S1.3.1 Emission spectra and energy levels

Photon Emission, Electromagnetic Spectrum, and Spectral Types

1. You're watching a fireworks display.
2. The vibrant reds, greens, and blues streaking across the night sky are not just beautiful—they are the result of electrons transitioning between energy levels in atoms, releasing light as photons.
3. But what exactly happens at the atomic level to create this spectacle?

To answer this, let’s dive into photon emission, the electromagnetic spectrum, and the differences between continuous and line spectra.

Photon Emission: The Source of Light

1. When atoms absorb energy—whether from heat, electricity, or other sources—their electrons can jump to higher energy levels, or "excited states."
2. However, these excited states are unstable, and the electrons soon return to lower energy levels, releasing the absorbed energy as photons of light.

The Mechanics of Photon Emission

1. Each photon emitted corresponds to a specific energy difference between two energy levels in the atom.
2. This energy is tied to the frequency of the emitted light through the equation:

$$E=h\cdot f$$

Where:

• $E$ is the energy of the photon (in joules, $J$),
• $h$ is Planck’s constant ($6.63\times {10}^{-34}\text{Js}$),
• $f$ is the frequency of the light (in hertz, $Hz$, or $s^{-1}=\frac{1}{\text{second}}$).

This process of photon emission is fundamental to understanding atomic spectra, which provide insights into the structure of atoms.

The Electromagnetic Spectrum: Wavelength, Frequency, and Energy

1. The light we see is just a small part of the electromagnetic spectrum, which includes all forms of electromagnetic radiation, from radio waves to gamma rays.
2. Each type of radiation is defined by its wavelength $\lambda$, frequency $f$, and energy $E$.

Relationships Between Wavelength, Frequency, and Energy

The speed of light $c$ is constant in a vacuum and relates wavelength and frequency as follows:

$$c=f\cdot \lambda$$

Where:

• $c=3.00\times {10}^8\text{m/s}$,
• $f$ is the frequency (Hz or $s^{-1}$),
• $\lambda$ is the wavelength $m$.

Energy is inversely proportional to wavelength and directly proportional to frequency:

$$E\propto \frac{1}{\lambda }\text{and}E\propto f$$

Visible Light: A Slice of the Spectrum

1. The visible spectrum ranges from approximately 400 nm (violet, high energy) to 700 nm (red, low energy).
2. Beyond this range lie ultraviolet radiation (shorter wavelengths, higher energy) and infrared radiation (longer wavelengths, lower energy).

Continuous vs. Line Spectra: The Fingerprints of Atoms

1. Not all light sources produce the same kind of spectrum.
2. By analyzing the light emitted or absorbed by a substance, we can uncover details about its atomic structure.

Continuous Spectrum

• Incandescent objects, such as the Sun or a filament bulb, produce this type of spectrum.
• When passed through a prism, it appears as a rainbow.

Emission and Absorption Line Spectrums

1. These lines are produced when electrons transition between quantized energy levels in an atom.
2. Each element has a unique line spectrum, acting like a "barcode" for identifying the element.

Emission Spectrum

Absorption Spectrum

• These dark lines correspond to wavelengths of light absorbed by an atom when electrons are excited to higher energy levels.

Reflection and Practice