Describe and give an everyday example of the Tyndall effect.

A solution of $6.85 \mathrm{~g}$ of a carbohydrate in $100.0 \mathrm{~g}$ of water has a density of $1.024 \mathrm{~g} / \mathrm{mL}$ and an osmotic pressure of $4.61 \mathrm{~atm}$ at $20.0^{\circ} \mathrm{C}$. Calculate the molar mass of the carbohydrate.

A quantity of $7.480 \mathrm{~g}$ of an organic compound is dissolved in water to make $300.0 \mathrm{~mL}$ of solution. The solution has an osmotic pressure of $1.43$ atm at $27^{\circ} \mathrm{C}$. The analysis of this compound shows that it contains $41.8$ percent $\mathrm{C}$, $4.7$ percent $\mathrm{H}, 37.3$ percent $\mathrm{O}$, and $16.3$ percent $\mathrm{N}$. Calculate the molecular formula of the compound.

A solution containing $0.8330 \mathrm{~g}$ of a polymer of unknown structure in $170.0 \mathrm{~mL}$ of an organic solvent was found to have an osmotic pressure of $5.20 \mathrm{mmHg}$ at $25^{\circ} \mathrm{C}$. Determine the molar mass of the polymer.

The molar mass of benzoic acid $\left(\mathrm{C}_6 \mathrm{H}_5 \mathrm{COOH}\right)$ determined by measuring the freezing-point depression in benzene is twice what we would expect for the molecular formula, $\mathrm{C}_7 \mathrm{H}_6 \mathrm{O}_2$. Explain this apparent anomaly.

A solution of $2.50 \mathrm{~g}$ of a compound having the empirical formula $\mathrm{C}_6 \mathrm{H}_5 \mathrm{P}$ in $25.0 \mathrm{~g}$ of benzene is observed to freeze at $4.3^{\circ} \mathrm{C}$. Calculate the molar mass of the solute and its molecular formula.

The elemental analysis of an organic solid extracted from gum arabic (a gummy substance used in adhesives, inks, and pharmaceuticals) showed that it contained $40.0$ percent $\mathrm{C}, 6.7$ percent $\mathrm{H}$, and $53.3$ percent $\mathrm{O}$. A solution of $0.650 \mathrm{~g}$ of the solid in $27.8 \mathrm{~g}$ of the solvent diphenyl gave a freezing-point depression of $1.56^{\circ} \mathrm{C}$. Calculate the molar mass and molecular formula of the solid. ( $K_{\mathrm{f}}$ for diphenyl is $8.00^{\circ} \mathrm{C} / \mathrm{m}$.)

Explain why we cannot use osmotic pressure to determine both molar mass and percent ionization for an unknown monoprotic acid.

Calculate the difference in osmotic pressure (in atm) at the normal body temperature between the blood plasma of a diabetic patient and that of a healthy adult. Assume that the sole difference between the two people is due to the higher glucose level in the diabetic patient. The glucose levels are $1.75$ and $0.84 \mathrm{~g} / \mathrm{L}$, respectively. Based on your result, explain why such a patient frequently feels thirsty.

Calculate the osmotic pressure of a $0.0500 \mathrm{MgSO}_4$ solution at $25^{\circ} \mathrm{C}$. (Hint: See Table 13.3.)

A $0.86$ percent by mass solution of $\mathrm{NaCl}$ is called "physiological saline" because its osmotic pressure is equal to that of the solution in blood cells. Calculate the osmotic pressure of this solution at normal body temperature $\left(37^{\circ} \mathrm{C}\right)$. Note that the density of the saline solution is $1.005 \mathrm{~g} / \mathrm{mL}$.

Both $\mathrm{NaCl}$ and $\mathrm{CaCl}_2$ are used to melt ice on roads and sidewalks in winter. What advantages do these substances have over sucrose or urea in lowering the freezing point of water?

At $25^{\circ} \mathrm{C}$, the vapor pressure of pure water is $23.76 \mathrm{mmHg}$ and that of seawater is $22.98 \mathrm{mmHg}$. Assuming that seawater contains only $\mathrm{NaCl}$, estimate its molal concentration.

What are the normal freezing points and boiling points of the following solution: $15.4 \mathrm{~g}$ of urea in $66.7 \mathrm{~mL}$ of water?