Textbook Question
Write balanced net ionic equations and the corresponding equilibrium equations for the stepwise dissociation of the triprotic acid H3PO4.
17. Acid and Base Equilibrium
Triprotic Acids and Bases
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Which of the following statements about the ionization of a polyprotic acid, such as H3PO4, is false?
A
Ka1 is much larger than Ka2 and Ka3<.sub>.
B
Ka1 = Kw
C
We can assume [H2PO4−] = [H3O+].
D
[HPO42−] ≈ Ka2
E
The protons are removed from the polyprotic acid in a stepwise fashion.
Verified step by step guidance1
Understand that a polyprotic acid, like H3PO4, can donate more than one proton (H+), and each ionization step has its own equilibrium constant (Ka).
Recognize that for polyprotic acids, the first ionization constant (Ka1) is typically much larger than the second (Ka2) and third (Ka3) because it is easier to remove the first proton than subsequent ones.
Evaluate the statement 'Ka1 = Kw'. Kw is the ion product of water, which is a constant at a given temperature (1.0 x 10^-14 at 25°C). This statement is false because Ka1 is not equal to Kw; Ka1 is specific to the acid's first ionization step.
Consider the statement '[H2PO4−] = [H3O+]'. This is a common assumption in the first ionization step of a polyprotic acid, where the concentration of the conjugate base formed is approximately equal to the concentration of H3O+ produced.
Analyze the statement '[HPO42−] ≈ Ka2'. This is incorrect because [HPO42−] is not approximately equal to Ka2; rather, it is the concentration of the species formed in the second ionization step, and Ka2 is the equilibrium constant for that step.
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