History of Atomic Theory - Comprehensive Study Notes

Ancient Greek Ideas

Based on observation and argument, Aristotle came to believe there are four fundamental elements: earth, ext{ } ext{wind}, ext{ } ext{fire}, ext{ } ext{water}.
Aristotle was a towering figure in philosophy; his ideas persisted for thousands of years.
Even as late as the 1800s, educated people were convinced that Aristotle was right.

Atomism

Democritus was a Greek philosopher born around 460 BCE.
Parmenides argued that change (in the natural world) was an illusion because it requires that something should come to be from nothing.
Democritus argued that there must be unchanging “material principles” (atoms) which can rearrange themselves to “form the changing world of appearances.”
His arguments went much deeper and sometimes came remarkably close to modern theory developed over 2000 years later.

The Greek Atom

Atomos is the Greek word for “uncuttable” or indivisible.
There are many conflicting reports of how Greek atomists envisioned atoms.
They were thought of as purely spherical and solid (unyielding).
They were thought to move through “void” which is “yielding.”
Atomists imagined clusters of atoms that held together with hook-like structures on their surfaces.
Weirdly enough, none of this is too far off from reality.

Robert Boyle

Born in 1627 as the 14th child of an English nobleman.
At Oxford, he was committed to the New Philosophy which put equal emphasis on observation and experiment with logic.
Conducted early work on gases and built vacuum pumps.
Conducted important chemical work and made many novel observations.
Was a committed alchemist his entire life and never stopped believing in transmutation.

Corpuscularism

Aristotle’s elements (earth, wind, fire, water) and Paracelsus’ elements (salt, sulfur, mercury) were being displaced by atomists in the 1600s.
Corpuscles (particles) were a primitive atomic idea that came out of philosophy about chemical experiments. They were never observed.
Boyle stated in his book, The Sceptical Chymist (1661), that elements were “certain primitive and simple, or perfectly unmingled bodies; which not being made of any other bodies, or one another, are the ingredients of which all those called perfectly mixt bodies are immediately compounded, and into which they are ultimately resolved.”

Groundbreaking Discoveries on Mass and Energy

Antoine Lavoisier (c. 1774) theorized that mass is neither created nor destroyed in ordinary chemical reactions; this is the Law of Conservation of Mass.
This discovery quickly became the basis for many developments in chemistry.

Joseph Proust

A French chemist who found that a given compound always contains the same proportions of elements by mass.
This is known as the Law of Definite Proportion. 1798 - 1806

John Dalton (1766-1844)

He reasoned that if elements were composed of tiny individual particles, a given compound should always contain the same combination of these particles (atoms).
Dalton also discovered the Law of Multiple Proportions: when two elements form a series of compounds, the ratios of the masses of the second element that combine with a fixed mass of the first element can always reduce to small whole numbers. 1808

Dalton’s Atomic Theory

Each element is made up of tiny particles called atoms.
The atoms of a given element are identical; the atoms of different elements are different in some fundamental way.
Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers of types of atoms.
Chemical reactions involve reorganization of the atoms, changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

First Table of Atomic Masses

Dalton organized the first table of atomic masses; mass is determined by being compared to a standard mass (weighing).
Many of his assumptions were later found incorrect, but his work constituted a very important step in characterizing the atom.
His work was not recognized for many years, but was instrumental to the work of Joseph Gay-Lussac and Amadeo Avogadro.

Avogadro’s Hypothesis

In 1811, Italian chemist Avogadro proposed that:
- at the same temperature and pressure, equal volumes of different gases contain the same number of particles.
Following this theory, we can conclude that the volume of a gas determines the number of molecules present; not the size of the gas’ individual particles. 1811

Interpretations of Avogadro’s Hypothesis

Gay-Lussac theorized that when two volumes of hydrogen react with one volume of oxygen, two volumes of water vapor are created.
This eventually was simplified into: two molecules of hydrogen + with one molecule of oxygen = one molecule of water.

The Electron

J.J. Thomson’s work marked the beginning of scientists understanding the composition of the atom.
He studied cathode rays and postulated that the negative stream of particles that came from cathode rays can be characterized as electrons.
\frac{e}{m} = -1.76 \times 10^{8} \ \mathrm{C/g}
(e represents the electrons’ charge in coulombs and m represents its mass in grams.)
1898 - 1903

Electron Charge

In an experiment involving oil drops, scientist Robert Millikan calculated the magnitude of the electron charge.
Using this and Thomson’s charge-to-mass ratio, Millikan calculated the mass of the electron to be m_e = 9.11 \times 10^{-31} \ \mathrm{kg}. 1909

Assumptions from Thomson’s Theory

Thomson surmised that all atoms must contain electrons because electrons can be produced from electrodes, which are made up of various kinds of metals.
Additionally, based on the previously confirmed proposition that atoms have neutral charge, Thomson concluded that a particle with a positive charge must also exist to cancel out the electrons’ negative charge.
With all his assumptions sorted out, he laid out the first visual model of the atom’s structure. He dubbed it the Plum Pudding Model.

The Plum Pudding Model

Description: negatively charged electrons embedded in a positively charged “pudding” medium; the electrons are like “plums” in the positive cake.
Visual hints in the diagram include a positively charged region surrounding the electrons.

The Nuclear Atom

In 1911, Ernest Rutherford carried out an experiment designed to test Thomson’s plum pudding model using alpha particles directed at a thin sheet of metal foil.
His results indicated that Thomson’s model was incorrect and that the atom’s structure could only be explained plausibly by the presence of a nucleus surrounded by electrons.
His work provided the foundation for the structure of the modern atomic model.

The Modern Atom

Atomic structure: the nucleus contains positively charged protons and neutral neutrons.
This tiny center is surrounded by a vast, negative electron cloud.
Different numbers of electrons and protons separate the atoms of different elements.
Atoms with the same number of electrons and protons but different number of neutrons are isotopes of the same element.

How to Read an Atomic Symbol

Atomic number (z) = the number of protons.
Mass number (a) = protons + neutrons.
Number of electrons is equal to the atomic number in an uncharged atom. In ions, the electron number varies depending on the change in charge.

Planck’s Findings: The Foundations of the Quantum Model

Through a series of experiments, Max Planck found that the energy of matter is not continuous; rather, it is quantized and appears in discrete units called quantums.
Therefore, energy exhibits the same properties as particles.
Planck’s Constant h is the constant that, when multiplied by the frequency of the electromagnetic radiation, characterizes a quantum.
c. 1900

The Photoelectric Effect

Einstein proposed that electromagnetic radiation can be viewed as a stream of particle-like components that he referred to as photons.
The characteristics of these photons are described in Einstein's award-winning analysis of their photoelectric effect.
After analyzing how photons interact, Einstein concluded that they do, in fact, have mass. 1921

The Atomic Spectrum of Hydrogen

When hydrogen samples are excited, they emit their excess energy in the form of light.
This light is emitted in discrete wavelengths, resulting in an emission spectrum.
This spectrum is considered a line spectrum: only certain energies are allowed for the electrons in the atom.
The visible light is emitted in discrete lines.
Because diffraction patterns can only be explained in terms of waves, electrons must have wavelengths.
This was tested in a 1927 experiment by C.J. Davidson and L.H. Germer involving nickel crystals. 1927

The Bohr Model

In response to hydrogen’s atom spectrum, Danish physicist Niels Bohr developed a quantum model for the hydrogen atom.
The structure places the hydrogen electron in a particular ring in which it orbits around the nucleus.
This finding violated some laws of classical physics, and, after ample experimentation, the technicalities of the model were ultimately proven wrong.
The ideas behind the model, however, paved the way for future, groundbreaking research on the atom. 1913

The Atom’s Quantum Mechanical Model

The unreliability of the Bohr Model prompted three early 20th century scientists – Heisenberg, de Broglie, and Schrödinger – to develop a new approach in characterizing the atom known as quantum mechanics.
Their early research was centered on how electrons bind to the nucleus in a standing wave, a position similar to how the strings of musical instruments vibrate to create a musical tone.

Wave Functions

Electrons’ position in relation to the nucleus and their movements about it can be described in orbital wave functions, which are a part of Schrödinger’s elaborate mathematical formula for standing waves.
Wave functions are the foundation of the quantum mechanical model of the atom.

An Overview of the Quantum Mechanical Model

An electron’s position in an orbital corresponds to its energy level.
The orbital lowest in energy is the 1s orbital. From there, energy levels in the continuing orbits increase with each orbit.
Although we know how electrons organize themselves within orbitals, we do not know their exact patterns of movements as they circle the nucleus.
This inevitable lack of knowledge led Heisenberg to develop his Uncertainty Principle, which states that the more accurately we know a particle’s position, the less accurately we know its momentum. 1920s

The Neutron (Chadwick)

James Chadwick (1891-1974) – The Existence of a Neutron. (publication: 1932)
Experimental setup (as depicted): vacuum pump, Be target, paraffin wax, radiation source from polonium, leads to a recoil in paraffin wax; gamma radiation could not cause this.
Conclusion: a neutral, very massive particle could cause the observed effects; this particle is the neutron.

The Neutron

Neutron has no charge, making it hard to detect.
It was the last part of the atom to be discovered.
The experiment involved bombarding beryllium with radiation from a polonium source; this radiation penetrated a lead shield.
The mysterious radiation could dislodge protons from paraffin wax with high recoil; gamma radiation could not.
Conclusion: a neutral particle that is very massive explains the observations; this is the neutron.