Chemistry

Absorption and Emission

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The Electromagnetic Spectrum

The EM spectrum in order of decreasing wavelength (increasing frequency and energy):

Radio waves → Microwaves → Infrared → Visible light → Ultraviolet → X-rays → Gamma rays

Light travels at $c = 3 \times 10^8 \text{ m/s}$ in a vacuum. This speed is always constant.

Wave-Particle Duality

Light exhibits both wave and particle properties — this is called wave-particle duality.

• As a wave: it has a wavelength $\lambda$ and a frequency $\nu$, related by: $\lambda = \frac{c}{\nu}$
• As a particle: it is a stream of photons, which are discrete packets (quanta) of energy: $E = h\nu$ where $h = 6.63 \times 10^{-34}$ J s is Planck's constant.

Combining these two equations gives: $E = \frac{hc}{\lambda}$

Quantisation

Quantisation is the process of splitting something into its smallest indivisible unit, called a quantum (plural: quanta). Planck showed that light can be quantised — it is made up of a smallest unit of energy that cannot be divided. The photon is the quantum of light (EM radiation). The electron is the quantum of charge.

Energy Levels and Spectra

Energy levels are discrete

Energy levels in an atom are discrete (quantised) — electrons can only exist at certain specific energy values. This means only certain frequencies of light can be absorbed or emitted. Every element has a unique set of energy levels, so every element has a unique absorption and emission spectrum. Always mention that energy levels are quantised when describing spectral evidence.

Absorption spectra

When light passes through a sample of an element, photons of specific wavelengths are absorbed. This excites electrons, causing them to jump up to a higher energy level. The wavelengths absorbed appear as black lines on an otherwise continuous coloured spectrum.

Emission spectra

An excited electron is unstable in its higher energy level. After some time it drops back down to a lower energy level, releasing energy as a photon of light. The wavelengths emitted appear as coloured lines on a black background.

There is a corresponding emission spectrum for every absorption spectrum — the dark lines in absorption match the coloured lines in emission for the same element.

When a particle absorbs or emits light, the energy absorbed or released equals the energy of the photon: $E = h\nu$.

Flame Colours

When elements are heated, electrons absorb energy and jump to higher energy levels. As they fall back down, they release specific wavelengths of visible light, producing characteristic colours:

Different elements have different electron arrangements, so electrons jump between different energy gaps, emitting different wavelengths of light. This is why each element gives a distinct colour. The colour is related to specific visible wavelengths, governed by $E = h\nu$.

The Lyman Series and Other Series

The Lyman series describes transitions where electrons fall down to the ground state ($n = 1$). Each line in the series corresponds to an electron dropping from a higher shell ($n = 2, 3, 4, \ldots$) back to $n = 1$.

Other series exist for electrons dropping to $n = 2$ (Balmer), $n = 3$ (Paschen), etc.

• The absorption series shows the same transitions but with arrows pointing upward (absorption rather than emission).
• The energy gap between shells decreases as shell number increases, so the lines in the spectrum get closer together at higher $n$.

Ionisation

Ionisation is where atoms lose or gain electrons; this always occurs in the gaseous state:

For ionisation from a +1 ion: $\ce{Sr^{+}(g) -> Sr^{2+}(g) + e^{-}}$

Ionisation enthalpies decrease down a group because more electron shells means a greater distance between the outer electron and the nucleus, plus more electron shielding — so the force of attraction is weaker and less energy is required to remove the electron.