In-Depth Notes on Solutions and Concentration

Dissolving and Recrystallization: At saturation point, a solute dissolves and some solid precipitates, establishing dynamic equilibrium.
Unsaturated Solutions: Have room for more solute to dissolve; adding more solute will continue to dissolve.
Supersaturated Solutions: Contains more solute than typically soluble; for example, 6 grams of solute in 100 grams of water.

Temperature Dependence:
- Solubility of solids generally increases with temperature (example: sugar dissolving better when water is heated).
- Heating increases solubility, allowing for supersaturation when cooled slowly, trapping solute between solvent molecules.
- Most ionic compounds show increased solubility with temperature rise (e.g., 200 grams of solute in 100 grams of water at higher temperatures).
Gases in Solutions:
- Solubility of gases decreases with increased temperature (e.g., soda goes flat quicker at higher temperatures due to gas escaping).

Henry's Law: Solubility of a gas in liquid is directly related to pressure: higher pressure equals higher gas solubility.
- Example: Pressurized vs. unpressurized soda can — pressurized has more dissolved CO2 due to the pressure.

General Rules:
- Ions that make compounds soluble include Group 1 elements (Li extsuperscript{+}, Na extsuperscript{+}, K extsuperscript{+}, etc.) and ammonium ion (NH extsubscript{4} extsuperscript{+}).
- Nitrates (NO extsubscript{3} extsuperscript{-}) and acetates (C extsubscript{2}H extsubscript{3}O extsubscript{2} extsuperscript{-}) are always soluble regardless of the cation.
Exceptions:
- Halides (Cl extsuperscript{-}, Br extsuperscript{-}, I extsuperscript{-}) are generally soluble except with Ag extsuperscript{+}, Pb extsuperscript{2+}, and Hg extsubscript{2} extsuperscript{2+}.
- Sulfates (SO extsubscript{4} extsuperscript{2-}) are generally soluble except with Ba extsuperscript{2+}, Pb extsuperscript{2+}, Ca extsuperscript{2+}, Sr extsuperscript{2+}, and Hg extsubscript{2} extsuperscript{2+}.

Mixing two soluble ionic compounds can sometimes produce an insoluble product (precipitate).
Net Ionic Equations: Focus on the ions that participate in the reaction. Spectator ions do not participate.

Concentration is the measure of solute in a solution.
Expressed as solute per solution in various forms:
- Mass/volume (g/mL), mass/mass (%), molarity (M = moles of solute/L of solution).
Dilution: Process used to lower a concentration by adding solvent; uses the formula M extsubscript{1}V extsubscript{1} = M extsubscript{2}V extsubscript{2}.

Solutions alter physical properties (e.g., boiling and freezing points) due to added solute; these changes are determined by the number of solute particles:
- Freezing Point Depression: For every mole of solute, freezing point decreases by 1.86 °C.
- Boiling Point Elevation: For every mole of solute, boiling point increases by 0.51 °C.

Solutions: Homogeneous mixtures where the solute is indistinguishable.
Colloids: Medium-sized particles, partially visible but cannot be filtered out (e.g., milk).
Suspensions: Large particles that settle out if left unstirred (e.g., muddy water, milkshake).

Vapor Pressure: The pressure of vapor in equilibrium with its liquid; decreases when solute is added.
Osmotic Pressure: The pressure required to stop osmosis; affects biological systems significantly.
Isotonic Solutions: Same concentration as body fluids; preventing cell swelling (hemolysis) or shrinking (crenation) when different concentrations are introduced.

IV Solutions: Normal saline (0.9% NaCl) and glucose solution mimic body fluid concentrations to prevent osmotic imbalance.
Hemolysis and Crenation:
- Hemolysis: Burst from excess water influx due to hypotonic solutions.
- Crenation: Shriveling from excessive water loss due to hypertonic solutions.
Dialysis: Mimics kidney function, using osmotic principles to remove waste from blood through a semi-permeable membrane.