CHEM 110 Final Exam Study Guide: Key Concepts and Principles

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Chemical Bonding

Sharing vs Transfer of Electrons

Chemical bonds form when atoms interact to achieve stable electron configurations. The way electrons are distributed determines the type of bond formed.

Sharing of electrons: Occurs in covalent bonds, where two nonmetals share one or more pairs of electrons.
Transfer of electrons: Occurs in ionic bonds, where one atom (usually a metal) donates electrons to another atom (usually a nonmetal), resulting in oppositely charged ions.
Example: In NaCl, sodium transfers an electron to chlorine, forming Na+ and Cl-.

Ionic vs Covalent Compounds

Ionic compounds: Composed of cations and anions held together by electrostatic forces. Usually form crystalline solids with high melting points.
Covalent compounds: Consist of molecules formed by shared electrons. Often have lower melting and boiling points and can exist as gases, liquids, or solids.
Example: H2O is covalent; NaCl is ionic.

Core vs Valence Electrons

Core electrons: Electrons in inner shells, not involved in bonding.
Valence electrons: Electrons in the outermost shell, responsible for chemical reactivity and bonding.
Example: Oxygen (atomic number 8) has 6 valence electrons.

Octet Rule

Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons, resembling noble gas configuration.
Exceptions: Hydrogen (2 electrons), Boron (6 electrons), expanded octets for elements in period 3 or higher.

Single, Double, and Triple Bonds

Single bond: One pair of shared electrons (e.g., H–H).
Double bond: Two pairs of shared electrons (e.g., O=O).
Triple bond: Three pairs of shared electrons (e.g., N≡N).

Lewis Dot Structures

Visual representations of valence electrons around atoms.
Used to predict bonding and molecular structure.

Molecular Structure and Polarity

Dipole Moment

Occurs when there is a separation of charge in a molecule due to differences in electronegativity.
Measured in Debye units (D).
Example: H2O has a significant dipole moment.

Electron and Molecular Geometries

Describes the 3D arrangement of atoms and electron pairs around a central atom.
Predicted using VSEPR theory.
Example: CH4 is tetrahedral; H2O is bent.

Polar vs Nonpolar Molecules

Polar molecules: Have an uneven distribution of charge (e.g., H2O).
Nonpolar molecules: Have an even distribution of charge (e.g., CO2).

Hydrophilic vs Hydrophobic

Hydrophilic: "Water-loving"; substances that dissolve well in water (usually polar).
Hydrophobic: "Water-fearing"; substances that do not dissolve well in water (usually nonpolar).

Intermolecular Forces and Properties

All Intermolecular Attractive Forces

London dispersion forces: Weak, present in all molecules.
Dipole-dipole interactions: Between polar molecules.
Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

Polarizability

The ease with which the electron cloud of a molecule can be distorted, affecting the strength of London dispersion forces.

Viscosity

A measure of a liquid's resistance to flow; higher intermolecular forces lead to higher viscosity.

Gases and Gas Laws

All Gas Laws

Boyle's Law: (at constant T and n)
Charles's Law: (at constant P and n)
Avogadro's Law: (at constant P and T)
Ideal Gas Law:

Solutions and Solubility

Homogenous Solutions

Mixtures with uniform composition throughout.
Example: Saltwater.

Insoluble vs Soluble

Soluble: Substances that dissolve in a solvent.
Insoluble: Substances that do not dissolve appreciably.

Chemical Composition and Analysis

Formula and Experimental Mass Analysis

Determining the mass percent of each element in a compound.
Formula mass: Sum of atomic masses in a formula unit.
Experimental mass analysis: Laboratory determination of composition.

Empirical vs Molecular Formula

Empirical formula: Simplest whole-number ratio of atoms in a compound.
Molecular formula: Actual number of atoms of each element in a molecule.
Example: Glucose: Empirical CH2O, Molecular C6H12O6.

Polyatomic Ions and Acids

Polyatomic Ions

Charged species composed of two or more atoms covalently bonded.
Examples: NO3- (nitrate), SO42- (sulfate).

Oxy Acids

Acids containing hydrogen, oxygen, and another element.
Example: H2SO4 (sulfuric acid).

Chemical Reactions

Oxidation/Reduction (Redox)

Involves transfer of electrons between species.
Oxidation: Loss of electrons.
Reduction: Gain of electrons.

Neutralization Reactions

Reaction between an acid and a base to produce water and a salt.
Example:

Decomposition and Combustion

Decomposition: A single compound breaks down into two or more simpler substances.
Combustion: A substance reacts with oxygen, releasing energy, usually producing CO2 and H2O.

Atomic Structure and Electron Configuration

Orbitals

Regions in an atom where electrons are likely to be found.
Types: s, p, d, f.

Electron Configuration

Describes the arrangement of electrons in an atom.
Example: Oxygen: 1s2 2s2 2p4

Orbital Diagram

Visual representation of electron configuration using boxes and arrows to show electron spin.

Acids, Bases, and Conjugate Pairs

Acid-Base Reactions

Acids donate protons (H+), bases accept protons.
Example:

Conjugate Pairs

Each acid has a conjugate base, and each base has a conjugate acid, differing by one proton.
Example: HCl (acid) and Cl- (conjugate base).

Thermochemistry

Study of energy changes during chemical reactions.
Key equation:
Where q is heat, m is mass, c is specific heat, and is temperature change.

Phase Changes

Evaporation vs Condensation

Evaporation: Liquid to gas phase change.
Condensation: Gas to liquid phase change.

Additional info: This guide covers all major topics typically found in an introductory college chemistry course, including atomic structure, bonding, chemical reactions, states of matter, solutions, acids and bases, and thermochemistry. Students should be familiar with definitions, examples, and basic calculations for each topic.