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Chemistry: Matter and Change

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Title: Chemistry: Matter and Change


1
(No Transcript)
2
Chapter Menu
Electrons in Atoms
Section 5.1 Light and Quantized Energy Section
5.2 Quantum Theory and the Atom Section 5.3
Electron Configuration
Click a hyperlink or folder tab to view the
corresponding slides.
Exit
3
Section 5-1
Section 5.1 Light and Quantized Energy
  • Compare the wave and particle natures of light.
  • Define a quantum of energy, and explain how it is
    related to an energy change of matter.
  • Contrast continuous electromagnetic spectra and
    atomic emission spectra.

radiation the rays and particles alpha
particles, beta particles, and gamma raysthat
are emitted by radioactive material
4
Section 5-1
Section 5.1 Light and Quantized Energy (cont.)
electromagnetic radiation wavelength frequency amp
litude electromagnetic spectrum
quantum Planck's constant photoelectric
effect photon atomic emission spectrum
Light, a form of electronic radiation, has
characteristics of both a wave and a particle.
5
Section 5-1
The Atom and Unanswered Questions
  • Recall that in Rutherford's model, the atoms
    mass is concentrated in the nucleus and electrons
    move around it.

6
Section 5-1
The Atom and Unanswered Questions
  • The model doesnt explain how the electrons were
    arranged around the nucleus.
  • The model doesnt explain why negatively charged
    electrons arent pulled into the positively
    charged nucleus.

7
Section 5-1
The Atom and Unanswered Questions (cont.)
  • In the early 1900s, scientists observed certain
    elements emitted visible light when heated in a
    flame.
  • Analysis of the emitted light revealed that an
    elements chemical behavior is related to the
    arrangement of the electrons in its atoms.

8
Section 5-1
The Wave Nature of Light
  • Visible light is a type of electromagnetic
    radiation, a form of energy that exhibits
    wave-like behavior as it travels through space.

9
Section 5-1
The Wave Nature of Light
  • All waves can be described by several
    characteristics.

10
Section 5-1
The Wave Nature of Light (cont.)
  • The wavelength (?) is the shortest distance
    between equivalent points on a continuous wave.

11
Section 5-1
The Wave Nature of Light (cont.)
  • The frequency (?) is the number of waves that
    pass a given point per second.
  • The amplitude is the waves height from the
    origin to a crest.

12
Section 5-1
The Wave Nature of Light (cont.)
13
Section 5-1
The Wave Nature of Light (cont.)
  • The speed of light (3.00 ? 108 m/s) is the
    product of its wavelength and frequency c ??.

14
Section 5-1
The Wave Nature of Light (cont.)
  • Sunlight contains a continuous range of
    wavelengths and frequencies.

15
Section 5-1
The Wave Nature of Light (cont.)
  • A prism separates sunlight into a continuous
    spectrum of colors.

16
Section 5-1
The Wave Nature of Light (cont.)
  • The electromagnetic spectrum includes all forms
    of electromagnetic radiation.

17
Section 5-1
The Particle Nature of Light
  • The wave model of light cannot explain all of
    lights characteristics.
  • Matter can gain or lose energy only in small,
    specific amounts called quanta.
  • A quantum is the minimum amount of energy that
    can be gained or lost by an atom.
  • Plancks constant has a value of 6.626 ? 1034 J
    ? s.

18
Section 5-1
The Particle Nature of Light (cont.)
  • The photoelectric effect is when electrons are
    emitted from a metals surface when light of a
    certain frequency shines on it.

19
Section 5-1
The Particle Nature of Light (cont.)
  • Albert Einstein proposed in 1905 that light has a
    dual nature.

20
Section 5-1
The Particle Nature of Light (cont.)
  • A beam of light has wavelike and particlelike
    properties.
  • A photon is a particle of electromagnetic
    radiation with no mass that carries a quantum of
    energy.

Ephoton h? Ephoton represents energy. h is
Planck's constant. ? represents frequency.
21
Section 5-1
Atomic Emission Spectra
  • Light in a neon sign is produced when electricity
    is passed through a tube filled with neon gas and
    excites the neon atoms.
  • The excited atoms emit light to release energy.

22
Section 5-1
Atomic Emission Spectra (cont.)
23
Section 5-1
Atomic Emission Spectra (cont.)
  • The atomic emission spectrum of an element is the
    set of frequencies of the electromagnetic waves
    emitted by the atoms of the element.
  • Each elements atomic emission spectrum is unique.

24
Section 5-1
Section 5.1 Assessment
What is the smallest amount of energy that can be
gained or lost by an atom? A. electromagnetic
photon B. beta particle C. quanta
D. wave-particle
  1. A
  2. B
  3. C
  4. D

25
Section 5-1
Section 5.1 Assessment
What is a particle of electromagnetic radiation
with no mass called? A. beta particle B. alpha
particle C. quanta D. photon
  1. A
  2. B
  3. C
  4. D

26
End of Section 5-1
27
Section 5-2
Section 5.2 Quantum Theory and the Atom
  • Compare the Bohr and quantum mechanical models of
    the atom.
  • Explain the impact of de Broglie's wave article
    duality and the Heisenberg uncertainty principle
    on the current view of electrons in atoms.
  • Identify the relationships among a hydrogen
    atom's energy levels, sublevels, and atomic
    orbitals.

atom the smallest particle of an element that
retains all the properties of that element, is
composed of electrons, protons, and neutrons.
28
Section 5-2
Section 5.2 Quantum Theory and the Atom (cont.)
ground state quantum number de Broglie
equation Heisenberg uncertainty principle
quantum mechanical model of the atom atomic
orbital principal quantum number principal energy
level energy sublevel
Wavelike properties of electrons help relate
atomic emission spectra, energy states of atoms,
and atomic orbitals.
29
Section 5-2
Bohr's Model of the Atom
  • Bohr correctly predicted the frequency lines in
    hydrogens atomic emission spectrum.
  • The lowest allowable energy state of an atom is
    called its ground state.
  • When an atom gains energy, it is in an excited
    state.

30
Section 5-2
Bohr's Model of the Atom (cont.)
  • Bohr suggested that an electron moves around the
    nucleus only in certain allowed circular orbits.

31
Section 5-2
Bohr's Model of the Atom (cont.)
  • Each orbit was given a number, called the quantum
    number.

32
Section 5-2
Bohr's Model of the Atom (cont.)
  • Hydrogens single electron is in the n 1 orbit
    in the ground state.
  • When energy is added, the electron moves to the n
    2 orbit.

33
Section 5-2
Bohr's Model of the Atom (cont.)
34
Section 5-2
Bohr's Model of the Atom (cont.)
35
Section 5-2
Bohr's Model of the Atom (cont.)
  • Bohrs model explained the hydrogens spectral
    lines, but failed to explain any other elements
    lines.
  • The behavior of electrons is still not fully
    understood, but it is known they do not move
    around the nucleus in circular orbits.

36
Section 5-2
The Quantum Mechanical Model of the Atom
  • Louis de Broglie (18921987) hypothesized that
    particles, including electrons, could also have
    wavelike behaviors.

37
Section 5-2
The Quantum Mechanical Model of the Atom (cont.)
  • The figure illustrates that electrons orbit the
    nucleus only in whole-number wavelengths.

38
Section 5-2
The Quantum Mechanical Model of the Atom (cont.)
  • The de Broglie equation predicts that all moving
    particles have wave characteristics.

39
Section 5-2
The Quantum Mechanical Model of the Atom (cont.)
  • Heisenberg showed it is impossible to take any
    measurement of an object without disturbing it.
  • The Heisenberg uncertainty principle states that
    it is fundamentally impossible to know precisely
    both the velocity and position of a particle at
    the same time.
  • The only quantity that can be known is the
    probability for an electron to occupy a certain
    region around the nucleus.

40
Section 5-2
The Quantum Mechanical Model of the Atom (cont.)
41
Section 5-2
The Quantum Mechanical Model of the Atom (cont.)
  • Schrödinger treated electrons as waves in a model
    called the quantum mechanical model of the atom.
  • Schrödingers equation applied equally well to
    elements other than hydrogen.

42
Section 5-2
The Quantum Mechanical Model of the Atom (cont.)
  • The wave function predicts a three-dimensional
    region around the nucleus called the atomic
    orbital.

43
Section 5-2
Hydrogen Atomic Orbitals
  • Principal quantum number (n) indicates the
    relative size and energy of atomic orbitals.
  • n specifies the atoms major energy levels,
    called the principal energy levels.

44
Section 5-2
Hydrogen Atomic Orbitals (cont.)
  • Energy sublevels are contained within the
    principal energy levels.

45
Section 5-2
Hydrogen Atomic Orbitals (cont.)
  • Each energy sublevel relates to orbitals of
    different shape.

46
Section 5-2
Hydrogen Atomic Orbitals (cont.)
47
Section 5-2
Section 5.2 Assessment
Which atomic suborbitals have a dumbbell shape?
A. s B. f C. p D. d
  1. A
  2. B
  3. C
  4. D

48
Section 5-2
Section 5.2 Assessment
Who proposed that particles could also exhibit
wavelike behaviors? A. Bohr B. Einstein
C. Rutherford D. de Broglie
  1. A
  2. B
  3. C
  4. D

49
End of Section 5-2
50
Section 5-3
Section 5.3 Electron Configuration
  • Apply the Pauli exclusion principle, the aufbau
    principle, and Hund's rule to write electron
    configurations using orbital diagrams and
    electron configuration notation.
  • Define valence electrons, and draw electron-dot
    structures representing an atom's valence
    electrons.

electron a negatively charged, fast-moving
particle with an extremely small mass that is
found in all forms of matter and moves through
the empty space surrounding an atom's nucleus
51
Section 5-3
Section 5.3 Electron Configuration (cont.)
electron configuration aufbau principle Pauli
exclusion principle Hund's rule valence
electrons electron-dot structure
A set of three rules determines the arrangement
in an atom.
52
Section 5-3
Ground-State Electron Configuration
  • The arrangement of electrons in the atom is
    called the electron configuration.
  • The aufbau principle states that each electron
    occupies the lowest energy orbital available.

53
Section 5-3
Ground-State Electron Configuration (cont.)
54
Section 5-3
Ground-State Electron Configuration (cont.)
  • The Pauli exclusion principle states that a
    maximum of two electrons can occupy a single
    orbital, but only if the electrons have opposite
    spins.
  • Hunds rule states that single electrons with the
    same spin must occupy each equal-energy orbital
    before additional electrons with opposite spins
    can occupy the same energy level orbitals.

55
Section 5-3
Ground-State Electron Configuration (cont.)
56
Section 5-3
Ground-State Electron Configuration (cont.)
  • Noble gas notation uses noble gas symbols in
    brackets to shorten inner electron configurations
    of other elements.

57
Section 5-3
Ground-State Electron Configuration (cont.)
  • The electron configurations (for chromium,
    copper, and several other elements) reflect the
    increased stability of half-filled and filled
    sets of s and d orbitals.

58
Section 5-3
Valence Electrons
  • Valence electrons are defined as electrons in the
    atoms outermost orbitalsthose associated with
    the atoms highest principal energy level.
  • Electron-dot structure consists of the elements
    symbol representing the nucleus, surrounded by
    dots representing the elements valence electrons.

59
Section 5-3
Valence Electrons (cont.)
60
Section 5-3
Section 5.3 Assessment
In the ground state, which orbital does an atoms
electrons occupy? A. the highest
available B. the lowest available C. the n 0
orbital D. the d suborbital
  1. A
  2. B
  3. C
  4. D

61
Section 5-3
Section 5.3 Assessment
The outermost electrons of an atom are called
what? A. suborbitals B. orbitals C. ground
state electrons D. valence electrons
  1. A
  2. B
  3. C
  4. D

62
End of Section 5-3
63
Resources Menu
Chemistry Online Study Guide Chapter
Assessment Standardized Test Practice Image
Bank Concepts in Motion
64
Study Guide 1
Section 5.1 Light and Quantized Energy
Key Concepts
  • All waves are defined by their wavelengths,
    frequencies, amplitudes, and speeds. c ??
  • In a vacuum, all electromagnetic waves travel at
    the speed of light.
  • All electromagnetic waves have both wave and
    particle properties.
  • Matter emits and absorbs energy in
    quanta.Equantum h?

65
Study Guide 1
Section 5.1 Light and Quantized Energy (cont.)
Key Concepts
  • White light produces a continuous spectrum. An
    elements emission spectrum consists of a series
    of lines of individual colors.

66
Study Guide 2
Section 5.2 Quantum Theory and the Atom
Key Concepts
  • Bohrs atomic model attributes hydrogens
    emission spectrum to electrons dropping from
    higher-energy to lower-energy orbits. ?E E
    higher-energy orbit - E lower-energy orbit E
    photon h?
  • The de Broglie equation relates a particles
    wavelength to its mass, its velocity, and
    Plancks constant. ? h / m?
  • The quantum mechanical model of the atom assumes
    that electrons have wave properties.
  • Electrons occupy three-dimensional regions of
    space called atomic orbitals.

67
Study Guide 3
Section 5.3 Electron Configuration
Key Concepts
  • The arrangement of electrons in an atom is called
    the atoms electron configuration.
  • Electron configurations are defined by the aufbau
    principle, the Pauli exclusion principle, and
    Hunds rule.
  • An elements valence electrons determine the
    chemical properties of the element.
  • Electron configurations can be represented using
    orbital diagrams, electron configuration
    notation, and electron-dot structures.

68
Chapter Assessment 1
The shortest distance from equivalent points on a
continuous wave is the A. frequency
B. wavelength C. amplitude D. crest
  1. A
  2. B
  3. C
  4. D

69
Chapter Assessment 2
The energy of a wave increases as ____.
A. frequency decreases B. wavelength decreases
C. wavelength increases D. distance increases
  1. A
  2. B
  3. C
  4. D

70
Chapter Assessment 3
Atoms move in circular orbits in which atomic
model? A. quantum mechanical model
B. Rutherfords model C. Bohrs model
D. plum-pudding model
  1. A
  2. B
  3. C
  4. D

71
Chapter Assessment 4
It is impossible to know precisely both the
location and velocity of an electron at the same
time because A. the Pauli exclusion principle
B. the dual nature of light C. electrons travel
in waves D. the Heisenberg uncertainty
principle
  1. A
  2. B
  3. C
  4. D

72
Chapter Assessment 5
How many valence electrons does neon have? A. 0
B. 1 C. 2 D. 3
  1. A
  2. B
  3. C
  4. D

73
STP 1
Spherical orbitals belong to which sublevel?
A. s B. p C. d D. f
  1. A
  2. B
  3. C
  4. D

74
STP 2
What is the maximum number of electrons the 1s
orbital can hold? A. 10 B. 2 C. 8 D. 1
  1. A
  2. B
  3. C
  4. D

75
STP 3
In order for two electrons to occupy the same
orbital, they must A. have opposite charges
B. have opposite spins C. have the same spin
D. have the same spin and charge
  1. A
  2. B
  3. C
  4. D

76
STP 4
How many valence electrons does boron contain?
A. 1 B. 2 C. 3 D. 5
  1. A
  2. B
  3. C
  4. D

77
STP 5
What is a quantum? A. another name for an atom
B. the smallest amount of energy that can be
gained or lost by an atom C. the ground state
of an atom D. the excited state of an atom
  1. A
  2. B
  3. C
  4. D

78
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CIM
Figure 5.11 Balmer Series Figure 5.12 Electron
Transitions Table 5.4 Electron Configurations
and Orbital Diagrams for Elements 110 Table 5.6
Electron Configurations and Dot Structures
102
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