Chemistry Study Guide: Electromagnetic Radiation & Atomic Models

Electromagnetic Radiation

Energy that travels in waves through space.

Wavelength (λ)

The distance between successive wave crests.

Frequency (ν)

The number of wave cycles per second (measured in Hertz, Hz).

Quantum

The minimum amount of energy that can be gained or lost by an atom.

Photon

A particle of light energy that carries a quantum of energy.

Atomic Emission Spectrum

The unique pattern of light emitted by an element when its electrons move between energy levels.

Speed of Light Equation

c = 3.00 × 10⁸ m/s (speed of light), λ = wavelength (m), ν = frequency (Hz).

Planck's Equation (Energy of a Photon)

E = hν, where E = energy (Joules), h = Planck's constant (6.626 × 10⁻³⁴ J·s), ν = frequency (Hz).

Radio Waves

Lowest energy, longest wavelength, lowest frequency.

Microwaves

Low energy, long wavelength, low frequency.

Infrared

Medium energy, medium wavelength, medium frequency.

Visible Light

Higher energy, shorter wavelength, higher frequency.

Ultraviolet (UV)

High energy, short wavelength, high frequency.

X-rays

Very high energy, very short wavelength, very high frequency.

Gamma Rays

Highest energy, shortest wavelength, highest frequency.

Bohr's Model of the Atom

Proposed that electrons move in fixed orbits (energy levels) around the nucleus.

Energy levels (n)

Electrons in higher orbits have more energy.

Quantum Mechanical Model

Electrons do NOT travel in fixed orbits but exist in regions called orbitals.

Electron Density Clouds

Orbital = 90% probability of finding an electron.

Heisenberg Uncertainty Principle

States that it is impossible to know both the position and velocity of an electron simultaneously.

Bohr Model Diagram

Depicts electrons in circular orbits around the nucleus.

Quantum Model Diagram

Shows probability clouds instead of defined orbits.

Wave Diagrams

Illustrate wavelength, frequency, and energy relationships.

Heisenberg Uncertainty Principle

Impossible to simultaneously know an electron's exact position and momentum.

Rutherford Model

Proposed that electrons orbit the nucleus.

Bohr's Model

Electrons exist in fixed orbits (energy levels) and can move between levels by absorbing or emitting energy.

Energy absorption in Bohr's Model

Causes electrons to jump to higher energy levels.

Atomic emission spectrum

When electrons return to lower levels, they emit energy as light.

Wavelength (λ)

Distance between wave peaks.

Frequency (ν)

Number of wave cycles per second (measured in Hz).

Speed of light (c)

3.00 × 10⁸ m/s.

Electromagnetic Spectrum

Arranges waves by energy and frequency.

Order of Electromagnetic Spectrum

From lowest to highest energy: Radio waves < Microwaves < Infrared < Visible light < Ultraviolet < X-rays < Gamma rays.

Hydrogen Emission Spectrum

Hydrogen emits light at specific wavelengths, creating a line spectrum.

Limitations of Bohr's Model

Only worked for hydrogen; failed for multi-electron atoms.

Quantum mechanical model

Developed by Erwin Schrödinger; electrons are in orbitals, not fixed paths.

Orbitals

Regions where electrons are likely to be found, described by different shapes (s, p, d, f).

Planck's constant (h)

6.626 × 10⁻³⁴ J·s.

Planck's Equation (Energy of a Photon)

E = hν, where E is energy (Joules) and ν is frequency (Hz).

Significance of quantum mechanics

Essential in modern atomic theory.

Chemical bonding

Affected by the shape of orbitals.

Quiz preparation tips

Be ready to explain Bohr's contributions and how his model improved upon Rutherford's.

Spectral lines

Created by electrons moving between energy levels.

Wave equations

Used to determine energy, wavelength, and frequency.

Electromagnetic spectrum ordering

Recognize how the electromagnetic spectrum is ordered based on energy.