ELECTRONS IN ATOMS

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ELECTRONS IN ATOMS

• Chapter 13
• CHEMISTRY

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Atomic Models

• John Dalton 1766-1844
• The atom is considered a solid indivisible mass.
• JJ Thomson 1856-1940
• Plum pudding model, nothing about number of
  electrons or protons, arrangements or loss of
  electrons.
• Ernest Rutherford 1871-1937
• Atom is mostly empty space with a dense nucleus.
• Niels Bohr 1885-1962
• Electrons travel in definite orbits around
  nucleus.

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Bohr Model

• Sometimes called the Planetary Model.
• 1. Electrons are arranged in concentric paths,
  orbitals, around the nucleus.
• 2. Electrons in a certain orbit have a fixed
  amount of energy and cannot lose energy and fall
  into the nucleus. (rungs on a ladder)
• 3. To move from one level to another, an
  electron must gain just the right of energy.

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Quantum

• The amount of energy required to move an electron
  from its present energy level to the next higher
  one.
• The higher on the ladder the more energy the
  electron has and the farther it is from the
  nucleus.
• Not all steps rungs are not always the same
  amount of energy. Steps become smaller as you
  move up the ladder.
• The energy of electrons is thus said to be
  quantized.

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The Quantum Mechanical Model

• 1926 Erwin Schrodinger
• Schrodingers Equation describes the energy and
  location of electrons in the hydrogen atom.
• This became the quantum mechanical model, truly a
  mathematically based model and not one based on
  physical models involving the motion of large
  objects.

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The Quantum Mechanical Model

• Energy of electrons is restricted to certain
  values.
• Does not define an exact path an electron takes
  around an nucleus.
• Estimates the probability of finding an electron
  in a certain location.
• The probability of finding an electron within a
  certain volume of space surrounding the nucleus
  can be represented by a fuzzy cloud.

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Electron Clouds

• Attempts to show probabilities as a fuzzy cloud
  are usually limited to the volume in which the
  electron is found 90 of the time.
• It is unclear just where a cloud ends, there is
  at least a slight chance of finding the electron
  a considerable distance from the nucleus.

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Atomic Orbitals

• Principal quantum number (n) designates energy
  levels that electrons occupy.
• N 1, 2, 3, 4, 5, .
• As the number increases, so does the energy.
• The average distance of the electron from the
  nucleus increases as n increases.
• Within each energy level, electrons occupy energy
  sublevels.

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Energy Sublevels

• The number of sublevels is the same as the
  principal quantum number.
• In energy level 1, there is one sublevel (1s).
• In energy level 2, there are two sublevels, (2s
  2p).
• s, p, d, and f, are used to designate different
  sublevels.
• These regions in which electrons are likely to be
  found are called atomic orbitals.

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Atomic Orbitals

• Letters denote the atomic orbitals.
• s orbitals are spherical
• p orbitals are dumbbell-shaped
• These orbitals have different orientations in
  space px, py, pz
• d orbitals are cloverleaf in shape (4 of 5).
• f orbitals are very complex and hard to
  visualize.
• Regions of the p and d orbitals close to the
  nucleus show a very low probability of finding
  the electron and are called nodes.

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Number of Electrons

• The number of electrons allowed in each energy
  level depends on the number of orbitals.
• Level 1 1 sublevel 1 orbital 2 electrons
• Level 2 2 sublevel 4 orbitals 8 electrons
• Level 3 3 sublevels 9 orbitals 18 electrons
• Level 4 4 sublevels 16 orbitals 32 electrons

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Electron Arrangement in Atoms

• Electron Configurations the ways in which
  electrons are arranged around the nuclei of
  atoms.
• Three rules control this
• Aufbau principle
• Pauli exclusion principle
• Hunds rule

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Aufbau Principle

• Electrons enter orbitals of lowest energy first
• 1. Orbitals within a sublevel are always of
  equal energy
• 2. s sublevel is always the lowest-energy
  sublevel
• 3. The range of energy levels within a
  principal energy level can overlap the energy
  levels of an adjacent principal level

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Pauli Exclusion Principle

• An atomic orbital may describe at most two
  electrons.
• One or two electrons may occupy an orbital.
• In order for two electrons to occupy the same
  orbital they must have opposite spins, clockwise
  counterclockwise. This is signified with up and
  down arrows.

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Hunds Rule

• When electrons occupy orbitals of equal energy,
  one electron enters each orbital until all
  orbitals contain one electron with parallel
  spins.
• Second electrons then add to each orbital so
  their spins are paired.

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Electron Configurations

• A convenient shorthand method is used to
• show the electron configuration of an atom.
• A number signifies the principal energy level, a
  letter to indicate the sublevel, and a
  superscript to indicate the number of electrons
  occupying that sublevel.
• 1s22s22p5
• The sum of superscripts equals the number of
  electrons in the atom.

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Exceptional Configurations

• Copper and chromium are highlighted on page 370
  in text.
• Points of Emphasis
• Configurations are to show most stable situation
  possible.
• Filled sublevels are most stable
• Half-filled sublevels are the next most stable
• Partially filled are not as stable as half-filled.

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Light and Atomic Spectra

• According to the wave model
• -light consists of electromagnetic waves
• -electromagnetic radiation includes radio waves
  to gamma rays
• -all electromagnetic waves travel at the speed of
  light in a vacuum
• 3.0 x 108 m/s

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Wave Terminology

• Wave cycle begins and ends at the wave origin.
• Amplitude is the waves height from origin to
  crest.
• Wavelength, (?, lambda), is the distance between
  crests.
• Frequency, (f), is the number of wave cycles to
  pass a certain point per unit time.

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Wave Terminology

• Frequency and wavelength are inversely related.
• c (speed of light) ?f
• Units are cycles per second or hertz (Hz). This
  is sometimes expressed simply as s-1.

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Electromagnetic Spectrum

• Visible window (10-7 and 1015 Hz range) is
  quite small.
• When sunlight passes through a prism or
  diffraction grating, the different wavelengths
  separate into different colors.
• ROY G. BIV
• Red, orange, yellow, green, blue, indigo, violet
• Red is longest ? and lowest frequency.

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Atomic Emission Spectra

• Element emits light when it is excited by the
  passage of an electric discharge through its
  vapor or gas.
• Atoms first absorb energy, and then lose the
  energy as they emit light.
• Emission spectra consist of bright lines.
• Each line corresponds to one exact frequency of
  light emitted by the atom.
• Each line corresponds to a specific amount of
  energy being emitted by the atom.

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Quantum Concept Photoelectric Effect

• Although classical physics says that atomic
  spectra should be continuous, we see only lines.
• Max Planck stated that as a substance heats up,
  changes in energy are small discrete units
  (quanta).
• Mathematically speaking
• E h f
• h is called Plancks constant 6.6262 x 10-34
  Js
• Energy absorbed or emitted is proportional to the
  frequency of the radiation.

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Quantum Concept Photoelectric Effect

• The size of an emitted or absorbed quantum
  depends on the size of the energy change.
• Small energy change involves the emission
  absorption of low frequency radiation.
• Large energy change involves high frequency
  radiation.
• Ordinary thermometers are not sensitive enough to
  measure these small changes.

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Quantum Concept Photoelectric Effect

• Albert Einstein proposed this same particle
  nature to light.
• Light quanta are called photons.
• The energy of photons is found using the same
  equation developed by Planck.
• Although this dual wave-particle behavior of
  light was difficult to accept, it does help to
  explain the photoelectric effect.

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Photoelectric Effect

• Classical physics could not explain this because
  shining a light on a photoelectric material
  should release electrons if only applied long
  enough.
• Only very specific frequencies of light seem to
  work.
• billiard ball vs ping pong ball vs golf ball

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An Explanation of Atomic Spectra

• Energy absorbed energy emitted
• Three groups of lines in hydrogen spectrum
  correspond to the transition of electrons from
  higher levels to lower levels.
• Lyman series level 1 ultraviolet
• Balmer series level 2 visible
• Paschen series level 3 infrared

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Quantum Mechanics

• Louis de Broglie
• If light can behave as wave and particle, can
  particles of matter behave as waves?
• De Broglies Equation
• ? h / mv
• Electrons wavelength vs baseballs wavelength

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Classical vs Quantum Mechanics

• Classical mechanics adequately describes the
  motions of bodies much larger than the atoms that
  they comprise. It appears that energy loss/gain
  can be in any amount.
• 2. Quantum mechanics describes the motions of
  subatomic particles and atoms as waves. These
  particles gain/lose energy in packages called
  quanta.

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Heisenberg Uncertainty Principle

• It is impossible to know exactly both the
  velocity and the position of a particle at the
  same time.
• Need photons to see things. Photons carry so
  much energy that when they strike an electron its
  velocity must change.

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• Hopefully this presentation has shed some light
  on the subject!!