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Chapter 15: Problem 83

What is the value of the equilibrium constant for the dissociation of acetic acid's proton? What does this value tell you?

Short Answer

Expert verified
The equilibrium constant K_a for the dissociation of acetic acid's proton is \(1.8 \times 10^{-5}\). This value indicates that acetic acid is a weak acid that does not dissociate completely in water.

Step by step solution

01

Write the dissociation equation for acetic acid

First, let's write the chemical equation for the dissociation of acetic acid (CH3COOH). Acetic acid releases a proton (H+) and forms the acetate ion (CH3COO-). The equation is: \[CH_3COOH \rightleftharpoons CH_3COO^- + H^+\]
02

Write the expression for the equilibrium constant, K_a

Next, we write the expression for the equilibrium constant, K_a, for the dissociation of acetic acid. The K_a is given by \[K_a = \frac{[CH_3COO^-][H^+]}{[CH_3COOH]}\] Here the square brackets represent the molar concentrations of each species in the solution at equilibrium.
03

Find the value of K_a

To find the value of K_a for acetic acid, we can consult a table of K_a values for different acids or look it up in reference materials. The K_a for acetic acid is \(1.8 \times 10^{-5}\).
04

Interpret the value of K_a

Now, let's interpret the value of K_a for acetic acid. A K_a value of \(1.8 \times 10^{-5}\) tells us that the acid strength of acetic acid is relatively weak. In general, a K_a value of less than \(10^{-4}\) suggests that the acid is weak and does not dissociate completely in water. In conclusion, the equilibrium constant K_a for the dissociation of acetic acid's proton is \(1.8 \times 10^{-5}\). This value indicates that acetic acid is a weak acid that does not dissociate completely in water.

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Most popular questions from this chapter

How would the Bronsted-Lowry theory explain that ammonia is a weak base?

Two students each dissolve 1 mole of acid in enough water to get 1 L of solution. Student \(A\) uses acetic acid, and student \(B\) uses hydrochloric acid. Do the solutions have the same pH? If not, which has the lower pH? Why should the two pH values be different?

The oxide ion, \(\mathrm{O}^{2-}\), present in sodium oxide \(\left(\mathrm{Na}_{2} \mathrm{O}\right)\) reacts violently with water to produce a highly basic solution. The hydride ion, \(\mathrm{H}^{-}\), in sodium hydride (NaH) does the same. (a) Write a balanced total ionic equation for the reaction of sodium oxide with water. (b) In terms of the Bronsted-Lowry definition, how are oxide and hydride similar? (c) What is it about the hydride and oxide ions that allow them to do what they do in water?

A solution is prepared by dissolving \(0.250\) mole of \(\mathrm{Ba}(\mathrm{OH})_{2}\) in enough water to get \(4.00 \mathrm{I}\). of solution. What are the \(\mathrm{OH}^{-}\) and \(\mathrm{H}_{3} \mathrm{O}^{+}\) molar concentrations?

Why is a pH of 7 equal to neutrality?
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