Consider a water sample in a closed system with Cr,c03-8x10-4 M and pH=8. a) What is the alkalinity of this water sample? b) How much acid is required to reduce the pH of the sample to 6 assuming no exchange of CO₂(g) with the atmosphere (i.e., the system stays closed)? c) What would be the pH of the sample if it came to equilibrium with the atmosphere (Pc02-10-3.5 atm)? Hint: recall that CO₂(g) → CO₂ (aq) has KH-10-¹.5 and [u,co;]= [co,(aq)].

Consider a water sample in a closed system with Cr,c03-8x10-4 M and pH=8. a) What is the alkalinity of this water sample? b) How much acid is required to reduce the pH of the sample to 6 assuming no exchange of CO₂(g) with the atmosphere (i.e., the system stays closed)? c) What would be the pH of the sample if it came to equilibrium with the atmosphere (Pc02-10-3.5 atm)? Hint: recall that CO₂(g) → CO₂ (aq) has KH-10-¹.5 and [u,co;]= [co,(aq)].

Chemistry & Chemical Reactivity

10th Edition

ISBN:9781337399074

Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Publisher:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Chapter20: Environmental Chemistry-earth's Environment, Energy, And Sustainability

Section20.7: Green Chemistry And Sustainability

Problem 2.1ACP: Assume that a sample of hard water contains 50. mg/L of Mg2+ and 150 mg/L of Ca2+, with HCO3 as the...

See similar textbooks

Similar questions

a) What mass of Fe(NH4)2(SO4)2•2H2O(s) is required to prepare a 500 mL solution containing 100 ppm (m:v) in Fe? b) Using this stock solution, what aliquot must be used to prepare calibration solutions, 100 mL volume, of the following concentrations: 0.100 ppm, 0.500 ppm, 2.00 ppm, 4.00 ppm, and 7.00 ppm.
6. (a) Aspirin is a weak acid with a pKa of 3.6 and MW 180 g/mol. The saturation solubility (S) of aspirin in water is 3.3 g /L and a saturated solution has pH equal to your answer to question 4(b) above. What is the solubility of the unionised form (So) of aspirin at this pH using units of mol/L and g/L. (b) Use the So value (just obtained) to calculate the saturation solubility (S) of aspirin in water at pH 6.6 using units of mol/L and g/L. (c) If you modified the aspirin suspension APF12 formula by replacing the purified water with pH 6.6 buffer, would the product be a suspension or solution?
(a) Find the pH of a 1.00 L solution prepared with 12.43 g of base (CH2OH)3CNH2 (MW =121.135 g/mol) plus 4.67 g of its conjugate acid (CH2OH)3CNH3Cl (MW = 157.596 g/mol).(b) How many mL of 0.500 M NaOH should be added to 10.0 g of (CH2OH)3CNH3Cl to give a pH of 7.60in a final volume of 250 mLConsider the pKa of (CH2OH)3CNH3+ = 8.072
One strategy for dealing with the acidification of lakes is periodically to add powered limestone (CaCO3) to them, resulting in pH of 10.95. Calculate the Ca2+ concentration in the lake. CO2 in atmosphere 370 ppm KH,CO2 = 0.033 mol/(L∙atm) Ka1 = 10-6.3 Ka2 = 10-10.3 Ksp,CaCO3 = 4.57 x 10-9
3.A sample is known to contain NaOH, Na2CO3 NaHCO3 or compatible mixture of these together with inert matter. With methyl orange, a 1.100 g sample requires 31.40 mL of HCl ( 1.00 mL is equivalent to 0.0140g CaO). With phenolphthalein indicator, the same weight of sample requires 13.30mL of the acid. What is the percent composition of the sample? ww ww
During the course of the titration of a 20.0 ml of solution containing 0.1 M of NaOH and 0.08 M of hydrazine (NH2-NH2) (Ka = 1.05 x 10) by a 0.2 M standard perchloric acid (HC104) titrant, calculate the pH of the titration mixture after (a) 5.0 ml; (b) 10.0 ml; and (c) 15.0 ml of HC1O4 titrant is added.
The solubility of salts that contain the conjugates of weak acids increase in acidic solution. Given the reaction of CUCN (Ksp = 3.47 x 10-20) with an acid in the form H3O+ shown below, find the equilibrium constant for this reaction given that Ka HCN = 4.9 x 10-10. CUCN(s) + H:O*(aq) = Cu"(aq) + HCN(aq) + H2O(1)
(vi) HCN is a weak acid with a Ka of 4.9 x 10-10. Write an equilibrium expression for this process and calculate the pH for a 0.4 mol L-1 solution of HCN.
Lakes have a natural buffering capacity, especially in regions where limestone gives rise to dissolved calcium carbonate. Write an equation for the effect of a small amount of acid rain containing sulfuric acid if it falls into alake containing carbonate (CO3-2 ) ions. Discuss how the lake will resist further pH changes. What happens if a large excess of acid rain is deposited?
A 0.1530 g of pure KIO3 was dissolved in 50 mL distilled water. This was then added with KI and 6N H₂SO4. After allowing the said solution to stand in the dark for 3 minutes, this was then titrated to end point with 31.35 mL sodium thiosulfate solution, with starch as the indicator. What is the normality of the sodium thiosulfate solution? Atomic weights: K = 39.1 1= 126.9 O= 16.0
Sn2 and N (starting concentration 10 mol/L) should be separated by fractionated precipitation using H₂S. Does a pH range exist allowing to completely precipitate one metal (i.e. the remaining concentration of the precipitated ion has to be 10 mol/L or less) whereas the other stays in solution at its initial concentration? Please calculate this pH range whereby pH values below zero are set to zero! Please illustrate the situation graphically by plotting pH versus in solution/precipitating/fully precipitated for the two ions (9) Solubility product (SnS): 10; (NIS): 1021; Acid constant (H2S): 10-20, r/Concentration of H2S in water: 0.1 mol/L
1. A 0.3516 g sample of commercial phosphate detergent was ignited at a heat to destroy the organic matter. The residue was taken up in hot HCI which converted P to H3PO4. The phosphate was precipitated as MGNH4PO4.6H2O by addition of Mg+ followed by aqueous NH3. After being filtered and washed, the precipitate was converted to Mg2P2O7 by ignition at 1000 OC. This residue weighed 0.2161 g. Calculate the percent P (30.974) in the sample? Ans 17.11%