Ch 9 - Solutions and Their Properties: GOB Chemistry Study Notes

Solutions and Their Properties

Introduction to Solutions

Solutions are homogeneous mixtures composed of two or more substances. In chemistry, understanding solutions is essential for studying reactions, biological systems, and industrial processes.

• Solute: The substance that is dissolved in a solution (present in lesser amount).

• Solvent: The substance that does the dissolving (present in greater amount). In aqueous solutions, water is the solvent.

• Examples: Salt water (NaCl in H2O), sugar in tea, air (a solution of gases).

Types of Solutions

Solutions can exist in all three physical states: solid, liquid, and gas.

Solute State	Solvent State	Example
Gas	Gas	Air (O2 in N2)
Gas	Liquid	Carbonated water (CO2 in H2O)
Liquid	Liquid	Alcohol in water
Solid	Liquid	Salt in water
Solid	Solid	Alloys (brass: Zn in Cu)

Solubility and Factors Affecting Solubility

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.

• "Like dissolves like": Polar solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents.

• Temperature: Solubility of solids in liquids generally increases with temperature; for gases, solubility decreases as temperature increases.

• Pressure: Affects the solubility of gases in liquids (Henry's Law).

Example: Sugar dissolves better in hot water than in cold water.

Electrolytes and Nonelectrolytes

When ionic compounds dissolve in water, they may dissociate into ions, making the solution conductive.

• Electrolyte: A substance that produces ions in solution and conducts electricity (e.g., NaCl, KBr).

• Nonelectrolyte: A substance that does not produce ions in solution (e.g., sugar, ethanol).

• Strong Electrolyte: Completely dissociates into ions (e.g., NaCl).

• Weak Electrolyte: Partially dissociates (e.g., acetic acid).

Solubility Rules for Ionic Compounds in Water

Solubility rules help predict whether an ionic compound will dissolve in water.

Compound Type	Solubility
All nitrates (NO3-), acetates (C2H3O2-), and most alkali metal salts	Soluble
Most chlorides, bromides, iodides	Soluble (except with Ag+, Pb2+, Hg22+)
Most sulfates (SO42-)	Soluble (except with Ba2+, Pb2+, Ca2+, Sr2+)
Most carbonates, phosphates, sulfides, oxides, hydroxides	Insoluble (except with alkali metals and NH4+)

Concentration Units

Concentration expresses the amount of solute in a given amount of solution or solvent.

• Mass/Volume Percent (% m/v): $\frac{\text{grams of solute}}{\text{mL of solution}} \times 100$

• Mass Percent (% m/m): $\frac{\text{grams of solute}}{\text{grams of solution}} \times 100$

• Volume Percent (% v/v): $\frac{\text{mL of solute}}{\text{mL of solution}} \times 100$

• Molarity (M): $\frac{\text{moles of solute}}{\text{liters of solution}}$

Example: A 0.9% NaCl solution contains 0.9 g NaCl per 100 mL solution.

Preparation of Solutions

Solutions of known concentration can be prepared by dissolving a measured amount of solute in solvent and diluting to a specific volume.

• Dilution Equation: $M_1V_1 = M_2V_2$

• Where $M_1$ and $V_1$ are the initial molarity and volume, $M_2$ and $V_2$ are the final molarity and volume.

Example: To prepare 250 mL of 0.5 M NaCl from 1.0 M NaCl, use $V_1 = \frac{M_2V_2}{M_1} = \frac{0.5 \times 250}{1.0} = 125$ mL.

Colligative Properties

Colligative properties depend on the number of solute particles in solution, not their identity.

• Boiling Point Elevation: Adding solute increases the boiling point of the solvent.

• Freezing Point Depression: Adding solute lowers the freezing point of the solvent.

• Osmosis: The movement of solvent through a semipermeable membrane from low to high solute concentration.

Equations:

• Boiling Point Elevation: $\Delta T_b = K_b m$

• Freezing Point Depression: $\Delta T_f = K_f m$

• Where $m$ is the molality of the solution.

Osmosis and Osmotic Pressure

Osmosis is crucial in biological systems, affecting cell volume and function.

• Osmotic Pressure ($\pi$): The pressure required to stop osmosis.

• Equation: $\pi = MRT$

• Where $M$ is molarity, $R$ is the gas constant, and $T$ is temperature in Kelvin.

• Isotonic Solution: Same osmotic pressure as cells; no net water movement.

• Hypotonic Solution: Lower osmotic pressure than cells; water enters cells (may cause hemolysis).

• Hypertonic Solution: Higher osmotic pressure than cells; water leaves cells (may cause crenation).

Suspensions, Colloids, and Solutions

Mixtures can be classified based on particle size and behavior.

Type	Particle Size	Appearance	Separation
Solution	< 1 nm	Clear	Does not separate
Colloid	1-1000 nm	Cloudy	Can be separated by semipermeable membrane
Suspension	> 1000 nm	Cloudy, settles	Can be separated by filtration

Example: Milk is a colloid; muddy water is a suspension; salt water is a solution.

Summary Table: Properties of Solutions, Colloids, and Suspensions

Property	Solution	Colloid	Suspension
Particle Size	< 1 nm	1-1000 nm	> 1000 nm
Settling	No	No	Yes
Filtration	No	No	Yes
Appearance	Clear	Cloudy	Cloudy

Key Equations and Concepts

• Concentration (molarity): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution: $M_1V_1 = M_2V_2$

• Percent by mass: $\%\,m/m = \frac{\text{g solute}}{\text{g solution}} \times 100$

• Percent by volume: $\%\,v/v = \frac{\text{mL solute}}{\text{mL solution}} \times 100$

• Osmotic Pressure: $\pi = MRT$

Applications in Health and Industry

• IV Solutions: Must be isotonic to prevent cell damage.

• Water Purification: Relies on filtration and osmosis principles.

• Food Science: Solutions, colloids, and suspensions are common in beverages, dairy, and processed foods.

Summary

Understanding solutions, their properties, and how to calculate concentrations is fundamental in chemistry and essential for applications in biology, medicine, and industry. Mastery of these concepts enables prediction of solubility, preparation of solutions, and comprehension of physiological processes such as osmosis.