QUESTION 1 Calculate the pH of a 0.126 M solution of ascorbic acid (CgHgO6). Ka1=8.00X10-5 and Ka2=1.60x10-12, (Report your answer to the thousandths place.)

If the Kb of a weak base is 8.7E-9, what is the Ka of its conjugate acid? Do not include any units with your answer.  Your answer should have two significant figures. For scientific notation use the same as that in the problem (i.e. 3.3E-2 is the same as 3.3 x 10-2).

Which is false? O CH3COOH is a weak acid H,PO4: HPO,2 is the principle weak acid/conjugate base pair responsible for maintaining a neutral pH in blood O H,PO4 is both a weak acid and a weak base NAOH is a strong base An amino acid is both a weak acid and a weak base

Using the constant below, calculate the molar concentration (M) of acid H3O* in solution if the concentration of OH- is 3.4 x 10-5. Kw = 1.0 × 10-14 M²= [H3O*][OH]. %3D O a) 2.9 x 10-10 O b) 2.0 x 10-10 O) 3,4 x 109 d) 2.4 x 10-10 e) 3.4 x 109

A weak base has a Kb = 1.5 x 109. What is the value of the pKa for the conjugate acid? Enter your answer, rounded to the correct number of significant figures, into the box below. Do not use scientific notation. If your answer is a decimal, please include a 0 before the decimal point. pka = type your answer...

What is the final pH of a 2.86 M solution of arsenic acid, H3AsO4, where Ka1 = 5.5 x 10^-3, Ka2 = 1.7 x 10^-7, and Ka3 = 5.1 x 10^-12 Give your final solution to two decimal places (i.e. x.xx)

[H+] = .0001 [H+] = 0.0030 [ОН-] — 8.6 е -8 РОН 3 4.9 pH = 7.3 %3D At 37 degrees C, which of these solutions can we say is nearly neutral?

Each row of the table below describes an aqueous solution at about 25 °C. Complete the table. That is, fill in any missing entries in third column. Be sure each entry you write includes the correct number of significant digits. [1,0] solution pH x10 -7 A 3.1 х 10 mol/L - 11 mol/L 1.2 x 10 В 0.0025 mol/L

A certain weak acid, HAHA, has a Ka�a value of 4.4×10−7.  Calculate the percent ionization of HAHA in a 0.10 M� solution. Express your answer as a percent using two significant figures.

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