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Atomic Structure

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Atomic Structure From Indivisible to Quantum Mechanical Model of the Atom Classical Model Democritus Dalton Thomson Rutherford Democritus Circa 400 BC Greek ... – PowerPoint PPT presentation

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Title: Atomic Structure


1
Atomic Structure
  • From Indivisible to Quantum Mechanical Model of
    the Atom

2
Classical Model
  • Democritus
  • Dalton
  • Thomson
  • Rutherford

3
Democritus
  • Circa 400 BC
  • Greek philosopher
  • Suggested that all matter is composed of tiny,
    indivisible particles, called atoms

4
Daltons Atomic Theory (1808)
  1. All matter is made of tiny indivisible particles
    called atoms.
  2. Atoms of the same element are identical. The
    atoms of any one element are different from those
    of any other element.
  3. Atoms of different elements can combine with one
    another in simple whole number ratios to form
    compounds.
  4. Chemical reactions occur when atoms are
    separated, joined, or rearrangedhowever, atoms
    of one element are not changed into atoms of
    another by a chemical reaction.

5
J.J. Thomson (1897)
  • Determined the charge to mass ratio for electrons
  • Applied electric and magnetic fields to cathode
    rays
  • Plum pudding model of the atom

6
Rutherfords Gold Foil Experiment (1910)
  • Alpha particles (positively charged helium ions)
    from a radioactive source was directed toward a
    very thin gold foil.
  • A fluorescent screen was placed behind the Au
    foil to detect the scattering of alpha (?)
    particles.

7
(No Transcript)
8
Rutherfords Gold Foil Experiment (Observations)
  • Most of the ?-particles passed through the foil.
  • Many of the ?-particles deflected at various
    angles.
  • Surprisingly, a few particles were deflected back
    from the Au foil.

9
Rutherfords Gold Foil Experiment (Conclusions)
  • Rutherford concluded that most of the mass of an
    atom is concentrated in a core, called the atomic
    nucleus.
  • The nucleus is positively charged.
  • Most of the volume of the atom is empty space.

10
Shortfalls of Rutherfords Model
  • Did not explain where the atoms negatively
    charged electrons are located in the space
    surrounding its positively charged nucleus.
  • We know oppositely charged particles attract each
    other
  • What prevents the negative electrons from being
    drawn into the positive nucleus?

11
Bohr Model (1913)
  • Niels Bohr (1885-1962), Danish scientist working
    with Rutherford
  • Proposed that electrons must have enough energy
    to keep them in constant motion around the
    nucleus
  • Analogous to the motion of the planets orbiting
    the sun

12
Planetary Model
  • The planets are attracted to the sun by
    gravitational force, they move with enough energy
    to remain in stable orbits around the sun.
  • Electrons have energy of motion that enables them
    to overcome the attraction for the positive
    nucleus

13
Think about satellites.
  • We launch a satellite into space with enough
    energy to orbit the earth
  • The amount of energy it is given, determines how
    high it will orbit
  • We use energy from a rocket to boost our
    satellite, what energy do we give electrons to
    boost them?

14
Electronic Structure of Atom
  • Waves-particle duality
  • Photoelectric effect
  • Plancks constant
  • Bohr model
  • de Broglie equation

15
Radiant Energy
  • Radiation ? the emission of energy in various
    forms
  • A.K.A. Electromagnetic Radiation
  • Radiant Energy travels in the form of waves that
    have both electrical and magnetic impulses

16
  • Electromagnetic Radiation ? radiation that
    consists of wave-like electric and magnetic
    fields in space, including light, microwaves,
    radio signals, and x-rays
  • Electromagnetic waves can travel through empty
    space, at the speed of light (c3.00x108m/s) or
    about 300million m/s!!!

17
Waves
  • Waves transfer energy from one place to another
  • Think about the damage done by waves during
    strong hurricanes.
  • Think about placing a tennis ball in your bath
    tub, if you create waves at one it, that energy
    is transferred to the ball at the other bobbing
  • Electromagnetic waves have the same
    characteristics as other waves
  •  

18
Wave Characteristics
Wavelength, ? (lambda) ? distance between
successive points
2m
10m
19
Wave Characteristics
  • Frequency, ? (nu) ? the number of complete wave
    cycles to pass a given point per unit of time
    Cycles per second

t5
t0
t0
t5
20
Units for Frequency
  • 1/s
  • s-1
  • hertz, Hz
  • Because all electromagnetic waves travel at the
    speed of light, wavelength is determined by
    frequency
  • Low frequency long wavelengths
  • High frequency short wavelengths

21
Waves
  • Amplitude ? maximum height of a wave

22
Waves
  • Node ? points of zero amplitude

23
Electromagnetic Spectrum
  • Radio TV, microwaves, UV, infrared, visible
    light all are examples of electromagnetic
    radiation (and radiant energy)
  • Electromagnetic spectrum entire range of
    electromagnetic radiation

24
Electromagnetic Spectrum
Frequency Hz
1024 1020
1018 1016 1014
1012 1010 108 106
Gamma Xrays UV
Microwaves FM AM
IR
10-16 10-9 10-8 10-6
10-3 100 102
Wavelength m
Visible Light
25
Notes
  • Higher-frequency electromagnetic waves have
    higher energy than lower-frequency
    electromagnetic waves
  • All forms of electromagnetic energy interact with
    matter, and the ability of these different waves
    to penetrate matter is a measure of the energy of
    the waves

26
What is your favorite radio station?
  • Radio stations are identified by their frequency
    in MHz.
  • We know all electromagnetic radiation(which
    includes radio waves) travel at the speed of
    light.
  • What is the wavelength of your favorite station?

27
Velocity of a Wave
  • Velocity of a wave (m/s) wavelength (m) x
    frequency (1/s)
  • c ??
  • c speed of light 3.00x108 m/s
  • My favorite radio station is 105.9 Jamming
    Oldies!!!
  • What is the wavelength of this FM station?

28
Wavelength of FM
  • c ??
  • c speed of light 3.00x108 m/s
  • ? 105.9MHz or 1.059x108Hz
  • ? c/? 3.00x108 m/s 2.83m
  • 1.059x1081/s

29
What does the electromagnetic spectrum have to do
with electrons?
  • Its all related to energy energy of motion(of
    electrons) and energy of light

30
States of Electrons
  • When current is passed through a gas at a low
    pressure, the potential energy (energy due to
    position) of some of the gas atoms increases.
  • Ground State the lowest energy state of an atom
  • Excited State a state in which the atom has a
    higher potential energy than it had in its ground
    state

31
Neon Signs
  • When an excited atom returns to its ground state
    it gives off the energy it gained in the form of
    electromagnetic radiation!
  • The glow of neon signs,is an example of this
    process

32
White Light
  • White light is composed of all of the colors of
    the spectrum ROY G BIV
  • When white light is passed through a prism, the
    light is separated into a spectrum, of all the
    colors
  • What are rainbows?

33
Line-emission Spectrum
  • When an electric current is passed through a
    vacuum tube containing H2 gas at low pressure,
    and emission of a pinkish glow is observed.
  • What do you think happens when that pink glow is
    passed through a prism?

34
Hydrogens Emission Spectrum
  • The pink light consisted of just a few specific
    frequencies, not the whole range of colors as
    with white light
  • Scientists had expected to see a continuous range
    of frequencies of electromagnetic radiation,
    because the hydrogen atoms were excited by
    whatever amount of energy was added to them.
  • Lead to a new theory of the atom

35
Bohrs Model of Hydrogen Atom
  • Hydrogen did not produce a continuous spectrum
  • New model was needed
  • Electrons can circle the nucleus only in allowed
    paths or orbits
  • When an e- is in one of these orbits, the atom
    has a fixed, definite energy
  • e- and hydrogen atom are in its lowest energy
    state when it is in the orbit closest to the
    nucleus

36
Bohr Model Continued
  • Orbits are separated by empty space, where e-
    cannot exist
  • Energy of e- increases as it moves to orbits
    farther and farther from the nucleus
  • (Similar to a person climbing a ladder)

37
Bohr Model and Hydrogen Spectrum
  • While in orbit, e- can neither gain or lose
    energy
  • But, e- can gain energy equal to the difference
    between higher and lower orbitals, and therefore
    move to the higher orbital (Absorption)
  • When e- falls from higher state to lower state,
    energy is emitted (Emission)

38
Bohrs Calculations
  • Based on the wavelengths of hydrogens
    line-emission spectrum, Bohr calculated the
    energies that an e- would have in the allowed
    energy levels for the hydrogen atom

39
Photoelectric Effect
  • An observed phenomenon, early 1900s
  • When light was shone on a metal, electrons were
    emitted from that metal
  • Light was known to be a form of energy, capable
    of knocking loose an electron from a metal
  • Therefore, light of any frequency could supply
    enough energy to eject an electron.

40
Photoelectric Effect pg. 93
  • Light strikes the surface of a metal (cathode),
    and e- are ejected.
  • These ejected e- move from the cathode to the
    anode, and current flows in the cell.
  • A minimum frequency of light is used. If the
    frequency is above the minimum and the intensity
    of the light is increased, more e- are ejected.

41
Photoelectric Effect
  • Observed For a given metal, no electrons were
    emitted if the lights frequency was below a
    certain minimum, no matter how long the light was
    shone
  • Why does the light have to be of a minimum
    frequency?

42
Explanation.
  • Max Planck studied the emission of light by hot
    objects
  • Proposed objects emit energy in small, specific
    amounts quanta
  • (Differs from wave theory which would say objects
    emit electromagnetic radiation continuously)
  • Quantum is the minimum quantity of energy that
    can be lost or gained by an atom.

43
Plancks Equation
  • E radiation Plancks constant x frequency of
    radiation
  • E h?
  • h Plancks constant 6.626 x 10-34 Js
  • When an object emits radiation, there must be a
    minimum quantity of energy that can be emitted at
    any given time.

44
Einstein Expands Plancks Theory
  • Theorized that electromagnetic radiation had a
    dual wave-particle nature!
  • Behaves like waves and particles
  • Think of light as particles that each carry one
    quantum of energy photons

45
Photons
  • Photons a particle of electromagnetic radiation
    having zero mass and carrying a quantum of energy
  • Ephoton h?

46
Back to Photoelectric Effect
  • Einstein concluded
  • Electromagnetic radiation is absorbed by matter
    only in whole numbers of photons
  • In order for an e- to be ejected, the e- must be
    struck by a single photon with minimum frequency

47
Example of Plancks Equation
  • CD players use lasers that emit red light with a
    ? of 685 nm. Calculate the energy of one photon.
  • Different metals require different minimum
    frequencies to exhibit photoelectric effect

48
Answer
  • Ephoton h?
  • h Plancks constant 6.626 x 10-34 Js
  • c ??
  • c speed of light 3.00x108 m/s
  • ? (3.00x108 m/s)/(6.85x10-7m)
  • ?4.37x10141/s
  • Ephoton (6.626 x 10-34 Js)(4.37x10141/s)
  • Ephoton 2.90 x 10-19J

49
Wave Nature of Electrons
  • We know electrons behave as particles
  • In 1925, Louis de Broglie suggested that
    electrons might also display wave properties

50
de Broglies Equation
  • A free e- of mass (m) moving with a velocity (v)
    should have an associated wavelength ? h/mv
  • Linked particle properties (m and v) with a wave
    property (?)

51
Example of de Broglies Equation
  • Calculate the wavelength associated with an e- of
    mass 9.109x10-28 g traveling at 40.0 the speed
    of light.
  • 1 J 1 kg m2/s2

52
Answer
  • C(3.00x108m/s)(.40)1.2x108m/s
  • ? h/mv
  • ? (6.626 x 10-34 Js) 6.06x10-12m
  • (9.11x10-31kg)(1.2x108m/s)
  • Remember 1J 1(kg)(m)2/s2

53
Wave-Particle Duality
  • de Broglies experiments suggested that e- has
    wave-like properties.
  • Thomsons experiments suggested that e- has
    particle-like properties
  • measured charge-to-mass ratio

54
Quantum mechanical model
  • SchrÖdinger
  • Heisenberg
  • Pauli
  • Hund

55
Where are the e- in the atom?
  • e- have a dual wave-particle nature
  • If e- act like waves and particles at the same
    time, where are they in the atom?
  • First consider a theory by German theoretical
    physicist, Werner Heisenberg.

56
Heisenbergs Idea
  • e- are detected by their interactions with
    photons
  • Photons have about the same energy as e-
  • Any attempt to locate a specific e- with a photon
    knocks the e- off its course
  • ALWAYS a basic uncertainty in trying to locate an
    e-

57
Heisenbergs Uncertainty Principle
  • Impossible to determine both the position and the
    momentum of an e- in an atom simultaneously with
    great certainty.

58
SchrÖdingers Wave Equation
  • An equation that treated electrons in atoms as
    waves
  • Only waves of specific energies, and therefore
    frequencies, provided solutions to the equation
  • Quantization of e- energies was a natural outcome

59
SchrÖdingers Wave Equation
  • Solutions are known as wave functions
  • Wave functions give ONLY the probability of
    finding and e- at a given place around the
    nucleus
  • e- not in neat orbits, but exist in regions
    called orbitals

60
SchrÖdingers Wave Equation
  • Here is the equation
  • Dont memorize this or write it down
  • It is a differential equation, and we need
    calculus to solve it
  • -h (?2 ? ) (?2? )( ?2? ) V? E?
  • 8(p)2m (?x2) (?y2) (?z2 )
  • Scary???

61
Definitions
  • Probability ? likelihood
  • Orbital ? wave function region in space where
    the probability of finding an electron is high
  • SchrÖdingers Wave Equation states that orbitals
    have quantized energies
  • But there are other characteristics to describe
    orbitals besides energy

62
Quantum Numbers
  • Definition specify the properties of atomic
    orbitals and the properties of electrons in
    orbitals
  • There are four quantum numbers
  • The first three are results from SchrÖdingers
    Wave Equation

63
Quantum Numbers (1)
  • Principal Quantum Number, n

64
Quantum Numbers
  • Principal Quantum Number, n
  • Values of n 1,2,3, ?
  • Positive integers only!
  • Indicates the main energy level occupied by the
    electron

65
Quantum Numbers
  • Principal Quantum Number, n
  • Values of n 1,2,3, ?
  • Describes the energy level, orbital size

66
Quantum Numbers
  • Principal Quantum Number, n
  • Values of n 1,2,3, ?
  • Describes the energy level, orbital size
  • As n increases, orbital size increases.

67
Principle Quantum Number
n6
n5
n4
n3
Energy
n2
n 1
68
Principle Quantum Number
  • More than one e- can have the same n value
  • These e- are said to be in the same e- shell
  • The total number of orbitals that exist in a
    given shell n2

69
Quantum Numbers (2)
  • Angular momentum quantum number, l

70
Quantum Numbers
  • Angular momentum quantum number, l
  • Values of l n-1, 0

71
Quantum Numbers
  • Angular momentum quantum number, l
  • Values of l n-1, 0
  • Describes the orbital shape

72
Quantum Numbers
  • Angular momentum quantum number, l
  • Values of l n-1, 0
  • Describes the orbital shape
  • Indicates the number of sublevel (subshells)
  • (except for the 1st main energy level, orbitals
    of different shapes are known as sublevels or
    subshells)

73
Orbital Shapes
  • For a specific main energy level, the number of
    orbital shapes possible is equal to n.
  • Values of l n-1, 0
  • Ex. Orbital which n2, can have one of two shapes
    corresponding to l 0 or l1
  • Depending on its value of l, an orbital is
    assigned a letter.

74
Orbital Shapes
  • Angular magnetic quantum number, l
  • If l 0, then the orbital is labeled s.
  • s is spherical.

75
Orbital Shapes
  • If l 1, then the orbital is labeled p.
  • dumbbell shape

76
Orbital Shapes
  • If l 2, the orbital is labeled d.
  • double dumbbell or four-leaf clover

77
Orbital Shapes
  • If l 3, then the orbital is labeled f.

78
Energy Level and Orbitals
  • n1, only s orbitals
  • n2, s and p orbitals
  • n3, s, p, and d orbitals
  • n4, s,p,d and f orbitals
  • Remember l n-1

79
Atomic Orbitals
  • Atomic Orbitals are designated by the principal
    quantum number followed by letter of their
    subshell
  • Ex. 1s s orbital in 1st main energy level
  • Ex. 4d d sublevel in 4th main energy level

80
Quantum Numbers (3)
  • Magnetic Quantum Number, ml

81
Quantum Numbers
  • Magnetic Quantum Number, ml
  • Values of ml l0-l

82
Quantum Numbers
  • Magnetic Quantum Number, ml
  • Values of ml l0-l
  • Describes the orientation of the orbital
  • Atomic orbitals can have the same shape but
    different orientations

83
Magnetic Quantum Number
  • s orbitals are spherical, only one orientation,
    so m0
  • p orbitals, 3-D orientation, so m -1, 0 or 1 (x,
    y, z)
  • d orbitals, 5 orientations, m -2,-1, 0, 1 or 2

84
Quantum Numbers (4)
  • Electron Spin Quantum Number,ms

85
Quantum Numbers
  • Electron Spin Quantum Number,ms
  • Values of ms 1/2 or 1/2
  • e- spin in only 1 or 2 directions
  • A single orbital can hold a maximum of 2 e-,
    which must have opposite spins

86
Electron Configurations
  • Electron Configurations arragenment of e- in an
    atom
  • There is a distinct electron configuration for
    each atom
  • There are 3 rules to writing electron
    configurations

87
Pauli Exclusion Principle
  • No 2 e- in an atom can have the same set of four
    quantum numbers (n, l, ml, ms ). Therefore, no
    atomic orbital can contain more than 2 e-.

88
Aufbau Principle
  • Aufbau Principle an e- occupies the lowest
    energy orbital that can receive it.
  • Aufbau order

89
Hunds Rule
  • Hunds Rule orbitals of equal energy are each
    occupied by one e- before any orbital is occupied
    by a second e-, and all e- in singly occupied
    orbitals must have the same spin

90
Electron Configuration
  • The total of the superscripts must equal the
    atomic number (number of electrons) of that atom.
  • The last symbol listed is the symbol for the
    differentiating electron.

91
Differentiating Electron
  • The differentiating electron is the electron that
    is added which makes the configuration different
    from that of the preceding element.
  • The last electron.
  • H 1s1
  • He 1s2
  • Li 1s2, 2s1
  • Be 1s2, 2s2
  • B 1s2, 2s2, 2p1

92
Orbital Diagrams
  • These diagrams are based on the electron
    configuration.
  • In orbital diagrams
  • Each orbital (the space in an atom that will hold
    a pair of electrons) is shown.
  • The opposite spins of the electron pair is
    indicated.

93
Orbital Diagram Rules
  • Represent each electron by an arrow
  • The direction of the arrow represents the
    electron spin
  • Draw an up arrow to show the first electron in
    each orbital.
  • Hunds Rule Distribute the electrons among the
    orbitals within sublevels so as to give the most
    unshared pairs.
  • Put one electron in each orbital of a sublevel
    before the second electron appears.
  • Half filled sublevels are more stable than
    partially full sublevels.

94
Orbital Diagram Examples
  • H ?_
  • 1s
  • Li ?? ?_
  • 1s 2s
  • B ?? ?? ? __ __
  • 1s 2s 2p
  • N ?? ?? ? ? ?_
  • 1s 2s 2p

95
Dot Diagram of Valence Electrons
  • When two atom collide, and a reaction takes
    place, only the outer electrons interact.
  • These outer electrons are referred to as the
    valence electrons.
  • Because of the overlaying of the sublevels in the
    larger atoms, there are never more than eight
    valence electrons.

96
Rules for Dot Diagrams
  • Xy

. .
Px orbital
S sublevel electrons
. .
Py orbital
Pz orbital
97
Rules for Dot Diagrams
  • Remember the maximum number of valence
    electrons is 8.
  • Only s and p sublevel electrons will ever be
    valence electrons.
  • Put the dots that represent the s and p electrons
    around the symbol.
  • Use the same rule (Hunds rule) as you fill the
    designated orbitals.

98
Examples of Dot Diagrams
  • H
  • He
  • Li
  • Be

99
Examples of Dot Diagrams
  • C
  • N
  • O
  • Xe

100
Summary
  • Both dot diagrams and orbital diagrams will be
    use full to use when we begin our study of atomic
    bonding.
  • We have been dealing with valence electrons since
    our initial studies of the ions.
  • The number of valence electrons can be determined
    by reading the column number.
  • Al 3 valence electrons
  • Br 7 valence electrons
  • All transitions metals have 2 valence electrons.
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