Solubility
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AP Chemistry › Solubility
A student adds magnesium sulfate ($\mathrm{MgSO_4}$), an ionic compound, to water. Water is highly polar and can form strong ion–dipole interactions. Is $\mathrm{MgSO_4}$ likely to be soluble in water, and why?
No; $\mathrm{MgSO_4}$ is insoluble because sulfates never dissolve in water.
Yes; it dissolves because the dissolving rate is fast, which means the solubility must be high.
No; $\mathrm{MgSO_4}$ is insoluble because water’s hydrogen bonding prevents ions from separating.
Yes; water’s polarity allows strong ion–dipole attractions that can overcome the ionic lattice.
No; ionic solids cannot dissolve in water because water molecules are neutral overall.
Explanation
This question tests knowledge of ionic compound solubility in polar solvents via ion-dipole interactions. Magnesium sulfate is an ionic compound, and water's high polarity enables strong ion-dipole attractions that hydrate and stabilize the Mg2+ and SO4 2- ions. These interactions are sufficient to overcome the lattice energy of MgSO4, allowing the ions to separate and dissolve. Hence, MgSO4 is soluble in water, as the polar solvent effectively surrounds and isolates the ions. Choice A is tempting but wrong, as it misstates that ionic solids cannot dissolve in neutral water molecules, overlooking ion-dipole forces that enable dissolution. For predicting ionic solubility, assess if the solvent can provide stabilizing interactions like ion-dipole to counter the lattice forces.
Carbon tetrachloride ($\mathrm{CCl_4}$) is a nonpolar molecular liquid (London dispersion forces). A student attempts to dissolve acetone ($\mathrm{(CH_3)_2CO}$), a polar molecule with a strong dipole (but no O–H bond), in $\mathrm{CCl_4}$. Is acetone likely to be highly soluble in $\mathrm{CCl_4}$, and why?
Yes; acetone will hydrogen-bond to $\mathrm{CCl_4}$ through chlorine atoms, increasing solubility.
No; acetone is polar while $\mathrm{CCl_4}$ is nonpolar, so solute–solvent attractions are relatively weak.
No; acetone is insoluble because it has a higher density than $\mathrm{CCl_4}$ and will sink.
Yes; acetone is polar, and polar solutes are most soluble in nonpolar solvents due to dipole alignment.
Yes; vigorous shaking increases solubility, so acetone will become highly soluble in $\mathrm{CCl_4}$.
Explanation
The skill tested is predicting solubility by matching intermolecular forces between polar and nonpolar substances. Acetone is a polar molecule with a significant dipole moment but no O-H for hydrogen bonding, while CCl4 is nonpolar with only London dispersion forces. The polarity mismatch leads to weak dipole-induced dipole interactions, which are not strong enough to favor mixing over the separate pure substances. Therefore, acetone is not highly soluble in CCl4, as the nonpolar solvent cannot adequately interact with the polar solute. Choice A is a tempting distractor, incorrectly stating that polar solutes are most soluble in nonpolar solvents due to dipole alignment, which confuses the 'like dissolves like' principle. A transferable strategy is to classify both solute and solvent as polar or nonpolar and expect high solubility only when they match.
Magnesium sulfate ($\mathrm{MgSO_4}$) is an ionic compound. A student places $\mathrm{MgSO_4(s)}$ into liquid diethyl ether ($\mathrm{C_4H_{10}O}$), which is only slightly polar and cannot hydrogen-bond as a donor. Is $\mathrm{MgSO_4}$ likely to be soluble in diethyl ether, and why?
No; ionic compounds dissolve only when the solvent is nonpolar so ions do not recombine.
No; ether is not polar enough to provide strong ion–dipole attractions to stabilize $\mathrm{Mg^{2+}}$ and $\mathrm{SO_4^{2-}}$.
Yes; shaking the mixture increases the dissolving rate, which makes $\mathrm{MgSO_4}$ soluble.
Yes; the oxygen in ether can hydrogen-bond strongly to ions and pull the lattice apart.
Yes; any solvent with an oxygen atom will dissolve ionic solids due to covalent bonding.
Explanation
The skill being tested is assessing the solubility of ionic compounds in weakly polar solvents lacking strong ion-solvating abilities. Magnesium sulfate (MgSO₄) has high-charge ions requiring strong solvation, but diethyl ether is only slightly polar and cannot donate hydrogen bonds. The weak dipole in ether provides insufficient ion-dipole attractions to stabilize Mg²⁺ and SO₄²⁻ compared to their lattice energy. Therefore, dissolution is unlikely due to inadequate solute-solvent interactions. A tempting distractor is choice A, which overstates ether's ability by claiming its oxygen enables strong hydrogen bonding to ions, ignoring the misconception that any oxygen-containing solvent can solvate ions effectively. Always quantify a solvent's polarity and hydrogen-bonding capacity when predicting ionic solubility to ensure accurate assessments.
A student adds hydrogen chloride gas ($\mathrm{HCl}$), a polar molecule that ionizes in water, to liquid water. Is $\mathrm{HCl}$ likely to be soluble in water, and why?
No; $\mathrm{HCl}$ is nonpolar because it is diatomic, so it cannot dissolve in polar water.
No; gases cannot dissolve in water because water molecules are held too tightly by hydrogen bonding.
Yes; $\mathrm{HCl}$ dissolves because bubbling increases the dissolving rate, which increases solubility.
Yes; $\mathrm{HCl}$ is polar and can interact strongly with water, and it forms ions that are stabilized by hydration.
No; $\mathrm{HCl}$ cannot dissolve because its molar mass is greater than water’s.
Explanation
This question tests understanding of polar gas solubility and ionization in water. Hydrogen chloride is polar and ionizes to H+ and Cl- in water, with ions stabilized by hydration and ion-dipole forces. These strong interactions enhance solubility beyond mere polarity. Thus, HCl is likely soluble in water. Choice A is tempting, incorrectly claiming gases cannot dissolve due to hydrogen bonding tightness, overlooking disruptions by compatible solutes. For gases that ionize, consider both molecular interactions and the stability of resulting ions in the solvent.
Benzoic acid ($\mathrm{C_6H_5CO_2H}$) has a polar carboxylic acid group capable of hydrogen bonding, but also a nonpolar aromatic ring. A student adds benzoic acid to water. Is benzoic acid likely to be highly soluble in water at room temperature, and why?
Yes; solubility depends mainly on mixing speed, so stirring makes it highly soluble.
No; the large nonpolar ring reduces overall polarity, so solubility in water is limited.
Yes; aromatic rings are polarizable, so they hydrogen-bond strongly to water.
Yes; any molecule with an O–H bond is highly soluble in water due to hydrogen bonding.
No; benzoic acid is insoluble because it has a higher molar mass than water.
Explanation
The skill being tested is analyzing solubility of molecules with both polar and nonpolar regions in polar solvents. Benzoic acid has a polar carboxylic acid group for hydrogen bonding but a large nonpolar aromatic ring that reduces overall polarity. The nonpolar ring leads to weaker interactions with water, limiting solubility despite the polar group. Therefore, benzoic acid is not highly soluble in water at room temperature. Choice A is tempting, overgeneralizing that any O-H bond guarantees high solubility, ignoring the impact of nonpolar portions. A transferable strategy is to consider the overall polarity of the molecule, weighing polar and nonpolar parts when predicting solubility in polar solvents.
A student tries to dissolve calcium carbonate, CaCO3(s), in water, H2O(l), at room temperature. CaCO3 is an ionic solid with relatively strong lattice energy; water is polar and can form ion–dipole interactions. Is CaCO3 likely to be very soluble in pure water, and why?
Yes; CaCO3 becomes very soluble if crushed into smaller pieces because particle size increases solubility.
No; CaCO3 is insoluble because it is denser than water, so it cannot dissolve and will always sink.
No; CaCO3 has strong ionic attractions in its lattice, so hydration by water is not sufficient to make it very soluble.
Yes; CaCO3 dissolves because carbonate can hydrogen bond strongly with water, breaking the lattice completely.
Yes; any ionic compound is highly soluble in water because ion–dipole forces always overcome lattice energy.
Explanation
This question tests understanding of ionic compound solubility and the balance between lattice energy and hydration energy. Calcium carbonate (CaCO3) has an exceptionally high lattice energy due to the +2 charge on Ca2+ and the -2 charge on CO32-, creating very strong electrostatic attractions in the solid. While water can form ion-dipole interactions with the ions, the hydration energy gained is insufficient to overcome the large lattice energy, resulting in very low solubility (Ksp ≈ 10^-9). The carbonate ion's large size and charge distribution also make it less effectively hydrated than smaller, simpler ions. Choice A incorrectly assumes all ionic compounds are highly soluble, ignoring that compounds with high lattice energies (like those with multiply charged ions) often have low solubility. The key insight is that solubility depends on the balance between lattice energy and hydration energy, not simply on whether a compound is ionic.
A student tries to dissolve carbon dioxide, CO2(g), in water, H2O(l), at room temperature. CO2 is linear and nonpolar overall (dispersion forces), while water is polar and hydrogen-bonding. Is CO2 expected to be highly soluble in water, and why?
Yes; CO2 becomes highly soluble if stirred because stirring increases the solubility of gases in liquids.
No; CO2 is nonpolar overall, so it has limited favorable interactions with water compared with water–water hydrogen bonding.
Yes; CO2 dissolves because water forms strong hydrogen bonds directly to the carbon atom in CO2.
Yes; CO2 is highly soluble because nonpolar gases always dissolve well in polar solvents.
No; CO2 cannot dissolve because gases are never soluble in liquids under any conditions.
Explanation
This question tests understanding of gas solubility in polar solvents. Carbon dioxide (CO2) is a linear molecule that, despite having polar C=O bonds, is nonpolar overall due to its symmetry where the bond dipoles cancel out. CO2 interacts primarily through weak London dispersion forces, which provide limited stabilization when dissolved in highly polar, hydrogen-bonding water. The weak CO2-water interactions cannot effectively compete with the strong water-water hydrogen bonds that must be disrupted to accommodate the gas molecules. Choice C incorrectly suggests water forms hydrogen bonds to carbon, but carbon lacks the high electronegativity and lone pairs needed for hydrogen bonding. The strategy is to consider molecular geometry when assessing polarity - linear molecules with identical terminal atoms like CO2 are nonpolar regardless of bond polarity, leading to low solubility in polar solvents.
A student adds glucose, C6H12O6(s), to water, H2O(l). Glucose has multiple –OH groups and is highly polar; water is polar and hydrogen-bonding. Is glucose likely to be soluble in water, and why?
No; glucose is a covalent compound, and covalent solutes are always insoluble in water.
No; glucose will dissolve only if stirred fast enough since stirring changes the equilibrium solubility.
Yes; glucose dissolves because its large molar mass increases its solubility in water.
No; glucose will not dissolve because it is a solid at room temperature, and solids cannot be soluble in liquids.
Yes; glucose can form many hydrogen bonds with water, leading to strong solute–solvent interactions.
Explanation
This question tests understanding of how molecular structure affects solubility in polar solvents. Glucose (C6H12O6) contains five hydroxyl (-OH) groups that can form extensive hydrogen bonds with water molecules, making it highly polar despite being a molecular compound. Each glucose molecule can form multiple hydrogen bonds as both donor and acceptor with surrounding water molecules, creating a stable hydration shell that makes dissolution energetically favorable. The numerous OH groups provide sufficient solute-solvent interactions to overcome the glucose-glucose interactions in the solid crystal. Choice B incorrectly assumes covalent compounds cannot dissolve in water, failing to recognize that polarity and hydrogen bonding capability, not bond type, determine solubility. The strategy is to count hydrogen bonding sites - molecules with multiple OH groups typically show high water solubility.
A student attempts to dissolve calcium carbonate ($\mathrm{CaCO_3}$), an ionic solid with a high lattice energy, in liquid water (polar). Is $\mathrm{CaCO_3}$ likely to be very soluble in water, and why?
Yes; any ionic compound dissolves completely in water because water is polar.
No; strong ionic attractions in the solid (high lattice energy) are not fully overcome by hydration, so solubility is low.
Yes; carbonate ions hydrogen-bond strongly with water, making $\mathrm{CaCO_3}$ highly soluble.
Yes; stirring increases the solubility of $\mathrm{CaCO_3}$ until all of it dissolves.
No; $\mathrm{CaCO_3}$ is insoluble because calcium is a metal and metals cannot dissolve in water.
Explanation
This question tests understanding of how lattice energy affects ionic solubility. Calcium carbonate has exceptionally high lattice energy due to the small, highly charged Ca²⁺ and CO₃²⁻ ions, which create very strong electrostatic attractions in the solid. While water can provide ion-dipole interactions for hydration, the hydration energy is insufficient to overcome CaCO₃'s large lattice energy, resulting in very low solubility (Ksp ≈ 10⁻⁹). This explains why limestone and chalk (forms of CaCO₃) persist in nature despite water exposure. Choice A incorrectly assumes all ionic compounds are highly soluble in water, but solubility depends on the balance between lattice energy and hydration energy—compounds with high lattice energies relative to hydration energies have low solubility. When predicting ionic solubility, consider both ion charges and sizes: small, highly charged ions typically form less soluble compounds.
A student mixes acetone, (CH3)2CO(l), with water, H2O(l). Acetone is polar due to its C=O group and can accept hydrogen bonds from water; water is polar and hydrogen-bonding. Is acetone likely to be soluble (miscible) in water, and why?
No; acetone cannot dissolve in water because it lacks an O–H bond and therefore cannot interact with water.
Yes; acetone is polar and can engage in dipole interactions and hydrogen bonding (as an acceptor) with water.
No; acetone is soluble only in nonpolar solvents because its methyl groups make it entirely nonpolar.
Yes; acetone will be miscible only if the mixture is stirred, since stirring determines miscibility.
No; acetone will not dissolve because its boiling point is lower than water’s, preventing mixing.
Explanation
This question tests understanding of polar molecular solubility and hydrogen bonding interactions. Acetone ((CH3)2CO) has a polar carbonyl group (C=O) with a significant dipole moment, making the molecule polar overall despite having two methyl groups. The oxygen atom in acetone has lone pairs that can accept hydrogen bonds from water molecules, while water can also interact with acetone through dipole-dipole forces. These favorable interactions allow acetone and water to be completely miscible, as the acetone-water interactions effectively replace water-water hydrogen bonds. Choice A incorrectly claims acetone cannot interact with water due to lacking an O-H bond, confusing hydrogen bond donation with hydrogen bond acceptance. The strategy is to identify polar functional groups and remember that molecules can participate in hydrogen bonding as acceptors even without O-H, N-H, or F-H bonds.