Solubility and Solution Chemistry Notes

Solubility

• Solubility basics
  • Solubility = the maximum amount of solute that can be dissolved in a given amount of solvent.
  • When one substance (solute) dissolves in another (solvent) it is soluble; example: table salt in water.
  • When a substance does not dissolve, it is insoluble; example: oil in water.
  • The tendency for substances to mix depends on nature and intermolecular forces between solute and solvent.
  • Entropy: a measure of energy dispersal (degree of randomness). Mixing two substances increases entropy.
  • Gases are always soluble in each other (conceptual, under appropriate conditions).
• Key concepts
  • Like dissolves like: a chemical will dissolve in a solvent with a similar structure.
  • Polar molecules and ionic compounds tend to be soluble in polar solvents; nonpolar molecules tend to be soluble in nonpolar solvents.
  • Solubility varies with temperature and pressure.
  • Miscible vs. immiscible: mutually soluble liquids are miscible (e.g., ethanol and water); oil and water are immiscible.
• Practical implications
  • Solubility depends on intermolecular attractive forces and the structure of solute/solvent.
  • Solubility curves and temperature/pressure effects help predict dissolution behavior.

Heat of Solution

• Heat effects during dissolution
  • For some solutes (e.g., NaOH), dissolution is exothermic: the container gets hot.
  • For others (e.g., NaNO3), dissolution is endothermic: the container gets cold.
• Three contributing processes (energetics)
  • ΔHsolve = ΔHsolute + ΔHsolvent + ΔHmix
  • IMF between solute molecules must be overcome: ΔH_solute > 0
  • IMF between solvent molecules must be overcome: ΔH_solvent > 0
  • IMF between solute and solvent must form new interactions: ΔH_mix < 0
  • The overall sign of ΔHsolve determines exothermic (ΔHsolve < 0) or endothermic (ΔH_solve > 0).
• Relative energy balance examples
  • If the energy released from solute–solvent interactions (ΔHmix) plus hydration overcomes the energy to break solute/solvent interactions, dissolution is exothermic (AHhydration < AHsolute → AHsolution < 0).
  • If breaking solute/solvent interactions costs more energy than is released by hydration, dissolution is endothermic (AHhydration > AHsolute → AH_solution > 0).
• Heat of hydration (ionic compounds)
  • Dissolution of ionic compounds involves overcoming lattice (crystal) energy (lattice energy) and forming ion–dipole interactions with solvent (e.g., water).
  • Hydration energy is the energy released when ions in gas form become solvated in water; this is typically negative (exothermic).
  • For ions in water, hydration energy is large and negative, contributing to dissolution when sufficient to overcome lattice energy.
• Key relationships for ionic dissolution
  • Ion–dipole interactions drive hydration and solvation;
  • If hydration energy is large in magnitude relative to lattice energy, dissolution is favored.

Solution Equilibrium

• Dissolution as an equilibrium process
  • Initially, ions are pulled out of the crystal lattice.
  • Over time, many ions are in the solvent, and some ions recombine to form solid.
  • At equilibrium, rate of dissolution equals rate of deposition; the solution is saturated.
• Saturation concepts
  • Saturated: no more solute dissolves at the given temperature; adding solute does not dissolve.
  • Unsaturated: more solute can dissolve.
  • Supersaturated: contains more dissolved solute than typically solvable at that temperature; can be formed by heating and then slowly cooling a saturated solution.
  • Supersaturated solutions can crystallize upon seeding with a small crystal.
• Temperature dependence of solubility
  • For most solids, solubility in a solid solvent increases as temperature increases.
  • Solubility curves (solubility vs. temperature) help classify a solution as saturated, unsaturated, or supersaturated:
  • On the curve: saturated
  • Below the curve: unsaturated
  • Above the curve: supersaturated
• Applications and concepts
  • Temperature-based purification: dissolve a solid in a hot solvent and crystallize upon cooling to remove impurities.
  • Gas solubility in liquids behaves differently with temperature (see Gas solubility section).
• Solubility curves (solids)
  • Examples show various salts (e.g., NaNO3, KNO3, etc.) with solubility plotted against temperature.

Solubility of Gases and Henry's Law

• Temperature effect for gases
  • Gas solubility in liquids generally decreases as temperature increases.
  • Practical example: aquatic life in warm ponds: higher water temperature reduces dissolved O2, impacting life.
• Pressure effect for gases
  • For solids: solubility is largely pressure-independent.
  • For gases: solubility increases with increased pressure above the liquid, due to more gas molecules being forced into solution until equilibrium with gas phase is reestablished.
• Henry's Law
  • Henry's law expresses the direct proportionality between the solubility of a gas and the partial pressure of the gas above the liquid:
  • $S = k<em>H \, P</em>{gas}$
  • Here, S is the gas solubility, Pgas is the partial pressure, and kH is the Henry's law constant for the specific gas–solvent system.
• Conceptual interpretation
  • The dynamic equilibrium exists between gas molecules trapped in the liquid by intermolecular forces and molecules that escape into the gas phase.
  • Increasing the gas pressure increases the amount dissolved until equilibrium is reached.

Concentration and Units

• Concentration: definition
  • Concentration represents the amount of solute dissolved in a given amount of solvent or solution.
  • Terms like diluted and concentrated describe relative amounts of solute.
• Common concentration units
  • Molarity: $M = \frac{n<em>{solute}}{V</em>{solution}}$ (moles per liter of solution)
  • Molality: $m = \frac{n<em>{solute}}{n</em>{solvent, ext{kg}}}$ (moles per kilogram of solvent)
  • Mole fraction: xi = rac{ni}{\sumj nj}
  • Mass percent (weight/weight): $w\% = \frac{m<em>{solute}}{m</em>{solution}} \times 100\%$
  • Parts per million (ppm): $ext{ppm} = \frac{m<em>{solute}}{m</em>{solution}} \times 10^6$
  • Parts per billion (ppb): $ext{ppb} = \frac{m<em>{solute}}{m</em>{solution}} \times 10^9$
• Dilution and conversion between units
  • Conversion between concentration units typically involves rewriting the given concentration as a ratio, isolating the solute and solvent parts, converting to the desired units, and applying definitions to compute the new concentration.

Making Solutions (Practical Lab Procedure)

• Example: Makes 100 mL of 1 M NaCl solution
  • Calculation: 1 M in 100 mL requires 0.1 mol of NaCl.
  • Weigh out 5.84 g NaCl (since 0.1 mol × 58.44 g/mol = 5.844 g ≈ 5.84 g).
  • Add the salt to a 100 mL volumetric flask.
  • Add water to dissolve the solid.
  • Add more water until the 100 mL mark is reached and mix until dissolved.
• Alternative description (for 1 L of 1 M):
  • Weigh 58.44 g NaCl, dissolve in water, and dilute to 1 L to obtain 1 M.

Converting Between Concentration Units (Procedure)

• Steps to convert units
  • 1) Write the given concentration as a ratio (solute : solvent or solution).
  • 2) Separate the numerator (solute amount) and denominator (volume or mass, depending on units).
  • 3) Convert the solute amount into the required unit (e.g., moles to grams or grams to moles).
  • 4) Convert the solution part into the required unit (e.g., liters to milliliters).
  • 5) Use the definitions to calculate the new concentration units.

Compare Solutions by Concentration (Osmolar Concepts)

• Osmotic relationships
  • Hyperosmotic: one solution has a higher total solute concentration than another.
  • Isoosmotic (isotonic): same total solute concentration in two solutions.
  • Hypoosmotic: one solution has a lower total solute concentration than another.
• Practical importance
  • IV fluids should be isosmotic with cells to avoid cell shrinkage or swelling.
  • Hyperosmotic environments cause water to move from cells to the surrounding medium, shrinking cells.
  • Hypoosmotic environments cause water to move into cells, potentially causing swelling or bursting.
• Real-world example: cucumber pickling
  • Cucumbers placed in concentrated salt and vinegar solutions lose water to the surrounding medium, increasing intracellular solute concentration and preserving the cucumber.

Miscellaneous Notes and Connections

• Temperature and purification strategies
  • For solids, dissolve in hot solvent and crystallize as the solution cools to remove impurities.
• Summary of key terms
  • Solute, solvent, solubility, miscible, immiscible, entropy, lattice energy, hydration energy, ion-dipole interactions, Henry's law, saturation, supersaturation, isotonic/isotonicity, osmolarity.
• Formulas to remember
  • Enthalpy of solution: $\Delta H<em>{solution} = \Delta H</em>{solute} + \Delta H<em>{solvent} + \Delta H</em>{mix}$
  • Exothermic if $\Delta H<em>{solution} < 0$; Endothermic if \Delta H{solution} > 0
  • Solubility dependence on gas pressure (Henry): $S = k<em>H P</em>{gas}$
  • Molarity: $M = \frac{n<em>{solute}}{V</em>{solution}}$
  • Molality: $m = \frac{n<em>{solute}}{m</em>{solvent, \text{kg}}}$
  • Mole fraction: $x<em>i = \frac{n</em>i}{\sum<em>j n</em>j}$
  • Mass percent: $w\% = \frac{m<em>{solute}}{m</em>{solution}} \times 100\%$
  • ppm: $ext{ppm} = \frac{m<em>{solute}}{m</em>{solution}} \times 10^6$
  • ppb: $ext{ppb} = \frac{m<em>{solute}}{m</em>{solution}} \times 10^9$