Quiz: Water, pH and Buffer Solutions — 16 questions

1. Which description correctly identifies the structure of a water molecule?

A. Two oxygen atoms and one hydrogen atom joined by covalent bonds at an angle of 120°

B. Two hydrogen atoms and one oxygen atom joined by covalent bonds at an angle of 104.5°

C. One hydrogen atom and two oxygen atoms joined by ionic bonds at an angle of 109.5°

D. One hydrogen atom and one oxygen atom joined by a hydrogen bond at an angle of 90°

Answer: (B)

Explanation

Water contains two hydrogen atoms covalently bonded to one oxygen atom, and its H–O–H angle is 104.5°. The molecule is not built from ionic or hydrogen bonds between its constituent atoms.

2. Why can a water molecule be globally neutral while still having partial charges?

A. The molecule contains unequal numbers of protons and electrons, producing a permanent net charge

B. The covalent bonds transfer complete electrons to oxygen, creating two fully charged ions

C. Oxygen attracts shared electrons more strongly, making oxygen partially negative and hydrogen partially positive

D. Hydrogen attracts shared electrons more strongly, making hydrogen partially negative and oxygen partially positive

Answer: (C)

Explanation

Oxygen is more electronegative than hydrogen, so electron density is distributed unevenly even though the molecule’s total charge is zero. This creates partial charges rather than complete ionic charges.

3. What defines a hydrogen bond between two water molecules?

A. A sharing of electron pairs between oxygen atoms in two neighboring molecules

B. An attraction between the partial positive charge of one molecule and the partial negative charge of another

C. A repulsion between similarly charged hydrogen atoms on adjacent molecules

D. A transfer of an electron from one water molecule to a neighboring water molecule

Answer: (B)

Explanation

A hydrogen bond is an electrostatic attraction between oppositely charged partial regions on different molecules. Covalent bonds, by contrast, involve sharing electrons within a molecule.

4. How does hydrogen bonding differ among ice, liquid water, and water vapor?

A. Ice lacks hydrogen bonds, liquid water has four permanent bonds, and vapor contains transient networks

B. Ice has four stable bonds per molecule, liquid water has short-lived fewer bonds, and vapor lacks intermolecular interactions

C. Ice has fewer short-lived bonds, liquid water has four stable bonds, and vapor has the strongest network

D. Ice and liquid water have identical stable networks, while vapor contains additional intermolecular bonds

Answer: (B)

Explanation

Each water molecule in ice participates in a stable crystalline network of four hydrogen bonds, whereas liquid water has fewer bonds that last roughly $10^{-8}$ to $10^{-11}$ seconds. Water vapor has no intermolecular interactions; it is not a stronger hydrogen-bonded state.

5. What happens when water interacts with the ions in a salt crystal?

A. Water links all ions into a continuous covalent network that remains intact in solution

B. Water converts the ions into neutral atoms before they can leave the crystal lattice

C. Water strengthens the crystal’s internal electrostatic attractions and prevents ions from separating

D. Water organizes oppositely charged ions and promotes dissolution of the crystal into separate ions

Answer: (D)

Explanation

Hydrogen bonds between water molecules and ions help organize the oppositely charged species and promote dissolution into ions. The dissolved ions then interact more weakly with one another than they did within the crystal.

6. Which group cannot form hydrogen bonds with water?

A. Hydroxyl group

B. Amine group

C. Methyl group

D. Carbonyl group

Answer: (C)

Explanation

Hydroxyl, carbonyl, carboxyl, and amine groups can form hydrogen bonds with water. A methyl group lacks the appropriate polar functionality for such bonding.

7. Which behavior best distinguishes hydrophilic molecules from hydrophobic molecules?

A. Hydrophilic and hydrophobic molecules both dissolve because water forms equivalent interactions with each type

B. Hydrophilic molecules form ionic crystals in water, whereas hydrophobic molecules become fully charged in solution

C. Hydrophilic molecules hydrogen-bond with water and tend to dissolve, whereas hydrophobic molecules separate from water

D. Hydrophilic molecules are nonpolar and separate from water, whereas hydrophobic molecules hydrogen-bond and dissolve

Answer: (C)

Explanation

Hydrophilic molecules can form hydrogen bonds with water and therefore tend to be soluble. Hydrophobic molecules are nonpolar, lack these interactions, and tend to separate from water.

8. Which expression correctly represents the ionic product of water at ordinary conditions?

A. $K_e=[\mathrm{H_2O}][\mathrm{OH^-}]=10^{-14}\,(\mathrm{mol/L})^2$

B. $K_e=\frac{[\mathrm{H^+}]}{[\mathrm{OH^-}]}=10^{-14}\,(\mathrm{mol/L})^2$

C. $K_e=[\mathrm{H_2O}][\mathrm{H^+}]=10^{-14}\,(\mathrm{mol/L})^2$

D. $K_e=[\mathrm{H^+}][\mathrm{OH^-}]=10^{-14}\,(\mathrm{mol/L})^2$

Answer: (D)

Explanation

The ionic product is defined as the product of the hydrogen-ion and hydroxide-ion concentrations, with a value of $10^{-14}\,(\mathrm{mol/L})^2$. Including water concentration would describe a different equilibrium expression rather than the ionic product.

9. What pH results when a solution has a hydrogen-ion concentration of $10^{-7}\,\mathrm{mol/L}$ and is pure water?

A. 10

B. 2

C. 14

D. 7

Answer: (D)

Explanation

Using $\mathrm{pH}=-\log_{10}[\mathrm{H^+}]$, a concentration of $10^{-7}\,\mathrm{mol/L}$ gives pH 7, and pure water has equal hydrogen-ion and hydroxide-ion concentrations. A pH of 2 would correspond to a much higher hydrogen-ion concentration.

10. A solution contains more hydroxide ions than hydrogen ions; how should it be classified?

A. Basic

B. Saturated

C. Neutral

D. Acidic

Answer: (A)

Explanation

A solution is basic when its hydroxide-ion concentration exceeds its hydrogen-ion concentration. Neutrality requires the two concentrations to be equal, while acidity requires more hydrogen ions.

11. What distinguishes a weak acid–base pair from a strong acid or base?

A. The strong substance reaches equal acid and base concentrations at equilibrium

B. The weak pair dissociates completely and contains one dominant species

C. The strong substance dissociates incompletely and retains both conjugate species

D. The weak pair dissociates incompletely and contains both acid and base species

Answer: (D)

Explanation

A weak acid–base pair dissociates incompletely, so acid and conjugate-base species coexist. Complete dissociation characterizes a strong acid or base rather than a weak pair.

12. For a weak acid–base pair with equal concentrations of acid and conjugate base, which relationship gives the pH?

A. $\mathrm{pH}=\mathrm{p}K_a+1$

B. $\mathrm{pH}=\mathrm{p}K_a$

C. $\mathrm{pH}=-\mathrm{p}K_a$

D. $\mathrm{pH}=\mathrm{p}K_a-1$

Answer: (B)

Explanation

The Henderson–Hasselbalch relationship is $\mathrm{pH}=\mathrm{p}K_a+\log\frac{[\mathrm{base}]}{[\mathrm{acid}]}$, so equal concentrations make the logarithmic term zero. Adding or subtracting one would require a tenfold concentration ratio.

13. What does the pKa represent for a weak acid and its conjugate base?

A. The pH at which their concentrations are equal

B. The pH at which the solution contains no conjugate base

C. The concentration at which hydroxide ions equal hydrogen ions

D. The pH at which the acid has fully dissociated

Answer: (A)

Explanation

The pKa is the pH where the weak acid and its conjugate base have equal concentrations. Full dissociation is associated with strong acids, not the defining meaning of pKa.

14. What is the defining property of a buffer solution?

A. Its pH changes only slightly when acid or base is added

B. Its pH remains exactly constant when acid or base is added

C. Its pH is always equal to 7 before any substance is added

D. Its pH changes more than pure water after acid or base is added

Answer: (A)

Explanation

A buffer resists pH changes, so adding an acid or base causes only a small pH variation. It does not maintain an exactly fixed pH, and its initial pH need not be neutral.

15. An acetate/acetic acid buffer starts with equal concentrations of both components and has a pKa of 4.85. What happens after adding $10^{-2}\,\mathrm{mol\,L^{-1}}$ HCl?

A. The acetate concentration becomes $40\times10^{-3}\,\mathrm{mol\,L^{-1}}$, the acetic acid concentration becomes $60\times10^{-3}\,\mathrm{mol\,L^{-1}}$, and the pH becomes 2.00

B. The acetate concentration becomes $60\times10^{-3}\,\mathrm{mol\,L^{-1}}$, the acetic acid concentration becomes $40\times10^{-3}\,\mathrm{mol\,L^{-1}}$, and the pH becomes 5.03

C. Both component concentrations remain $50\times10^{-3}\,\mathrm{mol\,L^{-1}}$, and the pH remains 4.85

D. The acetate concentration becomes $40\times10^{-3}\,\mathrm{mol\,L^{-1}}$, the acetic acid concentration becomes $60\times10^{-3}\,\mathrm{mol\,L^{-1}}$, and the pH becomes 4.67

Answer: (D)

Explanation

Hydrochloric acid converts some acetate into acetic acid, producing the stated concentrations and lowering the pH from 4.85 to 4.67. A pH of 2 would represent the much larger change expected for the same acid addition to pure water.

16. When is a buffer most effective at resisting pH changes?

A. When its pH is far above the pKa of its weak acid–base pair

B. When its pH is fixed at 7 regardless of the pair used

C. When its pH is close to the pKa of its weak acid–base pair

D. When its pH is far below the pKa of its weak acid–base pair

Answer: (C)

Explanation

Buffer effectiveness is greatest when the working pH is close to the pair’s pKa, because both weak-acid and conjugate-base forms are available. A pH far from the pKa produces a less balanced pair and weaker buffering.