CHEM 1405 Concept Review: Atomic Theory & Electronic Structure

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Experiments Leading Up to the Modern Atomic Theory

The Cathode Ray Experiments: Joseph John Thompson (J.J. Thompson) carried out a series of experiments with cathode rays. These rays were attracted by positively charged plates and repelled by negatively charged plates. The rays were also formed regardless of what material was used to make the electrode source. Thompson concluded that the “cathode rays” were actually tiny, negatively charged particles which were part of all atoms and named them electrons. He also determined the charge to mass ratio of these electrons.

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The Gold Foil Experiment: Ernest Rutherford and his assistants carried out an experiment in which thin gold foil was bombarded with alpha radiation. Rutherford expected the particles to pass right through, as it was believed that the positive charge in an atom was sort of amorphously spread over the whole atom. Instead, though most particles passed straight through, a number of particles were deflected at sharp angles. This led to the conclusion that an atom was composed of mostly empty space and that most of the mass and all the positive charge of an atom was contained in a very small dense core called the nucleus.

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Milliken’s Oil Drop Experiment: Robert Milliken devised an experiment where tiny droplets of oil were allowed to pass one at a time between two charged plates while being bombarded by X-rays. The ionizing radiation caused the droplets to take on charges that were whole number ratios of each other. The smallest increment of charge was found to be the charge of a single electron.

The Beryllium-Alpha Bombardment Experiment: James Chadwick used this experiment to discover the presence of uncharged particles in the nucleus which had a similar mass to the positively charged particles (protons) that had already been discovered by Rutherford and Eugen Goldstein. These newly discovered neutral particles were called neutrons.

Modern Atomic Theory

The current atomic theory states that an atom is composed of three fundamental particles; protons, neutrons, and electrons. The protons and neutrons combine to form a small but very dense mass in the center of the atom known as the nucleus. The electrons spin around the nucleus at blindingly fast speeds (close to the speed of light). Since the electrons are moving so fast, we speak of them as having a probability “cloud” where they are likely to be.

Electron Cloud Model.png

Subatomic Particles: The Components of an Atom

*Particle*	*Symbol*	*Charge*	*Mass (amu)*	*Location in Atom*
Proton	p or p⁺	+1	1.0073	in the nucleus
Neutron	n or n⁰	0	1.0087	in the nucleus
Electron	e^-	-1	5.486*10^-4	orbiting around the nucleus

Niels Bohr and Electron Arrangement

Niels Bohr used line spectra to develop a new theory about how electrons were arranged in an atom. He reasoned that only distinct lines were observed in line spectra because electrons were only capable of absorbing and releasing energy of specific amounts. Basically, the energy of electrons was quantized, meaning that electrons could only have certain specific amounts of energy. Bohr imagined these different amounts of energies as different energy levels and created the following model of the atom:

Bohr Model.png

Bohr was able to further deduce that the maximum number of electrons a particular energy level (n) could hold could be determined by the following formula:

$LaTeX: Maximum\:Number\:of\:Electrons\:\left(in\:Energy\:Level\right)=2n^2$ $Maximum\:Number\:of\:Electrons\:\left(in\:Energy\:Level\right)=2n^2$

Rules and Principles for Electron Behavior

Aufbau Principle: Electrons will always occupy lower energy orbitals before higher energy orbitals.

Hund’s Rule: Electrons will fill orbitals of equal energy (degenerate orbitals) with one electron with the same spin (parallel spin) in each orbital before any of those orbitals will gain a second electron.

Pauli Exclusion Principle: No two electrons can have the same set of 4 quantum numbers. (Basically, this means no two electrons can be in the same orbital with the same spin.)

Writing Electron Configurations

The periodic table can be used as a roadmap for electron configuration. Each “box” on the periodic table represents a place for one electron. The different subshells are represented by different “blocks” in the periodic table; the s-block, the p-block, the d-block, and the f-block. These blocks are color coded in the diagram below.

For a complete electron configuration:

Start at the beginning (1s) and fill the orbitals in order, going left to right until you reach the end of a period and then going to the next row. Indicate each the energy level, the subshell being filled, and how many electrons are filling it as you pass through each subshell. Represent the symbols in the following manner:

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Continue filling subshells until you reach the element for which you are determining the electron configuration. Remember that the f-orbitals are actually in between the 6s & 5d and the 7s & 6d orbitals. The subshell filling guide below on the right can help you check to make sure that you have filled the orbitals in the correct order. Also note that subshells drop 1 energy level when you pass into the d-block and 2 energy levels when you pass into the f-block.

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Example Configurations: O: 1s² 2s² 2p⁴

Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Orbital Diagrams

Orbital Diagram: a pictorial representation of electron configuration that uses boxes to represent the orbitals of a subshell and arrows to represent electrons occupying a subshell.

The different subshells in an atom have different numbers of orbitals and different shapes according to the type of subshell. The following table summarizes this information: