CH 11 Study Notes for College Physics and Chemistry 2e Chapter on Colloids
CHAPTER OUTLINE
11.1 The Dissolution Process
11.2 Electrolytes
11.3 Solubility
11.4 Colligative Properties
11.5 Colloids
11.1 THE DISSOLUTION PROCESS
Learning Objectives:
Describe the basic properties of solutions and how they form
Predict whether a given mixture will yield a solution based on molecular properties of its components
Explain why some solutions either produce or absorb heat when they form
Matter Classification
Pure substances
Elements
Compounds
Mixtures
Homogeneous
Heterogeneous
Solutions:
Definitions:
Solutions are homogeneous mixtures of two or more substances.
They form when there is sufficient attraction between the solute and solvent molecules.
Components:
Solvent: present in a larger amount
Solute: present in a smaller amount
Example of an Aqueous Solution:
NaCl(s)
ightarrow Na^+(aq) + Cl^-(aq)Note: Most chemical reactions occur in solutions; e.g., living systems
Characteristics of Solutions:
Solutions exist in all three physical states:
Solid solution: metal alloys (two or more metals)
Gas solution: Example: air (nitrogen, oxygen, etc.)
Liquid solution: Primarily discussing aqueous solutions
Example dissolving in water:
Sucrose: C{12}H{22}O{11}(s) ightarrow C{12}H{22}O{11}(aq)
Aluminum sulfate: Al2(SO4)3(s) ightarrow 2Al^{3+}(aq) + 3SO4^{2-}(aq)
Types of Solutes and Solvents
Table 9.1: Some Examples of Solutions
Gas Solutions: Air (O$2$ in N$2$)
Liquid Solutions:
Soda water (CO$2$ in H$2$O)
Vinegar (HC$2$H$3$O$2$ in H$2$O)
Seawater (NaCl in H$_2$O)
Solid Solutions: Brass (Cu + Zn), Steel (Fe + C)
Formation of Solutions:
Intermolecular attractive forces (IMF) influencing solution formation:
solute-solute
solvent-solvent
solute-solvent
Process Steps:
Step 1: Overcome solute-solute IMFs (energy consumed)
Step 2: Overcome solvent-solvent IMFs (energy consumed)
Step 3: Form solute-solvent attractive forces (energy released)
Energy Changes in Solution Formation
Relative magnitudes of energy changes determine the process type:
Exothermic: energy released
Endothermic: energy absorbed
No reaction: unable to occur
Examples & Applications
Instant cold pack: dissolving ammonium nitrate in water absorbs heat (endothermic process).
11.2 ELECTROLYTES
Learning Objectives:
Define and give examples of electrolytes
Distinguish between physical and chemical changes during dissolution for ionic and covalent electrolytes
Relate electrolyte strength to solute-solvent attractive forces
Electrolytes:
Classification: Electrolytes vs Nonelectrolytes
Example electrolytes crucial for body: sodium, potassium, chloride, bicarbonate
Strong Electrolytes:
Dissociate 100% in water (e.g., NaCl(s)
ightarrow Na^+(aq) + Cl^{-}(aq)), conduct electricity
Weak Electrolytes:
Dissociate only slightly (e.g., HF(aq)
ightleftharpoons H^+(aq) + F^-(aq))
Nonelectrolytes: Molecules that do not produce ions in water (e.g., CH4O(l) ightarrow CH4O(aq))
Conductivity of Aqueous Solutions
Some current flows:
Examples:
Nonelectrolyte: no current
Weak electrolyte: limited current
Strong electrolyte: extensive current
11.3 SOLUBILITY
Learning Objectives:
Describe temperature and pressure effects on solubility
State Henry’s Law and apply in gas solubility calculations
Explain degrees of solubility for liquid-liquid solutions
Solubility Definitions:
Solubility: maximum concentration of a solute achievable in a solvent at specific temperature.
Saturated: concentration = solubility
Unsaturated: concentration < solubility
Supersaturated: concentration > solubility
Factors Affecting Solubility:
Temperature: Solubility increases with temperature for solids and liquids; decreases for gases.
Pressure: Affects gas solubility (higher pressure = higher solubility).
Example of Gas Solubility:
Henry’s Law: Cg = kPg
Relationship: Gas solubility increases with increasing gas pressure.
Relation:
k = constant for gas-liquid system
Real World Applications:
Thermal pollution impact on ecosystems by decreasing dissolved oxygen levels.
11.4 COLLIGATIVE PROPERTIES
Learning Objectives:
Express solution concentrations in mole fraction and molality
Describe the effect of solute concentration on solution properties
Perform calculations for colligative properties
Explain distillation and osmosis processes
Concentration Units:
Molarity (M): M = rac{moles~solute}{volume~solution}
Mass percent: grams of solute in 100 grams of solution.
Mole fraction (X): X_A = rac{moles~A}{total~moles~solution}
Molality (m): moles of solute per kg of solvent.
Colligative Properties:
Properties dependent on solute concentration, not identity.
Include:
Vapor Pressure Lowering
Boiling Point Elevation: riangle Tb = m imes Kb
Freezing Point Depression: riangle Tf = m imes Kf
Osmotic Pressure: P_{osmotic} = MRT
Example Calculations:
Vapor pressure lowering explanation using Raoult’s Law
Boiling point elevation and freezing point depression applications
Practical Applications:
Distillation techniques used to separate mixtures (e.g., crude oil).
11.5 COLLOIDS
Learning Objectives:
Describe colloidal compositions and properties
List technological applications of colloids
Colloidal Systems:
Differences among solutions, colloids, and suspensions:
Suspensions: heterogeneous, large particle visibility, settle out over time
Solutions: homogeneous, invisible solutes, do not settle
Colloids: intermediate properties, larger particles that do not settle
Key Concepts:
Tyndall Effect: Colloidal particles scatter light, resulting in visibility (e.g., fog).
Colloidal Components:
Dispersed Phase: minor component
Dispersion Medium: major component
Application in soap/detergent emulsifiers explained.
Colloids in cleaning actions due to amphiphilic structure of detergents.
Historical Aspect:
Production of soaps from fats with bases demonstrated practical applications of colloids in everyday products.