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Where are the electrons ?

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Title: Where are the electrons ?


1
Where are the electrons ?
  • Rutherford found the nucleus to be in the center.
  • He determined that the atom was mostly empty
    space.
  • So, how are the electrons arranged in that space?
  • This was the shortcoming, of Rutherford's model
    of the atom, the electron position could not be
    explained.

2
What is the electromagnetic spectrum?
  • brainstorm

3
Electromagnetic radiation
  • Is a form of energy that exhibits wavelike
    behavior as it travels through space.
  • Electromagnetic spectrum all the forms of
    electromagnetic radiation.

4
Electromagnetic spectrum
5
Electromagnetic spectrum
  • Each line on the spectrum represents a certain
    wave frequency of radiation.
  • Each wave frequency is associated with a certain
    amount of energy.

6
Electromagnetic spectrum
  •         The elecromagnetic spectrum includes all
    forms of radiation, one of which is visible light
    -- the radiation to which our eyes are sensitive.
  •         Gamma-rays, X-rays, ultraviolet,
    infrared and radio waves are also forms of
    radiation. We divide the spectrum up according to
    the wavelength of the radiation.

7
Electromagnetic spectrum
8
Electromagnetic spectrum
9
A wave is described by
  • Wavelength - the distance between corresponding
    points on adjacent waves
  • Frequency- the number of waves that pass a given
    point in a specific time, usually one second.
    Unit is the hertz, (Hz).
  • http//www.colorado.edu/physics/2000/waves_particl
    es/

10
Wave description
  • Mathematically c speed 3 x108 m/s
  • Speed of light frequency x wavelength
  • Since c is constant
  • Frequency is inversely proportional to
    wavelength
  • All three, speed, wavelength, frequency are
    related

11
example
  • What is the frequency of a wave that has a
    wavelength of 200 nm?
  • 3 x 108/ 200 nm

12
1900s
  • Wave model of light was accepted

13
Early 1900s Photoelectric experiments
  • Experiment - Planck
  • Photoelectric effect refers to the emission of
    electrons from a metal when light shines on the
    metal.
  • Shined light on a metal varying the frequency of
    the light. Below a certain frequency the
    electrons were not emitted.
  • When the frequency was high enough there was
    enough energy to knock electrons loose from the
    metal.

14
Quantum
  • Planck suggested that an object emits energy in
    little packets of energy called a quantum.
  • Quantum minimum quantity of energy that can be
    lost or gained by an electron.
  • Energy that is given off is related to wave
    frequency.

15
Go to
  • http//www.colorado.edu/physics/2000/quantumzone/l
    ines2.html
  • http//www.colorado.edu/physics/2000/waves_particl
    es/

16
Planck in other words.
  • Proposed that there is a fundamental restriction
    on the amount of energy that an object emits or
    absorbs each piece of energy is a quantum.
  • E hv
  • h 6.6262 x 10-34 Js
  • v the frequency of the wave

17
Dual wave-particle nature of light
  • Einstein expanded on Plancks theory and said
    that electromagnetic radiation has a dual
    wave-particle nature.
  • While light exhibits wave like properties, it can
    also be thought of as a stream of particles.
  • Each particle carries a quantum of energy these
    particles are called photons.

18
Vocab
  • Photon particle of electromagnetic radiation
    having zero mass and carrying a quantum of
    energy.
  • Ephoton hv
  • Quantum the minimum quantity of energy that can
    be gained or lost by an electron.

19
Photoelectric effect explained
  • Electromagnetic radiation is only absorbed in
    whole number of photons.
  • In order for an electron to be bumped off, it
    must be struck by a photon of a certain minimal
    amount of energy.
  • The minimal energy corresponds to a minimal
    frequency

20
continued
  • Different metals require energy at different
    frequencies to exhibit a photoelectric effect.
  • Each metal has a certain required minimal level
    of energy required for the electrons to be
    knocked loose.

21
Ground state
  • The lowest possible energy level
  • Close to the nucleus
  • The lowest energy state of an atom

22
Excited state
  • When an atom absorbs energy and the electrons
    move to a higher energy level.
  • An atom has a higher potential energy than its
    ground state.

23
Like a ladder
  • Energy levels are like the rungs on a ladder, you
    can not step between the rungs. Electrons must
    jump from level to level, they can not reside
    between the levels

24
electrons
  • When an electron gains enough energy it jumps up
    an energy level ( rung on a ladder), and becomes
    excited.
  • It then immediately returns to its ground state.
  • The energy is released as a particular wavelength
    that corresponds to a particular color. A photon
    of radiation.

25
Line emission spectra
  • The spectra given off by the electrons of a
    certain element.
  • Each spectra is unique to the element or
    compound.
  • Line spectra are used to identify elements and/or
    compounds

26
Line emission spectra
  • The fact that specific frequencies are emitted
    indicates that the energy differences between an
    atoms energy states are fixed
  • Each line on the spectra represents a photon of
    energy.

27
Bohr model of atom
  • The electron circles the nucleus in an orbital.
  • When the electron gains enough energy ( a certain
    photon) it jumps to a higher level orbital.
  • When it returns to ground state it emits the
    energy (photon). The frequency of the photon is
    seen in the spectra.

28
Bohr successfully calculated the hydrogen line
spectra
  • Bohrs model of the hydrogen atom accounted
    mathematically for the energy of each of the
    transitions of the Lyman, Balmer and Paschen
    spectral series.
  • Bottom line Bohrs model of the H atom, with
    the line spectra, could be mathematically
    supported.
  • summary and beyond
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