Determine the pH of a solution that is prepared by dissolving 8 g of Ba(OH)2 in 3.8 L water at 25°C. Assume that the volume of the solvent does not change upon addition of the solute. Atomic Weights: Ba = 137.3 g/mol; O = 16.00 g/mol; H = 1.008 g/mol.

A diprotic acid, H, A, has acid dissociation constants of K₁1 = 4.76 x 10-4 and K₁2 = 3.59 x 10-12. Calculate the pH and molar concentrations of H₂A, HA, and A2- at equilibrium for each of the solutions. A 0.205 M solution of H₂A. pH = [HA] = A 0.205 M solution of NaHA. pH= [HA] = M M [H₂A] = [A²-] = [H₂A] = [A²-] = M M M M

Consider the following data on some weak acids and weak bases: name hydrocyanic acid hydrofluoric acid acid solution 0.1 M KCN 0.1 M HONH3Br 0.1 M NaCl formula HCN 0.1 M C₂H5NHCI HF Ra X K 4.9 × 10 - 10 6.8 × 10 pH Use this data to rank the following solutions in order of increasing pH. In other words, select a '1' next to the solution that will have the lowest pH, a '2' next to the solution that will have the next lowest pH, and so on. 4 ✓ choose one 1 (lowest) 2 3 4 (highest) choose one S base 3 name Kb formula C5H₂N 1.7 x 10 -9 pyridine - 8 hydroxylamine HONH₂ 1.1 × 107

For the acid HA, the acid dissociation constant, Ka, is 1.1 x 105. After 48 mL of 0.9 M NaOH is added to 48 mL of a 1.8 M HA solution, what will the pH of the resulting solution be?

A chemist must prepare 350.0 mL of hydrobromic acid solution with a pH of 1.80 at 25 °C. He will do this in three steps: • Fill a 350.0 mL volumetric flask about halfway with distilled water. • Measure out a small volume of concentrated (7.0M) stock hydrobromic acid solution and add it to the flask. • Fill the flask to the mark with distilled water. Calculate the volume of concentrated hydrobromic acid that the chemist must measure out in the second step. Round your answer to 2 significant digits. mL x10 Explanation Check O 2022 McGraw Hill LLC, All Rights Reserved. Terms of Use Privacy Center Acce: 80°F Cloudy DELL Esc F2 F3 F4 F5 F6 F7 F8 F9 F10 F11 F12 PrtScr Insert Delete F1 "R A # $4 & Backspace 1 3 4 6 80 E Y *00

The pH of an aqueous solution of sulfurous acid, H2SO3, that has a concentration of 0.60 mol/L and dissociation constants of Ka1 1.5 x 102 and Ke = 6.3 x 10 8. %3D

Atmospheric CO2 levels have risen by nearly 20% over the past 40 years from 315 ppm to 380 ppm Part A Given that the average pH of clean, unpolluted rain today is 5.4, determine the pH of unpolluted rain 40 years ago. Assume that carbonic acid (H2CO3) formed by the reaction of CO2 and water is the only factor influencing pH. CO2(g)+NO11) = HgCO(aq) Express your answer using one decimal place. pH = 5.4 ✓ Correct Part B Previous Answers What volume of CO2 at 25 °C and 1.0 atm is dissolved in a 13.0-L bucket of today's rainwater? Express your answer using two significant figures. Η ΕΠΙ ΑΣΦ Vco₂ = 0 ? Submit Previous Answers Request Answer mL 1015

Calculate the concentrations of all species in a 1.01 M Na₂SO3 (sodium sulfite) solution. The ionization constants for sulfurous acid are Kal = 1.4 x 10-2 M and K₁2 = 6.3 x 10-8M. [Na+] = [HSO3] = [OH-] = M [so-] = M [H₂SO3] = M [H+] = M M M

A solution is prepared at 25 °C that is initially 0.69M in dimethylamine ((CH3)2NH), a weak base with K = 5.4 × 10-4, and 2.4M in dimethylammonium ((CH3)2NH₂CI). Calculate the pH of the solution. Round your answer to 2 decimal places. chloride pH = ☐

A solution is prepared at 25 °C that is initially 0.15M in propanoic acid (HC₂H-CO₂), a weak acid with K=1.3 × 105, and 0.18M in potassium propanoate (KC₂H,CO₂). Calculate the pH of the solution. Round your answer to 2 decimal places. pH = 0 X S