Introduction to Chemistry: Key Concepts and Study Guide

Chapter 9: Atomic Structure and Electromagnetic Radiation

Electromagnetic Radiation

Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. It is fundamental to understanding atomic structure and chemical behavior.

• Light is embodied in an electromagnetic field and travels at $3.0 \times 10^8$ m/s (c).

• Wavelength (λ): The distance between successive crests of a wave.

• Amplitude: The height of the wave, related to intensity.

• Frequency (ν): The number of wave cycles per second.

• c = λ × ν: The speed of light equation.

• Short wavelength = high frequency; long wavelength = low frequency.

• High frequency = high energy; low frequency = low energy.

• Energy of a photon: $E_{\text{quantum}} = h \nu = 6.626 \times 10^{-34} \text{J·s}$

• Types of electromagnetic radiation: Gamma, X-rays, UV, visible (violet – red), IR (heat), radar/microwave, radio/TV.

Continuous Spectrum vs. Discrete Spectrum

A continuous spectrum contains all wavelengths, while a discrete spectrum contains only specific wavelengths, as seen in atomic emission spectra.

• Continuous spectrum: All colors/wavelengths present.

• Discrete spectrum: Only certain wavelengths present, characteristic of elements.

Schrödinger Theory (Quantum Mechanical Theory)

The quantum mechanical model describes electrons in terms of probabilities and energy levels, rather than fixed orbits.

• Shell (n = 1, 2, 3...): Energy level; higher n means larger and higher energy shell.

• Subshells (l = 0, 1, 2...): Shape of probability region (s, p, d, f).

• Orbitals: Regions of space where electrons are likely to be found.

• Number of orbitals per shell:

  • n = 1: 1 orbital

  • n = 2: 4 orbitals

  • n = 3: 9 orbitals

  • n = 4: 16 orbitals

• Orbital orientation: Each orbital has a specific orientation in space.

• Pauli Exclusion Principle: 2 electrons can occupy each orbital, but must have opposite spins.

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom or ion, based on quantum mechanics.

• Use quantum numbers and the Aufbau principle to fill orbitals.

• Be able to write electron configurations for atoms and ions.

• Apply the Pauli exclusion principle and Hund’s rule.

Chapter 10: Chemical Bonding

Ionic Bonding

Ionic bonding involves the transfer of electrons from one atom to another, resulting in the formation of ions.

• Ionic bond: Electrostatic attraction between cation (positive ion) and anion (negative ion).

• Formula unit: The smallest unit of an ionic compound.

Covalent Bonding

Covalent bonding involves the sharing of electrons between atoms to achieve stability.

• Non-polar covalent bond: Equal sharing of electrons (atoms have similar electronegativity).

• Polar covalent bond: Unequal sharing of electrons (difference in electronegativity).

• Use electronegativity chart to determine bond type.

Bond Properties

Bonds have specific lengths and strengths, which depend on the atoms involved and the type of bond.

• Bond length: Optimal distance between nuclei.

• Bond strength: Increases with number of electrons shared (double/triple bonds stronger than single).

Lewis Dot Structures

Lewis dot structures represent the valence electrons in molecules and help predict bonding and molecular shape.

• Recall that carbon gets "4" bonds; hydrogen gets "1" bond.

• Exceptions: Boron (often less than octet), odd number of electrons (e.g., NO), resonance structures.

• Resonance: Multiple valid Lewis structures for a molecule.

Geometry and the VSEPR Theory

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the shapes of molecules based on electron pair repulsion.

• Determine geometry from Lewis dot structure.

• Associate shapes with electron pair arrangements.

Overall Molecular Polarity

Molecular polarity depends on the symmetry and distribution of charge within the molecule.

• Symmetrical molecule: Non-polar.

• Non-symmetrical molecule: Polar.

• Look for lone pairs and central atom arrangement.

Chapter 11: States of Matter and Gas Laws

Properties of Matter

Matter can exist in different states and has various physical properties.

• Identify properties of matter (mass, volume, density, etc.).

Kinetic Theory of Gases

The kinetic theory explains the behavior of gases in terms of particle motion.

• Gases consist of particles in constant, random motion.

• Diffusion: Movement of particles from high to low concentration.

• Effusion: Escape of gas through a small hole.

Graham’s Law of Effusion

Graham’s law relates the rate of effusion to the molar mass of gases.

• $\text{Rate}_1/\text{Rate}_2 = \sqrt{M_2/M_1}$

Gas Pressure and Measurement

Gas pressure is the force exerted by gas particles on the walls of a container.

• Measured in atmospheres (atm), pascals (Pa), or torr.

• Barometer: Instrument for measuring atmospheric pressure.

Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

• Boyle’s Law: $P V = k$ (at constant temperature, pressure is inversely proportional to volume).

• Charles’s Law: $V/T = k$ (at constant pressure, volume is directly proportional to temperature).

• Avogadro’s Law: $V/n = k$ (volume is directly proportional to number of moles).

• Combined Gas Law: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

• Ideal Gas Law: $P V = n R T$

Chapter 12: Intermolecular Forces and States of Matter

Properties of Condensed States

Condensed states include liquids and solids, which have stronger intermolecular forces than gases.

• Identify properties of liquids and solids.

Intermolecular Forces

Intermolecular forces are attractions between molecules that affect physical properties.

• London Dispersion Forces: Present in all molecules and atoms; increase with mass and surface area.

• Dipole-Dipole Forces: Occur in polar molecules; permanent dipole.

• Hydrogen Bonding: Strongest type; occurs when H is attached to N, O, or F and interacts with N, O, or F in another molecule.

Properties Observed Due to Intermolecular Forces

Intermolecular forces influence many physical properties of substances.

• Surface tension: Stronger forces lead to higher surface tension.

• Viscosity: Resistance to flow; stronger forces increase viscosity.

• Boiling point: Stronger forces result in higher boiling points.

• Melting point: Stronger forces result in higher melting points.

• Vapor pressure: Stronger forces result in lower vapor pressure.

Changes in States of Matter

Changes in state involve energy changes and are classified as endothermic or exothermic.

• Exothermic: Releases energy (condensation, freezing).

• Endothermic: Absorbs energy (melting, vaporization, sublimation).

• Phase changes:

  • Condensation (gas to liquid)

  • Freezing (liquid to solid)

  • Melting (solid to liquid)

  • Vaporization (liquid to gas)

  • Sublimation (solid to gas)

Additional info: Solid arrangements (crystal structures) are not required for the exam.