Chapter 9: Solutions – General Chemistry Study Notes

Chapter 9: Solutions

9.1 Solutions

Solutions are homogeneous mixtures composed of two or more substances. The major component is called the solvent, and the minor component is called the solute. Homogeneity means the mixture is uniform throughout, with no visible separation between components.

• Solvent: The substance present in the largest amount; it dissolves the solute.

• Solute: The substance present in a smaller amount; it is dissolved by the solvent.

• Example: Salt water is a solution where water is the solvent and salt is the solute.

9.2 Electrolytes and Nonelectrolytes

Electrolytes are substances that produce ions when dissolved in water, enabling the solution to conduct electricity. Nonelectrolytes do not produce ions and thus do not conduct electricity.

• Strong Electrolytes: Dissociate completely in water, producing many ions (e.g., NaCl).

• Weak Electrolytes: Dissociate only partially, producing few ions (e.g., HF).

• Nonelectrolytes: Dissolve as molecules, do not produce ions (e.g., CH3OH).

• Conductivity: Directly proportional to the number of ions present in solution.

• Example: NaCl in water conducts electricity; sugar in water does not.

9.3 Solubility

Solubility is the maximum amount of solute (in grams) that can dissolve in 100 grams of water at a given temperature. Solutions can be classified based on the amount of solute dissolved.

• Unsaturated Solution: Contains less than the maximum amount of solute; more solute can dissolve.

• Saturated Solution: Contains the maximum amount of solute; excess solute remains undissolved.

• Example: Adding sugar to tea until no more dissolves creates a saturated solution.

Soluble and Insoluble Ionic Compounds

Not all ionic compounds are soluble in water. Solubility rules help predict which compounds will dissolve.

Positive Ions	Negative Ions	Exceptions
Li+, Na+, K+, Rb+, Cs+, NH4+	NO3-, ClO3-	None
	Cl-, Br-, I-	Except with Ag+, Pb2+, Hg22+
	SO42-	Except with Ba2+, Pb2+, Ca2+, Sr2+, Hg22+

Additional info: Ionic compounds not containing at least one of these ions are usually insoluble.

9.4 Solution Concentrations

Concentration expresses the amount of solute in a given quantity of solution. Common units include mass percent, volume percent, and molarity.

• Mass Percent (w/w%): $\text{mass percent (m/m)} = \frac{\text{g of solute}}{\text{g of solute} + \text{g of solvent}} \times 100\%$

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

• Mass/Volume Percent (m/v%): $\text{mass/volume percent (m/v)} = \frac{\text{g of solute}}{\text{mL of solution}} \times 100\%$

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

• Example: Mixing 15.0 g Na2CO3 and 235 g H2O yields a 6.00% (w/w) Na2CO3 solution.

9.5 Dilution of Solutions

Dilution involves adding solvent to decrease the concentration of a solution. The amount of solute remains constant before and after dilution.

• Dilution Equation: $M_1 V_1 = M_2 V_2$

• Stock Solution: A concentrated solution used to prepare more dilute solutions.

• Example: Diluting 0.50 L of 6.0 M HCl to 1.0 L results in a final concentration of 3.0 M.

Water as a Polar Solvent

Water is a polar molecule due to its bent geometry and polar O—H bonds. This allows water molecules to form hydrogen bonds, making water an excellent solvent for ionic and polar substances.

• Polarity: Uneven distribution of electron density creates partial charges.

• Hydrogen Bonding: Attraction between the partial positive charge on hydrogen and partial negative charge on oxygen in adjacent molecules.

• Example: Water dissolves salt (NaCl) due to its polarity.

Why Some Solutions Conduct Electricity

Electrical conductivity in solutions depends on the presence of mobile ions. The more ions present, the greater the conductivity.

• Mobile Electrolytes: Ions that move freely in solution.

• Direct Proportionality: Conductivity increases with ion concentration.

• Example: NaCl solution conducts electricity; pure water does not.

Solution, Colloid & Suspension

Mixtures can be classified based on particle size and separation methods.

Mixture	Particle Size	Settling	Separation
Solution	Small, < 1 nm	No	Cannot be separated by filtration
Colloid	Large, 1 nm–1 μm	No	Cannot be separated by common filtration
Suspension	Large, > 1 μm	Yes	Can be separated by filtration or centrifugation

Osmosis

Osmosis is the movement of solvent (usually water) through a semipermeable membrane from a region of lower solute concentration to higher solute concentration.

• Semipermeable Membrane: Allows passage of solvent but not solute.

• Result: The solution with higher solute concentration increases in volume.

• Example: Water entering a plant cell by osmosis.

Osmotic Pressure

Osmotic pressure is the pressure required to prevent the flow of solvent into a more concentrated solution through a semipermeable membrane.

• Importance: Maintains cell integrity and fluid balance in biological systems.

Reverse Osmosis

Reverse osmosis occurs when pressure greater than the osmotic pressure is applied, forcing solvent to move from higher to lower solute concentration. This process is used in desalination to obtain pure water from seawater.

• Application: Water purification in desalination plants.

Isotonic, Hypotonic, and Hypertonic Solutions

These terms describe the relative solute concentrations of solutions compared to body fluids, affecting cell volume and function.

• Isotonic Solution: Same osmotic pressure as body fluids; cells retain normal shape (e.g., 0.90% NaCl).

• Hypotonic Solution: Lower solute concentration; water enters cells, causing swelling and possible bursting (hemolysis).

• Hypertonic Solution: Higher solute concentration; water leaves cells, causing shrinkage (crenation).

Dialysis

Dialysis is the process where solvent and small solute particles pass through an artificial membrane, while larger particles are retained. It is used medically to remove waste from blood (hemodialysis).

• Artificial Membrane: Selectively permeable to small molecules.

• Application: Removal of urea and other wastes in kidney failure patients.

Calculating Solution Quantities and pH

Quantitative calculations are essential for preparing solutions and understanding their properties.

• Moles and Mass: Use molarity and volume to calculate moles and mass of solute.

• pH Calculation: pH is a measure of acidity, defined as:

$\text{pH} = -\log[\text{H}_3\text{O}^+]$

• Example: If $[\text{H}^+] = 8.5 \times 10^{-4} \text{ M}$, then $\text{pH} = -\log(8.5 \times 10^{-4})$

Additional info: These calculations are fundamental for laboratory work and understanding chemical equilibria in solutions.