The Evolution of the Atomic

TheoryBy John, Caroline, and Gabi

Greek Observations• The ancient Greeks took a great interest in

science, even prior to the discovery of the empirical method.

• Pre-Socratic philosopher Demokritos questioned whether matter is continuous or infinitely divisible into smaller pieces– He found matter was indeed divisible and coined

the term “atomos” (atom), using it to refer to the smallest pure building block of matter.

– This is considered the first universal atomic theory.

Discovery of Chemical Elements

• Robert Boyle, author of the Skeptical Chymist, is the father of two major quantitative sciences: Physics and Chemistry.

• His revolutionary ideas about the existence and properties of chemical elements are considered his greatest contributions.– He proposed that a substance was

scientifically an element unless it could be broken down into two or more simpler substances

c. 1661

Groundbreaking Discoveries on Mass and Energy

• Antoine Lavoisier, a French chemist, theorized that mass is neither created nor destroyed in ordinary chemical reactions.

• This discovery led the Law of Conservation of Mass which quickly became the basis for many developments in chemistry.

c. 1774

Groundbreaking Discoveries on Mass and Energy

• Joseph Proust, a Frenchmen, 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 Theory1. Each element is made up of tiny particles

called atoms.2. The atoms of a given element are identical;

the atoms of different elements are different in some fundamental way.

3. 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.

4. 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 by the size of the gas’ individual particles.

1811

Interpretations of

Avogadro’s Hypothesis

• Following 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.

• e/m = - 1.76 x (10)^8 C/g–Where 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 9.11 x 10-31 kilograms.

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

Negatively charged electrons are randomly distributed “plums” on the positive cloud of cake.

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.

1911

The Modern Atom• Atomic structure

– The nucleus contains positively charge protons and neutral neutrons.

– This comparably 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 considered 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 = 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 let off in various wavelengths, which results in an emission spectrum.

• This spectrum is considered a line spectrum. In a line spectrum, only certain energies are allowed for the electrons in the atom. The visible light is only 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 the discovery of hydrogen’s

atom spectrum, Danish physicist Niels Bohr developed a quantum model for the hydrogen atom.

• The structure of his model places hydrogen’s electron in a particular ring in which it orbits around the nucleus.

• This finding violated some laws of classical physicals, and, after ample experimentation, the technicalities of the model were ultimately proven wrong. The ideas behind the model itself, 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 Schrodinger – 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 Schrodinger’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 wave function of 1s orbital. From there on, energy level in the continuing orbits increases 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 in which he proposes that the more accurately we know a particle’s position, the less accurately we know its momentum.

1920s