Chapter 13: Solutions – General Chemistry Study Notes

Chapter 13: Solutions

Topics Covered

• Review of solutions and dissolving

• Factors that influence solubility

• Expressing solution concentration

• Colligative properties of solutions

Review of Solutions and Dissolving

Mixtures and Entropy

Most substances in our environment are mixtures, and the process of mixing increases entropy (disorder). When a solute dissolves in a solvent, the system moves toward a state of uniform concentration, maximizing entropy.

• Entropy: A measure of disorder or randomness in a system. Mixing increases entropy.

• Example: Dissolving NaCl in water results in a uniform distribution of Na+ and Cl- ions among H2O molecules.

Examples of Solutions

State of Solute	State of Solvent	Example
Gas	Gas	Air (O2 in N2)
Gas	Liquid	Club Soda (CO2 in H2O)
Gas	Solid	Catalytic Converter (CO into Pt)
Liquid	Gas	Water Vapor in Air
Liquid	Liquid	Vodka (ethanol in water)
Liquid	Solid	Dental Amalgams (Hg in Ag and other metals)
Solid	Gas	Naphthalene in Air
Solid	Liquid	Seawater
Solid	Solid	Brass or other alloys (Zn in Cu)

Solubility and Its Dependence on Temperature and Pressure

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature and pressure. It is commonly expressed in grams of solute per 100 grams of water.

Substance Dissolved	283 K	293 K	313 K	333 K	353 K
Potassium nitrate	21	32	62	106	167
Sodium chloride	36	36	36	37	37
Potassium chloride	35	40	54	45	66
Ammonium chloride	24	37	41	55	66

Types of Solutions

• Saturated: Contains the maximum amount of solute that can dissolve at a given temperature.

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

• Supersaturated: Contains more solute than is stable at a given temperature; formed under special conditions.

How Does a Solution Form?

• For ionic compounds (e.g., NaCl): $\text{NaCl}_{(s)} \rightarrow \text{Na}^+_{(aq)} + \text{Cl}^-_{(aq)}$

• For molecular compounds (e.g., glucose): $\text{C}_6\text{H}_{12}\text{O}_6_{(s)} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6_{(aq)}$

• Solvation: The process by which solvent molecules surround and interact with solute ions or molecules.

Factors Affecting Solubility

Nature's Tendency Toward Mixing and Intermolecular Forces (IMFs)

• Solubility depends on the natural tendency toward mixing (increasing entropy) and the types of intermolecular forces (IMFs) present.

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

Types of Intermolecular Attractions

• Dispersion forces: Present in all molecules, especially nonpolar ones (e.g., hexane).

• Dipole-dipole interactions: Occur between polar molecules (e.g., acetone, chloroform).

• Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F (e.g., ethanol, water).

• Ion-dipole forces: Occur between ions and polar molecules (e.g., Na+ in water).

Predicting Solubility Patterns

Solubility in water increases with the presence of polar groups (e.g., -OH) and decreases with increasing nonpolar hydrocarbon chains.

Effect of Temperature

• Ionic Compounds: Solubility generally increases with temperature (e.g., KNO3 in water).

• Gases: Solubility decreases with increasing temperature (e.g., CO2 in water).

Expressing Concentrations of Solutions

Solution Concentration Terms

Unit	Definition	Units
Molarity (M)	amount solute (mol) / volume solution (L)	mol/L
Molality (m)	amount solute (mol) / mass solvent (kg)	mol/kg
Mole fraction (X)	amount solute (mol) / total amount (mol)	None
Percent by mass (%)	mass solute / mass solution × 100%	%
Parts per million (ppm)	mass solute / mass solution × 106	ppm
Parts per billion (ppb)	mass solute / mass solution × 109	ppb

Preparing Solutions

• To prepare a solution of given molarity, dissolve the calculated mass of solute in a volumetric flask and add solvent up to the desired volume.

• To dilute a solution: $C_1V_1 = C_2V_2$ or $M_1V_1 = M_2V_2$

• Where $C_1$ and $M_1$ are the concentration/molarity of the stock solution, $V_1$ is the volume of stock solution, $C_2$ and $M_2$ are the concentration/molarity of the dilute solution, and $V_2$ is the volume of dilute solution.

Sample Calculations

• Calculate molarity, molality, and mole fraction using given mass, volume, and density data.

• Use the relationships:

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

  • Molality: $m = \frac{\text{moles solute}}{\text{kg solvent}}$

  • Mole fraction: $X = \frac{\text{moles component}}{\text{total moles}}$

Colligative Properties of Solutions

Definition and Types

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

• Vapor pressure lowering

• Boiling point elevation

• Freezing point depression

• Osmotic pressure

Vapor Pressure Lowering

• The vapor pressure of a solution is lower than that of the pure solvent due to solute-solvent interactions and the blocking of solvent surface area by solute particles.

• Raoult's Law: $P_{\text{solution}} = X_{\text{solvent}} P^\circ_{\text{solvent}}$ Where $X_{\text{solvent}}$ is the mole fraction of the solvent and $P^\circ_{\text{solvent}}$ is the vapor pressure of the pure solvent.

Boiling Point Elevation

• Adding a nonvolatile solute raises the boiling point of the solvent.

• Equation: $\Delta T_b = i K_b m$ Where $\Delta T_b$ is the boiling point elevation, $i$ is the van't Hoff factor, $K_b$ is the boiling point elevation constant, and $m$ is the molality.

Freezing Point Depression

• Adding a solute lowers the freezing point of the solvent.

• Equation: $\Delta T_f = i K_f m$ Where $\Delta T_f$ is the freezing point depression, $i$ is the van't Hoff factor, $K_f$ is the freezing point depression constant, and $m$ is the molality.

Colligative Properties of Electrolytes

• Electrolyte solutions show greater colligative effects because they dissociate into multiple ions.

• van't Hoff factor (i): $i = \frac{\text{moles of particles in solution}}{\text{moles of formula units dissolved}}$

• For nonelectrolytes, $i = 1$; for NaCl, $i \approx 2$; for CaCl2, $i \approx 3$.

Osmosis and Osmotic Pressure

• Osmosis: The movement of solvent from a region of higher solvent concentration to lower solvent concentration through a semipermeable membrane.

• Osmotic pressure (\( \Pi \)): The minimum pressure required to stop osmosis.

• Equation: $\Pi = MRT$ Where $M$ is molarity, $R$ is the gas constant (0.08206 L·atm/mol·K), and $T$ is temperature in Kelvin.

• For electrolytes: $\Pi = iMRT$

Applications

• IV fluids must be isotonic with blood to prevent cell damage due to osmosis.

Sample Exercises

• Predicting solubility patterns based on molecular structure and IMFs.

• Calculating solution concentrations (molarity, molality, mole fraction, mass %).

• Using Raoult's Law to determine vapor pressure lowering.

• Calculating boiling point elevation and freezing point depression for various solutions.

• Determining osmotic pressure and its implications for biological systems.

Additional info: These notes are based on standard General Chemistry curriculum and include all major concepts, equations, and applications relevant to Chapter 13: Solutions.