Chapter 11 Notes

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Solutions:

Solutions are homogeneous mixtures of two or more pure substances

Two substances that don't mix are a heterogeneous mixture

In a solution, the solute is dispersed uniformly throughout the solvent

How Does a Solution Form?

As a solution, the solvent pulls the solute particles apart and surrounds, or solvates, them

If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal

The ions are solvated (surrounded by the solvent)

If the solvent is water, the ions are hydrated

The IMF here is ion-dipole

Lattice Energy:

Lattice energy (U): the energy released when one mole of an ionic compound forms from its free ions in the gas phase

U = K(Q1-Q2)/d

K is a proportionality constant depending on the lattice structure

Lattice energy depends on:

Ionic charge

Ionic radius

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Energy changes in a Solution:

The enthalpy change of the overall process depends on ∆H for each of these steps

The Born-Haber cycle can be used to determine the enthalpy of hydration of an ionic compound

Solubility:

When solutions are formed, intermolecular forces are key

Polar substances dissolve in a polar solvent, while nonpolar substances dissolve in a nonpolar substance

Like dissolves like

Nonpolar substances won't dissolve in polar substances, and vice versa

Factors That Affect Solubility:

Structure effects

Pressure effects

Temperature effects (aqueous solutions only)

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The Three Factors That Affect Solubility:

Structure effects: solubility is favored when the solute and the solvent have similar polarities; "like dissolves like"

Polar solutes dissolve in polar solvents; nonpolar in nonpolar

Example: Vitamins A, D, E, and K are fat-soluble while B and C are water-soluble

The stronger the IMF between solute and solvent, the more likely the solute will dissolve

Ions in water also have ion-dipole forces

Example: ethanol in water

Amphiphilic: a molecule that has both a polar and nonpolar region, allowing it to dissolve nonpolar molecules in polar solvents

Emulsions: a type of colloid that contains tiny liquid droplets randomly dispersed in a liquid solvent

Glucose (which has hydrogen bonding) is very soluble in water

Cyclohexane (which has only dispersion forces) is not water-soluble

Pressure effects: this only affects gas solubility in a solvent

Example: carbonation in a bottle of soda

To ensure the large amount of CO2 is dissolved in water, soda is bottled at high temperatures

"Fizzing" occurs as CO2 (g) molecules escape from the liquid solvent because they are no longer soluble at atmospheric pressure (1 atm)

Larger molecules have stronger dispersion forces

Increasing pressure above solution forces more gas to dissolve

Temperature effects: consider solubilities of solid, liquid, and gas solutes separately

Solid solute: most solids become more soluble as temperature increases

Liquid solute: mos liquids become more soluble as temperature increases

Gas solute: most gases become less soluble as temperature increases

Why? It's because as the temperature increases, dissolved gas particles have a greater tendency to escape the liquid, thus decreasing their solubility in the solvent

Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature