Electrons in Atoms

1
Chapter 5

• Electrons in Atoms

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What we know so far about the atom

• Atoms have a nucleus ( made of protons and
  neutrons) surrounded by fast moving electrons.
• All atoms positive charge and virtually all its
  mass is concentrated in the nucleus.
• Always the number of protons in the nucleus
  equals the number of electrons.

3
Light and quantized energy

• Light , a form of electromagnetic radiation, has
  characteristics of both wave and particle.
• An electromagnetic radiation is a form of energy
  that exhibits wavelike behavior as it travels
  through space.
• Examples of electromagnetic radiation
  microwaves, X rays and waves that carry radio and
  television programs

4
Characteristics of waves

• Wavelength ( represented by Greek letter lambda)
  is the shortest distance between 2 equivalent
  points.
• Frequency ( represented by the Greek letter nu)
  is the number of waves that passes a given point
  per second.
• The amplitude of a wave is the waves height from
  the origin to crest, or from the origin to
  trough. Wavelength and frequency do not affect
  the amplitude of a wave.

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How we make electromagnetic waves

• Electromagnetic waves
• are formed when an
• electric field (shown with
• blue arrows) couples
• with a magnetic field
• (shown with red arrows).

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Electromagnetic wave relationship

• C ( speed of light in vacuum) wavelength X
  frequency

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Electromagnetic waves have different wavelength
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Electromagnetic spectrum

• The EM spectrum includes all forms of
  electromagnetic radiation, with the only
  differences in the type of radiation being their
  wavelength and frequencies.

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Practice problems

• Independent work
• Textbook page 140 1,2,3 and 4

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The particle nature of light

• As mentioned before light , a form of
  electromagnetic radiation, has characteristics of
  both wave and particle.
• When objects are heated they emit glowing light.
  Max Planck discovered that matter can gain or
  lose energy in small amounts called quanta. A
  quantum is the minimum amount of energy that can
  be gained or lost by an atom.

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The photoelectric effect

• In the photoelectric effect electrons, called
  photoelectrons, are emitted from a metals
  surface when light of a certain frequency, or
  higher than a certain frequency shines on a
  surface.

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Formulas Need to know

• Energy of a
• quantum Planck constant X frequency
• A photon is a massless particle that carries a
  quantum of energy
• Energy of a
• photon Planck constant X frequency

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Practice problems

• Independent work
• Textbook page 143 5,6,7

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Atomic emission spectra

• The atomic emission spectrum of an element is a
  set of frequencies of the electromagnetic waves
  emitted by atoms of the element.
• Each elements atomic emission spectrum is unique
  and can be used to identify an element or
  determine whether the element is part of an
  unknown compound.

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Atomic emission spectrum for hydrogen, barium
and iron
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Comparison of atomic emission spectra

• http//jersey.uoregon.edu/vlab/elements/Elements.h
  tml

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Quantum theory and the atom

• Scientists tried to explain the relationship
  between atomic structure, electrons and atomic
  emission spectra
• Wavelike properties of electrons help relate
  atomic emission spectra, energy states of atoms
  and atomic orbitals
• .

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5.2 Bohrs model of the atom

• The lowest allowable energy state of an atom is
  called its ground state.
• When an atom gains energy, it is said to be in an
  excited state.

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Example for the atom of hydrogen

• Ground state Excited state

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Bohrs description of the hydrogen atom

• The electron in a hydrogen atom moves around the
  nucleus only in certain allowed orbits. The lower
  the electrons orbit, the higher the energy state
  or energy level. Due to the electrons attraction
  to the nucleus.

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The hydrogen line spectrum

• In order to complete his calculations Bohr
  assigned a number n, called a quantum number, to
  each orbit.

Atoms orbit Quantum number Orbit radius nm Corresponding atomic energy level Relative energy
First n1 0.0529 1 E1
Second n2 0.212 2 E24E1
Third n3 0.476 3 E39E1
Fourth n4 0.846 4 E416E1
Fifth n5 1.32 5 E525E1
Sixth n6 1.90 6 E636E1
Seventh n7 2.59 7 E749E1
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The hydrogen line spectrum-cont.

• Bohr suggested that the hydrogen atom is in the
  ground state, also called the first energy level,
  when its single electron is in n1.
• When energy is added from an outside source the
  electron moves to a higher energy orbit, such as
  n2, so the atom is now in an excited state. When
  the atom is in an excited state, the electron can
  drop from the higher-energy orbit to a lower
  energy orbit, so as a results of the transition
  an photon is emitted.
• E photonE higher-energy orbit - E lower-energy
  orbit

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The hydrogen line spectrum
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The hydrogen line spectrum
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Limits of Bohrs model

• The model explained the hydrogens spectral lines
  but not for other elements
• Bohrs model did not account for the chemical
  behavior of the atom
• In later experiments it was proven that electrons
  in atoms do not move around the nucleus in
  circular orbits

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The quantum mechanical model of the atom

• Louis De Broglie proposed that just as light had
  been shown to have both a wave and a particle
  aspect, so might matter.
• His now famous wave equation, indicated that an
  electron in a Bohr orbit racing around the atom's
  nucleus would possess a wavelength of the right
  dimension to form standing waves.

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De Broglies equation

• wavelength planks constant / mass of particle
  x velocity

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Heisenberg uncertainty principle

• states that it is fundamentally impossible to
  know precisely both the velocity and the position
  of a particle at the same time.

"The more precisely the POSITION is
determined,the less precisely the MOMENTUM is
known"