Atomic Structure and Electromagnetic Radiation: Foundations of Modern Chemistry

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Atomic Structure and Electromagnetic Radiation

Overview

This section introduces the fundamental experiments and concepts that led to our modern understanding of atomic structure and the nature of electromagnetic radiation. Key experiments by J.J. Thomson and Rutherford, as well as the properties of light and atomic spectra, are discussed.

J.J. Thomson’s Cathode Ray Tube (CRT) Experiment

Discovery of the Electron

Cathode rays are streams of negatively charged particles (later called electrons) emitted from the cathode in a vacuum tube.
Thomson showed that these rays were deflected by electric and magnetic fields, indicating they carried a negative charge.
He measured the charge-to-mass ratio of the electron, showing it was much smaller than that of a hydrogen ion.

Conclusion: Atoms contain tiny, negatively charged particles (electrons), suggesting atoms are divisible and have internal structure.

Thomson’s Atomic Model: Proposed the "plum pudding" model, where electrons are embedded in a positively charged sphere.

Millikan’s Oil Drop Experiment

Measurement of the Electron’s Charge

Millikan measured the charge of individual oil droplets suspended in an electric field.
He determined the fundamental unit of electric charge: .
Combined with Thomson’s charge-to-mass ratio, the mass of the electron was calculated: .

Rutherford’s Alpha Particle Scattering Experiment

Discovery of the Nucleus

Alpha particles were directed at a thin gold foil.
Most particles passed through, but some were deflected at large angles.
Conclusion: Atoms have a small, dense, positively charged nucleus containing most of the mass, with electrons surrounding it.

Comparison of Subatomic Particles:

Particle	Relative Mass	Relative Charge
Electron	1	-1
Proton	1836	+1
Neutron	1839	0 (neutral)

Properties of a Wave

Wavelength and Frequency

Wavelength (λ): The distance between two consecutive crests or troughs in a wave (measured in meters).
Frequency (ν): The number of wave cycles passing a point per second (measured in Hz or s-1).
Relationship: where is the speed of light ( m/s).

Example: Calculate the frequency of blue light with a wavelength of 440 nm:

Electromagnetic Radiation (Light)

Electromagnetic Spectrum

Includes all types of electromagnetic waves, from gamma rays to radio waves.
Visible light is a small portion of the spectrum (approx. 400–700 nm).
Amplitude: The height of the wave, related to the intensity of the light.

Line Spectrum and Atomic Spectra

Continuous vs. Line Spectrum

Continuous spectrum: Contains all wavelengths in a given range (e.g., sunlight).
Line spectrum: Contains only specific wavelengths, characteristic of the element (atomic spectra).

Atomic Emission Spectrum: Produced when excited electrons in atoms fall to lower energy levels, emitting light at specific wavelengths.

Atomic Absorption Spectrum: Produced when atoms absorb specific wavelengths of light, corresponding to energy differences between levels.

Photoelectric Effect

Particle Nature of Light

When light of sufficient frequency strikes a metal surface, electrons are ejected.
The number of electrons depends on the light’s intensity, but their energy depends on the light’s frequency.
Explained by Einstein’s photon theory: light consists of particles (photons) with energy .

Summary Table: Key Experiments and Discoveries

Experiment	Scientist	Discovery
Cathode Ray Tube	J.J. Thomson	Electron, charge-to-mass ratio
Oil Drop	Millikan	Electron charge
Alpha Scattering	Rutherford	Nucleus, nuclear model
Photoelectric Effect	Einstein	Photon, quantization of light

Additional info:

Some diagrams and images referenced in the notes are not included here, but their content has been described in the text.
Examples and calculations are based on standard textbook problems for General Chemistry.