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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

Development of Atomic Models


  • Greek philosopher

  • 400 BC

  • “Atomos” concept

Development of atomic models

John dalton early 1800 s
John Dalton (early 1800’s)

  • Coined the term “atom”.

Dalton s atomic theory
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.

Development of atomic models

Development of atomic models

Dalton’s Atomic Theory called

“Hard Spheres Model”

Thomson experiments
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 Rays were negatively charged

Cathode Ray Tube and Magnet

They bent toward (+) plate

Development of atomic models

Cathode Rays were particles

They couldn’t pass through matter.

Jj is awesome
JJ is Awesome

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

  • Discovered first subatomic particle (electron).

What about the positive
What about the Positive?

  • But…matter is neutral.

  • Therefore:

    • A positive charge must exist to balance the negative.

Plum pudding model
Plum Pudding Model

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


Brian Cox:

Thompson and Discovery of Electron

Ernest rutherford 1908
Ernest Rutherford (1908)

  • Physicist who worked in new field of radioactivity.

Found 3 different types of radiation
Found 3 Different Types of Radiation

  • Used magnetic field to isolate three types of radiation.

  • Alpha (α)

  • Beta (β)

  • Gamma (γ)

Charges of radiation
Charges of Radiation

  • The radiation had different charges.

Identify the charge each type of radiation has.

Gold foil experiment
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

Development of atomic models


Most alpha particles passed straight

through gold foil.


Atom’s volume is mostly empty space.

Development of atomic models


A few alpha particles

deflected at an angle

or bounced back.


Atoms have a very

small, dense positively

charged nucleus.

Development of atomic models

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
The Nuclear Model

Rutherford’s Model is called the “Nuclear Model”

Brian Cox: Rutherford and the Nucleus

Comparison to thomson
Comparison to Thomson

  • Positively charge only contained in nucleus.

  • Negatively particles scattered outside nucleus.

  • Charge is not disbursed evenly.

Niels bohr 1913
Niels Bohr (1913)

  • Came up with the “Planetary Model”

Bohr s theory
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
Energy Levels

  • Different energy levels can contain different numbers of electrons.

How many per level
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
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
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 (cations)

(+) ions are smaller

Lost electron(s)

Ions anions
(-) Ions (anions)

(-) ions are larger

Gained electron(s)

How did bohr come up with his model
How Did Bohr Come Up With His Model?

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

What are spectral lines
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”)

What causes spectral lines
What Causes Spectral Lines?

Jumping Electrons!!

Video of Line Spectra of Hydrogen

Jumping electrons
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
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 electrons1
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

Development of atomic models

  • 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
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
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: Argon 2-8-8, Neon 2-8

Valence electrons1
Valence Electrons

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

  • Ex:

    • Al 2-8-3 3 valence

    • Ca 2-8-8-2 2 valence

    • F 2-7 7 valence

Lewis dot diagrams
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 2-8-3 has 10 kernel electrons

      and 3 valence electrons

The nature of light
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
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.

Development of atomic models

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?

Development of atomic models

Relationship of Frequency, Wavelength and Energy of colored line

Good overview videos
Good Overview Videos line

  • Crash Course: History of Atomic Theory


  • Quantum Mechanics and the Bohr Model


Calculating the energy of a spectral line honors
Calculating the Energy of a Spectral Line (HONORS) line


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 honors1
Calculating the Energy of a Spectral Line (HONORS) line


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