Development of atomic models
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Development of Atomic Models. Democritus. Greek philosopher 400 BC “ Atomos ” concept. Can matter can be divided forever? Eventually, a piece would be “indivisible” “ Atomos ,” meaning “not to be cut , ” is smallest piece of matter. John Dalton (early 1800’s).

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Development of Atomic Models

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Development of Atomic Models


  • Greek philosopher

  • 400 BC

  • “Atomos” concept

  • Can matter can be divided forever?

  • Eventually, a piece would be “indivisible”

  • “Atomos,” meaning “not to be cut,” is smallest piece of matter

John Dalton (early 1800’s)

  • Coined the term “atom”.

Dalton’s Atomic Theory

  • Matter made of tiny indivisible particles called “atoms”.

  • Atoms of one element are alike, and different from atoms of other elements.

Page from Dalton’s Journal

  • Compounds form when different atoms combine in fixed proportions.

  • Chemical reactions involverearrangement of atoms.

  • Atoms can’t be created or destroyed, but are conservedin chemical reactions.

Dalton’s Atomic Theory called

“Hard Spheres Model”

JJ Thomson (1897)

Thomson’ Experiments

  • Studied “cathode rays” (electric current) in a “Crooke’s Tube”.

  • Fluorescent screen, shows how cathode ray behaved in a magnetic field.

Lets draw a typical Crooke’s Tube in our notes.

Cathode Rays were negatively charged

Cathode Ray Tube and Magnet

They bent toward (+) plate

Cathode Rays were particles

They couldn’t pass through matter.

JJ is Awesome

  • Concluded the negative “cathode ray” particles came from within atoms.

  • Discovered first subatomic particle (electron).

What about the Positive?

  • But…matter is neutral.

  • Therefore:

    • A positive charge must exist to balance the negative.

Plum Pudding Model

Atoms are positively charged spheres with negatively charged particles scattered throughout.


Brian Cox:

Thompson and Discovery of Electron

Ernest Rutherford (1908)

  • Physicist who worked in new field of radioactivity.

Found 3 Different Types of Radiation

  • Used magnetic field to isolate three types of radiation.

  • Alpha (α)

  • Beta (β)

  • Gamma (γ)

Charges of Radiation

  • The radiation had different charges.

Identify the charge each type of radiation has.

Gold Foil Experiment

  • Shot alpha particles, at very thin piece of gold foil.

  • Alpha particles have a positive charge, and a mass of 4 amu

  • Fluorescent screen shows where the particles went.

Rutherford Gold Foil


Most alpha particles passed straight

through gold foil.


Atom’s volume is mostly empty space.


A few alpha particles

deflected at an angle

or bounced back.


Atoms have a very

small, dense positively

charged nucleus.

Nucleus is extremely small compared to the size of the atom as a whole.

Deflections happened rarely (1/8000).

Modern Example of Gold Foil Experiment in Action

The Nuclear Model

Rutherford’s Model is called the “Nuclear Model”

Brian Cox: Rutherford and the Nucleus

Comparison to Thomson

  • Positively charge only contained in nucleus.

  • Negatively particles scattered outside nucleus.

  • Charge is not disbursed evenly.

Niels Bohr (1913)

  • Came up with the “Planetary Model”

Bohr’s Theory

  • Electrons circle nucleus in specific energy levels or “shells”.

  • The higher the “energy level,” the higher the electron’s energy.

Energy Levels

  • Different energy levels can contain different numbers of electrons.

How many per level?

  • n = the number of the energy level

    2n2 = maximum number of electrons an energy level can hold.

    Ex: Level 3 can hold 2(3)2 = 18 electrons

Draw a Bohr Atom

  • Ex: The Fluorine Atom (F)

    • Protons = 9

    • Neutrons = 10

    • Electrons = 9

    • How many energy levels do you draw?

    • How many electrons in each level?

Human Bohr Model

Draw a Bohr Ion

  • They only difference is that one or more electrons gets added or taken out of the outer energy level.

  • Ex: The Magnesium Ion (Mg+2)

    • Protons = 12

    • Neutrons = 12

    • Electrons = 10

(+) Ions (cations)

(+) ions are smaller

Lost electron(s)

(-) Ions (anions)

(-) ions are larger

Gained electron(s)

How Did Bohr Come Up With His Model?

  • Studied the spectral lines emitted by various elements (especially Hydrogen)

What are Spectral Lines?

  • Energy gets absorbed by an atom causing it to emit a unique set of colored lines.

  • Used to identify what elements are present in a sample. (elemental “Fingerprint”)

Spectral Lines are Different for Each Element

Answer: 1

What Causes Spectral Lines?

Jumping Electrons!!

Video of Line Spectra of Hydrogen

Jumping Electrons

Electrons normally exist in the lowest energy level possible called the “ground state”. (stable)

“Ground state” e- configurations are written on the periodic table for each element.

Ex: Aluminum is 2-8-3

Calcium is 2-8-8-2

An Electron Gets “Excited”

Electrons can absorb a photon (or “quanta”) of energy and “jump up” to a higher energy level farther from the nucleus.

This is called the “excited state”. (unstable)

Jumping Electrons

  • They quickly “fall back down” to the ground state. (stable)

  • They emit a photon (or “quanta”) of energy that corresponds to how far they jumped.

Spectral Lines

  • This photon of energy is seen as a spectral line!

  • Each spectral line corresponds to a specific photon of energy that is released.

    Model Of Hydrogen Atom and Electrons Jumping


Absorb Energy

Jump Up

Emit Energy

Fall Down

Excited vs Ground State

  • Periodic table lists ground state electron configurations for neutral atoms.

    • To recognize an “excited state” configuration, count the electrons and see if the configuration matches the one on the table.

  • Ex: 2-8-7-3 = 20 electrons

    • Calcium (atomic # 20) is 2-8-8-2

    • So this must be showing one of the ways calcium could be in the excited state.

Valence Electrons

  • Electrons in highest occupied energy level.

  • Involved in forming bonds with other atoms.

  • Atoms are most stable when they obtain a “stable octet” of 8 valence electrons

  • Noble Gases: (Group 18)

    • Have stable octet already and are “inert” and unreactive

      • Ex: Argon2-8-8, Neon 2-8

Valence Electrons

  • Look at the last number in the atom’s electron configuration to determine the number of valence electrons.

  • Ex:

    • Al2-8-33 valence

    • Ca2-8-8-22 valence

    • F2-77 valence

Lewis Dot Diagrams

  • Shows the number of valence electrons an atom has as “dots” around the atom’s symbol.

Phosphorus is 2-8-5


  • Nucleus and non-valence electrons

  • Inner part of atom not involved directly in reactions

  • Ex:

    • Al 2-8-3has 10 kernel electrons

      and 3 valence electrons

The Nature of Light

  • Study of light has provided important information about the structure of atoms.

  • Dual Nature of Light:

    • behaves as both waves and as particles (depending on what type of experiment is being performed.)

  • Speed of Light: all light waves travel at the same velocity

    • C = 3.0 x 108 meters/sec

Electromagnetic Spectrum

  • Spectral lines can come from all areas of the EM Spectrum.

  • Lines of visible colors make up only a small part of the spectrum.

  • EM waves carry different amounts of energy based upon their wavelength and frequency.

Wavelength (λ): distance between two peaks of a wave

Frequency (γ): number of peaks that pass per second. (Hertz (Hz) or cycles/sec)

Which wave has higher energy?

Relationship of Frequency, Wavelength and Energy of colored line

Good Overview Videos

  • Crash Course: History of Atomic Theory


  • Quantum Mechanics and the Bohr Model


Calculating the Energy of a Spectral Line (HONORS)


If you know the wavelength of the spectral line you can find it’s frequency.

c = λ x ү

c = the speed of light = 3 x 108 meters/sec

λ = wavelength (in meters)

ү = frequency of the wave

Calculating the Energy of a Spectral Line (HONORS)


Using the frequency find the energy of the line (in Joules)

E = h x ү

E = energy in Joules

h = Planck's constant = 6.63 × 10-34 kg x m2 / sec

ү = frequency of the wave

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