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Arrangement of Electrons in Atoms

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No medium is required. Electromagnetic Radiation ... If the electron is circling the nucleus then why doesn't it emit radiation and ... – PowerPoint PPT presentation

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Title: Arrangement of Electrons in Atoms


1
Arrangement of Electrons in Atoms
2
Development of a New Atomic Model
3
Light as a wave.
4
Electromagnetic Radiation
  • A form of energy
  • Exhibits wavelike characteristics
  • No medium is required

5
Electromagnetic Radiation
  • All Electromagnetic Radiation travels at the same
    rate
  • 3.00 x 108 m/s

6
Electromagnetic Radiation
7
Electromagnetic Radiation
  • Wavelength
  • Frequency
  • Amplitude
  • Crest
  • Trough

8
Wavelength and Frequency
  • c ??
  • c speed of light
  • ? wavelength
  • ? frequency

9
Light as a particle.
10
Problems with classical physics.
  • Blackbody radiation.
  • The photoelectric effect.
  • Stability of the atom.

11
Blackbody radiation.
  • Heating a piece of metal changes the color from
    red, to orange, to white.
  • Classical physics should see all colors
    (energies)
  • Could not explain why only certain colors showed
    up.

12
Photoelectric Effect
  • Yellow light shone on certain metals
  • Nothing happens
  • Really bright yellow light (more energy)
  • Nothing happens

13
Photoelectric Effect
  • Blue light shone on certain metals
  • Electrons given off
  • Really bright blue light
  • More electrons given off but same speed

14
Stability of the atom.
  • Why isnt the electron attracted to the nucleus?
  • If the electron is circling the nucleus then why
    doesnt it emit radiation and fall into the
    nucleus?

15
Classical Physics revised.
16
Blackbody Radiation
  • Max Planck energy delivered in packets,
    quantum.
  • If energy is delivered in packets, then
    blackbodies gain energy in packets.
  • Only packets of certain frequencies are delivered.

17
Energy and Frequency
  • Energy and frequency are related.

18
Energy and Frequency
  • E h?
  • E Energy, Joules (J)
  • h Planck constant, 6.626x10-34 Js
  • ? Frequency, Hertz (Hz)

19
The Photoelectric Effect
20
Einsteins Theory
  • Einstein theorized the dual nature of light.
  • Used the two-slit experiment to verify the
    particle and wave nature of light.

21
Photoelectric Effect
  • Einstein proposed that light should be both a
    particle and a wave.
  • Coined the term photon (a quantum of light).
  • Only certain photons have enough energy

22
Why talk about waves?
  • De Broglie determined that all matter behaves as
    waves.
  • Electrons are matter.
  • Therefore, electrons behave as waves and as
    particles both.

23
Stability of the Atom
  • Explained by Niels Bohr using experiments by
    others.
  • Looked at the emission line spectrum published by
    Lyman, Balmer, and Paschen

24
Line spectrum.
  • Hydrogen only gave a line spectrum when passed
    through a prism.
  • Thought that any amount of energy should excite
    the Hydrogen and therefore a continuous spectrum
    should be given.

25
Lyman, Balmer, Paschen
  • Investigated Hydrogen gas excited in a tube.
  • Passed light through a prism.
  • The H2 light emitted a line emission spectrum.
  • The line emission spectrum is the Lyman, Balmer,
    Paschen Series.

26
Bohr Model of the Atom
  • Bohr proposed that electrons exist on quantized
    orbits
  • The orbits have a defined amount of energy

27
Bohr Model of the Atom
  • Ground State lowest energy state
  • Excited State gained energy
  • Must gain energy to get to an excited state

28
Bohr Model of the Atom
  • Explains Lyman, Balmer, Paschen
  • Electron gains energy and jumps orbits
  • Electron is unstable
  • Electron drops back to ground state, giving off
    energy (light)

29
Bohr Model of the Atom
  • Since the orbits have only certain energies, only
    certain wavelengths exist.
  • Also mathematically verified theory.

30
Bohr Model of the Atom.
  • Proposed Orbits
  • When the electron is in an orbit, it has a
    defined amount of energy.
  • The electron can not exist between orbits.

31
Bohr Model of the Atom
  • The electron must absorb energy to move orbits.
  • The electron emits the same amount of energy as
    it absorbs.
  • The energy difference is the difference between
    energy levels.

32
Bohr Model of the Atom
  • 1st orbit is a certain distance from the nucleus.
  • The higher the orbit, the more energy.

33
Bohr Model of the Atom
  • Orbits are not equally spaced
  • Since the orbit has inherent energy, the electron
    has energy from the beginning.
  • Only explains single electron atoms

34
The Quantum Model of the Atom
35
Electron as a wave De Broglie
  • All matter behaves as waves.
  • Electron exhibits wavelike properties.

36
Electron as a wave De Broglie
  • The electron is a wave confined to the nucleus.
  • The electron can be bent (diffraction) like a
    wave.
  • The electron can cancel itself out (interference)
    like a wave

37
Electron as a wave Heisenberg
  • Heisenberg Uncertainty Principle
  • It is impossible to know the position and
    velocity of an electron at the same time.
  • Once you measure it, you change it.

38
Electron as a wave Schrodinger
  • Developed Schrodinger wave equation.
  • Only waves of specific energy solves the equation

39
Electron as a wave Schrodinger
  • Paved the way for Quantum Theory
  • Quantum Theory describes mathematically the
    movement of very small things.
  • Solving Schrodinger wave equation gives quantum
    numbers.

40
Quantum Model and Electron Configuration.
41
Quantum Model
  • Has Orbitals
  • An orbital is a fuzzy region of probable electron
    placement (90)
  • Orbitals have distinct shapes dependent upon
    amount of energy of the electron.

42
Quantum Model
  • Orbital shapes are given symbols
  • s, p, d, f, etc.

43
Quantum Model has rules for electron placement.
44
Aufbau Principle
  • Lazy mans rule.
  • Aufbau Principle - electron will occupy the
    lowest energy orbital available.

45
Hunds Rule
  • Summer camp rule.
  • Orbitals of equal energy are each occupied by one
    electron before any orbital is occupied by a 2nd
    electron.

46
Pauli Exclusion Principle
  • No two electrons in an atom can have the same 4
    quantum numbers
  • Quantum numbers describe the placement of
    electrons in an orbital
  • Only allows for 2 electrons per orbital shape

47
Electron configuration.
  • Use the aforementioned rules to determine
    electron configuration notation.

48
Orbital Notation
  • Uses arrows and lines to show electrons and
    orbitals

49
Energy Levels
  • Highest occupied energy level is the highest
    energy level with an electron.
  • Ex. 1s22s22p63s23p64s23d7
  • Inner shell electrons are all of the electrons
    not on the outside.
  • Ex. 1s22s22p63s23p64s23d7

50
Noble Gas Configuration
  • A Noble gas configuration has the outermost s and
    p orbitals filled.
  • Example 1s22s22p6, 1s22s22p63s23p6
  • Noble Gas Shortcut - use the Noble gas before the
    element (except He).
  • Example Br Ar4s23d104p5
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