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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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democritus
Democritus
  • Greek philosopher
  • 400 BC
  • “Atomos” concept
slide3
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
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.
slide7
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.
slide8
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

http://youtu.be/XU8nMKkzbT8

They bent toward (+) plate

slide12

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.

yummy
Yummy…

Brian Cox:

Thompson and Discovery of Electron

http://youtu.be/IdTxGJjA4Jw

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

http://www.youtube.com/watch?v=5pZj0u_XMbc

slide21
Observation:

Most alpha particles passed straight

through gold foil.

Conclusion:

Atom’s volume is mostly empty space.

slide22
Observation:

A few alpha particles

deflected at an angle

or bounced back.

Conclusion:

Atoms have a very

small, dense positively

charged nucleus.

http://www.kentchemistry.com/moviesfiles/Units/AtomicStructure/Rutherford3.htm

slide23

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

http://youtu.be/XBqHkraf8iE

the nuclear model
The Nuclear Model

Rutherford’s Model is called the “Nuclear Model”

Brian Cox: Rutherford and the Nucleus

http://youtu.be/wzALbzTdnc8

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

http://www.youtube.com/watch?v=PLpZfJ4rGts

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

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

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

http://www.youtube.com/watch?v=QI50GBUJ48s

slide42
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

http://www.upscale.utoronto.ca/PVB/Harrison/BohrModel/Flash/BohrModel.html

remember
REMEMBER

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

kernel
Kernel
  • 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.
slide51
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?

slide52

Relationship of Frequency, Wavelength and Energy of colored line

http://employees.oneonta.edu/viningwj/sims/plancks_equation_s.html

good overview videos
Good Overview Videos
  • Crash Course: History of Atomic Theory
  • http://www.youtube.com/watch?v=thnDxFdkzZs
  • Quantum Mechanics and the Bohr Model
  • http://www.youtube.com/watch?v=-YYBCNQnYNM
calculating the energy of a spectral line honors
Calculating the Energy of a Spectral Line (HONORS)

STEP 1:

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)

STEP 2:

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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