chemistry

1
chemistry
Chapter 5
2
Models of the Atom
5.1

• The scale model shown is a physical model.
  However, not all models are physical. In fact,
  several theoretical models of the atom have been
  developed over the last few hundred years. You
  will learn about the currently accepted model of
  how electrons behave in atoms.

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The Development of Atomic Models
5.1

• The Development of Atomic Models
• What was inadequate about Rutherfords atomic
  model?

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The Development of Atomic Models
5.1

• Rutherfords atomic model could not explain the
  chemical properties of elements.
• Rutherfords atomic model could not explain why
  objects change color when heated.

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The Development of Atomic Models
5.1

• The timeline shoes the development of atomic
  models from 1803 to 1911.

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The Development of Atomic Models
5.1

• The timeline shows the development of atomic
  models from 1913 to 1932.

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The Bohr Model
5.1

• The Bohr Model
• What was the new proposal in the Bohr model of
  the atom?

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The Bohr Model
5.1

• Bohr proposed that an electron is found only in
  specific circular paths, or orbits, around the
  nucleus.

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The Bohr Model
5.1

• Each possible electron orbit in Bohrs model has
  a fixed energy.
• The fixed energies an electron can have are
  called energy levels.
• A quantum of energy is the amount of energy
  required to move an electron from one energy
  level to another energy level.

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The Bohr Model
5.1

• Like the rungs of the strange ladder, the energy
  levels in an atom are not equally spaced.
• The higher the energy level occupied by an
  electron, the less energy it takes to move from
  that energy level to the next higher energy level.

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The Quantum Mechanical Model
5.1

• The Quantum Mechanical Model
• What does the quantum mechanical model determine
  about the electrons in an atom?

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The Quantum Mechanical Model
5.1

• The quantum mechanical model determines the
  allowed energies an electron can have and how
  likely it is to find the electron in various
  locations around the nucleus.

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The Quantum Mechanical Model
5.1

• Austrian physicist Erwin Schrödinger (18871961)
  used new theoretical calculations and results to
  devise and solve a mathematical equation
  describing the behavior of the electron in a
  hydrogen atom.
• The modern description of the electrons in atoms,
  the quantum mechanical model, comes from the
  mathematical solutions to the Schrödinger
  equation.

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The Quantum Mechanical Model
5.1

• The propeller blade has the same probability of
  being anywhere in the blurry region, but you
  cannot tell its location at any instant. The
  electron cloud of an atom can be compared to a
  spinning airplane propeller.

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The Quantum Mechanical Model
5.1

• In the quantum mechanical model, the probability
  of finding an electron within a certain volume of
  space surrounding the nucleus can be represented
  as a fuzzy cloud. The cloud is more dense where
  the probability of finding the electron is high.

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Atomic Orbitals
5.1

• Atomic Orbitals
• How do sublevels of principal energy levels
  differ?

17
Atomic Orbitals
5.1

• An atomic orbital is often thought of as a region
  of space in which there is a high probability of
  finding an electron.
• Each energy sublevel corresponds to an orbital of
  a different shape, which describes where the
  electron is likely to be found.

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Atomic Orbitals
5.1

• Different atomic orbitals are denoted by letters.
  The s orbitals are spherical, and p orbitals are
  dumbbell-shaped.

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Atomic Orbitals
5.1

• Four of the five d orbitals have the same shape
  but different orientations in space.

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Atomic Orbitals
5.1

• The numbers and kinds of atomic orbitals depend
  on the energy sublevel.

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Atomic Orbitals
5.1

• The number of electrons allowed in each of the
  first four energy levels are shown here.

22
Atomic Orbitals

• Animation 5
• Observe the characteristics of atomic orbitals.

23
Electron Arrangement in Atoms
5.2

• If this rock were to tumble over, it would end up
  at a lower height. It would have less energy than
  before, but its position would be more stable.
  You will learn that energy and stability play an
  important role in determining how electrons are
  configured in an atom.

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Electron Configurations
5.2

• Electron Configurations
• What are the three rules for writing the electron
  configurations of elements?

25
Electron Configurations
5.2

• The ways in which electrons are arranged in
  various orbitals around the nuclei of atoms are
  called electron configurations.
• Three rulesthe aufbau principle, the Pauli
  exclusion principle, and Hunds ruletell you how
  to find the electron configurations of atoms.

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Electron Configurations
5.2

• Aufbau Principle
• According to the aufbau principle, electrons
  occupy the orbitals of lowest energy first. In
  the aufbau diagram below, each box represents an
  atomic orbital.

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Electron Configurations
5.2

• Pauli Exclusion Principle
• According to the Pauli exclusion principle, an
  atomic orbital may describe at most two
  electrons. To occupy the same orbital, two
  electrons must have opposite spins that is, the
  electron spins must be paired.

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Electron Configurations
5.2

• Hunds Rule
• Hunds rule states that electrons occupy orbitals
  of the same energy in a way that makes the number
  of electrons with the same spin direction as
  large as possible.

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Electron Configurations
5.2

• Orbital Filling Diagram

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Electron Configurations

• Simulation 2
• Fill atomic orbitals to build the ground state of
  several atoms.

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for Conceptual Problem 1.1
Problem Solving 5.9 Solve Problem 9 with the help
of an interactive guided tutorial.
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Exceptional Electron Configurations
5.2

• Exceptional Electron Configurations
• Why do actual electron configurations for some
  elements differ from those assigned using the
  aufbau principle?

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Exceptional Electron Configurations
5.2

• Some actual electron configurations differ from
  those assigned using the aufbau principle because
  half-filled sublevels are not as stable as filled
  sublevels, but they are more stable than other
  configurations.

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Exceptional Electron Configurations
5.2

• Exceptions to the aufbau principle are due to
  subtle electron-electron interactions in orbitals
  with very similar energies.
• Copper has an electron configuration that is an
  exception to the aufbau principle.

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Physics and the Quantum Mechanical Model
5.3

• Neon advertising signs are formed from glass
  tubes bent in various shapes. An electric current
  passing through the gas in each glass tube makes
  the gas glow with its own characteristic color.
  You will learn why each gas glows with a specific
  color of light.

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Light
5.3

• Light
• How are the wavelength and frequency of light
  related?

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Light
5.3

• The amplitude of a wave is the waves height from
  zero to the crest.
• The wavelength, represented by ? (the Greek
  letter lambda), is the distance between the
  crests.

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Light
5.3

• The frequency, represented by ? (the Greek letter
  nu), is the number of wave cycles to pass a given
  point per unit of time.
• The SI unit of cycles per second is called a
  hertz (Hz).

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Light
5.3

• The wavelength and frequency of light are
  inversely proportional to each other.

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Light
5.3

• The product of the frequency and wavelength
  always equals a constant (c), the speed of light.

44
Light
5.3

• According to the wave model, light consists of
  electromagnetic waves.
• Electromagnetic radiation includes radio waves,
  microwaves, infrared waves, visible light,
  ultraviolet waves, X-rays, and gamma rays.
• All electromagnetic waves travel in a vacuum at a
  speed of 2.998 ? 108 m/s.

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Light
5.3

• Sunlight consists of light with a continuous
  range of wavelengths and frequencies.
• When sunlight passes through a prism, the
  different frequencies separate into a spectrum of
  colors.
• In the visible spectrum, red light has the
  longest wavelength and the lowest frequency.

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Light
5.3

• The electromagnetic spectrum consists of
  radiation over a broad band of wavelengths.

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Light

• Simulation 3
• Explore the properties of electromagnetic
  radiation.

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5.1
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5.1
50
5.1
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for Sample Problem 5.1
Problem-Solving 5.15 Solve Problem 15 with the
help of an interactive guided tutorial.
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Atomic Spectra
5.3

• Atomic Spectra
• What causes atomic emission spectra?

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Atomic Spectra
5.3

• When atoms absorb energy, electrons move into
  higher energy levels. These electrons then lose
  energy by emitting light when they return to
  lower energy levels.

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Atomic Spectra
5.3

• A prism separates light into the colors it
  contains. When white light passes through a
  prism, it produces a rainbow of colors.

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Atomic Spectra
5.3

• When light from a helium lamp passes through a
  prism, discrete lines are produced.

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Atomic Spectra
5.3

• The frequencies of light emitted by an element
  separate into discrete lines to give the atomic
  emission spectrum of the element.

Mercury
Nitrogen
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An Explanation of Atomic Spectra
5.3

• An Explanation of Atomic Spectra
• How are the frequencies of light an atom emits
  related to changes of electron energies?

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An Explanation of Atomic Spectra
5.3

• In the Bohr model, the lone electron in the
  hydrogen atom can have only certain specific
  energies.
• When the electron has its lowest possible energy,
  the atom is in its ground state.
• Excitation of the electron by absorbing energy
  raises the atom from the ground state to an
  excited state.
• A quantum of energy in the form of light is
  emitted when the electron drops back to a lower
  energy level.

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An Explanation of Atomic Spectra
5.3

• The light emitted by an electron moving from a
  higher to a lower energy level has a frequency
  directly proportional to the energy change of the
  electron.

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An Explanation of Atomic Spectra
5.3

• The three groups of lines in the hydrogen
  spectrum correspond to the transition of
  electrons from higher energy levels to lower
  energy levels.

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An Explanation of Atomic Spectra

• Animation 6
• Learn about atomic emission spectra and how neon
  lights work.

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Quantum Mechanics
5.3

• Quantum Mechanics
• How does quantum mechanics differ from classical
  mechanics?

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Quantum Mechanics
5.3

• In 1905, Albert Einstein successfully explained
  experimental data by proposing that light could
  be described as quanta of energy.
• The quanta behave as if they were particles.
• Light quanta are called photons.
• In 1924, De Broglie developed an equation that
  predicts that all moving objects have wavelike
  behavior.

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Quantum Mechanics
5.3

• Today, the wavelike properties of beams of
  electrons are useful in magnifying objects. The
  electrons in an electron microscope have much
  smaller wavelengths than visible light. This
  allows a much clearer enlarged image of a very
  small object, such as this mite.

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Quantum Mechanics

• Simulation 4
• Simulate the photoelectric effect. Observe the
  results as a function of radiation frequency and
  intensity.

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Quantum Mechanics
5.3

• Classical mechanics adequately describes the
  motions of bodies much larger than atoms, while
  quantum mechanics describes the motions of
  subatomic particles and atoms as waves.

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Quantum Mechanics
5.3

• The Heisenberg uncertainty principle states that
  it is impossible to know exactly both the
  velocity and the position of a particle at the
  same time.
• This limitation is critical in dealing with small
  particles such as electrons.
• This limitation does not matter for
  ordinary-sized object such as cars or airplanes.

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Quantum Mechanics
5.3

• The Heisenberg Uncertainty Principle