Calculate the K_a of a monoprotic weak acid if a 0.50 M solution has 1.14% ioniz

Question

Calculate the K_a of a monoprotic weak acid if a 0.50 M solution has 1.14% ionization. a) 6.57 times 10^-5 b) 3.84 times 10^-6 c) 2.66 times 10^-4 d) 2.59 times 10^-5 e) 7.49 times 10^-5. How many grams of barium hydroxide would you need to prepare 400 mL of a solution with pH = 12.48? a) 5.17 g b) 4.14 g c) 2.58 g d) 2.07 g e) 1.04 g Calculate the K_b for diethylamine, (C_2H_5)_2NH, if the pH of a 0.975 M (C_2H_5)_2NH solution is 12.55. a) 4.84 times 10^-4 b) 9.66 times 10^-10 c) 3.91 times 10^-26 d) 7.74 times 10^-26 e) 1.34 times 10^-3 Which of the following is true for a 0.15 M solution of a weak acid, HA, with K_a = 1.4 times 10^-4? a) The acid would be 3.1% dissociated. b) The equilibrium concentration of the acid, HA, would be 0.145 M. c) The solution would have a pH = 2.34 d) The equilibrium concentration of the conjugate base would be 4.58 times 10^-3 M. e) All of the above are true.

Explanation / Answer

Q25.

find mass of Ba(OH)2 for

V = 400 mL = 0.4 L and

pH = 12.48

then pOH = 14-pH = 14-12.48 = 1.52

pOH = 1.52

[OH-] = 10^-pOH = 10^-1.52 = 0.030199 M

now

Ba(OH)2 --> Ba+2 + 2OH-

so

Ba = 1/2*[OH-] = 1/2*0.030199 = 0.0150995 M of Ba2+

then

V = 0.4 L

mol = MV = (0.0150995)*0.4 = 0.0060398 mol

mass = mol*Mw = (0.0060398)(171.34) = 1.03485 g

choose E

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