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Chemistry I Electrons in Atoms Chapter 5

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Chemistry I Electrons in Atoms Chapter 5 46.) What is the Pauli Exclusion Principle? A maximum of two electrons can occupy an orbital and they must have opposite spins. – PowerPoint PPT presentation

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Title: Chemistry I Electrons in Atoms Chapter 5


1
Chemistry IElectrons in AtomsChapter 5
2
Rutherfords nuclear model did not provide enough
detail about how electrons occupy the space
around the nucleus. In this chapter we will learn
how electrons are arranged around the nucleus and
how that arrangement effects chemical behavior.
3
In what ways did many scientists
findRutherfords model to be incomplete?
4
1.)In what ways did many scientists
findRutherfords model to be incomplete?
  • It did not explain
  • 1.How the electrons were arranged around the
    nucleus
  • 2.Why they were not pulled into the nucleus
  • 3. The differences in chemical behavior of the
    different elements.

5
2.)In the early 1900s what did scientists
observe about certain elements when heated in a
flame?
  • Every element gives off specific wavelengths of
    light that can be used like a fingerprint to
    identify the element.

6
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7
spectroscope
8
  • http//jersey.uoregon.edu/vlab/elements/Elements.h
    tml

9
An elements chemical behavior is related to the
arrangement of electrons in its atoms.
10
Wave Nature of Light
11
What is electromagnetic radiation?
12
3.)What is electromagnetic radiation?
  • A type of wave that is part electrical and
    magnetic energy acting at right angles.
  • Visible light is a type of electromagnetic
    radiation.

13
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14
What are the 4 characteristics of waves?
15
4.) What are the 4 characteristics of waves?
  • 1.) wavelength distance from crest to crest
  • 2.) frequency- of waves that pass a given point
    per second
  • 3.) Amplitude height of wave from the normal
    resting to wave crest
  • 4.) speed the speed of light is a constant 3.0
    x 108 m/s.

16
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17
5.)What unit do we use to express frequency?
  • Hertz (Hz)
  • The hertz unit is 1/second ( the inverse of a
    second)

18
6.) What is the speed of light?
  • The speed of light is a constant 3.0 x 108 m/s
  • The formula for the speed of light is C ? v
  • ? wavelength (measured in meters)
  • v frequency (measured in 1/sec)
  • The unit for speed is m/s (distance/time)

19
6 continued
  • When wavelength (?) increases then frequency (v)
    decreases because the speed of light ( c) is a
    constant 3.0X108m/s.
  • Also
  • When ? decreases, V increases

20
7.) List types of electromagnetic radiation from
the lowest to the highest in energy ( page 139)
21
List types of electromagnetic radiation from the
lowest to the highest in energy ( page 140)
  • Radio waves ?decreasing wavelength
  • Microwaves ?increasing frequency
  • Infrared
  • Visible
  • Ultraviolet (U.V.)
  • X rays
  • gamma

22
Particle Nature of Light
  • When an object is heated only certain
    wavelengths of light are emitted. The instrument
    used to separate light into its different
    wavelengths is called a spectroscope. The pattern
    of wavelengths is called the spectrum.

23
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24
  • http//jersey.uoregon.edu/vlab/elements/Elements.h
    tml

25
  • When a substance is heated the spectrum of light
    that is given off can be used to identify the
    substance.

26
  • http//jersey.uoregon.edu/vlab/elements/Elements.h
    tml

27
  • The wave model of light could not explain why
    different substances emit particular wavelengths
    of light when heated

28
8.)What are some problems with the wave model of
the atom?
  • It doesnt explain why only certain wavelengths
    of light are emitted when an object is heated.
  • It also does not explain the photoelectric
    effect.( we will discuss the photoelectric effect
    in just a little bit)

29
Who is Max Planck?
30
9.)Who is Max Planck?
  • A German physicist who, in 1900, began searching
    for the reason that only certain wavelengths of
    light are given off for every element.

31
10.)What did Planck conclude?
  • Matter can gain or lose energy in only small
    specific amounts called quanta.

32
As metal is heated, it glows. In fact, it glows
different colors as it gets hotter (white hot
being the hottest). By studying glowing metals,
Max Planck discovered that only certain
wavelengths of light are emitted at each specific
temperature.
33
11.) What is a quantum?
  • The minimum amount of energy that can be gained
    or lost by an atom.
  • The amount of energy it takes for a electron to
    get from one energy level to the next.

34
12.) If energy can only absorbed in quanta,
specific amounts, why does energy appear to
continuous to us?
  • b/c quanta are extremely small.

35
13.) What is Plancks equation that demonstrates
that energy is related to the frequency of
radiation?
  • Equation
  • E hv
  • E energy of a quantum
  • h- plancks constant
  • v- frequency

36
14.) What is Plancks constant?
  • 6.62 x 10-34 Js

37
15.) How are energy and frequency related?
  • They are directly related
  • As frequency increase, energy increases.

38
What is the photoelectric effect?
39
What is the photoelectric effect?
40
  • http//www.tutorvista.com/content/physics/physics-
    iv/radiation-and-matter/photoelectric-effect-and-c
    ell.php

41
16.)What is the photoelectric effect?
  • When light of a certain frequency shines on the
    surface of a metal, electrons are ejected. Blue
    light always results in electrons being ejected.
    Red light of any brightness does not eject
    electrons.

42
17.) In 1905 how did Albert Einstein explain the
photoelectric effect?
  • He proposed that all electromagnetic radiation
    (including visible light) is both wavelike and
    particle like in nature.
  • Continued next slide

43
17.) In 1905 how did Albert Einstein explain the
photoelectric effect?
  • He suggested that while a beam of light has many
    wavelike characteristics, it is also like a
    stream of particles called photons.

44
18.) What is a photon?
  • A photon is a particle of electromagnetic
    radiation with no mass that carries a quantum of
    energy.

45
19.)How did Einstein use Plancks idea of quantum
energy to explain the photoelectric effect?
  • Planck proposed that Ehv. Energy is equal to his
    constant times the frequency of radiation.
  • Einstein proposed that since blue light has a
    larger frequency, it has a larger quantum of
    energy.
  • Continued next slide

46
19.) How did Einstein use Plancks idea of
quantum energy to explain the photoelectric
effect?
  • E hv for blue light is more than Ehv for red
    light because blue light has a larger frequency
    than red light.

47
Practice Problems page 143
48
20.)What is the atomic emission spectrum?
  • The atomic emission spectrum of an element is the
    set of frequencies of (light) electromagnetic
    waves emitted ( given off) by atoms to that
    element when it is heated.

49
21.) How does the atomic emission spectrum for
one element compare to another element?
  • Each is unique and can be used to identify the
    element like a fingerprint.

50
22.) Are atomic emission spectrums of elements
continuous or distinct individual wavelengths of
light?
  • They are distinct, individual wavelengths of
    light.

51
23.) In a line spectrum which line has the
highest energy?
  • The line farthest to the blue end of the
    spectrum. The line corresponding to the shortest
    wavelength/ highest frequency.

52
24.) How does the quantitization of energy (
energy only exists in specific sizes) help
explain line spectrum?
  • The energy of the photon of light that is emitted
    is tied to a specific frequency/color of light
  • E hv Each frequency (v) corresponds to a
    specific color

53
Questions page 145
54
Section 2 Quantum Theory and the Atom
55
  • Scientists concluded that light was both wave and
    particle. Scientists were able to understand
    atomic structure, electrons and atomic emission
    spectra better because of a better understanding
    of light.

56
  • Scientists wanted to know why atomic emission
    spectrums are not continuous but are
    discontinuous.

57
23.)Who was Niels Bohr?
  • A Danish physicist who worked in Rutherfords
    laboratory. He proposed the quantum model of the
    atom that helped explain why atomic emission
    spectrums only contain certain frequencies of
    light.

58
25.)What did Niels Bohr propose?
  • He proposed that a hydrogen atom has only certain
    allowable energy levels.
  • The lowest level is called the ground state.
  • continued

59
25.)What did Niels Bohr propose?
  • All other levels are called excited states.
  • Electrons move around the nucleus in only certain
    allowed orbits.
  • Bohr labeled these orbits.
  • continued

60
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61
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62
25.)What did Niels Bohr propose?
  • The orbit closest to the nucleus is labeled n1
    and was lowest in energy.

63
26.) How did Niels Bohr explain line spectrum
with the quantization of energy for electrons. (
How did Bohr explain that certain wavelengths of
light are released when an element is heated by
proposing that electrons can only be at certain
energy levels.)
64
26
  • He said that when an element is at ground state (
    the lowest energy level) no energy (light) is
    released. But when an electron becomes excited,
    it jumps up to another energy level.

65
26
  • When the electron returns to the lower energy
    level the energy is released. The energy (light)
    that is released is equal to the difference
    between energy levels.

66
26
  • The amount of energy that is released is equal to
    a particular frequency of light.

67
27.)Do we still accept Bohrs model?
  • Bohr was correct about his idea of quantization
    of energy and his basic explanation of atomic
    emission spectrum. However he could only predict
    the atomic emission spectrum of hydrogen.

68
27.)Do we still accept Bohrs model?
  • His basic model of the model is considered to be
    incorrect. Today we do not believe that the
    electrons are in orbits so that we can predict
    their positions.

69
The Quantum Mechanical Model of the Atom
70
28.)Who was Louis de Broglie?
  • A French physics student that proposed an idea
    that eventually accounted for the fixed energy
    levels in Bohrs model.

71
29.)What did De Broglie propose?
  • He proposed that all matter moves in waves. The
    larger the mass the smaller the wavelength. Only
    very small objects have noticeable wavelengths.

72
30.)What was De Broglies equation?
  • ? h/ mv
  • DeBroglies equation predicts a large mass will
    have a small wavelength (?) only a very small
    mass will have a noticeable wavelength.

73
Page 149
74
If electrons move in waves then only electrons
with matching wavelengths can be accepted into an
energy level and when an electron is released to
a lower energy level particular wavelengths of
energy are released.
75
29.)What did Heisenberg propose?
  • It is impossible to make a measurement of an
    object without disturbing it. For example, a
    thermometer changes the temperature of an object.

76
30.)What is the Heisenberg uncertainty principle?
  • It is impossible to know precisely both the
    velocity and position of a particle at the same
    time.

77
The Schrodinger Wave Equation
78
32.) The Austrian physicist Edwin Schrodinger
derived an equation that treated the electron as
a wave. What was significant about his equation?
  • It worked equally well for atoms of other
    elements unlike Bohrs model which worked only
    for hydrogen.

79
33.) What is the name of the model proposed in
which electrons are treated as waves?
  • The quantum mechanical model of the atom.

80
34.)What is the quantum mechanical model of the
atom?
  • It limits the electrons energy to certain values
    but makes no attempt to describe the path of the
    electron around the nucleus

81
35.) What is an atomic orbital?
  • In the quantum mechanical model of the atom, the
    orbital is the region in which we expect to find
    an electron 90 of the time.
  • Orbitals are described by electron clouds.

82
Hydrogens Atomic Orbital
83
36.)What are principal quantum numbers?
  • Principal quantum numbers represent the relative
    sizes and energies of the atomic orbitals.
  • As n increases, the size and energy level
    increases.

84
37.) What are sublevels?
  • Each principle energy level has sublevels. The
    number of sublevels is equal to the n level.
  • n 1 has one sublevel
  • n 2 has two sublevels
  • etc.

85
38.) What type of sublevels are there?
  • The sublevels are named s, p, d, and f

86
Sublevel Shape energy level
  • s spherical lowest
  • p dumbell
  • d cloverleaf
  • f complex highest

87
39.)How many electrons does an orbital contain?
  • two

88
40.) How many orbitals does each sublevel
contains?
  • Sublevel orbitals e-
  • s 1 2
  • p 3 6
  • d 5 10
  • f 7 14

89
n 4n 3n 2n 1
s p d f
s p d
s p
s
90
Principal Quantum Sublevel types orbitals Total of orbitals Total electrons
1 s 1 1 2
2 s p 1 3 4 8
3 s P d 1 3 5 9 18
4-7 s p d f 1 3 5 7 16 32

91
Section 3 Electron ConfigurationsThe arrangement
of electrons in atoms follows a few very specific
rules.
92
Ground State Electron Configurations
93
41.)What is an electron configuration?
  • The arrangement of electrons in an atom

94
42.) What is ground state electron configuration?
  • The most stable, lowest energy arrangement of
    electrons.

95
43.) What are the three rules that govern ground
state electron configurations.
  • 1.) Aufbau principle
  • 2.) Hunds Rule
  • 3.) Pauli Exclusion Principle

96
44.)What is the Aufbau principle?
  • Electrons occupy the lowest energy orbital
    available.
  • German for building up

97
45.) In general what is the order of increasing
energy among energy levels and sublevels?
  • In general n1 is lower than n2 . The order of
    increasing order of sublevels is s, p, d, f.
  • Use the diagram

98
46.) What is the Pauli Exclusion Principle?
  • A maximum of two electrons can occupy an orbital
    and they must have opposite spins.
  • The way to indicate two electrons with opposite
    spin is

??
99
47.) What is the Hunds Rule?
  • Single electrons with the same spin must occupy
    equal energy orbitals before additional electrons
    with opposite spin can occupy the same orbital.

100
48.)What are the two ways that electron
configurations can be described?
  • Orbital diagrams using boxes
  • Mg
  • 1s 2s 2p 3s

??
??
??
??
??
??
101
49.) What is noble gas configuration?
  • Electron configurations that are used to show
    just the valence electrons. The full inner core
    orbitals are represented by the noble gas symbol
    with the lower atomic number and the electrons in
    the valence shell.
  • K 1s22s22p63s23p64s1
  • K Ar 4s1

102
50.) Exceptions to predicted configurations.
  • Cr Ar4s13d5
  • Cu Ar4s13d10
  • It is more stable to borrow an electron from
    the s orbital and put it in the d orbital

103
Pg 160 practice problems
104
51.) Which electrons determine the chemical
properties of an element?
  • The valence electrons/ the outermost electrons.

105
52. How does the number of valence electrons
compare to the family the element is in?
  • Group 1 1 valence electron
  • Group 2 2 valence electrons
  • Group 13 3 valence electrons
  • Group 14 4 valence electrons
  • Group 15 5 valence electrons
  • Group 16 - 6 valence electrons
  • Group 17 7 valence electrons
  • Group 18 8 valence electrons except for
    hydrogen which only has two.
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