General Chemistry I: Comprehensive Study Guide

Chapter 1: Introduction to Chemistry

1.1 Classification of Matter

Chemistry studies the composition, structure, and properties of matter. Matter can be classified in several ways:

• States of Matter: Solid, liquid, gas

• Pure Substances vs. Mixtures:

  • Element: A substance that cannot be broken down into simpler substances.

  • Compound: A substance composed of two or more elements chemically combined.

  • Mixtures: Physical combinations of substances. Homogeneous mixtures have uniform composition (solutions), while heterogeneous mixtures do not.

1.2 Atomic Structure and Molecules

• Atomic Model: Atoms are the basic units of matter, consisting of protons, neutrons, and electrons.

• Molecules: Two or more atoms bonded together.

1.3 Chemical and Physical Properties and Changes

• Chemical Properties: Describe a substance's ability to undergo chemical changes (e.g., flammability).

• Physical Properties: Can be observed without changing the substance (e.g., melting point, density).

• Chemical Changes: Result in new substances (e.g., rusting iron).

• Physical Changes: Do not alter the chemical identity (e.g., melting ice).

1.4 Units and Measurements

• SI Units: Standard units for scientific measurements (meter, kilogram, second, mole, kelvin, ampere, candela).

• Prefixes: Used to indicate multiples or fractions of units (e.g., kilo-, centi-, milli-).

• Avogadro's Number: $N_A = 6.022 \times 10^{23}$ (number of particles in a mole).

• Dimensional Analysis: A method for converting between units using conversion factors.

1.5 Density and Physical Properties

• Density: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Used to identify substances and convert between mass and volume.

Chapter 2: Atomic Structure

2.1 Subatomic Particles

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles orbiting the nucleus.

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

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

2.2 Ions

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

• Charge Calculation: $\text{Charge} = \text{number of protons} - \text{number of electrons}$

Chapter 3: The Periodic Table

3.1 Organization and Trends

• Groups: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows.

• Main Group/Transition Elements: Main group (s and p blocks), transition elements (d block).

• Metals, Nonmetals, Metalloids: Classified by properties and position on the table.

• Special Groups: Alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides.

Chapter 8: Quantum Theory and Atomic Structure

8.1 Duality of Matter

• All matter exhibits both wave-like and particle-like properties.

8.2 Properties of Waves

• Wavelength ($\lambda$): Distance between two consecutive peaks.

• Frequency ($\nu$): Number of waves passing a point per second.

• Speed of Light ($c$): $c = \lambda \nu$

• Relationship: As wavelength increases, frequency decreases (inverse relationship).

8.3 Electromagnetic Spectrum

• Includes all types of electromagnetic radiation, from radio waves to gamma rays.

• Visible light is a small portion of the spectrum.

8.4 Energy of Photons

• Planck's Equation: $E = h\nu$

• Rydberg Equation: $\Delta E = -R_H \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$

• De Broglie Wavelength: $\lambda = \frac{h}{mv}$

8.5 Quantum Numbers and Atomic Orbitals

• Principal Quantum Number ($n$): Energy level (positive integer).

• Angular Momentum Quantum Number ($l$): Shape of orbital (0 = s, 1 = p, 2 = d, 3 = f).

• Magnetic Quantum Number ($m_l$): Orientation of orbital.

• Spin Quantum Number ($m_s$): Electron spin (+1/2 or -1/2).

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing.

8.6 Electron Configurations

• Describes the arrangement of electrons in an atom.

• Use the Aufbau principle (fill lowest energy orbitals first).

• Example: Carbon (Z=6): 1s2 2s2 2p2

Periodic Trends

Trends in the Periodic Table

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electron Affinity: Generally becomes more negative across a period.

• Electronegativity: Increases across a period, decreases down a group.

• Metallic Character: Increases down a group, decreases across a period.

Sample Table: Electromagnetic Spectrum

Type	Wavelength (nm)	Frequency (Hz)	Energy
Gamma Rays	< 0.01	> 1019	Highest
X-Rays	0.01 - 10	1016 - 1019	High
Ultraviolet	10 - 400	1015 - 1016	Moderate
Visible	400 - 700	4.3 × 1014 - 7.5 × 1014	Moderate
Infrared	700 - 106	1012 - 4.3 × 1014	Low
Microwave	106 - 109	109 - 1012	Lower
Radio	> 109	< 109	Lowest

Key Equations and Constants

• Avogadro's Number: $N_A = 6.022 \times 10^{23}$

• Speed of Light: $c = 2.99792458 \times 10^8$ m/s

• Planck's Constant: $h = 6.62607 \times 10^{-34}$ J·s

• Rydberg Constant: $R_H = 2.17987 \times 10^{-18}$ J

• Energy of a Photon: $E = h\nu$

• Rydberg Equation: $\Delta E = -R_H \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$

• De Broglie Wavelength: $\lambda = \frac{h}{mv}$

Study and Exam Tips

• Practice problems from the textbook and worksheets.

• Understand and memorize key equations and constants.

• Bring a scientific calculator and required materials to the exam.

• Show all work for calculations; partial credit is given for correct steps.

Additional info: This guide is based on a course syllabus and exam review for a General Chemistry I college course. It covers foundational topics, key equations, and periodic trends essential for first-semester chemistry students.