Chapter 15

If pure water has both hydronium ions (acid) and hydroxide ions (base) in it, how can it be neutral?

Based solely on concentrations, when is an aqueous solution judged to be acidic? Give two answers to this question.

Based solely on concentrations, when is an aqueous solution judged to be basic? Give two answers to this question.

True or false? Even in a very basic aqueous solution, there are some \(\mathrm{H}_{3} \mathrm{O}^{+}\) ions present. Explain your answer.

Consider the following statement: "As the \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration in an aqueous solution increases, the \(\mathrm{OH}^{-}\) concentration must decrease." (a) Why is this true? (b) Is there any way to stop the \(\mathrm{OH}^{-}\) concentration from decreasing as you increase the \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration?

True or false? In an aqueous solution at \(25^{\circ} \mathrm{C}\), you will always get the same number when you multiply the equilibrium \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration by the equilibrium OH \(^{-}\) concentration. Explain your answer.

An aqueous solution has an \(\mathrm{H}_{3} \mathrm{O}\) ' concentration of \(1.0 \mathrm{M}\). What is the OH concentration? Is this solution acidic or basic? Justify your answer.

An aqueous solution has an OH concentration of \(1.0 \times 10^{-11} \mathrm{M}\). What is the \(\mathrm{H}_{3} \mathrm{O}\) concentration? Is this solution acidic or basic? Justify your answer

A solution is prepared by dissolving \(2.50\) moles of \(\mathrm{LiOH}\) in enough water to get \(4.00 \mathrm{~L}\) of solution. What are the \(\mathrm{OH}^{-}\) and the \(\mathrm{H}_{3} \mathrm{O}^{+}\) molar concentrations?

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?

A solution is prepared by dissolving \(2.40 \mathrm{~g}\) of \(\mathrm{Mg}(\mathrm{OH})_{2}\) in enough water to get \(4.00 \mathrm{~L}\) of solution. What are the \(\mathrm{OH}^{-}\) and \(\mathrm{H}_{3} \mathrm{O}^{+}\) molar concentrations? (Hint: You need to calculate the molar mass of \(\mathrm{Mg}(\mathrm{OH})_{2}\).)

A solution is prepared by dissolving \(2.00\) moles of \(\mathrm{HNO}_{3}\) in enough water to get \(800.0 \mathrm{~mL}\) of solution. What are the \(\mathrm{H}_{3} \mathrm{O}^{+}\) and the \(\mathrm{OH}\) molar concentrations?

A solution has an \(\mathrm{OH}\) concentration of \(10^{-11} \mathrm{M}\). (a) Is this solution basic or acidic? Explain how you know. (b) Without using a calculator, explain how you could quickly determine the \(\mathrm{H}_{3} \mathrm{O}\) concentration of this solution. What is that concentration?

What is a base-10 logarithm?

Without using a calculator, what is the base- 10 logarithm of \(10^{-34}\) ?

Without using a calculator, what is the base-10 logarithm of \(10^{13}\) ?

Without using a calculator, what is the base- 10 logarithm of \(10^{0}\) ? Of 1?

The base- 10 logarithm of 60 is a number: (a) Between \(-2\) and \(-1\) (b) Between \(-1\) and 0 (c) Between 0 and 1 (d) Between 1 and 2 (e) Between 2 and 3

The base-10 logarithm of \(0.73\) is a number: (a) Between \(-2\) and \(-1\) (b) Between \(-1\) and 0 (c) Between 0 and 1 (d) Between 1 and 2 (e) Between 2 and 3

True or false? As a solution's acidity increases, its pH decreases.

Write the mathematical definition of \(\mathrm{pH}\), and explain what fundamental change there would be in the pH scale if you left the negative out.

Why is a pH of 7 equal to neutrality?

What is the \(\mathrm{pH}\) of a solution whose hydronium ion concentration is \(0.0010 \mathrm{M?}\) Is the solution acidic or basic?

What is the \(\mathrm{pH}\) of a solution whose \(\mathrm{H}_{3} \mathrm{O}\) concentration is \(10^{6} \mathrm{M}\) ? Is the solution acidic or basic?

What is the \(\mathrm{pH}\) of a solution whose \(\mathrm{H}_{3} \mathrm{O}\) concentration is \(6.40 \times 10^{-9} \mathrm{M} ?\) Is the solution acidic or basic?

What is the \(\mathrm{pH}\) of a solution whose \(\mathrm{OH}^{-}\) concentration is \(10^{-14} \mathrm{M} ?\) Is the solution acidic or basic?

What is the \(\mathrm{pH}\) of a solution whose \(\mathrm{OH}\) concentration is \(2.0 \times 10^{-3} \mathrm{M}\) ? Is the solution acidic or basic?

Solution A has a \(\mathrm{pH}\) of 3 . Solution \(\mathrm{B}\) has a \(\mathrm{pH}\) of 6 . Which solution is more acidic, and by how much?

What is the \(\mathrm{pH}\) of a solution whose \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration is \(10.0 \mathrm{M}\) ? Is this solution acidic or basic?

The \(\mathrm{pH}\) of a solution is 4 . (a) What is the \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration? (b) What is the \(\mathrm{OH}^{-}\) concentration? (c) Is this solution acidic or basic?

The pH of a solution is 8 . (a) What is the \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration? (b) What is the OH concentration? (c) Is this solution acidic or basic?

The pH of a solution is \(-1\). What are the \(\mathrm{H}_{3} \mathrm{O}^{+}\) and OH concentrations? Is the solution acidic or basic?

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?

Aniline, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\), is a weak base, with a lone pair of electrons on the nitrogen atom. (a) According to Bronsted and Lowry, what must aniline do to act as a base? (b) Why can aniline act as a base? (c) What are the molecular formula and charge of the conjugate acid of aniline?

Can a weak acid and its conjugate base ever have the same charge? Explain.

Knowing that aniline (Problem \(15.158\) ) is a weak base, is its conjugate acid a weak acid or a strong acid?

Nitric acid, \(\mathrm{HNO}_{3}\), is a very strong acid. Solutions of sodium nitrate, \(\mathrm{NaNO}_{3}\), contain lots of nitrate ions. Would you expect such a solution to be acidic, basic, or neutral? Explain. (Hint: Think about the conjugate base of \(\mathrm{HNO}_{3}\).)

Why is the conjugate base of a weak acid like acetic acid often referred to as a salt of the acid?

If \(\mathrm{CI}^{-}\) is the conjugate base of \(\mathrm{HCl}\), why isn't an aqueous solution of \(\mathrm{NaCl}\) acidic?

What do we mean when we say that a solution has buffering ability?

Does pure liquid water have any buffering ability? Explain.

What is the general recipe for making a buffer? Explain the function of each ingredient.

Name one biological system in which control of \(\mathrm{pH}\) (buffering ability) is important, and name the buffer that is in control.

Can a buffer resist \(\mathrm{pH}\) changes for any added amount of strong acid or base? Explain.

Does a mixture of carbonic acid, \(\mathrm{H}_{2} \mathrm{CO}_{3}\), a weak acid, and sodium bicarbonate, \(\mathrm{NaHCO}_{3}\), in water constitute a buffer? If no, explain why. If yes, explain why and use chemical equations to show what happens when either \(\mathrm{OH}\) or \(\mathrm{H}_{3} \mathrm{O}^{+}\) is added to the solution.

When a strong acid is added to a buffer, the \(\mathrm{pH}\) changes a little bit. (a) Does the pH increase or decrease? (b) Why does the \(\mathrm{pH}\) change at all? Why doesn't the buffer hold the \(\mathrm{pH}\) constant?

One way to make an acetic acid buffer is to mix substantial amounts of acetic acid and its salt, sodium acetate, in water. Another way to make the same buffer is to add a substantial amount of acetic acid to water and then add half as much \(\mathrm{NaOH}\) to the water. Explain how and why this latter procedure makes a buffer.

It is possible to make two completely different buffers using the dihydrogen phosphate ion, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) (a) In one buffer, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) serves as the weak acid. What is the conjugate weak base? (b) Write the equations that show how the buffer in part (a) works when either \(\mathrm{H}_{3} \mathrm{O}^{+}\) or \(\mathrm{OH}^{-}\) is added. (c) In the other buffer, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) serves as the weak base. What is the conjugate weak acid? (d) Write the equations that show how the buffer in part (c) works when either \(\mathrm{H}_{3} \mathrm{O}^{+}\) or \(\mathrm{OH}^{-}\) is added.

Suppose \(2.0\) moles of sodium acetate, \(\mathrm{NaO}_{2} \mathrm{CCH}_{3}\), are dissolved in some water and then \(1.0 \mathrm{~L}\) of a \(1.0 \mathrm{M} \mathrm{HCl}\) solution is added. (a) Write the chemical reaction that occurs. (b) After the reaction, what are the predominant species in solution? How many moles of each species are there? (c) Is the resulting solution a buffer? If yes, explain why.

A buffer works by replacing added strong acid with weak acid. Explain how.