Electrons in Atoms

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Chapter 5

• Electrons in Atoms

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5.1 Light and Quantized Energy

• Electromagnetic Radiation-a form of energy that
  exhibits wavelike behavior as it travels through
  space
• Radio waves
• Microwaves
• Infrared
• Visible light
• Ultraviolet light
• X-rays
• Gamma rays

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• All electromagnetic waves travel at the speed of
  light in a vacuum (3.00 x 108 m/s).
• Crest-the highest point of a wave
• Trough-the lowest point of a wave
• Wavelength(?)-the distance from crest to crest
• Frequency(?)-the number of waves that pass a
  certain point in one second
• Amplitude-the height from the origin in to the
  crest

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• Wavelength and frequency have an inverse
  relationship.
• c ? x ?

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Electromagnetic Spectrum
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Review

• Draw a wave and label the parts including crest,
  amplitude, frequency, wavelength, trough.
• Which waves travel at the speed of light in a
  vacuum?
• As the wavelength decreases, what happens to the
  energy and frequency of the wave?

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• Plancks Theory
• There is a fundamental restriction on the amount
  of energy that an object emits or absorbs
• energy emitted or absorbed is restricted to
• pieces(quantum)
• Quantum-fixed amount
• ? h x ?
• ? energy
• ? frequency
• h Plancks constant 6.6262 x 10-34 J-s

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• Einstein used Plancks equation to explain the
  photoelectric effect.
• Einstein proposed that light consist of quanta of
  energy that behave like tiny particles (photons).

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Review

• Explain the photoelectric effect.
• What is Plancks equation?

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• Atomic Emission Spectrum/Line Spectrum-contains
  only certain colors, or wavelengths. Distinct
  lines appear when you pass light through a prism.
• All elements emit light when they are vaporized
  in an intense flame or when electricity passes
  through their gaseous state.
• Atoms absorb then release energy in the form of
  light.
• Every element emits light containing only certain
  wavelengths( line spectrum, color).
• The energy of radiation increases with the
  frequency.

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Review

• Describe the process of observing an atomic
  emission spectrum.

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5.2 Quantum Theory and the Atom

• Neils Bohr Atomic Model
• Described the atom as electrons moving around the
  nucleus in well defined orbits .
• The smaller the orbit, the lower the energy.
• The larger the orbit, the higher the energy.
• n1, n2, n3, n4
• Ground State-the lowest energy state of an atom.
• Bohr tried to explain the Hydrogen spectrum
  decided that e- could only exist in specific
  energy levels. These levels are quantized.

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• When electrons in the ground state gain energy,
  they can move a greater distance from the nucleus
  to an excited state.
• When the electrons lose the gained energy they
  fall back to the ground state and release the
  energy in the form of radiation.
• ROYGBIV
• Lower energy ---------------------? Higher Energy
• A quantum of energy the amount of energy
  needed to move an electron from one energy level
  to the next level.

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• Louis de Broglie - predicted all moving particles
  have wave characteristics.
• Heisenberg Uncertainty Principle it is
  impossible to know precisely the velocity and
  position of a particle at same time.

Energy Level n 1 n 2 n 3 n 4
of e- 2 8 18 32
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• Erwin Schrödinger
• Quantum Mechanical Model of the Atom
• An atomic model in which the electrons are
  treated as waves. Rather than describing the
  electron orbiting the nucleus, he describes the
  probability of finding an electron in a given
  place (electron cloud).
• Atomic Orbital-a three dimensional region around
  the nucleus. 4 different orbitals s, p, d, f.
• Ex.
• Carbon

4 e-
6 p 6 n
2 e-
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• Principal Energy Levels
• Energy levels of electrons are designated by
  principal quantum numbers (n).
• n1, n2, n3, n4
• The principal energy levels have sublevels.
• The principal energy levels quantum number is
  equal to the number of sublevels.
• These sublevels are called s, p, d, and f.
• n1 has 1 sub level
• 1s
• n2 has 2 sub levels
• 2s 2p
• n3 has 3 sub levels
• 3s, 3p, 3d

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• The sublevels are divided into atomic orbitals.
• Each sublevel has a certain number of orbitals.
• Each orbital can have a maximum of 2 electrons
  in it.
• All s sublevels have 1 orbital.
• All p sublevels have 3 orbitals.
• All d sublevels have 5 orbitals.
• All f sublevels have 7 orbitals.

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• Principal Energy Levels ( n1, n2, n3)
• Sub Levels (s, p, d, f)
• Orbitals (s1 orbital, p3 orbitals, d5
  orbitals, f7 orbitals)
• Increasing energy
• (increasing distance from nucleus)
• ???????????????
• Energy level n 1 2 3 4
• Max e- 2 8
  18 32

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Sublevels
Type of Sublevel of orbitals of electrons
s 1 2
p 3 6
d 5 10
f 7 14
n 1 1 sublevel s n 2 2 sublevels s, p n
3 3 sublevels s, p, d Electron spin spin up or
spin down.
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Review

• What is the maximum principal energy level?
• Name the sublevel orbitals the maximum number
  of electrons each orbital can hold.

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5.3 Electron Configurations

• 1s
• 2s 2p
• 3s 3p 3d
• 4s 4p 4d 4f
• 5s 5p 5d 5f
• 6s 6p 6d
• 7s 7p

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Electron configuration-the way in which electrons
are arranged around the nucleus.

• 1s22s22p63s2
• 1s22s22p63s23p64s23d104p6

Principal Energy Level
12 Mg 24.02
of electrons
Sublevel
36 Kr 83.80
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• Rules for Orbital Diagrams
• Aufbau principle-electrons enter orbitals of
  lower energy first.
• Pauli exclusion principle-an atomic orbital has
  atmost two electrons with opposite spin.
• Hunds rule-when electrons occupy orbitals of
  equal energy, one electron enters each orbital
  until all the orbitals contain one electron with
  parallel spins.
• Mg ?? ?? ?? ?? ?? ??
• 1s 2s 2p
  3s

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• Valence Electrons-electrons in an atoms
  outermost orbitals
• Group Group Name Valence e-
• 1 Alkali Metals 1
• 2 Alkaline Earth Metals 2
• 3-12 Transition Metals
  varies
• 13 Boron Group 3
• 14 Carbon Group 4
• 15 Nitrogen Group
  5
• 16 Oxygen Group 6
• 17 Halogens 7
• 18 Noble Gases
  8

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Abbreviated or Noble Gas Configuration

• Ca 1s22s22p63s23p6 4s2
• Ar 1s22s22p63s23p6
• Ca Ar 4s2

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• Electron Dot Structures
• Elements symbol represents the atomic nucleus
  and inner electrons
• Dots are used to represent the valence electrons
• Li O C Mg Ne

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Review

• Write the electron configuration for Chlorine in
  the standard form and in the noble gas
  configuration.
• What are valence electrons?
• What are the rules for electron configuration?
• Draw the Lewis dot structure for Na.