# High School Chemistry/Electrons and Light

In the early 1900s, scientists observed that different elements made flames of different colors. This small observation shows that an element's behavior is related to the arrangement of electrons!

## Part 1: Introduction

Electromagnetic Radiation travels in repetitive wave motions. Wavelength (λ) is measured in meters or nanometers = distance from two consecutive peaks or troughs, while frequency (v) is measured in hertz or hz = numbers of waves that pass a given point in one second (waves per second).

Visible light is a form of electromagnetic radiation (a form of energy that moves in waves), other types of EM radiation are radio waves, x-rays, microwaves. All forms of EM radiation together form the electromagnetic spectrum. All forms of EM radiation travel at the speed of light (C): 3 x 108 m/s.

Relationships between Wavelength and Frequency
Wavelength Frequency
↑λ ↓v
↓λ ↑v

The visible spectrum: Sunlight and white light consists of light with a continuous range of wavelength & frequency (400nm - 700nm). Each color of light has a unique wavelength and frequency. The part of the spectrum that contains harmful wavelengths/frequencies are the ones with short wavelengths and high frequencies.

## Part 2: Mathematics

The mathematical relationship between wavelength and frequency:
C=λv C = speed of light (3 x 108 m/s) The mathematical relationship between energy and frequency:
E=hv E = energy (jowls)
Planck's constant = 6.626 x 10-34 j/s

Max Planck proposed the theory that energy is not continuous, but quantized. EMR travels in waves and is made of energy-carrying particles known as photons, each photon carries a quantum of energy. Matter (atoms) can again or lose energy only in small, specific amounts called quanta. These quanta of energy are gained or lost when electrons move from one energy level to another energy level. Note that the number of energy levels an electron falls correspond to the color of light produced.

### Bohr Model/Energy levels

Bohr Model
• Electrons exist in energy levels/orbits.
• Energy levels closest to the nucleus are lowest in energy.
• Electrons can't be found in the space between levels.
Energy levels (evidence)

When atoms are heated/electrified, they absorb energy, causing electrons to "jump" to outer energy levels (the excited state). As the electrons lose this energy, they fall back to their lower energy level (ground state), excess energy is emitted as some form of EM radiation.

If energy levels did not exist, an infinite energy difference would be available to the naked eye and all colors of light, including white light, would always be seen. Since only a few bands of color are produced (when electrons are in the excited state), then only a few energy levels must be possible.

### Other Imp. Electron Knowledge

Albert Einstein found that light is made of photons, which are small particles/pieces ("quanta"). A photon's energy depends on its wavelength/frequency (remember their relationship)... so, a minimum frequency of light is needed for the ejection of photons from metals (for them to look shiny; Photoelectric Effect).

Heisenberg's Uncertainty Principle: It is impossible to know both the velocity and position of an electron; merely the act of observations changes what you are observing.

Prisms, Spectroscopes, and 3-D Glasses

Different wavelengths of light are refracted, or bent, when passing through a solid medium and separated into different colors. White light gives a continuous spectrum.

### The Quantum Mechanical Model of the Atom

From Schrodinger:

1. the "fly in the room" theory
2. mathematically-based model; this describes the probability of finding an electron in a certain area of the atom
3. region where electrons might be found are called orbitals
4. any one orbital can hold only 2 electrons
5. Orbitals come in several shapes and sizes - many of each size and shape found in each atom

Electron configuration is a map listing the orbital where each electron in an atom is found.

### Orbital Diagrams

1. Aufbau - start at the lowest energy level and go up
2. Hund's rule - when electrons occupy equal energy orbitals, one electron enters each orbital before any orbital gets a second electron (1 electron in each orbital before doubling up).
3. Pauli exclusion principle - one orbital can have at most 2 electrons; when 2 electrons the same orbital, they must have opposite spins

In terms of

Mo and Cr
• 4s1 and 3d5
• 5s1 and 4d5
Cu and Ag
• 4s1 and 3d10
• 5s1 and 4d10