Preview
IB DP Physics — Unit E: Nuclear and Quantum Physics | SL & HL | E.1 Structure of the Atom
📋 What You Will Learn
- Observe and describe the phenomenon of atomic spectra using real-world examples such as neon signs, sodium street lamps, and starlight
- Explain why atoms emit or absorb only specific, discrete wavelengths of light — and why this is fundamentally different from classical predictions
- Define energy levels, ground state, excited state, ionisation energy, and all related notation before applying them in equations
- Apply the photon energy equation \( E = hf \) and the energy level transition equation \( \Delta E = hf \) to calculate emitted and absorbed photon frequencies and wavelengths
- Distinguish between continuous spectra, emission line spectra, and absorption line spectra, and interpret spectral diagrams
- Analyse energy level diagrams and connect spectral lines to specific electron transitions
- Evaluate the historical significance of atomic spectra in the development of the quantum model of the atom
🌍 Real-World Context Look up at a sodium street lamp on a dark night and you will notice something unusual: the light it casts is not white, but a distinctive deep yellow-orange. Inside a neon sign, the glowing gas produces a vivid red-orange glow. A hydrogen discharge tube in a laboratory emits a faint pinkish-violet light. When this light is passed through a prism or diffraction grating, something remarkable happens — instead of a smooth rainbow of colours, you see only a handful of sharp, bright lines at very specific colours, separated by complete darkness. Each element produces its own unique pattern of lines, as distinctive as a fingerprint. This observation — that atoms emit light only at specific wavelengths — was one of the great mysteries of nineteenth-century physics, and its explanation would ultimately require an entirely new theory of matter.
🔍 Guiding Question Why do atoms emit and absorb light only at specific, discrete wavelengths, and what does this tell us about the internal structure of the atom?
Physical Setup: Observing Atomic Spectra
To study the light emitted by atoms, physicists use a simple but powerful apparatus. A gas discharge tube — a sealed glass tube containing a low-pressure gas — is connected to a high-voltage source. The electric current excites the atoms inside the tube, causing them to emit light. This light is then directed through a narrow slit and passed through a diffraction grating or prism, which separates the light by wavelength. The separated light is projected onto a screen or captured by a spectrometer.
What appears on the screen is not a continuous rainbow. Instead, only a few narrow, bright lines appear at specific positions — and each position corresponds to a specific wavelength of light. This pattern is called an emission line spectrum. The positions of the lines are unique to the element in the tube: hydrogen always produces the same set of lines, helium always produces its own different set, and so on.
Notice in the diagram that only four lines appear for hydrogen in the visible range — not a continuous band of colour. The dark gaps between the lines are regions where hydrogen emits no light at all. This is the defining feature of a line spectrum, and it is in stark contrast to the continuous spectrum produced by a hot solid or liquid, which shows all wavelengths blending smoothly into one another.
Three Types of Spectra
Before we can explain why line spectra exist, we need to clearly distinguish between the three types of spectra that physicists observe. Each arises from a different physical situation, and each carries different information about matter and light.
📖 Key Definition — Continuous Spectrum A continuous spectrum contains all wavelengths of electromagnetic radiation within a given