Atomic Structure

1
Atomic Structure

• Chapter 7
• Describe the properties of electromagnetic
  radiation.
• Understand the origin of light from excited atoms
  and its relationship to atomic structure.
• Describe the experimental evidence for
  wave-particle duality.
• Describe the basic ideas of quantum mechanics.
• Define the three quantum numbers and their
  relationship to atomic structure.

2
Electromagnetic Radiation

• Radiation is _____________!
• List forms of electromagnetic radiation
• _______________ ___________
• _______________ ___________
• Maxwell Theory (1831-1879) describe all forms of
  radiation in terms of ________
• ________________________________.
• Einstein Theory (1879-1955) light has
  _______________________________.

3
Wave Properties
wavelength
Visible light
Ultraviolet radiation
4
Electromagnetic Radiation
Frequency hertz (s-1) Speed wavelength (m) x
frequency (s-1) c l x v
5
What is the frequency of orange light, which has
a wavelength of 625 nm?
Students should be familiar with conversion of
units and conversion between l and v.
6
The Visible Spectrum of Light

• Long wavelength --gt ______ frequency
• _____ energy
• Short wavelength --gt _____ frequency
• _____ energy

7
Energy and Frequency

• Max Planck (1858-1947) the energy of a vibrating
  systems is proportional to the frequency of
  vibration.
• The proportionality constant
• h Plancks constant
• 6.6260693 x 10-34 J s
• E h v

8
Radiation given off by a Heated Body

• Planck solved the ___________________.
• Vibrations are _________ only vibrations with
  specific frequencies are allowed.
• There is a distribution of vibrations in a object.

9
Quantization of Energy

• An object can gain or lose energy by absorbing or
  emitting radiant energy in QUANTA.
• Energy of radiation is proportional to frequency.
• Light with large l (small v) has a _____ E.
• Light with a short l (large v) has a ____ E.
• E h v

10
Photoelectric Effect

• Experiment demonstrates the _______
  _____________________________.

No e- observed until light of a certain minimum
E is used.
11
Photoelectric Effect

• Classical theory said that E of ejected
• electron should increase with increase
• in light frequencynot observed!
• No e- observed until light of a certain
• minimum E is used.
• If the frequency is above the minimum,
• the number of e- ejected depends on
• light intensity.
• Einstein explained the photoelectric effect
  light consists of __________ particles called
  PHOTONS _______________.
• The energy of each photon is proportional to the
  ______________of radiation (Plancks relation).
• The greater the intensity of light, the more
  photons are available to strike per unit of time.

12
Show that the energy of a mol of blue photons (l
400 nm) is higher than the energy of a mol of
red photons (l685 nm)
13
Using Plancks Equation
v c/l
E h v
E h v h c
l
(wavenumber)

• As frequency (v) increases, energy (E)
  __________.
• As wavelength (l) decreases, energy (E) _________.

Students should be familiar with frequency,
wavelength, and energy calculations.
14
What is the color of light when its frequency is
6.0 x 1014 s-1?
15
Photosynthesis

• Chlorophylls absorb blue and red light and
  carotenoids absorb blue-green light, but green
  and yellow light are not effectively absorbed by
  photosynthetic pigments in plants therefore,
  light of these colors is either reflected by
  leaves or passes through the leaves. This is why
  plants are green.

16
Spectrum of White Light
17
Spectrum of Excited Hydrogen Gas

• Excited atoms emit light of only certain
  wavelengths. Evidence of ____________________.
• Line Emission Spectra of Excited Atoms.
• The wavelengths of emitted light depend on
  ______________________________.

18
Which Mathematical Expression represents the
Regular Patterns of Emission?

• Johann Balmer (1825-1898) and Johannes Rydberg
  (1854-1919) developed an equation
• Rydberg equation to calculate the
  _________________
• __________________
• __________________.
• Rydberg constant R
• R 1.0974 x 107 m-1

when n gt 2 n 3 , l red line n 4 , l green
line, Etc. Balmer Series
19
Atomic View of the Early 20th Century

• An electron (e-) traveled about the nucleus in
  an orbit.
• 1. Any orbit should be possible and so is any
  energy.
• 2. But a charged particle moving in an electric
  field should emit energy.
• End result should be matter self-destruction!

20
Bohr Model

• Niels Bohr (1885-1962) connected the observation
  of the spectra of excited atoms with the quantum
  ideas of Planck and Einstein.
• Based on Rutherfords work electrons are
  arranged in space outside the atom.
• Bohr model shows electrons moving in a circular
  orbit around the nucleus.
• Bohr postulated
• 1.- An electron could occupy only __________
  ___________or energy levels in which it is
  stable.
• 2.-The energy of the electron in the atom is
  ______________.

21
Atomic Spectra and Bohr
R h c n 2
Potential energy of electron in the nth level
En -

• n ? ___________ quantum number
• n is a _________________ having values of 1, 2, 3
  and so on.
• The energy of attraction between oppositely
  charged bodies (negative electron and positive
  nuclear proton) has a negative value. The value
  becomes more negative as the bodies move closer
  together (Coulombs law).
• As the value of n increases, the energy becomes
  less negative, the distance of the electron from
  the nucleus increases.

22
Atomic Spectra and Bohr

• Only orbits where n integral number are
  permitted.
• If e-s are in quantized energy states, then ?E
  of states can have only certain values. This
  explain sharp line spectra.

23
Ground State and Excited State

• Ground state The state of an atom in which all
  electrons are in the ______________________.
• Excited state The state of an atom in which at
  least one electron is ______________________
  ____________________.

24
CC alculate DE for an e- of the H atom falling
from high energy level (n 2) to low energy
level (n 1).
25
Atomic Spectra and Bohr

• The amount of energy that must be absorbed by the
  atom so that an electron can move from the first
  to the second energy state is 3/4RhC or 984
  kJ/mol of atoms no more or less energy levels
  in the H atom are quantized only certain
  amounts of energy may be absorbed or emitted.
• When an electron falls from a level of higher n
  to one of lower n, ________ energy. The negative
  sign indicates energy is _________, 984 kJ must
  be _______ per mole of H atoms.
• The energy ________ is observed as ______ This
  is the source of the lines observed in the
  emission spectrum of H atoms. The basic
  explanation holds for the spectra of other
  elements.

26
Atomic Spectra and Bohr
1
1
(
)
-
?E Efinal Einitial -R h c
n2final
n2initial

• The origin of atomic spectra is the movement of
  _________ between quantized energy states.
• Electron is excited from a lower energy state to
  a higher one Energy is ________.
• Electron moves from a higher energy state to a
  lower one Energy is _________.

27
Electronic Transitions in an Excited H Atom

• If electrons move from energy states n gt1 to the
  n 1 state emission lines have energies in the
  UV region (Lyman series).
• If electrons move from energy states n gt2 to the
  n 2 state emission lines have energies in the
  VIS region (Balmer series).
• If electrons move from energy states n gt3 to the
  n 3 state emission lines have energies in the
  IR region.

28
Calculate the wavelength of the photon emitted if
an electron in the H atom moves from n 4 to n 2
29
Flaws in Bohrs Theory

• Bohrs model of the atom explained only the
  spectrum of H atoms and of other systems having
  one electron (such as He).
• The idea that electrons are particles moving
  about the nucleus with a path of fixed radius,
  like that of the planets about the sun, is no
  longer valid.

30
Wave Mechanics

• Louis de Broglie (1892-1987) proposed that all
  moving objects have _______ _________________(1924
  ).
• For light (1) E mc2
• (2) E h v h c / l

31
Wave Mechanics Calculate the Broglie Wavelength
l h
m v

• Baseball (115 g) at 100 mph
• e- with velocity 1.9 x 108 cm/sec
• It is possible to observe wave-like properties
  only for particles of extremely __________, such
  as protons, neutrons, and electrons.

32
The Uncertainty Principle

• Erwin Schrödinger, 1887-1961 developed
  ________________or ______________.
• Werner Heisenberg, 1901-1976 The uncertainty
  principle it is impossible to fix both the
  ______________ electron in an atom and its
  ________ with any degree of certainty.
• Max Born, 1882-1970 if the energy of an
  electron in an atom is known with a small
  uncertainty, there will be large uncertainty in
  its position in the space about the atom's
  nucleus.
• We can assess only the likelihood, or
  probability, of finding an electron with a given
  energy within a given region of space.

33
Schrödinger's Wave Functions

• The behavior of the electron in the atom is best
  described as a standing wave In a vibrating
  string, only certain vibrations can be observed
  only certain wave functions are allowed for the
  electron in the atom.
• Each wave function (?) is associated with an
  allowed energy value, En, for the electron.
• Then, from 1 and 2, the energy of the electron is
  quantized only certain values of energy.

Wave motionwave length and nodes
4. In contrast to Bohrs theory quantization is
imposed as a postulate.
34
Schrödinger's Wave Functions

• 5. The is related to the probability of finding
  the electron within a given region of space
  _______________.
• 6. Energy is known precisely position is given
  by a probability. The region of space in which an
  electron of a given energy is most probably
  located is called its _______________.
• 7. The solution to the Schrödinger's equation,
  for an electron, in a 3-D space, are 3 integer
  numbers quantum numbers n, l, and ml. These
  numbers have only certain combination of values.

35
Quantum numbers

• n, Principal quantum number 1, 2, 3,
• Determines the ________ of the electron. Also
  related to size of orbital.
• En - Z2h R / n2
• Electrons with the same n value are in the same
  electron ______ or same electron _________.
• l, Angular Momentum quantum number 0, 1, 2, 3,
  , n-1
• Determines the ______ at which electrons
  circulate about the nucleus. Related to orbital
  __________.
• Electrons with the same l value are in the same
  _______ and have the same orbital _____
  (______). All orbitals in the same subshell have
  the same ___________.
• ml, Magnetic quantum number 0, 1, 2, 3,
  , l
• Determines the _____________ of the orbital
  motion of the electron. (Clockwise or
  counterclockwise). Related to ___________ in
  space of the orbitals within a subshell, this
  gives the ___________ of orbitals in a subshell.
• See Table 7.1 (p 319)

36
Quantum numbers and Orbitals

• Number of subshells in a shell n
• Number of orbitals in a subshell 2l 1
• Number of orbitals in a shell n2
• l 0 (s) l 1 (p) l 2 (d) l 3 (f)
• Name of orbital value of n and letter code for
  l
• If n1 l n-1 0 ml 0
• Only 1 subshell (s) only 1 orbital (1s)
• If n2 l 0, 1 ml 1, 0, -1
• There are 2 subshells (s and p)
• 4 orbitals (the 2s, and three 2p (3 orientations)

37
Orbitals

• Electron orbitals are probabilities represented
  as ____________________.

38
Orbitals
surface density plot or radial distribution plot

• For the s orbital, the probability of finding an
  electron is the same at the same distance from
  the nucleus the 1s orbital is ____________ in
  shape.
• Quantum mechanics electron has wave properties
  the maximum amplitude of the electron wave
  occurs at 0.053 nm from the nucleus.
• Bohrs radius 0.059 nm

39
Orbitals

• The p orbitals have 1 nodal surface zero
  probability of finding an electron.
• Number of nodal surfaces value of l
• There are three p orbitals in each p subshell ml
  1, 0, -1
• Refer to orbitals according to the axes along
  which the lobes lie px, py, pz

40
Orbitals

• The d five orbitals, l2 have 2 nodal surfaces
  (may not be flat).
• What type of orbital is designated n 4, l 3,
  ml -3?
• a. 4s
• b. 4p
• c. 4d
• d. 4f
• e. none

41
Orbitals
Students should be familiar with definitions of
quantum numbers and orbital types.
42
Practice

• Which of the following represent valid sets of
  quantum numbers?
• n3, l3, ml 1
• n5, l1
• n6, l5, ml1
• n4, l3, ml-4

43
Remember

• Go over all the contents of your textbook.
• Practice with examples and with problems at the
  end of the chapter.
• Practice with OWL tutors.
• W ork on your assignment for Chapter 7.
• Practice with the quiz on the cd or online
  service.