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Abnormal Molar Masses

Use the van't Hoff factor when a solute splits apart or joins together in solution.

6 micro-lessons

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Dissociation and Association

Why some molar masses come out wrong: when solutes split apart or join together.

Abnormal Molar Masses

When we dissolve a solute in a solvent, its molecules don't always stay intact. Depending on the chemical interactions, they might dissociate (break apart) or associate (clump together).

Dissociation: Ionic compounds like $KCl$ split into multiple ions ( $K^+$ and $Cl^-$ ). This increases the total number of particles in solution, amplifying colligative properties and resulting in an experimentally calculated molar mass that is lower than the true value.
Association: In solvents with low dielectric constants (like benzene), molecules like acetic acid can form dimers via hydrogen bonding. This decreases the total number of particles, reducing colligative properties and resulting in an experimentally calculated molar mass that is higher than the true value.

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Example 1.12: Association (Dimerization)

Calculates the percentage association of benzoic acid in benzene.

Problem

2 g of benzoic acid ( $C_6H_5COOH$ ) dissolved in 25 g of benzene shows a depression in freezing point equal to 1.62 K. Molal depression constant for benzene is $4.9 \text{ K kg mol}^{-1}$ . What is the percentage association of acid if it forms dimer in solution?

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Example 1.13: Dissociation of a Weak Acid

Calculates the van't Hoff factor and dissociation constant for acetic acid.

Problem

0.6 mL of acetic acid ( $CH_3COOH$ ), having density $1.06 \text{ g mL}^{-1}$ , is dissolved in 1 litre of water. The depression in freezing point observed for this strength of acid was $0.0205^\circ\text{C}$ . Calculate the van't Hoff factor and the dissociation constant of acid. (Assume $K_f = 1.86 \text{ K kg mol}^{-1}$ )

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Identifying the van't Hoff Factor

Find the van't Hoff factor for salts that fully dissociate.

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Assuming complete dissociation, what is the expected van't Hoff factor ( $i$ ) for potassium sulfate ( $K_2SO_4$ )?

Practice: Freezing Point with i

Multi-step problem using the van't Hoff factor.

Solve the multi-step problem by breaking down the calculation.

Exercise 1.32 Adaptation Calculate the depression in the freezing point of water when $10 \text{ g}$ of $CH_3CH_2CHClCOOH$ ( $M = 122.5 \text{ g mol}^{-1}$ ) is added to $250 \text{ g}$ of water.

Given: $K_a = 1.4 \times 10^{-3}$ , $K_f = 1.86 \text{ K kg mol}^{-1}$ .

Calculate Molality

Find the initial molality of the acid before considering dissociation.

State Equilibrium & Ka formula

What formula connects Ka to the degree of dissociation (alpha)?

Find Degree of Dissociation (alpha)

Substitute the values to solve for alpha (you can assume 1-alpha ≈ 1 for simplicity).

Calculate the van't Hoff Factor (i)

Convert your degree of dissociation into the van't Hoff factor.

Final Freezing Point Depression

Use ΔTf = i * Kf * m to find the final answer.

← previousColligative Properties: Freezing Point A...next →Chapter Review: Solutions

Dissociation and Association

Why some molar masses come out wrong: when solutes split apart or join together.

Abnormal Molar Masses

When we dissolve a solute in a solvent, its molecules don't always stay intact. Depending on the chemical interactions, they might dissociate (break apart) or associate (clump together).

Dissociation: Ionic compounds like $KCl$ split into multiple ions ( $K^+$ and $Cl^-$ ). This increases the total number of particles in solution, amplifying colligative properties and resulting in an experimentally calculated molar mass that is lower than the true value.
Association: In solvents with low dielectric constants (like benzene), molecules like acetic acid can form dimers via hydrogen bonding. This decreases the total number of particles, reducing colligative properties and resulting in an experimentally calculated molar mass that is higher than the true value.

1 / 3

Example 1.12: Association (Dimerization)

Calculates the percentage association of benzoic acid in benzene.

Problem

1 / 5

Example 1.13: Dissociation of a Weak Acid

Calculates the van't Hoff factor and dissociation constant for acetic acid.

Problem

1 / 5

Identifying the van't Hoff Factor

Find the van't Hoff factor for salts that fully dissociate.

1 / 5

Assuming complete dissociation, what is the expected van't Hoff factor ( $i$ ) for potassium sulfate ( $K_2SO_4$ )?

Practice: Freezing Point with i

Multi-step problem using the van't Hoff factor.

Solve the multi-step problem by breaking down the calculation.

Exercise 1.32 Adaptation Calculate the depression in the freezing point of water when $10 \text{ g}$ of $CH_3CH_2CHClCOOH$ ( $M = 122.5 \text{ g mol}^{-1}$ ) is added to $250 \text{ g}$ of water.

Given: $K_a = 1.4 \times 10^{-3}$ , $K_f = 1.86 \text{ K kg mol}^{-1}$ .

Calculate Molality

Find the initial molality of the acid before considering dissociation.

State Equilibrium & Ka formula

What formula connects Ka to the degree of dissociation (alpha)?

Find Degree of Dissociation (alpha)

Substitute the values to solve for alpha (you can assume 1-alpha ≈ 1 for simplicity).

Calculate the van't Hoff Factor (i)

Convert your degree of dissociation into the van't Hoff factor.

Final Freezing Point Depression

Use ΔTf = i * Kf * m to find the final answer.