Arrangement of Electrons in Atoms

1
Arrangement of Electrons in Atoms

• Development of a New Atomic Model

2
Wave Description Of Light

• Electromagnetic Radiation
• form of energy that exhibits wavelike behavior as
  it travels through space.
• EX visible light, X-ray, Ultraviolet and
  inferred light, microwaves, and radio waves.
• Travels at a constant speed of 3.0 x 108 m/s
• Electromagnetic Spectrum All the electromagnetic
  radiation form the ES. (fig 4-1, p. 92)

3
Electromagnetic Spectrum
4
Wave Calculations

• Wavelength (?) - distance between two peaks .
  Measured in meters
• Frequency (v) - number of peaks that pass a point
  each second.
• Hz Hertz s-1
• c ? v       
• where c 3.0 x 108 m/s

5
Is light really a wave?

• Max Planck did experiments with light-matter
  interactions where light did not act like a wave
• Photoelectric Effect - emission of electrons from
  a metal when light shines on the metal.
• Only emitted at certain energies wave theory
  said any energy should do it.
• Led to the particle theory of light

6

• Planck suggested that objects emit energy in
  specific amounts called QUANTA
• Quantum - minimum quantity of energy that can be
  lost or gained by an atom.
• led Planck to relate the energy of an electron
  with the frequency of EMR            
• E hv           
• E Energy (J, of a quantum of radiation)
• v frequency of radiation emitted
• h Plancks constant (6.626 x 10-34 Js)

7
Equation Practice

• What is the frequency of yellow light with a
  wavelength of 548 nm?

8
Equation Practice

• What is the wavelength of blue light with a
  frequency of 4.60 x 1023 Hz?

9
Equation Practice

• What is the energy of magenta light with a
  wavelength of 691 nm?

10

• leads to Einsteins dual nature of light (EMR
  behaves as both a wave and a particle)
• Photon - particle of EMR having zero mass and
  carrying a quantum of energy.

11
Hydrogen Emission Spectrum

• Ground State - Lowest energy state of electron.
• Excited State - higher energy than ground state.
• Bright-line Spectrum (emission spectrum)
• Series of specific light frequencies emitted by
  elements
• "spectra are the fingerprints of the elements"

12
The Development of A New Atomic Model

• Rutherfords model was an improvement over
  previous models, but still incomplete.
• Where exactly are electrons located?
• What prevented the electrons from being drawn
  into the nucleus?

13
Bohr Model Of H Atom

• Bohr explained how the electrons stay in the
  cloud instead of slamming into the nucleus
• Definite orbits paths
• The greater the distance from the nucleus, the
  greater the energy of an electron in that shell.

14

• Electrons start in lowest possible level - ground
  state.
• Absorb energy - become excited and shift upward.
• Dropping back down - emits photons (packets of
  energies equal to the previously absorbed
  energy).
• Hydrogen Emission Spectrum

15
Quantum Model of the Atom

• Bohrs model was great, but it didnt answer the
  question why?
• Why did electrons have to stay in specific
  orbits?
• Why couldnt the electrons exist anywhere within
  the electron cloud?
• Louis de Broglie pointed out that electrons act
  like waves
• Using Plancks equation (Ehv), dB proved that
  electrons can have specific energies and that
  Bohrs quantized orbits were actually correct

16
Heisenberg Uncertainty Principle

• Impossible to determine both the exact location
  and velocity of an electron

17
Schrodinger Wave Equation

• He gave more support to Bohrs quantized energy
  levels
• Quantum theory describes the wave properties of
  electrons using mathematical equations
• Disproved Bohrs train tracks within those
  energy levels