Chemistry 101 Final Exam Review Checklist
This is a guide for study. There may be additional material not listed here that will appear on the exam. Also study your
text notes, lecture notes, quizzes, exams, and problems.
Chapter 1
Classification of Matter: Mixtures (heterogeneous and homogeneous) and Pure Substances (compounds and
elements)
Physical and Chemical Changes
Physical and Chemical Properties
Chapter 2
Uncertainty: Accuracy vs. Precision, Significant Figures in calculations and measurements
Metric Prefixes
Exponential or Scientific Notation
Unit or Dimensional Analysis
Chapter 3
Modern Atomic Theory and Structure
Isotopes: Atomic Number, Mass Number, Isotope Symbol (example, 1H)
Electron Configurations of Ground State Atoms and Ions (1s2s2p3s3p4s3d4p5s….)
Periodic Behavior of Elements: Valence Electrons, Core Electrons
Shapes of Orbital (s, p, d)
Quantized Energy
Chapter 4
Ionic Compounds
Ionic Bond Formation
The Formula Unit
Naming Ionic Compounds: Binary Type I and II
Polyatomic Cations and Anions
Chapter 5
Molecular or Covalent Compounds
Covalent Bond Formation
Molecular Formula
Coordinate Covalent Bond
Lewis Dot Structures for both Covalent and Ionic Compounds
Total Number of Valence Electrons
Skeleton Structure
Octet Rule
Molecular Geometry and the VSEPR Model
Linear-2 electron groups around central atom
Trigonal Planar-3 electron groups around central atom (0 and 1 lp)
Tetrahedral -4 electron groups around central atom (0, 1, 2 lp)
Bond Polarity: Electronegativity Di fferences
Molecular Polarity: Combine Molecular Geometry and Bond Polarity
Naming Molecular Compounds: Binary Type III
Chapter 6
Stoichiometry
Balancing equations
The Mole
Mass-Mole Conversions
Mole Ratios from Balanced Equations
Limiting Reactants (How many bikes can you make with 57 tires and 149 frames?)
% Yield
Reaction Types: Precipitation, Acid-Base, Redox (Oxidation Numbers)
Complete and Net Ionic Reactions
Chapter 7
Enthalpy of a Reaction: ∆ H ≈ Energy Added to Break Reactants Bonds – Energy Released to Form Product Bonds
Exothermic (∆
∆ H < 0, P are stable relative to R) vs. Endothermic (∆
∆ H > 0, R are stable relative to P)
Entropy: Disorder (∆
∆ S > 0), Order ( ∆ S < 0)
Gibbs Free Energy: ∆ G = ∆ H - T∆
∆S
Spontaneous: ∆ G < 0, exergonic
Non-Spontaneous: ∆ G < 0, endergonic
Nature favors spontaneous reactions that are give stable products with high entropy
Every reaction has a barrier (Ea, activation energy) that depends on the frequency, energy, and orientation of
collisions). The Ea controls the Rate of Reaction.
Reaction Coordinate Diagram
Rate of Reaction affected by Temperature, Concentration of Reactants and Products, and Catalyst
All Reactions are Reversible in Theory but not in Practice, example CH4 + 2O2 ------> CO2 + 2H2O
Truly Reversible Reactions are at Equilibrium for a Given Temperature
At Equilibrium the Forward and Reverse Rates are Equal, the Product and Reactant Concentrations are Constant,
and the Equilibrium State is Dynamic
Write t he Equilibrium Expression (K) for a Balanced Equation
If K>1, then ?
If K<1, then?
If K = 1, then ?
LeChâtelier’s Principle
Chapter 8
Intermolecular Forces (IMF): London Dispersion, Dipole-Dipole, H-Bond
Determining IMF for a Particular Molecule based on the Molecule’s Polarity (see Chapter 5)
Chapter 9
The Solution Process: A Trade Off between the Energy Added to Separate Pure Solute and Solvent Particles and
the Energy Gained by Mixing Solute and Solvent Particles AND Entropy Always Increases for Solutions
“Like Dissolves Like”: Similar IMF between Solute and Solvent
Solubility is Affected by Temperature and Pressure
Concentration: Molarity (moles of solute per liter of solution), % by mass (w/w), % by volume (v/v), % by
mass/volume (w/v), ppm
Dilution: M1V1 = M2V2
Chapter 10
See the Handout and Focus on 10.1 -10.13 and 10.15