Unit 4 - Electrons in Atoms

Introduction:

Unit 4 concerns electrons in atoms and the behavior of light in terms of the wave-particle duality described by DeBroglie. This chapter will use a limited amount of formulas and math to compare frequency, wavelength and energy. It concludes with several days spent learning electron configuration for both elements and ions.

Student Objectives & Assignments

4.1

a. In the early 20th century, light was determined to have a dual wave-particle nature.

b. Quantum theory was developed to explain observations such as the photoelectric effect and the line-emission spectrum of hydrogen.

c. Quantum theory states that electrons can exist only at specific atomic energy levels.

d. When an electron moves from one main energy level to a main energy level of lower energy, a photon is emitted. The photon’s energy equals the energy difference between the two levels.

e. An electron in an atom can move from one main energy level to a higher main energy level only be absorbing an amount of energy exactly equal to the difference between the two levels.

4.2

a. In the early 20th century, electrons were determined to have a dual wave-particle nature.

b. The Heisenberg uncertainty principle states that it is impossible to determine simultaneously the position and velocity of an electron or any other particle.

c. Quantization of electron energies is a natural outcome of the Schrodinger wave equation, which describes the properties of an atom’s electrons.

d. An orbital, a three-dimensional region around the nucleus, shows the region in space where an electron is most-likely to be found.

e. The four quantum numbers that describe the properties of electrons in atomic orbitals are the principal quantum number, the angular momentum quantum number, the magnetic quantum number and the spin quantum number.

4.3

a. The ground-state electron configuration of an atom can be written by using the Aufbau principle, Hund’s rule and the Pauli exclusion principle.

b. Electron configurations can be depicted by using different types of notation. In this book, three types of notation are used: orbital notation, electron-configuration notation and noble-gas notation. c. Electron configurations of some atoms, such as chromium, deviate from the predictions of the Aufbau principle, but the ground-state configuration that results is the configuration with the minimum possible energy.

Ohio Department of Education Standards:

Atomic Structure:

The physical science syllabus included properties and locations of protons, neutrons and electrons, atomic number, mass number, cations and anions, isotopes and the strong nuclear force that hold the nucleus together. In this course, the historical development of the atom and the positions of electrons are explored in more detail.

Atomic models are constructed to explain experimental evidence and make predictions. The changes in the atomic model over time exemplify how scientific knowledge changes as new evidence emerges and how technological advancements like electricity extend the boundaries of scientific knowledge. Thompson’s study of electrical discharges in cathode-ray tubes led to the discovery of the electron and the development of the plum pudding model of the atom. Rutherford’s experiment, in which he bombarded gold foil with α-particles, led to the discovery that most of the atom consists of empty space with a relatively small, positively charged nucleus. Bohr used data from atomic spectra to propose a planetary model of the atom in which electrons orbit the nucleus, like planets around the sun. Later, Schrödinger used the idea that electrons travel in waves to develop a model in which electrons travel randomly in regions of space called orbitals (quantum mechanical model).

Based on the quantum mechanical model, it is not possible to predict exactly where electrons are located but there is a region of space surrounding the nucleus in which there is a high probability of finding an electron (electron cloud or orbital). Data from atomic spectra (emission and absorption) gives evidence that electrons can only exist at certain discrete energy levels and not at energies between these levels. Atoms are usually in the ground state where the electrons occupy orbitals with the lowest available energy. However, the atom can become excited when the electrons absorb a photon with the precise amount of energy (indicated by the frequency of the photon) to move to an orbital with higher energy. Any photon without this precise amount of energy will be ignored by the electron. The atom exists in the excited state for a very short amount of time. When an electron drops back down to the lower energy level, it emits a photon that has energy equal to the energy difference between the levels. The amount of energy is indicated by the frequency of the light that is given off and can be measured. Each element has a unique emission and absorption spectrum due to its unique electron configuration and specific electron energy jumps that are possible for that element. Being aware of the quantum mechanical model as the currently accepted model for the atom is important for science literacy as it explains and predicts subatomic interactions, but details should be reserved for more advanced study.

Electron energy levels consist of sublevels (s, p, d and f), each with a characteristic number and shape of orbitals. The shapes of d and f orbitals will not be assessed in high school. Orbital diagrams and electron configurations can be constructed to show the location of the electrons in an atom using established rules. However, the names of these rules will not be assessed. Valence electrons are responsible for most of the chemical properties of elements. In this course, electron configurations (extended and noble gas notation) and orbital diagrams can be shown for any element in the first three periods.

Although the quantum mechanical model of the atom explains the most experimental evidence, other models can still be helpful. Thinking of atoms as indivisible spheres is useful in explaining many physical properties of substances, such as the state (solid, liquid or gas) of a substance at room temperature. Bohr’s planetary model is useful to explain and predict periodic trends in the properties of elements.

Periodic Table:

In the physical science syllabus, elements are placed in order of increasing atomic number in the periodic table such that elements with similar properties are placed in the same column. How the periodic table is divided into groups, families, periods, metals, nonmetals and metalloids was also in the physical science syllabus. In chemistry, with more information about the electron configuration of elements, similarities in the configuration of the valence electrons for a particular group can be observed. The electron configuration of an atom can be written from the position on the periodic table. The repeating pattern in the electron configurations for elements on the periodic table explain many of the trends in properties across periods or down columns including atomic radii, ionic radii, first ionization energies, electronegativities and whether the element is a solid or gas at room temperature. Additional ionization energies, electron affinities and periodic properties of the transition elements, lanthanide and actinide series is reserved for more advanced study.

Note: Quantum numbers and equations of de Broglie, Schrödinger and Plank are beyond the scope of this course.

Helpful Video Links:

(NOTE: I don't consider you dummies, this video was just as explanatory for myself as you. Quantum mechanics is an extremely abstract concept and I'd be shocked if you understood it from class or the book without an explanation).