Physical Properties of Solutions – Lecture Review

24 question-and-answer flashcards covering the key concepts from the lecture on solution formation, intermolecular forces, energetics, ideal vs non-ideal behavior, and factors affecting solubility.

What is the basic definition of a solution?

A homogeneous mixture formed when one substance (solute) disperses uniformly throughout another (solvent).

What two general factors determine whether substances will form a solution?

(1) The natural tendency of substances to mix and spread into larger volumes (entropy). (2) The types and strengths of intermolecular interactions involved.

Why do gases mix spontaneously to form solutions?

Gas molecules have negligible intermolecular forces, so their random motion leads to spontaneous mixing and increased entropy.

How does entropy influence the formation of solutions?

Mixing increases the randomness (entropy) of molecules, which thermodynamically favors spontaneous solution formation.

Under what condition will liquids or solids not mix spontaneously?

When strong intermolecular forces or physical barriers restrain the molecules.

What must happen to solute–solute interactions for a solution to form?

They must be overcome (separated) so solute particles can disperse in the solvent.

Name the three enthalpy components of the solution process.

ΔHsolute (separation of solute), ΔHsolvent (separation of solvent), and ΔHmix (mixing of particles).

Are ΔHsolute and ΔHsolvent endothermic or exothermic?

Endothermic (ΔH > 0) because energy is required to separate particles.

Is ΔHmix endothermic or exothermic, and why?

Exothermic (ΔH < 0) because energy is released when solute and solvent particles attract each other during mixing.

How is ΔHsoln (overall enthalpy of solution) calculated conceptually?

ΔHsoln = ΔHsolute + ΔHsolvent + ΔHmix.

What defines an ideal solution?

All molecular interactions (solute–solute, solvent–solvent, solute–solvent) are equal, resulting in ΔHsoln ≈ 0 and predictable properties.

Give an example of an ideal solution mentioned in the lecture.

Toluene dissolved in benzene.

What characterizes a non-ideal exothermic solution (second scenario)?

Un-like intermolecular forces are stronger than like forces; ΔHmix releases more heat than absorbed, so ΔHsoln < 0. Example: acetone and chloroform.

What characterizes a non-ideal endothermic solution (third scenario)?

Un-like intermolecular forces are weaker than like forces but still allow mixing; ΔHsoln > 0. Example: carbon disulfide and acetone.

When will a heterogeneous mixture form instead of a solution (fourth scenario)?

When intermolecular forces between unlike molecules are much weaker than those between like molecules, preventing significant dissolution. Example: water and octane.

What is hydration in the context of ionic solutions?

The process where an ion is surrounded and stabilized by a cluster of water molecules, releasing hydration energy.

How does hydration energy influence the solubility of ionic solids?

If hydration energy exceeds the lattice energy required to separate ions, the ionic solid is likely to dissolve in water.

Describe the stages leading to a saturated solution.

(a) Only dissolution occurs initially; (b) crystallization rate increases over time; (c) saturation is reached when dissolution and crystallization rates are equal.

What is meant by ‘colligative properties’ as listed in the topic overview?

Properties of solutions that depend only on the number of solute particles, not their identity (e.g., boiling-point elevation, freezing-point depression).

Name two concentration units commonly used to describe solutions (from overall course context).

Molarity (M) and molality (m).

Why are ideal and non-ideal solutions treated differently in calculations?

Ideal solutions obey simple laws (like Raoult’s law); non-ideal solutions deviate due to unequal intermolecular forces, requiring correction factors or activity coefficients.

What role do phase diagrams play in understanding solutions?

They illustrate the equilibrium phases (solid, liquid, gas) of a system at various temperatures and pressures, aiding in predicting solubility and phase transitions.

List two factors (besides IMF) that affect solubility.

Temperature and pressure (especially for gases).

Why can gases have their solubility strongly influenced by pressure?

According to Henry’s law, the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the solution.