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Models of the Atom Section 13.1


Title: Models of the Atom Section 13.1 Author: System User Last modified by: System User Created Date: 10/8/2004 4:19:20 PM Document presentation format – PowerPoint PPT presentation

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Title: Models of the Atom Section 13.1

Models of the AtomSection 13.1
  • The story of how the atomic theory has evolved
    over time.

John Dalton
  • The atom is a solid indivisible mass.
  • He had several key ideas
  • All elements are composed of tiny indivisible
    particles called the atom.
  • Atoms of the same element are identical.
  • Atoms of different elements are different.
  • Atoms combine chemically with one another in
    simple whole-number ratios.
  • During chemical reactions, atoms are separated,
    joined, or rearranged. Atoms are never created
    nor destroyed.

JJ Thomson
  • Plum-pudding model.
  • I call it the Chocolate Chip Cookie Model
  • The atoms has negatively charged electrons stuck
    into a lump of positively charged material, like
    chocolate chips stuck in a cookie dough.
  • Did not address protons and neutrons.

Ernest Rutherford
  • Discovered the nucleus.
  • Showed that most of an atoms mass is
    concentrated in a small, positively charged
    region called the nucleus.
  • Electrons resided on the outside.
  • Did not address how electrons were arranged.

Neils Bohr
  • Electrons are arranged on concentric circular
    paths, or orbits around the nucleus.
  • Solar system model or planetary model.
  • Gave us the idea of definite energy levels.

Quantum Mechanical Model Our Currently Accepted
  • Erwin Schrodinger
  • Primarily a mathematical model using quantum
  • It addresses probabilities of finding an
    electron at any instant in an area called
    electron clouds.
  • Introduced the ideas of Principal Energy Levels
    and Sublevels of energies.
  • The electron clouds take certain shapes,
    represented by the s,p,d,f subatomic orbitals.

Principal Energy Levels
  • Just like the Bohr model, the Quantum Mechanical
    Model designates energy levels of electrons by
    means of principal quantum numbers
  • Principal Energy Levels refers to a major region
    where electrons are most likely to be found.
  • They are assigned values in order of increasing
    energy 1, 2, 3, etc.

  • Within each principal energy level, the electrons
    occupy energy sublevels.
  • The number of sublevels within each principal
    energy level is the same as the principal quantum
  • How many sublevels does the 4th principal energy
    level have?

Atomic Orbitals
  • The regions in which electrons are likely to be
    found are called atomic orbitals.
  • Letters denote the atomic orbitals
  • S-shape orbitals are spherical
  • P-shape orbitals are hour-glass shapes
  • D-shape orbitals have clover-leaf shapes
  • Draw an example of each into your notes.

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Exploring further
  • The lowest principal energy level has only one
    sublevel, called 1s.
  • The second principal energy level has 2
    sublevels, the 2s and 2p. The 2p is higher in
    energy and consists of three p orbitals.

The Electron Pyramid
  • The s orbitals have 1 spatial orientation,
    therefore can hold 2 electrons
  • The p orbitals have 3 spatial orientations,
    therefore can hold 6 electrons
  • The d orbitals have 5 spatial orientations,
    therefore can hold 10 electrons
  • The f orbitals have 7 spatial orientations,
    therefore can hold 14 electrons.

Electrons Fill following 3 simple rules
  • Aufbau principle Electrons enter the lowest
    energy level first.
  • Pauli Exclusion Principle An atomic orbital may
    describe at most 2 electrons, both spinning in
    opposite directions.
  • Hunds Rule When electrons occupy orbitals of
    equal energy, one electron enters each orbital
    until all the orbitals contain one electron with
    parallel spins.

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Exceptional Electron Configurations
  • Chromium and Copper have exceptional electron
  • They fill their d sublevel completely, leaving
    their 4s partially filled.
  • Much more stable this way!
  • Write them correctly into your notepacks

Physics and the Quantum Mechanical Model
  • This section studies the electron as a property
    of light.
  • Electrons travel as waves and are made of
    particles of light called photons
  • According to the wave model, light consists of

Electromagnetic Spectrum
  • This form of energy includes
  • Gamma rays
  • X-rays
  • Ultraviolet rays
  • Visible light
  • Infrared rays
  • Radar
  • FM
  • TV
  • Shortwave
  • AM

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Electromagnetic Spectrum
  • Every element emits light when it is excited by
    the passage of electric discharge through its gas
    or vapor.
  • The atoms first absorb energy, then lose the
    energy as they emit light.

Electromagnetic Spectrum
  • Electrons are said to move from their GROUND
    STATE (lowest energy level) to and EXCITED STATE
    (higher energy level).
  • When the electron falls back to its lower energy,
    it emits a PHOTON of energy, and can be seen in
    the visible spectrum.

Electromagnetic Spectrum
  • Passing the light emitted by an element through a
    prism gives the ATOMIC EMISSION SPECTRUM of the
  • Because each atom has a unique electron
    arrangement, each atom emits a unique wavelength
    during this process. This wavelength falls
    within the visible spectrum.

Kernel Structures
  • The kernel is a structure used to shorten an
    electron configuration.
  • A kernel is an inert gas symbol in brackets that
    stands in place of all of the filled orbitals
    contained in the inert gas.
  • Examples

Honors The Quantum Concept and the
Photoelectric Effect
  • Electrons travel as waves around the nucleus of
    an atom. Lets review the concept of wave

  • Amplitude the height of a wave from its origin.
  • Wavelength the distance between the crests l
  • Frequency the number of wave cycles to pass a
    given point per unit of time. n
  • Hertz a per second unit for n.

Speed of light, c
  • The frequency and wavelength are inversely
    related as shown by this relationship
  • c ln
  • The speed of light is a constant
  • C 3.0 E 8 meters/second
  • Examples

Honors Photoelectric Effect
  • By studying black body radiation, German
    physicist Max Planck described mathematically
    that the amount of radiant energy (E) absorbed or
    emitted by a body is proportional to the
    frequency of the radiation.
  • E h x n
  • h Plancks constant, 6.63 E -34 J-s

Albert Einstein
  • Nobel Prize Winner!!
  • In 1905, Albert Einstein proposed that light
    could be described as quanta of energy that
    behave as if they were particles he called
  • In the PHOTOELECTRIC EFFECT, metals eject
    electrons when light shines on them.
  • Photoelectric cells
  • Automatic doors at Meijer!

(Honors) Quantum Mechanics and Matter Waves
  • In 1924, Louis De Broglie derived an equation
    that described the wavelength of a moving
    particle, such an electron.
  • l h/mv
  • M mass (in kg)
  • Mass x velocity momentum
  • h is Plancks constant

Matter Waves
  • If the mass of an electron is 9.11 E -28 grams
    and moving nearly at the speed of light, an
    electron has a wavelength of about 2 E -10 cm.

  • De Broglies prediction that matter would exhibit
    both wave and particle properties is summarized
    in the following two statements
  • 1. Classical mechanics adequately describes the
    motions of bodies much larger than the atoms that
    they comprise.
  • 2. Quantum mechanics describes the motions of
    subatomic particles and atoms as waves. These
    particles gain or lose energy in packages called

Honors Quantum Numbers
  • 4 Quantum numbers are used to describe a single
    electrons position within an atom.
  • 1. Principal quantum number (n) size and energy
    of an orbital.
  • Has integer values gt0

Quantum numbers
  • 2. Angular momentum quantum number l
  • shape of the orbital.
  • integer values from 0 to n-1
  • l 0 is called s
  • l 1 is called p
  • l 2 is called d
  • l 3 is called f
  • l 4 is called g

Values of l 0 1 2 3
Letter used s Sharp p Princi-pal d Diffuse f Funda-mental
Quantum numbers
  • 3. Magnetic quantum number (m l)
  • integer values between - l and l, including
  • Describes the orientation of the orbital in
  • 4. Electron spin quantum number (m s)
  • Can have 2 values.
  • either 1/2 or -1/2

14.1 Classification of the Elements
  • By Electron Configuration

Classifying Elements by Electron Configuration
  • Of the three major subatomic particles, the
    ELECTRON plays the most significant role in
    determining the properties of an element.
  • The arrangement of elements in the PERIODIC TABLE
    depends on these properties.

Elements can be classified into 4 categories
  • The Noble Gases
  • These are elements in which the outermost s and
    p sublevels are filled.
  • Write for Helium, Neon, Argon, Krypton

Elements can be classified into 4 categories
  • The representative elements
  • In these elements, the outermost s and p
    sublevel is only partially filled.
  • Write for Lithium, Sodium, Potassium, Carbon,
    Silicon, Germanium

Elements can be classified into 4 categories
  • The transition metals
  • These are metallic elements in which the
    outermost s sublevel and nearby d sublevel
    contain electrons.
  • Write for Zinc and Zirconium.

Elements can be classified into 4 categories
  • The inner transition metals
  • These are metallic elements in which the
    outermost s sublevel and nearby f sublevel
    generally contain electrons.

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14.2 Periodic Trends
  • Atomic radius ½ the distance between the nuclei
    of two like atoms in a diatomic molecule.

Group Trends
  • Atomic size generally increases as you go down a
    group on the periodic table.
  • Adding additional energy levels!

Periodic Trends
  • Atomic size generally decrease as you move from
    left to right across a period.
  • Same energy level -
  • increasing nuclear charge pulls electrons closer
    to nucleus.

Ionization Energy
  • An ion a charged atom that results from either
    losing or gaining an electron.
  • Ionization Energy The energy required to
    overcome the attraction of the nuclear charge and
    remove an electron from a gaseous atom.
  • (The ease of losing an electron and forming a 1

Ionization Energy
  • First ionization energy the energy needed to
    remove the first electron from an atom.
  • Second ionization energy the energy needed to
    remove the second electron from an atom, etc.

Ionization Energy
  • Group Trends The first ionization energy
    generally decreases as you move down a group on
    the periodic table.
  • The size of the atoms increases, so the outermost
    electron is farther from the nucleus and will be
    more easily removed.

Ionization Energy
  • Periodic Trends For the representative
    elements, the first ionizatoin energy generally
    increases as you move from left to right across a
  • Increasing nuclear charge makes it more difficult
    to remove an electron.

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Ionic Size
  • The atoms of METALLIC elements have low
    ionization energies. They form POSITIVE ions
  • By contrast, the atoms of NONMETALLIC elements
    readily form NEGATIVE ions.

Trends in Ionic Size
  • Positive ions are always smaller than the neutral
    atoms from which they form.
  • They lose their outer shell electrons
  • Negative ions are always larger than the neutral
    atoms from which they form.
  • This is because the effective nuclear attraction
    is less for an increased number of electrons.

Trends in Electronegativity
  • Electronegativity the tendency for the atoms of
    the element to attract electrons when they are
    chemically combined with atoms of another
  • Electronegativity generally DECREASES as you go
    down a group.
  • As you go across a period from left to right, the
    electronegativity of the representative elements

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  • The electronegativity of cesium, the least
    electronegative element is 0.7
  • The electronegativity of fluorine, the most
    electronegative element, is 4.0
  • Electronegativity values help predict the type of
    bonding that can exist between atoms in
    compounds, either IONIC OR COVALENT bonds.

Summary of Periodic Trends
  • Using page 406, create a summary of periodic
    trends into your notes.
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