Atomic Theory: Development, Models, and Subatomic Particles, Study Guides, Projects, Research of Chemistry

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Atomic Theory

September 2023

ADDU-JHS Grade 9 Science (Chemistry) Sunday Hamoy

Development of Atomic Theory

https://www.youtube.com/watch?v=xazQRcSCRaY

Dalton’s “Solid Sphere” Atomic Theory (1803)

1. All matter consists of tiny particles called “atoms”. (Democritus called them “atomos”.)
2. Atoms of one element can neither be subdivided nor changed into atoms of any other element.
3. Atoms can neither be created nor destroyed. (Supports Law of Conservation of Mass)
4. Atoms of the same element are identical in mass, size, and other properties.
5. In compounds, atoms of different elements combine in simple, whole number ratios. (Law of Multiple Proportions). In a given compound, the numbers of atoms of each of its elements are always present in the same ratio (Law of Definite Proportions).
6. Atoms of one element differ in mass and other properties from atoms of other elements.

Electron Discovery (1897): J.J. Thomson’s Cathode Ray Experiment Deflection region Drift region Displacement

Anodes / collimators Cathode Volts

• Thomson observed an electric field deflecting cathode rays towards a positive plate.

Ernest Rutherford’s “Nuclear Model” (1911)

• Most of the α particles were not deflected, so the atom must be mostly empty space.
• Fewer α particles were deflected, so the positive objects being struck occupy small volumes only.
• Because the α particles were deflected so much, the positive objects being struck must have big masses.
• Small volume and big mass = high density
• This small, high-density, positively-charged area is the nucleus.
• The existence of the proton was proved by Rutherford in 1920, when he showed that the Hydrogen nucleus (a proton) was also present in other atoms.

Niels Bohr’s Planetary Model of the Atom (1913)

• Electrons are only stable in certain allowed orbits, or “stationary

states”. (Atomic orbital – a particular spatial distribution for an

electron)

• Electrons dropping down from one stable orbit to another would

radiate a single discrete packet of radiation, in the form of a photon of

light. (Albert Einstein, Mathematical Law of the Photoelectric Effect /

“light is both a wave and a particle”, 1905)

• Only certain atomic energies are possible, hence only certain

frequencies of electromagnetic radiation or photons are emitted – i.e.

“quantized” energy. (Max Planck, 1900)

Niels Bohr’s Planetary Model of the Atom (1913)

• Bohr’s theory successfully explained the regularities in the pattern of

light emitted from hot hydrogen gas in the laboratory and in the

atmospheres of stars. Heated hydrogen emits characteristic blue, red

and violet light, and a photon of each of those colors of light

corresponds to the energy difference between the different allowed

orbits.

• Direct experimental evidence for the existence of such discrete states

was obtained in 1914 by the German physicists James Franck and

Gustav Hertz.

Planetary vs. Nuclear Model

• The planetary model incorporates into the classical mechanics description of the atom, Planck’s ideas of quantization and Einstein’s finding that light consists of photons whose energy is proportional to their frequency.
• Bohr assumed that the electron orbiting the nucleus would not normally emit any radiation (the stationary state hypothesis), when moving within the same orbit. It would emit or absorb a photon if it moved to a different orbit. The energy absorbed or emitted would reflect differences in the orbital energies.

Atomic Emission Spectrum (or Line Spectrum)

• The Atomic Emission Spectrum is light composed of a few colors or

wavelengths emitted by the electrons of an atom; it is a kind of

atomic fingerprint that is useful in identifying elements.

• Bohr’s planetary model explains how each element emits its own

characteristic atomic emission spectrum:

• Electrons do not give off energy when they stay at their orbits with

defined energy levels.

• When energy is absorbed, an electron jumps from the ground

state to higher-energy excited state. This period of excitation is

very brief and the electron soon jumps back to a lower energy

level, not necessarily the ground state, along with the emission of

light of specific wavelength

Bohr’s Model:

Energy Levels of

Electron Orbits

Horizontal lines show the relative energy of orbits in the Bohr model of the H-atom. Vertical arrows depict the energy of photons absorbed (left) or emitted (right) as electrons move between these orbits.

Question:

Who proposed the

“charge cloud” atomic

model?

a) Erwin Schrodinger b) Louis de Broglie c) Max Planck d) Werner Heisenberg

Heisenberg Uncertainty Principle (1927)

“At every moment the electron has only an inaccurate position and an

inaccurate velocity, and between these two inaccuracies there is this

uncertainty relation.”

Heisenberg’s Support for Bohr’s Model

• Correction: Bohr’s model violated the Heisenberg uncertainty

principle by trying to specify simultaneously both the position (e.g.

orbit of a particular radius) and energy (i.e. related to momentum) of

the electron.

• Correction: Given its mass and wavelike nature, the electron could not

possibly orbit the nucleus in a well-defined circular path. Bohr’s model

was only able to predict the most probable radius of the electron in

the hydrogen atom.

Schrödinger’s Wave Functions (1926)

• Schrödinger developed wave mechanics, a mathematical technique

that describes the relationship between the motion of a particle that

exhibits wavelike properties (e.g. electron) and its allowed energies.

• Each wave function is a solution to Schrödinger’s equation, and is

associated with a particular energy described by a set of quantum

numbers (e.g. energy profile of electrons in an atom).

Schrödinger’s Support for Bohr’s Model

• Supported: Energy of an electron is quantized.
• Supported: Bohr only assumed the idea of quantization.

Schrödinger’s wave functions proved quantization is a natural

consequence of the wavelike behavior of an electron.