Electrons in Atoms

1
Electrons in Atoms

• Light Quantized Energy
• Quantum Theory the Atom
• Electron Configuration

2
Unanswered ?s

• Rutherford
• Lacked detail
• How electrons occupy space?
• Why electrons not pulled into nucleus?
• Early 1900s
• Certain chemical emit visible light when heated
  in flame
• Related to the arrangement of electrons

3
Wave Nature of Light

• Electromagnetic radiation
• Form of energy that exhibits wave like behavior
  as it travels through space
• Visible light, microwaves, X-rays, radio waves

4
Wave Nature of Light

• Wavelength
• The shortest distance between equivalent points
  on a continuous wave
• Greek letter Lamda (?)
• Expressed in meters, centimeters, or nanometers

5
Wave Nature of Light

• Frequency
• The number of waves that pass a given point per
  second
• Greek letter nu (?)
• Expressed in Hertz (Hz)
• One hertz is equal to one wave per second

6
Wave Nature of Light

• Amplitude
• Waves height from the origin to a crest, or from
  the origin to the trough

7
Speed of Light

• 3.00 x 108 m/s in a vacuum
• Given its own value
• c
• Product of wavelength (?) frequency (?)
• c ??

8
? ?

• Inversely related
• As one quantity decreases the other increases

9
Electromagnetic Spectrum

• EM spectrum
• Encompasses all forms of electromagnetic
  radiation
• Only differences in types of radiation is their
  frequency and wavelength

10
EM Spectrum
11
Particle Nature of Light

• Seeing light as a wave helps explain many things
• Does not describe how light interacts with matter
• New model is needed to address issues of phenomena

12
HOT!!!!

• WHY?
• As temperature increases
• Contains a greater amount of energy
• Different colors different wavelengths

13
Max Planck

• German physicist
• Studied light emitted from heated objects
• What did he find?
• Matter can gain or lose energy only in small
  amounts
• Called quanta

14
Quantum

• Quantum
• The minimum amount of energy that can be gained
  or lost by an atom
• Idea revolutionary to many and disturbing to many
  more
• Believed that energy could be absorbed and
  emitted in varying quantities
• Water temperature increases in infinitesimal
  steps as molecules absorb quanta
• Steps are so small, seems to rise continuously

15
Emitted light from glowing objects

• Planck proposed that light energy was quantized
• Energy of a quantum is related to frequency of
  emitted radiation
• Equantum hv
• E Energy
• v frequency
• h Plancks constant
• (6.626 x 10-34 J s)

16
Plancks Theory

• For a given frequency, matter can emit or absorb
  energy only in whole-number multiples
• Matter can only have certain amounts of energy,
  quantities between do not exist

17
Photoelectric Effect

• Electrons, called photoelectrons, are emitted
  from a metals surface when light of a certain
  frequency shines on the surface

18
Photoelectric Effect

• Einstein proposed that electromagnetic radiation
  has both wavelike and particlelike natures
• Light may have many wavelike characteristics, it
  also can be thought of as a stream of tiny
  particles, or bundles of energy
• Photon
• Particle of electromagnetic radiation with no
  mass and that carries a quantum of energy

19
Photons Energy

• Ephoton hv
• Einstein energy of a photon of light must have
  a certain minimum, or threshold, value to cause
  the ejection of a photoelectron

20
Atomic Emission Spectrum

• Neon lights??
• Cannot be explained by wave model
• Atomic Emission Spectrum
• Of an element is the set of frequencies of the
  electromagnetic waves emitted by atoms of the
  element

21
Atomic Emission Spectrum
22
Atomic Emission Spectrum

• Can be used to identify an element
• Or determine if that is part of an unknown
  compound
• Ephoton hv
• Not predicted by laws of classical physics
• Expected to observe emission in continuous series
  of colors as excited electrons lost energy

23
Quantum Theory

• Behavior of light
• Only by wave-particle model
• Good but
• Still need to explain atomic structure,
  electrons, atomic emission

24
Bohr Model of Atom

• Hydrogen emission discontinuous not continuous
• Why?
• Niels Bohr
• Quantum model for hydrogen
• Correctly predicted the frequencies of the lines
  of hydrogen emission

25
Bohrs Model

• Building on Planck Einsteins ideas of
  quantized energy (certain values allowed)
• Hydrogen only has certain allowable energy states
• Ground state
• The lowest allowable energy state of an atom

26
Bohr goes further

• Relates atoms energy state to the motion of the
  electron within the atom
• Single electron in H, moves around the nucleus in
  only certain allowed circular orbits

27
Bohrs Model

• Smaller electrons orbit in lowest energy states,
  larger electrons orbit in highest energy states
• Assigned a quantum number, n, to each orbit and
  calculated orbits radius
• n 1, radius 0.0529 nm
• n 2, radius 0.212 nm
• n 3, radius 0.476 nm

28
Hydrogen atom in ground state

• First energy level
• n 1
• In ground state atom does not radiate energy
• Added energy from outside source
• electron moves to a higher-energy orbit
• n 2
• Atom in excited state can drop from a
  higher-energy level to a lower-energy orbit

29
Hydrogen atom in ground state

• Higher to lower energy orbit
• Atom emits a photon corresponding to the
  difference between the energy levels associated
  with the two orbits
• E Ehigher-energy orbit Elower-energy orbit
  Ephoton hv

30
Hydrogen atom in ground state

• Only certain atomic energies are possible
• Only certain frequencies of EM radiation can be
  seen
• Four electron transmission accounts for visible
  lines in hydrogens atomic emission spectrum

31
Hydrogen atom in ground state

• Visible lines
• Balmer series
• Ultraviolet light
• Electrons drop into the n 1 orbit
• Lyman series
• Infrared light
• Electrons drop into the n 3 orbit
• Paschen series

32
PROBLEM!

• Bohrs models worked great for H
• Fails to explain any other element
• Did not fully account for chemical behavior
• Quantized energy levels
• Laid the ground work for atomic models to come
• Movement of electrons not completely understood
• Much evidence shows that electrons do not move in
  circular orbits

33
Quantum Mechanical Model

• After Bohr
• How were electrons arranged in the atom?
• Louis de Broglie
• French graduate student in physics
• Idea to explain the fixed energy levels of Bohr

34
Electrons as waves

• Bohrs quantized electron orbits had
  characteristics similar to those of waves
• De Broglie
• Only whole numbers of waves are allowed in a
  circular orbit of fixed radius
• Light contain both wave and particle
  characteristics

35
Electrons as waves

• If an electron has wave-like motion is
  restricted to circular orbits of fixed radius
• The electron is allowed certain possible
  wavelengths, frequencies, and energies
• de Broglie equation
• ? h / mv
• Predicts that all moving particles have wave
  characteristics

36
Heisenberg Uncertainty Principle

• It is impossible to make any measurement on an
  object without disturbing the object
• Helium balloon in dark room
• Photons help to determine position?
• Photon has same energy as electron
• The interaction between the two particles changes
  both the wavelength of the photon and the
  position and velocity of the electron

37
Heisenberg Uncertainty Principle

• States that it is fundamentally impossible to
  know precisely the velocity and position of a
  particle at the same time
• The act of observing the electron produces a
  significant, unavoidable uncertainty in the
  position motion of the electron

38
Erwin Schrödinger

• Austrian physicist
• Furthered the wave-particle theory
• New model for Hydrogen applied equally well to
  all elements
• Where Bohrs failed

39
Quantum Mechanical Model

• The atomic model in which electrons are treated
  as waves
• Called the wave mechanical model of the atom, or
• Quantum Mechanical Model of the Atom
• Limits an electrons energy to certain values
• Makes no attempt to describe the electrons path
  around the nucleus

40
Quantum Mechanical Model

• Equation too complex to be discussed here
• Solution is known as a wave function
• This is related to probability of finding the e-
  within a particular volume of space surrounding
  the nucleus
• Gives higher probability to find location of e-
• Atomic Orbital
• A three-dimensional region around the nucleus
• Describes the electrons probable location

41
Quantum Mechanical Model

• Picture as a fuzzy cloud in which the density
  of the cloud at a given point is proportional to
  the probability of finding the electron at that
  point.

42
Hydrogens Atomic Orbitals

• Orbital does not have an exact definite size
• Principal Quantum Numbers (n)
• Indicate the relative size and energies of atomic
  orbitals
• As n increases, the orbital becomes larger
• The electron spends more time further away from
  the nucleus, the atoms energy increases

43
Principal Energy Levels

• The major energy levels of an atom
• Lowest principal energy level assigned n 1
• Ground state
• Up to 7 energy levels have been detected for the
  H atom
• Giving n values ranging from 1 to 7

44
Energy Sublevels

• The energy levels contained within a principal
  energy level
• Level 1 single sublevel
• Level 2 2 sublevels
• Level 3 3 sublevels
• and so on
• Stadium Seating

45
Sublevels

• Labeled according to shape of atoms orbitals
• s, p, d, f
• s orbitals are spherical
• p orbitals are dumbbell
• d f orbitals are not all same shape

46
Sublevels s p

• Principal energy level 1
• 1s
• Spherical 1s orbital
• Principal energy level 2
• 2s 2p
• Larger spherical 1s 3 p orbitals
• 2px, 2py, 2pz
• x, y, z??

47
Sublevel d

• Principal energy level 3
• 3s, 3p, 3d
• Each d sublevel has 5 orbital of equal energy
• dxy, dxz, dyz, dx2-y2, dz2
• 4 have identical shape but different orientations
• 5th d orbital

48
Sublevel f

• Principal energy level 4
• 4f
• Consists of 7 f orbital of equal energy

49
(No Transcript)
50
Hydrogen only has 1 e-??

• At any given time, an electron can occupy just
  one orbit
• Mostly in ground state 1s
• Gain a quantum of energy
• Excites electron to 2s, one of 2p, or any other
  unoccupied orbitals

51
Electron Configuration

• The arrangement of electrons in an atom
• Low-energy systems more stable
• Aufbau principle
• Pauli exclusion principle
• Hunds rule

52
Aufbau Principle

• States that each electron occupies the lowest
  energy orbital available
• All orbitals related to an energy sublevel are of
  equal energy
• In a mutli-electron atom, the energy sublevels
  with a principal energy level have different
  energies
• In order of increasing energy, the sequence of
  energy sublevels within a principle energy level
  is s, p, d, f
• Orbitals related to energy sublevels within one
  principal energy level can overlap orbitals
  releated to energy sublevels within another
  principle level

53
Arsenic??
54
Pauli Exclusion Principle

• States that a maximum of two electrons may occupy
  a single atomic orbital, but only if the
  electrons have opposite spins
• Each electron can only spin in one of two
  directions

55
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 orbital

56
Orbital Diagram Electron Configuration Notation

• Represent an atoms electron configuration
• Orbital diagram
• Box for each of the atoms orbital
• Arrows in boxes represent electrons

57
Electron Configuration Notation

• This method designates the principle energy level
  and energy sublevel associated with each of the
  atoms orbitals and includes a superscript
  representing the number of electrons in the
  orbital
• Carbon
• 1s22s22p2
• Neon
• 1s22s22p6

58
Electron Configuration Notation

• Noble-gas notation
• Use noble-gas from previous period
• Electron configuration for energy level being
  filled
• Chlorine
• 1s22s22p63s23p5
• Ne3s23p5

59
Valence Electrons

• Electrons in the atoms outermost orbitals
• Generally those orbitals associated with the
  atoms highest principle energy level
• Sulfur
• S Ne3s23p4
• 6 valence electrons
• Cesium
• Cs Xe6s1
• 1 valence electron

60
Electron-dot Structure

• Consists of the elements symbol
• Represents the atomic nucleus and inner-level
  electrons
• Surrounded by dots
• Represents the atoms valence electrons
• Silicon
• Fluorine
• Calcium