C5: The Development of Atomic Theory, Foundations of Quantum Mechanics and Electron Behavior

Democritus

Proposed that matter is made of tiny indivisible particles called atomos (Greek for "uncuttable").

Democritus' Belief

Believed atoms differed in shape and size depending on the substance.

Aristotle

Rejected atomic theory and believed matter was made up of four basic elements: earth, air, fire, and water.

Alchemy Era

Practiced between ancient times and the 1600s, with goals including turning base metals into gold.

Alchemy Contributions

Contributed to early chemical procedures like distillation and metal extraction.

Robert Boyle

Published The Sceptical Chymist (1661) and disproved Aristotle's theory.

Boyle's Emphasis

Emphasized quantitative experiments.

Boyle's Definition of Elements

Defined elements as substances that cannot be broken down further.

Dalton's Atomic Theory

All matter is made of indivisible atoms; all atoms of an element are identical; different elements have different atoms.

Dalton's Ratios

Atoms combine in simple ratios to form compounds.

Chemical Reactions

Chemical reactions rearrange atoms but do not create or destroy them.

Cathode Rays

A stream of particles traveling from the cathode to the anode, discovered through Crookes Tube experiments.

Crookes Tube Setup

A sealed glass tube connected to a vacuum pump containing a cathode, anode, and metal cross.

Jean Baptiste Perrin's Observation

Discovered that a magnet deflected the cathode ray, concluding the beam was made of charged particles.

J.J. Thomson's Breakthrough

Modified Crookes tube to focus the cathode ray and observed that the beam bent toward the positive plate.

Charge-to-Mass Ratio

Determined the charge-to-mass ratio of the particles as 1.76 × 10⁸ C/g, which was 2000 times greater than the ratio for a hydrogen ion.

Discovery of the Electron

Concluded that particles had the same charge as a hydrogen ion but a mass 2000× smaller.

Universality of the Electron

Thomson found electrons are in every atom through repeated experiments with different cathode metals and gases.

Thomson's Announcements

Publicly announced the discovery of "corpuscles" (electrons) on April 29, 1897.

Thomson's Nobel Prize

Awarded the Nobel Prize in 1906 for work on the conduction of electricity through gases.

Significance of Thomson's Discovery

Atoms are not indivisible—they are made of smaller particles with opposite electrical charges.

Plum Pudding Model

Proposed the atom as positively charged "pudding" with negatively charged electrons scattered throughout.

Ernest Rutherford

Initially supported Thomson's "plum pudding" model of the atom and later conducted research on radioactivity.

Alpha (α)

Helium nuclei emitted with a speed of ~16,000 km/s, penetrating a few cm of air and stopped by aluminum foil.

Beta (β)

Electrons emitted with a speed of ~200,000 km/s, penetrating a few mm of aluminum and stopped by thin lead.

Gamma (γ)

High-energy EM radiation emitted with a speed of ~300,000 km/s (light), penetrating very high: 30 cm lead or 2 km air.

Gold Foil Experiment

Conducted in 1909 at the University of Manchester by Ernest Rutherford, Hans Geiger, and Ernest Marsden using thin gold foil and alpha particles.

Plum Pudding Model

A model that expected alpha particles to pass straight through gold foil with minimal deflection.

Rutherford's Conclusion

The atom has a tiny, dense, positively charged center called the nucleus, with electrons orbiting in surrounding empty space.

Nucleus

Contains virtually all of the atom's mass but very little of its volume.

Atomic Number (Z)

The number of protons in the nucleus of an atom, determining the element's identity and equal to the number of electrons in a neutral atom.

Mass Number (A)

Total number of protons + neutrons in an atom's nucleus, calculated using the formula: Mass Number (A) = Number of Protons + Number of Neutrons.

Isotopes

Atoms of the same element with the same number of protons but different numbers of neutrons.

Carbon-12

An isotope of carbon with 6 protons and 6 neutrons, making up 98.89% of carbon.

Carbon-13

An isotope of carbon with 6 protons and 7 neutrons, making up 1.11% of carbon.

Carbon-14

An isotope of carbon with 6 protons and 8 neutrons, making up less than 0.01% of carbon.

Mass Spectrometry

A technique used to determine the isotopic composition, masses, and relative abundances of elements based on the deflection of charged particles in a magnetic field.

Neon Isotopes

Isotopes of neon that separate inside the mass spectrometer, including Neon-20, Neon-21, and Neon-22.

Robert Boyle

Scientist known for early gas laws and promoting the idea of elements as substances that can't be broken down.

J.J. Thomson

Discovered the electron and proposed the plum pudding model.

Ernest Rutherford

Discovered the nucleus and conducted the gold foil experiment.

James Chadwick

Discovered the neutron in 1932.

Frederick Soddy

Coined the term 'isotope' and worked with Rutherford on radioactivity.

Electron (e⁻)

A subatomic particle with a charge of -1, located outside the nucleus, with a relative mass of ~0 (1/1836 amu) and an absolute mass of 9.109 × 10⁻²⁸ g.

Proton (p⁺)

A subatomic particle with a charge of +1, located in the nucleus, with a relative mass of 1 amu and an absolute mass of 1.673 × 10⁻²⁴ g.

Neutron (n⁰)

A subatomic particle with no charge, located in the nucleus, with a relative mass of 1 amu and an absolute mass of 1.675 × 10⁻²⁴ g.

Frederick Soddy

Coined the term "isotope"; worked with Rutherford on radioactivity.

William Crookes

Developed the Crookes tube, observed cathode rays.

Jean Baptiste Perrin

Demonstrated cathode rays were negatively charged particles.

Hans Geiger & E. Marsden

Conducted gold foil experiment with Rutherford.

Marie & Pierre Curie

Discovered radioactive elements like radium and polonium.

Robert Millikan

Conducted the oil drop experiment; measured the charge of an electron.

John Dalton

Proposed the first modern atomic theory.

Antoine Lavoisier

Known as the father of modern chemistry; formulated the law of conservation of mass.

Joseph Priestley

Discovered oxygen (credited by Lavoisier).

Joseph Proust

Proposed the law of definite proportions.

Carl Scheele

Discovered several elements including oxygen (independently).

Julius Plücker

Early experiments with discharge tubes that led to the study of cathode rays.

Rutherford's Model Challenge

Classical physics said accelerating charges (like orbiting electrons) should emit energy, leading to atom collapse.

Planck's Quantum Theory

Energy is not continuous—it is quantized; energy is absorbed or emitted in discrete packets called quanta.

Planck's Formula

E = hv, where E is energy of a quantum, h is Planck's constant (6.626 × 10⁻³⁴ J·s), and v is frequency of radiation.

Einstein & the Photoelectric Effect

Proposed that light consists of particles called photons; explained why only light of a certain frequency can knock electrons off a metal surface.

Niels Bohr and the Hydrogen Atom

Used quantum ideas to fix Rutherford's model and explain the hydrogen line spectrum.

Bohr's Postulates

Electrons orbit the nucleus only in specific, allowed circular paths (energy levels) with fixed energy.

Ground State

n = 1 (lowest energy).

Excited States

n = 2, 3, 4... (higher energy).

Energy Emission

Electrons emit energy only when moving between levels, equal to the difference in energy: ΔE = E_high - E_low.

Bright-Line Spectrum

When hydrogen gas is excited, it emits light at specific wavelengths, revealing a bright-line spectrum.

Quantized

It means it can only exist in specific, discrete amounts—not in any value in between.

Examples of Quantized

Energy in atoms, Money, Stairs.

Energy Absorption under Planck's Theory

Energy is absorbed or emitted in whole-number multiples of hv only (like 1hv, 2hv, 3hv).

Balmer Series

Electron falls to n = 2 from higher energy levels, producing visible light.

n = 3 to n = 2 Transition

Produces red light (~656 nm).

n = 4 to n = 2 Transition

Produces blue-green light (~486 nm).

n = 5 to n = 2 Transition

Produces blue light (~434 nm).

n = 6 to n = 2 Transition

Produces violet light (~410 nm).

Lyman Series

Electrons fall to n = 1, producing ultraviolet radiation.

Paschen Series

Electrons fall to n = 3, producing infrared radiation.

Hydrogen's Bright-Line Spectrum

Consists of four distinct colored lines—red, blue-green, blue, and violet—on a black background.

Electron Transition Mechanism

Electrons drop from higher levels (n = 3, 4, 5, 6) to n = 2, releasing energy as photons.

Ultraviolet Region Transitions

Transitions to n = 1 generate ultraviolet radiation.

Rydberg Constant (RH)

RH = 1.097 × 10^7 m−1.

De Broglie's Idea

Proposed that electrons have wave properties, similar to light.

De Broglie's Equation

λ = h / mv, where λ is wavelength, h is Planck's constant, m is mass, and v is velocity.

Planck's Constant (h)

h = 6.626 × 10−34 J·s.

Wave Behavior of Electrons

Electrons exhibit wave behavior, confirmed by diffraction patterns in experiments.

Impact of De Broglie's Theory

Provided a foundation for quantum mechanics and explained electron orbits.

Baseball Wavelength Example

Wavelength = 1.24 × 10−34 m, negligible and undetectable.

Electron Wavelength Example

Wavelength = 3.3 × 10−10 m, significant compared to atom size.

Stable Electron Orbits

Only orbits where a whole number of electron wavelengths fit are stable.

Bohr's Model Limitations

Worked only for hydrogen and failed for multi-electron atoms.

Wave-Particle Duality

Quantum mechanics treats particles like electrons as both waves and particles.

Diffraction Pattern Experiment

Davisson and Germer's experiment showed electrons acting like waves.

Energy Release Mechanism

Each electron drop releases energy in the form of a photon of specific wavelength.

Visible Spectrum Lines

Produced by specific electron transitions to n = 2.

Matrix Mechanics

Created by Werner Heisenberg in 1925, treating the electron as a quantum particle using complex math (matrix algebra).

Wave Mechanics

Developed by Erwin Schrödinger in 1926, treating the electron as a wave using a wave equation (simpler math for many cases).

Heisenberg's Uncertainty Principle

Core idea: You cannot know both the exact position and velocity of a particle at the same time. The more precisely you know one, the less precisely you know the other.

Uncertainty in Position (Δx)

The uncertainty in the position of a particle.

Uncertainty in Velocity (Δv)

The uncertainty in the velocity of a particle.