General Chemistry I: Midterm Exam 1 Study Guide

Chapter 1: Units, Measurements, and Scientific Notation

Units of Measurement

Understanding units is fundamental in chemistry for expressing quantities and performing calculations.

• Base Units: Standard units such as kilogram (kg), meter (m), second (s), mole (mol), etc.

• Metric Prefixes: Used to express multiples or fractions of base units (e.g., kilo-, centi-, milli-, micro-, nano-).

• Scientific Notation and Significant Figures: Scientific notation is used to express very large or small numbers. Significant figures indicate the precision of a measurement.

• Density: Defined as mass per unit volume. Formula: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

Chapter 2: Atomic Theory and Structure

Atomic Models

Atomic models have evolved to explain the structure and behavior of atoms.

• Key Contributors: Thomson (discovered electron), Rutherford (nuclear model), Bohr (quantized orbits), Schrödinger (quantum mechanics).

• Modern Model: Electrons exist in orbitals defined by quantum mechanics.

Nuclear Symbols and Isotopes

• Nuclear Symbol: Notation showing atomic number and mass number.

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

• Atomic Mass: Weighted average mass of all isotopes of an element.

• Avogadro's Number: $6.022 \times 10^{23}$ particles per mole.

Chapter 3: Quantum Mechanics and Atomic Structure

Kinetic Energy

• Formula: $KE = \frac{1}{2}mv^2$

Electromagnetic Radiation

• Relationship between wavelength, frequency, and speed of light: $c = \lambda \nu$ where $c$ is the speed of light ($3.00 \times 10^8$ m/s), $\lambda$ is wavelength, and $\nu$ is frequency.

• Energy of a Photon: $E = h\nu$ where $h$ is Planck's constant ($6.626 \times 10^{-34}$ J·s).

Photoelectric Effect

• Equation: $KE = h\nu - \phi$ where $\phi$ is the work function of the material.

Bohr Model of the Atom

• Absorption vs. Emission: Electrons absorb energy to move to higher energy levels and emit energy when returning to lower levels.

de Broglie Wavelength

• Equation: $\lambda = \frac{h}{mv}$ where $\lambda$ is wavelength, $h$ is Planck's constant, $m$ is mass, and $v$ is velocity.

Heisenberg Uncertainty Principle

• Equation: $\Delta x \cdot \Delta p \geq \frac{h}{4\pi}$ This principle states that the position and momentum of a particle cannot both be precisely determined at the same time.

Quantum Mechanics

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

• Quantum Numbers: Describe the properties of atomic orbitals (n, l, ml, ms).

• Rules for Quantum Numbers: n, l, ml, ms must follow specific allowed values.

Chapter 4: Electron Configuration and Periodic Trends

Electron Configuration and Exceptions

• Electron Configuration: Distribution of electrons among atomic orbitals.

• Exceptions: Some atoms/ions have electron configurations that differ from the expected order due to stability considerations.

• 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.

• Aufbau Principle: Electrons occupy the lowest energy orbitals first.

Periodic Properties and Trends

• Periodic Table Trends: Atomic radius, ionization energy, electron affinity, metallic character.

• Classification: Metals, nonmetals, metalloids; grouped by similar properties.

Chapter 5: Chemical Bonding and Molecular Structure

Lewis Structures

• Lewis Structures: Diagrams showing the bonding between atoms and the lone pairs of electrons in a molecule.

Bond Polarity and Electronegativity

• Bond Polarity: Describes the distribution of electron density in a bond (nonpolar, polar covalent, ionic).

• Electronegativity: A measure of an atom's ability to attract electrons in a chemical bond.

• Electronegativity Chart: Provided on the exam for reference.

Additional Info

• Nomenclature: Moved to Midterm Exam II (not covered in this exam).