Introduction to Chemistry: Exam 2 Study Guide

Chapter 4: Atoms and Elements

Fundamental Concepts of Atoms

Atoms are the basic building blocks of matter. Understanding their structure and properties is essential for studying chemistry.

• Atoms: The smallest unit of an element that retains its chemical properties.

• Subatomic Particles: Atoms are composed of protons (positively charged), neutrons (neutral), and electrons (negatively charged).

• Atomic Number (Z): The number of protons in the nucleus of an atom; determines the element's identity.

• Mass Number (A): The sum of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: Carbon-12 and Carbon-14 are isotopes of carbon, differing in their number of neutrons.

• Use the periodic table to identify elements by their atomic number and symbol.

• Calculate atomic mass from natural abundances and isotopic masses.

Formula:

$\text{Atomic Mass} = \sum (\text{Fractional Abundance} \times \text{Isotope Mass})$

Chapter 5: Molecules and Compounds

Chemical Formulas and Compound Classification

Chemical compounds are formed when atoms combine in fixed ratios. Their composition is represented by chemical formulas.

• Chemical Formula: Shows the types and numbers of atoms in a compound (e.g., H2O).

• Empirical Formula: Shows the simplest whole-number ratio of atoms.

• Molecular Formula: Shows the actual number of each type of atom in a molecule.

• Ionic Compounds: Composed of positive and negative ions held together by electrostatic forces.

• Molecular Compounds: Composed of molecules formed by covalent bonds between nonmetals.

• Acids: Substances that release hydrogen ions (H+) in solution.

Example: NaCl is an ionic compound; CO2 is a molecular compound.

• Recognize and name chemical compounds using systematic nomenclature.

• Calculate formula mass by summing the atomic masses of all atoms in the formula.

Formula:

$\text{Formula Mass} = \sum (\text{Number of Atoms} \times \text{Atomic Mass})$

Chapter 10: Chemical Bonding

Lewis Structures and Molecular Geometry

Chemical bonding involves the sharing or transfer of electrons between atoms. Lewis structures help visualize bonding and electron arrangement.

• Lewis Structure: Diagram showing valence electrons as dots around atoms.

• Electron Geometry: The spatial arrangement of electron groups around a central atom.

• Molecular Geometry: The arrangement of atoms (excluding lone pairs) in a molecule.

• Bond Angle: The angle between adjacent bonds in a molecule.

• Predict molecular shapes and polarity using VSEPR theory.

Table: Electron Geometry, Bond Angle, and Molecular Geometry

# of e- groups	Electron geometry	Bond angle	Molecular Geometry
2	linear	180	linear
3	trigonal planar	120	Trigonal planar (0 lone pairs); Bent (1 lone pair)
4	tetrahedral	109.5	Tetrahedral (0 lone pairs); Trigonal pyramidal (1 lone pair); Bent (2 lone pairs)

Formula:

$\text{Polarity:} \quad \text{A molecule is polar if it has an uneven distribution of electron density.}$

Organic Chemistry and Intermolecular Forces

Structural Formulas and Hydrocarbons

Organic chemistry focuses on carbon-containing compounds. Hydrocarbons are classified based on the types of bonds between carbon atoms.

• Structural Formula: Shows how atoms are connected in a molecule.

• Alkanes: Hydrocarbons with only single bonds.

• Alkenes: Hydrocarbons with at least one double bond.

• Alkynes: Hydrocarbons with at least one triple bond.

• Aromatic Hydrocarbons: Contain benzene rings.

• Identify common functional groups (e.g., alcohols, ketones, carboxylic acids).

Example: Ethene (C2H4) is an alkene; Benzene (C6H6) is aromatic.

Intermolecular Forces

Intermolecular forces are attractions between molecules that affect physical properties such as boiling and melting points.

• Dispersion Forces (London Forces): Weak forces present in all molecules due to temporary dipoles.

• Dipole-Dipole Forces: Attractions between polar molecules.

• Hydrogen Bonds: Strong dipole-dipole interactions involving H bonded to N, O, or F.

• Ion-Dipole Forces: Attractions between ions and polar molecules.

Comparison Table: Types of Intermolecular Forces

Type	Strength	Occurs Between
Dispersion	Weakest	All molecules
Dipole-Dipole	Intermediate	Polar molecules
Hydrogen Bond	Strong	H bonded to N, O, or F
Ion-Dipole	Strongest	Ions and polar molecules

Formula:

$\text{Boiling Point} \propto \text{Strength of Intermolecular Forces}$

Additional info: Some explanations and examples have been expanded for clarity and completeness.