The first four chapters of our textbook described chemical reactions and we were able to observe the chemical and physical properties of these reactions. We could also predict the amount of reactants consumed and products formed in each of the reactions we studied.

In Chapters 1-4 we didn’t really explain answer any “why?” questions, we simply showed the results and then related them to the properties of matter. In order to answer some of the “why?” questions in chemistry we need to use quantum mechanics to describe the arrangements of electrons in atoms, which chemists and physicists refer to as the electronic structure of atoms.

**6.1 Electronic Structure of Atoms Overview**

In order to understand the electronic structure of atoms we must understand the wave-like nature of light. We encounter this property of light in our everyday lives when discussing the various regions of the electromagnetic spectrum.

**6.1 Electromagnetic Raiation**

Now that we have a relationship established between energy, frequency, and wavelength we can inter-convert between each of these values.

**6.1 Energy, Frequency, Wavelength Example Problem**

Depending on the experimental circumstances, radiation can have either wave-like or particle-like properties. This was very puzzling to scientists in the early 1900s as they could not come up with a theory describing radiation under all circumstances. Eventually, the wave-particle duality was established. It was also proved that matter can emit or absorb energy in specific amounts, or we can say that energy is quantized.

**6.2 Wave-Particle Duality**

Many theories were developed describing the wave-particle duality of radiation, and Balmer was able to fit the lines of the hydrogen spectrum mathematically, but no one was able to theoretically explain the line spectra of a hydrogen atom. That was until Niels Bohr proposed his model of the hydrogen atom.

**6.3 Line Spectra and the Bohr Model Part 1**

**6.3 Line Spectra and the Bohr Model Part 2**

Louis de Broglie and Werner Heisenberg continued to study the wave-like nature of electrons and their experiments opened the door for a new branch of science which we currently refer to as Quantum Mechanics.

**6.4 The Wave Behavior of Matter**

In 1926 Erwin Schrodinger developed a mathematical equation that represents the energy of the electrons in an atom. These energies are the same in the Bohr model, but it improves the Bohr model in the fact that it accounts for the electrons orbiting the nucleus in atomic orbitals, rather than circular orbits. The Schrodinger is the foundation of Quantum Mechanics.

**6.5 Quantum Mechanics**

Solving the Schrodinger equation requires advanced calculus, but the results of these calculations give us some practical results. We are able to pictorially represent orbitals from this equation and in this course you are not required to solve the Schrodinger equation mathematically, but you will need to know how to identify the atomic orbitals.

**6.5 Atomic Orbitals**

Once the atomic orbitals are known, we need to determine how the electrons in an atom will occupy these orbitals. Electrons will always occupy orbitals in the lowest energy configuration possible, and we need to determine which orbitals have the lowest energy.

**6.5 Energies of Atomic Orbitals**

Now that we have a grasp on the relative energies of the atomic orbitals we can start populating them with electrons. The two methods we use to indicate how electrons fill in the atomic orbitals are electron configurations and orbital block diagrams.

**6.8 Electron Configurations and Block Diagrams Overview**

**6.8 Electron Configuration and Orbital Block Diagram for Elements 1-18**

**6.8 Electron Configuration and Orbital Block Diagram for Elements 19-30**

In addition to determining the electron configurations for atoms, we can also determine the electron configurations for ions.

**6.8 Electron Configuration of Ions**

Now that we have written out the electron configurations for atoms and ions, we need to distinguish each electron from each other. Quantum numbers are used to distinguish each electron from another, much like a ticket to a sporting event or concert distinguishes one seat from another.

**6.5 Quantum Numbers**

**6.5 Assigning Quantum Numbers to N**

**6.5 Assigning Quantum Numbers to Cr**

I’m not in your class, but these videos have been very helpful. thanks a lot!