U2 L1 Video Lesson

Early Atomic Theories

Scientific models of the atom have changed as new experimental evidence became available. Each model explained some observations, but later evidence revealed its limitations.

Learning Goals

Students should be able to:

• Explain how atomic models developed.
• Describe experiments that provided evidence about atomic structure.
• Explain the basic quantum-mechanical model of the atom.

Development of Atomic Theory

• Early Greek ideas: Some Greek philosophers viewed the universe as one continuous, unchanging whole. Democritus proposed instead that matter consists of tiny particles he called atomos—an early idea resembling atoms.
• Dalton’s atomic theory: Dalton proposed that matter is made of atoms; atoms of each element have characteristic properties; atoms combine in simple whole-number ratios to form compounds; and chemical reactions rearrange atoms. Some details—especially the idea that atoms cannot be divided—were later revised after the discovery of subatomic particles.
• Thomson and the electron: Experiments with low-pressure gas-discharge tubes produced cathode rays. In 1897, J. J. Thomson showed that these rays consist of negatively charged particles, later called electrons. Their charge-to-mass ratio was consistent across different gases and electrode materials, indicating that electrons are components of atoms.
• Thomson’s plum-pudding model: Thomson pictured electrons embedded in a diffuse sphere of positive charge, like fruit in a pudding. Rutherford’s experiment later challenged this model.
• Rutherford’s nuclear model: In the gold-foil experiment, alpha particles from a radioactive source were directed at very thin gold foil, and their paths were observed on a detector.
  • Most particles passed through, suggesting that atoms are mostly empty space.
  • A small number were sharply deflected or bounced back, suggesting that most of an atom’s positive charge and mass are concentrated in a tiny, dense central nucleus.
  • Rutherford’s model placed electrons around the nucleus.
• Bohr’s model: Studies of light emitted by energized gases suggested that electrons can occupy only certain allowed energy levels. Bohr proposed that electrons move between these levels by absorbing or releasing specific amounts of energy. His model explained hydrogen’s spectrum well but did not adequately explain more complex atoms.

Atomic Spectra and Electron Energy

• Electrons normally occupy their lowest available energy state, called the ground state.
• When an atom absorbs energy, an electron can move to a higher-energy excited state.
• As the electron falls to a lower level, it releases energy as a photon of light. The photon’s energy—and therefore its wavelength and color—depends on the energy difference between the levels.
• Each element has its own allowed energy levels and produces a characteristic emission spectrum. Spectra can therefore help identify elements.
• An emission spectrum has bright lines produced by energized, low-pressure gas. An absorption spectrum has dark lines where a cooler gas absorbs particular wavelengths from a continuous source. A continuous spectrum contains an uninterrupted range of wavelengths.
• Spectrometers use components such as a slit and prism to separate light into wavelengths for observation or measurement.
• The Bohr and Rydberg equations relate energy levels and electron transitions to photon energy and wavelength.
• Clarification: For hydrogen’s visible emission lines, electrons fall to the n = 2 level; hydrogen’s ground state is n = 1. Violet light has higher energy and frequency than red light, while red light has a longer wavelength.

Quantum-Mechanical Model

Bohr’s fixed circular orbits were eventually replaced by a quantum-mechanical description:

• Electrons are treated as having wave-like behavior and quantized energies.
• Rather than specifying an exact path for an electron, the model describes the probability of finding it in different regions around the nucleus.
• These regions are called orbitals or electron clouds. They vary in size, shape, and orientation; they are not simply circular shells at progressively greater distances.
• The Heisenberg uncertainty principle means that an electron’s exact position and speed cannot both be known simultaneously. The model therefore predicts likely locations rather than a precise electron trajectory.
• Bohr’s model remains useful for introducing energy levels and explaining hydrogen’s spectrum, while the quantum model better accounts for more complex atoms and their spectra.

Speakers and Sources Featured

• Speaker: One lesson narrator/teacher.
• Scientists and historical sources discussed: Democritus, John Dalton, J. J. Thomson, Ernest Rutherford, Niels Bohr, Albert A. Michelson, Arnold Sommerfeld, Louis de Broglie, Clinton Davisson, Erwin Schrödinger, and Werner Heisenberg.
• Experiments and evidence discussed: Cathode-ray tube experiments, Rutherford’s gold-foil experiment, observations of gas-emission spectra, and spectroscopy.