$$pH = -\log[1 \times 10$$^{-4}$$] = 4$$

$$[OH^-] = 1 \times 10$$^{-3}$$ M

Question 1

Which of the following best describes why water is considered a polar molecule, as shown in the diagram?

Accepted Answer

Because the oxygen atom is more electronegative, causing a partial negative charge near oxygen and a partial positive charge near hydrogen.

Question 2

A solution has a hydrogen ion concentration of $$1 	imes 10$$^{-4}$$ M. What is its pH?

Accepted Answer

$$pH = -\log[1 	imes 10$$^{-4}$$] = 4$$

Question 3

Which of the following best explains the function of a buffer in biological systems?

Accepted Answer

A buffer maintains a constant pH by absorbing excess H$$^+ or OH$^-$$ ions.

Question 4

If you add 0.1 M NaOH to pure water, what will happen to the pH compared to adding the same amount to a phosphate buffer?

Accepted Answer

The pH of pure water will change much more rapidly than the phosphate buffer.

Question 5

Which equation correctly represents the dissociation of water?

Accepted Answer

$$H_2O \leftrightarrow H^+ + OH^-$$

Question 6

A solution has a pH of 11. What is the hydroxide ion concentration $$[OH^-]$$?

Accepted Answer

$$[OH^-] = 1 	imes 10$$^{-3}$$ M

Question 7

Why is it important for blood to contain buffers?

Accepted Answer

To maintain homeostasis by keeping blood pH within a narrow range.

Question 8

Which of the following solutions required the greatest volume of NaOH to change its pH by 3 units, according to the experiment?

Accepted Answer

Milk

Question 9

What is the correct formula for calculating pH?

Accepted Answer

$$pH = -\log[H^+]$$

Question 10

Which of the following best describes what happens when a strong acid is added to a buffered solution?

Accepted Answer

The buffer absorbs the excess H$$^+$$ ions, minimizing the pH change.

Question 11

If the pH of a solution decreases from 7 to 4, how does the hydrogen ion concentration change?

Accepted Answer

It increases by 1000 times.

Question 12

Which of the following is NOT a safety precaution when working with strong acids and bases in the lab?

Accepted Answer

Mixing bases with any available chemical to observe reactions.

Question 13

Which of the following best explains why the phosphate solution acted as the best buffer in the experiment?

Accepted Answer

It contains a mixture of weak acids and bases that can absorb or release H$$^+$$ ions.

Question 14

Which of the following statements about the pH scale is correct?

Accepted Answer

A solution with pH 9 is more basic than a solution with pH 7.

Question 15

Which of the following best describes what happens to the pH when a small amount of strong base is added to an unbuffered solution?

Accepted Answer

The pH changes rapidly.

Question 16

Which of the following is the correct chemical equation for the bicarbonate buffer system in blood?

Accepted Answer

$$CO_2 + H_2O$$ \leftrightarrow H_2CO_3 \leftrightarrow H^+ + $$HCO_3$$^-$$

Buffers, pH, and the Chemical Context of Life: General Biology Lab Study Notes

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