7: Electronic Structure and Periodic Properties
The study of chemistry must at some point extend to the molecular level, for the physical and chemical properties of a substance are ultimately explained in terms of the structure and bonding of molecules. This module introduces some basic facts and principles that are needed for a discussion of organic molecules.
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- 7.1: Light and Atomic Structure
- Light and other forms of electromagnetic radiation move through a vacuum with a constant speed, c. This radiation shows wavelike behavior, which can be characterized by a frequency, ν, and a wavelength, λ, such that c = λν. Light is an example of a travelling wave. Other important wave phenomena include standing waves, periodic oscillations, and vibrations. Standing waves exhibit quantization, since their wavelengths are limited to discrete integer multiples of some characteristic lengths.
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- 7.2: Development of Quantum Theory
- Macroscopic objects act as particles. Microscopic objects (such as electrons) have properties of both a particle and a wave. but their exact trajectories cannot be determined. The quantum mechanical model of atoms describes the 3D position of the electron in a probabilistic manner according to a mathematical function called a wavefunction, often denoted as ψ. The squared magnitude of the wavefunction describes the distribution of the probability of finding the electron in a particular region in
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- 7.3: Electronic Structure of Atoms (Electron Configurations)
- The relative energy of the subshells determine the order in which atomic orbitals are filled. Electron configurations and orbital diagrams can be determined by applying the Pauli exclusion principle (no two electrons can have the same set of four quantum numbers) and Hund’s rule (whenever possible, electrons retain unpaired spins in degenerate orbitals). Electrons in the outermost orbitals, called valence electrons, are responsible for most of the chemical behavior of elements.
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- 7.4: Periodic Variations in Element Properties
- Electron configurations allow us to understand many periodic trends. Covalent radius increases as we move down a group because the n level (orbital size) increases. Covalent radius mostly decreases as we move left to right across a period because the effective nuclear charge experienced by the electrons increases, and the electrons are pulled in tighter to the nucleus. Anionic radii are larger than the parent atom, while cationic radii are smaller.
Contributors and Attributions
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Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110 ).