Atomic emission spectra

IB Chemistry SL· Structure 2: Electron Configuration· 15 min read

1. Continuous vs Line Emission Spectra★★☆☆☆⏱ 4 min

When light from different sources is passed through a prism, it produces two distinct types of spectra. Hot solid or liquid sources produce a continuous range of wavelengths with no gaps. Light emitted from excited gaseous atoms produces a spectrum with only discrete, separated lines.

📘 Definition

Line emission spectrum

A set of discrete coloured lines on a black background, produced when light from excited gaseous atoms is split by a prism

Example:

Hydrogen gas excited by high voltage produces a line spectrum with four visible lines

📐 Worked Example

Distinguish between the emission spectrum from a hot incandescent light bulb filament and excited gaseous neon.

1
The hot filament is a solid, so it produces a continuous spectrum with all visible wavelengths and no gaps.
2
Excited gaseous neon contains isolated neon atoms with electrons transitioning between quantized energy levels, so it produces a discrete line emission spectrum with only specific visible wavelengths.

2. Origin of Emission Lines★★☆☆☆⏱ 5 min

Atomic emission spectra provide direct experimental evidence that electron energy levels in atoms are quantized, meaning electrons can only occupy specific fixed energy levels. When an atom absorbs energy, an electron jumps from a lower energy level to a higher energy level (creating an excited state). When the electron drops back to a lower energy level, it emits a photon with energy exactly equal to the difference between the two levels:

\Delta E = E_{higher} - E_{lower} = h\nu = \frac{hc}{\lambda}

📘 Definition

Quantized energy

The property that electrons can only exist at specific, discrete energy levels, not any energy between these levels

📐 Worked Example

An electron in hydrogen drops from n=4 to n=2, with an energy difference of $4.09 \times 1 0^{- 19}$ J. Calculate the wavelength of the emitted photon ( $h = 6.626 \times 1 0^{- 34}$ Js, $c = 3.00 \times 1 0^{8}$ ms⁻¹).

1
Rearrange the energy-wavelength relationship to solve for $λ$ :
2
$\Delta E = \frac{hc}{\lambda} \rightarrow \lambda = \frac{hc}{\Delta E}$
3
Substitute the given values and calculate:
4
$\lambda = \frac{(6.626 \times 10^{-34} \text{ Js})(3.00 \times 10^8 \text{ ms}^{-1})}{4.09 \times 10^{-19} \text{ J}} = 4.86 \times 10^{-7} \text{ m} = 486 \text{ nm}$
5
486 nm falls in the visible region of the electromagnetic spectrum.

3. Hydrogen Emission Series★★★☆☆⏱ 5 min

Hydrogen has the simplest emission spectrum, with groups of lines called series that correspond to electron transitions that end on the same lower energy level ( $n_{f ina l}$ ). Each series falls in a different region of the electromagnetic spectrum, as shown in the table below:

Series Name	Final n Level	Spectral Region
Lyman	1	Ultraviolet
Balmer	2	Visible
Paschen	3	Infrared

📐 Worked Example

A hydrogen electron transition emits a photon with wavelength 122 nm, in the ultraviolet region. What is the final n level of the transition?

1
Ultraviolet emission from hydrogen corresponds to the Lyman series of transitions.
2
All transitions in the Lyman series end on the lowest energy level, n=1.
3
Answer: Final n level = 1

4. Emission Spectra as Evidence★★★☆☆⏱ 4 min

Before the development of the Bohr model, scientists predicted electrons could have any energy, so excited atoms would produce a continuous emission spectrum. The observation of discrete line spectra directly contradicted this, and provided the first clear evidence that electron energy is quantized.

5. Common Pitfalls

Wrong move:

Claims emission lines are produced when electrons jump from lower to higher energy levels.

Why:

Jumping up requires absorbing energy, so this produces absorption lines, not emission lines.

Correct move:

Emission lines are produced when excited electrons drop from higher to lower energy levels, emitting a photon equal to the energy difference.

Wrong move:

Forgets to convert wavelength from nanometers to meters when substituting into energy calculations.

Why:

Planck's constant and the speed of light use SI units (meters for distance), so mixing units gives an incorrect order of magnitude.

Correct move:

Always convert nanometers to meters by dividing by $1 0^{9}$ before calculating energy.

Wrong move:

Confuses the final n level for hydrogen emission series, e.g. says Balmer series ends at n=1.

Why:

Series are grouped by the lower (final) energy level after the transition, not the starting level.

Correct move:

Use the mnemonic: Lyman n=1 (UV), Balmer n=2 (Visible), Paschen n=3 (IR).

Wrong move:

Gets a negative wavelength from a negative $Δ E$ for emission.

Why:

$Δ E = E_{f ina l} - E_{ini t ia l}$ is negative for emission because energy is lost from the atom, but wavelength is always positive.

Correct move:

Use the magnitude (absolute value) of the energy difference when calculating photon wavelength.

Wrong move:

Claims all elements have the same emission spectrum because all electrons are identical.

Why:

Each element has a unique nuclear charge and unique electron arrangement, so energy differences between levels are unique.

Correct move:

Each element has a unique line emission spectrum that can be used to identify it.

6. Quick Reference Cheatsheet

Concept	Key Fact/Formula
Origin of emission lines	Electron drops from high to low energy, emits photon: $Δ E = h c / λ$
Continuous vs line	Continuous = all wavelengths, no gaps; Line = discrete, specific wavelengths
Hydrogen Balmer series	Transitions end at n=2, produces visible lines
Key evidence	Line spectra prove electron energy is quantized
Unit conversion	1 nm = $1 \times 1 0^{- 9}$ m

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2025 · 1
Identify hydrogen emission series by region
2024 · 2
Calculate energy of electron transition
2023 · 1
Explain origin of line emission spectra

Going deeper

What's Next

Atomic emission spectra are the experimental foundation for our understanding of quantized electron energy levels, which is the core of the modern model of the atom. Mastery of this concept is required to understand how electron configuration is built, and how different atomic properties like ionization energy relate to electron arrangement. The ideas you learned here about quantized energy also underpin spectroscopic techniques you will encounter in analytical chemistry topics later in the course. Next, you will build on this knowledge to explore how electrons are arranged in orbitals and how we write full electron configurations for multi-electron atoms.