1. History of the Atomic Model
Once upon a time…
Some info from

2. Dmitri Mendeleev organizes our elements
In 1869, a Russian scientist named Dmitri Mendeleev, drew up our first periodic table of elements arranged according to increasing atomic masses. (this was before knowing of protons, neutrons and electrons, he just intuitively believed that chemical properties were related to their masses)
The table revealed a periodic relationship among the elements’ properties. This means there were repeating patterns of related chemical properties.
He also left holes in it where he felt new elements would be discovered later on.

3. Mendeleev’s Original 1869 Table
Mendeleev’s table as published in 1869, with many gaps and uncertainties.
A later 1898 Version
The table, rotated ninety degrees, as shown in textbooks in 1898 when Marie Curie discovered Radium. Mendeleev promptly added Ra to his table also.

4. Henry Mosely fine-tunes the periodic table
In 1913, Henry Mosely (an Englishmen) discovered a method to calculate the atomic number (the positive charge of the nucleus) through the bombardment of electrons creating a unique X-ray spectrum of every atom. This process was called X-ray diffraction.
In 1914, Henry Mosley arranged the periodic table today according to an element’s atomic number…and we still do it this way today!
This method lead to better periodicity and explained why argon, with an atomic mass of amu (atomic number of 18), can come before potassium with the atomic mass of amu (atomic number of 19).
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5. Democrotis with first mention of atoms (460B.C. - 370B.C.)
The greek philospher Democrotis was among the first to suggest the existence of atoms…a discontinuous theory of matter.
The word atom comes from “atomos” meaning indivisible or indestructible.
Atoms were thought of as the building blocks (the Legos) of everything.
Aristotle (another greek philospher around 350BC) believed in a thought that all matter was made of only 4 elements: earth, wind(air), fire and water.
Aristotle

6. John Dalton’s Solid Sphere Model
In 1803, John Dalton was the first to show laboratory evidence of the existence of atoms.
Dalton’s model of the atom was a solid sphere with no subatomic particles…the atom was indivisible!
Dalton also proposed the first atomic theory which stated that
1. all things are made up of atoms and
2. all atoms of the same element are identical (later determined this not to be true…isotopes)
3. these atoms can combine in simple whole number ratios to form compounds.

7. Thomson’s Discovery of the Electron, shatters Solid Sphere Model
He used a cathode ray tube to discover the presence of electrons…which were even smaller than atoms!

8. JJ Thomson’s Plum Pudding Model
In 1898 JJ Thomson proposed The Plum Pudding Model for the structure of an atom, shattering the solid sphere model.
The model pictured a sphere of positive electricity (“pudding”) studded with negatively charged electrons (the “plums”).
Later this model proved to be wrong by his student E. Rutherford.

9. Earnest Rutherford’s Gold Foil Experiment
His expected results.
He shot alpha particles (+ charged) through the atoms of gold in a thin sheet of foil, and studied how they were scattered by the atoms. Oddly, some of the particles “bounced” back off the foil but most when straight through.
The circles are atoms of gold, the dots are the nucleus and the arrows are the alpha particles (positively charged electromagnetic radiation).
The actual results…Shocking!

10. Rutherford Nuclear Atom
Rutherford (a student of Thomson) set out to prove Thomson’s model correct, however ended up proving it wrong.
By 1911, Rutherford had theorized a new model with a positively charged dense central nucleus surround by a cloud of negative electrons.
Although very, very tiny the nucleus contains 99.9% of an atoms mass surrounded mostly by empty space.
Rutherford’s Model

11. Millikan’s Oil Drop Experiment
When he sprayed oil droplets into a chamber and bombarded them with X-rays to place a negative charge on them, the charged droplets were attracted to the positive plate. Changing the strength of the electrical field offset the attraction and allowed Millikan to determine the charge of the electrons.
He measured the charge of the electrons (1.60 X coulombs) and later used this to calculate their extremely small mass (9.11 x grams).

12. Bohr Model of the Atom Bohr Planetary Model
The Bohr model has protons and neutrons in the nucleus surrounded by orbitals with definite electron paths, kind of like a planetary model.
Atoms in their normal, energetically stable state are said to be in their ground state.
The greater the distance of the orbitals from the nucleus, the greater is the energy of the electrons in that shell (because they are overcoming the electrostatic attraction towards the nucleus).
Bohr Planetary Model

13. Neils Bohr
Neils Bohr studied light emitted from excited (electrically stimulated) gas atoms called atomic emission spectrums (bright-line spectrums). It turns out that every element emits a unique spectrum (combination) of wavelengths when excited. (or bright-line spectrum).
This emission spectrum can be used as a type of fingerprint for determining what elements are present. Atomic emission spectrums are used today to determine what elements are present in far away stars.
He determined that this must mean that each electron must jump exact, discrete amounts within set orbitals within an atom. This contributed to his discovery that there are quantum energy levels (or orbital locations) for electrons. Therefore, an electron in a given orbital has a definite amount of energy.

14. Mechanism of Bright-line Sprectra
1st Step atoms absorb energy….Radiant energy, heat energy, or electrical energy absorbed by the atom causes the electron to move (“jump”) from its lower-energy orbit ground state to a higher-energy orbit (an excited state).
2nd Step Immediately after, the atoms release energy… radiant energy (or photons) are emitted in the form of visible light or X-ray or UV rays when the electron moves from a higher-energy orbit back to its lower-energy orbit.
1st
2nd

15. This lead to Quantum Theory of Atomic Structure
Atoms and molecules could emit (or absorb) energy only in discrete quantities called “quanta”, like small packages or bundles, there were not unlimited energy contents for particles.
Observations of electron behavior in the 1920’s were inconsistent with the planetary model because electrons moved like a particle and yet also like a wave.
Researches such as Albert Einstein (1905), Louis de Broglie (1924) and Erwin Schrodinger (1926) devised a mathematical equation to describe the wave-particle behaviors of electrons.

16. Quantum-mechanical Model of the Atom
It is also known as the Shrodinger Model or Electron-Cloud Model.
Like the Bohr model, this model restricts the energy of electrons to certain levels. Unlike the Bohr model, The Q-M model does not show the definite paths of electrons; it shows the most probable location of an electron as a cloud (region of space)
An electron’s location at any time can not be pinpointed exactly, yet you have a good idea of where it should be.
Also, each energy level contains subshells called orbitals (these are the s, p, d and f orbitals).

17. Quantum mechanical Model
It is very hard to draw. Here are 2 attempts.

18. Photoelectric Effect
This is the theory that light radiation is both a particle and a wave (Albert Einstein worked on this).
In other words it is made of photons (tiny bundles of energy) that behave like waves in their movement.
A high-intensity, low-frequency light will not emit electrons from a metal, but a low-intensity, high-frequency beam will. Depends on the frequency of the light not its intensity.
Electronic eye in a store is an example of this effect.
Einstein’s main conclusion was that all matter exhibits wave properties.

19. Wave-Particle Duality Cartoon

20. Wave-Particle Dualtity
Wave/particle duality is the possession by physical entities (such as light and electrons) of both wavelike and particle-like characteristics. On the basis of experimental evidence, the German physicist Albert Einstein first showed (1905) that light, which had been considered a form of electromagnetic waves, must also be thought of as particle-like, or localized in packets of discrete energy (see the photoelectric effect).
The French physicist Louis de Broglie proposed (1924) that electrons and other discrete bits of matter, which until then had been conceived only as material particles, also have wave properties such as wavelength and frequency. Later (1927) the wave nature of electrons was experimentally established. An understanding of the complementary relation between the wave aspects and the particle aspects of the same phenomenon was announced in 1928.

21. Max Planck’s Eqtn  E= h· f
E is the energy emitted or absorbed
h is Planck’s constant 6.63 x J-s or J/Hz
f is the frequency of radiation in Hz or 1/s
Energy is always emitted in multiples of h f (h f, 2h f, 3 h f ) from an atom.
This is what tells us there are discrete energy levels within electron orbitals.
Wavelength and frequency of light waves when multiplied by each other will always equal a constant, c, which is the speed of light 3.0 x 108 m/s2. So c = wavelength x frequency.

22. SPECTRAL LINES
Spectral lines (emission spectra) are a result of transition of electrons between energy levels.
Their frequency is related to the energy spacing between levels using Planck’s relationship (E=hf )
The frequencies present give off an atomic fingerprint called a bright-line spectra.
We can use the color (determined by the frequency) of light emitted to identify the metal atoms present.

23. Flame Test
We will be doing a Flame Test Lab to identify the presence specific metals in our compounds.
We will be adding heat energy to metal ionic compounds. The metal ions will absorb energy and electron(s) will “jump” up to a higher energy level. This is called an “excited state”.
This excited state is highly unstable, so instantly (within a millisecond) the electron(s) will return to the lower energy level and emit a photon of light energy in return.
Each metal ion will emit a different frequency of light (dif. color) related to the energy levels that the electron moved within.

24. Other emission spectra uses
An emission spectrum (or bright-line spectrum) is a more quantitative way to determine the identity of atoms. This method looks through a spectroscope at the specific multiple wavelengths that are released.
This method can be used to determine the identities of atoms that are in stars in the sky.