GENERAL CHEMISTRY 1

General Overview

The following notes provide an in-depth overview of key concepts in general chemistry, including atomic structure, chemical bonding, quantum mechanics, and the periodic table. This summary is organized by themes and topics to facilitate easier studying and understanding.

Atomic Structure and Discoveries

• Atomic Model Timeline: The development of atomic theory has a significant history:

  • 1896: J.J. Thomson discovered the electron and proposed the concept of protons.

  • 1909: Ernest Rutherford established the nuclear model of the atom, identifying the nucleus as a dense central core.

  • 1913: Henry Moseley determined the atomic number, thus refining the organization of the periodic table.

  • 1913: Niels Bohr introduced the Bohr model, which depicted electrons in fixed orbits.

  • 1932: James Chadwick discovered the neutron, completing the modern understanding of atomic structure.

Fundamental Particles

Charge and Mass of Subatomic Particles

Atomic Number and Isotopes

• The atomic number (Z) refers to the number of protons in an atom, distinguishing one element from another.

• Isotopes are variants of the same element with different numbers of neutrons.

• The mass number (A) is the sum of protons and neutrons in an atomic nucleus:mass number = protons + neutrons.

Bohr’s Theory of the Atom

• Introduction (1913): Niels Bohr proposed that electrons move in fixed circular orbits around the nucleus, which are quantized levels of energy.

• Postulates of Bohr's Theory:

  1. Electrons occupy specific orbits with no energy lost while in orbit.

  2. Electrons can transition between energy levels by absorbing or emitting photons, as given by the equation ΔE = hν.

  3. The angular momentum of electrons in these orbits is quantized.

• Bohr successfully calculated radii of hydrogen's orbits and their corresponding energy levels, aligning with observed spectral lines.

Hydrogen Spectrum

• An electron in the hydrogen atom resides in its ground state (n=1) and may be excited to higher levels (n=2, 3, etc.) upon energy absorption.

• Excited electrons emit energy in photon form when transitioning back to lower energy states, giving rise to the observed spectral series:

  • Lyman Series: Emissions to n=1.

  • Balmer Series: Emissions to n=2.

  • Paschen Series: Emissions to n=3, etc.

Limitations of Bohr’s Model

• Bohr's model has limitations, including:

  • It cannot explain the Zeeman effect (magnetic field influences on atomic spectra).

  • It fails to explain the Stark effect (electric field effects).

  • It violates the Heisenberg Uncertainty Principle, especially in larger atoms.

Quantum Mechanics

Quantum Numbers Overview

• Four quantum numbers describe the properties of electrons in atoms:

  1. Principal Quantum Number (n): Indicates the shell (energy level).

  2. Azimuthal Quantum Number (l): Defines the shape of the orbital.

  3. Magnetic Quantum Number (m): Specifies the orientation of the orbital.

  4. Spin Quantum Number (s): Indicates the electron's spin direction.

Details of Quantum Numbers

Principal Quantum Number 'n'

• The principal quantum number indicates the major energy level and is denoted by positive integers (n = 1, 2, 3,...).

• Each energy level can hold a maximum of 2n² electrons.

Azimuthal Quantum Number 'l'

• Determines the shape of orbitals (s, p, d, f):

  • l = 0 for s (spherical), l = 1 for p (dumbbell), l = 2 for d (complex shapes), and l = 3 for f.

• Values range from 0 to n-1, signifying available sublevels.

Magnetic Quantum Number 'm'

• Presents the orientation of the orbital, allowing multiple orientations for each shape:

  • For p orbitals (l=1), m can take values -1, 0, +1 (px, py, pz orientations).

  • For d orbitals (l=2), m ranges from -2 to +2, corresponding to different orientations.

Spin Quantum Number 's'

• Represents the intrinsic spin of the electron which can be either +1/2 or -1/2.

• Electrons with parallel spins (same sign) repel, while electrons with opposite spins can coexist in the same orbital.

Summary

These notes cover various facets of atomic theory from foundational discoveries to the quantum mechanical description of atomic structure, including limitations of historical models. Understanding these concepts is crucial for grasping the complexities of chemical bonding and the behavior of atoms in various chemical contexts.