H5O2.
Determine the pH of (a) A 0.10 M CH3COOH solution (ICE) (b) A solution that is 0.10 M CH3COOH and 0.10 M CH3COONa. Ka for CH3COOH = 1.8 X 10-5.
Which is the best solution to buffer at pH = 10.50? (A) A buffer solution containing 0.25 M H2CO3 and 0.45 M Na2CO3; (Ka (H2CO3) = 4.3 x 10–7) (B) A buffer solution containing 0.25 M NaHCO3 and 0.45 M Na2CO3; (Ka (HCO3–) = 5.6 x 10–11) (C) A buffer solution containing 0.12 M NaHCO3 and 0.22 M Na2CO3; (Ka (HCO3–) = 5.6 x 10–11) (D) A buffer solution containing 0.12 M NH4Cl and 2.2 M NH3; (Ka (NH4+) = 5.6 x 10–10)
2. Determine [H3O+] and [OH] concentrations (in M to two decimal places) for each: (a) Pure water at 10 °C (K = 0.293 x 10-14) W (b) An aqueous solution with a pH of 3.23 A: 5.41x108; 5.41×10-8 A: 5.89x104; 1.70×10-11
Consider the weak bases below and their Kp values: NH3 Ky = 1.8 x 10-5 C2H5NH2 Kp = 5.6 x 10-4 C5H5N Kp = 1.7 x 10-9 Arrange the conjugate acids of these weak bases in order of increasing acid strength. O CsH;NH< NH,*< C2H5NH3* C2H5NH;*= NH,= C5H5NH* C2H5NH3*< NH4*< C5H5NH* O NH,*< C,H5NH;*
culate the pH of a buffer solution 1) Calculate the pH of a solution prepared by dissolving 2.05 g of sodium acetate, CH3COONa, in 85.0 mL of 0.10 Macetic Assume the volume change upon dissolving the sodium acetate is negligible. K₂ of CH3COOH is acid, CH3COOH(aq). 1.75 x 10-5. pH =
Calculate the pH of a buffer solution containing 0.67 M C3H,COOH and 0.33 M NaC3H,COO. The Ka of C3H,COOH is 1.514 x 10-5.
Given that Ka's for hydrofluoric acid (HF) and boric acid (H3BO3) are 6.3 x 10-4 and 5.4 x 10-10, respectively, calculate the pH of the following solutions: (a) The mixture from adding 50 mL 0.2 M HF to 50 mL 0.5 M sodium borate (NaH2BO3). (b) The mixture from adding an additional 150 mL 0.2 M HF to the solution in (a), i.e., a total of 200 mL 0.2 M HF was added to 50 mL 0.5 M NaH2BO3.