Models of the Atom Section 13.1 - PowerPoint PPT Presentation

1 / 56
About This Presentation
Title:

Models of the Atom Section 13.1

Description:

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

Number of Views:138
Avg rating:3.0/5.0
Slides: 57
Provided by: Syste377
Category:

less

Transcript and Presenter's Notes

Title: Models of the Atom Section 13.1


1
Models of the AtomSection 13.1
  • The story of how the atomic theory has evolved
    over time.

2
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.

3
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.

4
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.

5
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.

6
Quantum Mechanical Model Our Currently Accepted
Model
  • Erwin Schrodinger
  • Primarily a mathematical model using quantum
    mechanics
  • 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.

7
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.

8
Sublevels
  • 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
    number.
  • How many sublevels does the 4th principal energy
    level have?

9
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.

10
(No Transcript)
11
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.

12
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.

13
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.

14
(No Transcript)
15
(No Transcript)
16
Exceptional Electron Configurations
  • Chromium and Copper have exceptional electron
    configurations.
  • They fill their d sublevel completely, leaving
    their 4s partially filled.
  • Much more stable this way!
  • Write them correctly into your notepacks

17
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 RADIATION.

18
Electromagnetic Spectrum
  • This form of energy includes
  • Gamma rays
  • X-rays
  • Ultraviolet rays
  • Visible light
  • Infrared rays
  • Radar
  • FM
  • TV
  • Shortwave
  • AM

19
(No Transcript)
20
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.

21
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.

22
Electromagnetic Spectrum
  • Passing the light emitted by an element through a
    prism gives the ATOMIC EMISSION SPECTRUM of the
    element.
  • Because each atom has a unique electron
    arrangement, each atom emits a unique wavelength
    during this process. This wavelength falls
    within the visible spectrum.

23
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

24
Honors The Quantum Concept and the
Photoelectric Effect
  • Electrons travel as waves around the nucleus of
    an atom. Lets review the concept of wave
    mechanics

25
Definitions
  • 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.

26
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

27
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

28
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
    PHOTONS.
  • In the PHOTOELECTRIC EFFECT, metals eject
    electrons when light shines on them.
  • Photoelectric cells
  • Automatic doors at Meijer!

29
(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

30
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.

31
  • 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
    quanta.

32
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

33
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

34
Values of l 0 1 2 3
Letter used s Sharp p Princi-pal d Diffuse f Funda-mental
35
Quantum numbers
  • 3. Magnetic quantum number (m l)
  • integer values between - l and l, including
    zero.
  • Describes the orientation of the orbital in
    space.
  • 4. Electron spin quantum number (m s)
  • Can have 2 values.
  • either 1/2 or -1/2

36
14.1 Classification of the Elements
  • By Electron Configuration

37
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.

38
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

39
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

40
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.

41
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.

42
(No Transcript)
43
14.2 Periodic Trends
  • Atomic radius ½ the distance between the nuclei
    of two like atoms in a diatomic molecule.

44
Group Trends
  • Atomic size generally increases as you go down a
    group on the periodic table.
  • Adding additional energy levels!

45
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.

46
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
    charge)

47
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.

48
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.

49
Ionization Energy
  • Periodic Trends For the representative
    elements, the first ionizatoin energy generally
    increases as you move from left to right across a
    period.
  • Increasing nuclear charge makes it more difficult
    to remove an electron.

50
(No Transcript)
51
Ionic Size
  • The atoms of METALLIC elements have low
    ionization energies. They form POSITIVE ions
    easily.
  • By contrast, the atoms of NONMETALLIC elements
    readily form NEGATIVE ions.

52
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.

53
Trends in Electronegativity
  • Electronegativity the tendency for the atoms of
    the element to attract electrons when they are
    chemically combined with atoms of another
    element.
  • 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
    INCREASES.

54
(No Transcript)
55
Electronegativity
  • 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.

56
Summary of Periodic Trends
  • Using page 406, create a summary of periodic
    trends into your notes.
Write a Comment
User Comments (0)
About PowerShow.com