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

Development of Atomic Models


Democritus

Democritus

  • Greek philosopher

  • 400 BC

  • “Atomos” concept


Development of atomic models

  • 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.


Page from dalton s journal

Page from Dalton’s Journal


Development of atomic models

  • 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.


Development of atomic models

Dalton’s Atomic Theory called

“Hard Spheres Model”


Jj thomson 1897

JJ Thomson (1897)


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


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.


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


Development of atomic models

Observation:

Most alpha particles passed straight

through gold foil.

Conclusion:

Atom’s volume is mostly empty space.


Development of atomic models

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


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

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”)


Spectral lines are different for each element

Spectral Lines are Different for Each Element


Development of atomic models

Answer: 1


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


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

    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: Argon2-8-8, 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:

    • Al2-8-33 valence

    • Ca2-8-8-22 valence

    • F2-77 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 2-8-3has 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

  • 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?


Development of atomic models

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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