Electronic Structure and Periodic Trends

1
Chapter 9

• Electronic Structure and Periodic Trends

2
Homework

• Assigned Problems (odd numbers only)
• Questions and Problems 9.1 to 9.71 (begins on
  page 258)
• Additional Questions and Problems 9.81 to
  9.115 (page 284-286)
• Challenge Questions 9.119, 9.121, 9.123 (page
  286)

3
Electromagnetic Radiation

• Matter is anything that has mass and occupies
  space.
• Nearly all changes that matter undergoes requires
  the absorption or release of energy
• Energy is the capacity to do work
• The process of moving matter against an opposing
  force.
• Forms of energy include heat, electrical, and
  light
• One way energy is transmitted through space is by
  Electromagnetic Radiation
• Transmits from one place to another in the form
  of a wave
• Given off by atoms when they have been excited by
  any form of energy

4
Electromagnetic Radiation

• Light (radiant) energy, which is visible and
  invisible
• Classified into types according to the frequency
  of the wave
• Sunlight, visible light, radio waves, microwaves
  (ovens), X-rays, and heat from a fire (infrared),
  are all forms of this radiant energy
• These forms of radiant energy exhibit the same
  wavelike characteristics

5
Wavelength and Frequency

• Electromagnetic radiation is radiant (light)
  energy that travels in waves at the speed of
  light
• The waves have three basic characteristics
  wavelength, frequency, and speed
• The highest point on the wave is a peak
• Wavelength (l distance between neighboring
  peaks)
• generally measured in nanometers (1 nm 10-9 m)
• Velocity (v how fast the wave is moving)
• c speed of light
• 3.00 x 108 m/s
• Amplitude (how tall the waves are)
• Frequency (u the number of waves that pass a
  point in a given time)
• generally measured in Hertz (Hz),
• 1 Hz 1 wave/sec 1 sec-1

6
Waves

• c u x l

frequency
wavelength
frequency
C speed of light
wavelength
7
Electromagnetic Spectrum

• Classified by wavelength
• Lower energy (longer wavelength, lower frequency)
  
• Higher energy (shorter wavelength, higher
  frequency)
• Radiowaves AM/FM/TV signals, cell phones, low
  frequency and energy
• Microwaves Microwave ovens and radar
• Infrared (IR) Heat from sunlight, infrared lamps
  for heating
• Visible The only EM radiation detected by the
  human eye
• ROYGBIV
• Ultraviolet Shorter in wavelength than visible
  violet light, sunlight
• X-rays Higher in energy than UV
• Gamma rays Highest in energy, harmful to cells

8
Wavelengths of EM Radiation
9
Atomic Spectra and Energy Levels

• When white light passes through a prism it
  produces a continuous rainbow of colors from (red
  to violet)
• From red to violet the wavelength becomes shorter

10
Atomic Spectra and Energy Levels

• When an element is heated (strontium and barium)
  light is produced
• If this light is passed through a prism, it does
  not produce a continuous rainbow, only certain
  colors

11
Atomic Spectra and Energy Levels

• Only specific colors are produced in the visible
  region. This is called a bright-line spectrum
• Each line produced is a specific color, and thus
  has a specific energy
• Each element produces a unique set of lines
  (colors) which represents energy associated with
  a specific process in the atom

12
Light Energy and Photons

• Scientists associated the lines of an atomic
  spectrum with changes in an electrons energy
  (Bohr Model)
• An electron in a higher energy state will return
  to a lower energy state
• The energy that is given off (emitted)
  corresponds to the energy difference between the
  higher and lower energy states
• The light emitted behaves like a stream of small
  particles called photons

13
Electron Energy Levels

• Electrons possess energy they are in constant
  motion in the large empty space of the atom
• The arrangement of electrons in an atom
  corresponds to an electrons energy
• The electron resides outside the nucleus in one
  of seven fixed energy levels
• Energy levels are quantized Only certain energy
  values are allowed

14
Light Energy and Photons

• The energy of a photon is related by the equation
• The energy of a photon is directly proportional
  to its frequency
• The energy of a photon is inversely proportional
  to its wavelength

E h?
15
Electron Energy Levels

• The different lines in an atomic spectrum are
  associated with changes in an electrons energy
• Each electron resides in a specific E level
  called its principal quantum number (n, where
  n1, n2)
• Electrons closer to nucleus have lower energy
  (lower n values)
• Electrons farther from the nucleus have higher
  energy (higher n values)

16
Electron Energy Levels

• Electrons can be excited to a higher E level
  with the absorption of E
• The energy absorbed is equal to the difference
  between the two E states
• When an electron loses E and falls to a lower E
  level, it emits EM radiation (photon)

17
Electron Energy Levels

• If the EM radiation wavelength is in the visible
  spectrum a color is seen

18
Energy Levels of Hydrogen The Bohr Model

• In 1913 Bohr developed a quantum model based on
  the emission spectrum for hydrogen
• The proposal was based on the electron in
  hydrogen moving around the nucleus in a circular
  orbit

19
Energy Levels of Hydrogen/The Bohr Model

• The Bohr atom

nucleus
20
Energy Levels of Hydrogen/The Bohr Model

• The Bohr atom has several orbits with a specific
  radius and specific energy
• Each orbit or energy level is identified by n
  the principal quantum number
• Electrons can be excited to a higher energy
  level with absorption of energy
• The energy absorbed and released is equal to the
  energy difference between the two states

nucleus
21
Energy Levels of Hydrogen/The Bohr Model

• The energy levels calculated by the Bohr model
  closely agreed with the values obtained from the
  hydrogen emission spectrum
• The Bohr model did not work for other atoms
• Energy levels were OK but another model was
  needed to describe the location of the electron
  about the nucleus
• Shrodinger in 1926 (DeBroglie, Heisenberg)
  developed the more precise quantum mechanical
  model
• The quantum (wave) mechanical model is the
  current theory of atomic structure

22
Quantum Mechanical Model

• The electron is treated not as a particle but as
  a wave bound to the nucleus
• The electron does not move around the nucleus in
  a circular path (orbit)
• Instead, the electron is found in orbitals. It
  is not an circular path for the electron
• An orbital indicates the probability of finding
  an electron near a particular point in space
• An orbital is a map of electron density in 3-D
  space
• Each orbital is characterized by a series of
  numbers called quantum numbers

23
Electron Energy Levels

• The energy of an electron and its distances from
  the nucleus can be grouped into levels
• Principal quantum number n is the major energy
  level in the atom
• It has values of n 1, 2, 3, etc.
• As n increases the size of the principal energy
  level (shell) increases

Principal E level electron capacity 2n2
24
Electron Sublevels

• All electrons in a principal E level (shell) do
  not have the same energy
• Each principal level is divided into 1, 2, 3, or
  4 sublevels (subshells)
• An E level contains the same number of sublevels
  (s, p, d, and f) as its own pr. energy level
  number

of sublevels in a principal E level n
25
Electron Sublevels

• The order of the increasing energy for sublevels
  (within an E level)
• The sublevels with the lowest to highest energy
• s sublevel (holds up to 2 electrons)
• p sublevel (holds up to 6 electrons)
• d sublevel (holds up to 10 electrons)
• f sublevel (holds up to 14 electrons)

s
Lowest energy
Highest energy
26
Orbitals

• The third term used to describe electron
  arrangement about the atomic nucleus (shells,
  subshells) is the orbital
• Since the electron location cannot be known
  exactly, the location of the electron is
  described in term of probability, not exact paths
• Region in space around the nucleus where there is
  a high (90) probability of finding an electron
  of a specific energy

27
Orbitals

• Orbital shapes are 3-D regions where the highest
  probability exists
• Each orbital is represented by four quantum
  numbers
• Orbitals within the same sublevel differ mainly
  in orientation
• Orbitals of the same type, but in different E
  levels (i.e. 1s, 2s, 3s) have the same general
  shape, but differ in size

28
s-Orbitals

• Only one type of orbital
• Spherical in shape
• The larger the energy level, the larger the
  sphere
• Holds two electrons

29
s-Orbitals
30
p-Orbitals

• Can only occur in n2 or higher levels
• Are dumb-bell in shape
• Three sub-orbitals (px, py and pz) each holding 2
  electrons for a total of 6 electrons in a
  p-orbital

31
p-Orbitals
32
d-Orbitals

• Five possible d-orbitals
• Odd shapes
• Only possible in n3 and larger energy levels
• Holds a total of 10 electrons

33
d-Orbitals
34
f-Orbitals

• Seven possible types of f-orbital
• Shapes very difficult, so dont have to know
• Can hold a total of 14 electrons
• Only possible for energy levels n4 and higher

35
Writing Orbital Diagrams and Electron
Configurations

• To show how the electrons are distributed in the
  E levels within an atom
• Orbital diagrams
• Electron configurations
• The most stable arrangement of electrons is one
  where the electrons are in the lowest energy
  sublevels possible

36
Writing Orbital Diagrams and Electron
Configurations

• The most stable arrangement of electrons is
  called ground-state electronic configuration
• The most stable, lowest E arrangement of the
  electrons
• The GS configuration for an element with many
  electrons is determined by a building-up process

37
Writing Orbital Diagrams and Electron
Configurations

• For the building-up process, begin by adding
  electrons to specific E levels beginning with the
  1s sublevel
• Continue in the order of increasing sublevel
  energies

1s?2s ?2p ?3s ?3p ?4s ?3d ?4p ?5s ?4d ?etc.
38
Orbital Diagram

• The notation illustrating the electron
  arrangement in terms of which energy levels and
  sublevels are occupied
• Uses the building-up principal
• Hunds Rule When electrons are placed in a set
  of orbitals of equal energy, the orbitals will be
  occupied by one electron each before pairing
  together

39
Notation

• Draw a box for each orbital
• Use an arrow up or down to represent an electron
• Only one up and one down arrow is allowed in a box

1s
2s
2p
40
Filling of Orbitals

• In General
• Begin filling from the lowest to the highest
  energy level
• If there are more than one sub-orbital possible,
  electrons will spread out first instead of
  doubling up
• Once each sub-orbital is filled with one
  electron, they will double up, but MUST have
  opposite spins (Hunds Rule)

41
Orbitals Review

• s-orbitals
• Only one per n
• Can hold two electrons for a total of 2 electrons
  in an s-orbital
• p-orbitals
• Three per n
• Can each hold two electrons for a total of 6
  electrons in a p-orbital

42
Orbitals Review

• d-orbitals
• Five per n
• Can each hold two electrons for a total of 10
  electrons in a d-orbital
• f-orbitals
• Seven per n
• Can each hold two electrons for a total of 14
  electrons in an f-orbital

43
Orbital Diagram

• hydrogen
• Only one electron
• Occupies the 1s orbital
• helium
• Two electrons
• Both occupy the 1s orbital
• lithium
• Three electrons
• Two occupy the 1s orbital, one occupies the 2s
  orbital

1s
1s
1s
2s
44
Electron Configurations and the Periodic Table

• No need to memorize the filling order of the
  electron, use the periodic table
• The atomic numbers are in order of increasing
  sublevel
• Can build-up atoms by reading across the
  periods from left to right
• By following a path of increasing atomic number
  and note the various subshells as they are
  encountered
• Each box in the table (across a period) is an
  increase in one electron

45
Electron Configurations and the Periodic Table

• The elements are arranged by increasing atomic
  number
• The periodic table is divided into sections based
  on the type of subshell (s, p, d, or f) which
  receives the last electron in the build up
  process
• Different blocks on the periodic table correspond
  to the s, p, d, or f sublevels

46
Electron Configurations and the Periodic Table

• s-block elements (Groups 1A and 2A) gain their
  last electron in an s-sublevel
• p-block elements (Groups 3A to 8A) gain their
  last electron in a p-sublevel
• d-block elements (transition metals) gain their
  last electron in a d-sublevel. First appear
  after calcium (element 20)
• d-sublevel is (n-1) less than the period number
• f-block elements are in the two bottom rows of
  the periodic table
• f-sublevel is (n-2) less than the period number

47
Subshell Filling Order
(n-1)d
np
(n-2) f
ns
48
Writing Electronic Configurations Using Sublevel
Blocks

• Locate the element, the number of electrons is
  equal to the atomic number
• Lowest energy sublevel fills first, then the next
  lowest following a path across each period
• The configuration of each element builds on the
  previous element
• The p, d, or f sublevels must completely fill
  with electrons before moving to the next higher
  sublevel

49
Electron Configuration Example 1

• Write the complete electron configuration for
  chlorine
• Chlorine is atomic number 17 (on the periodic
  table) so the neutral atom has 17 electrons
• Writing sublevel blocks in order up to chlorine
  gives

1s22s22p63s23px
50
Electron Configuration Example 1
(n-1) d
np
(n-2) f
ns
51
Electron Configuration Example 1
1s
2s
2p
3s
3p
52
Electron Configuration Example 2

• Write the complete electron configuration for
  calcium
• Calcium is atomic number 20 (on the periodic
  table) so the neutral atom has 20 electrons
• Writing sublevel blocks in order up to calcium
  gives

1s22s22p63s23p64sx
53
Electron Configuration Example 2
(n-1) d
np
(n-2) f
ns
54
Electron Configuration Example 2
1s
2s
2p
3s
3p
4s
55
Electron Configs Examples
56
Periodic Trends of the Elements

• Per. Table Graphically represents the behavior
  of the elements
• Elements are arranged by increasing atomic number
• In the periodic table, elements with similar
  properties occur at regular intervals
• The arrangement of electrons and not the mass
  that determines chemical properties of the
  elements

57
Periodic Trends of the Elements/Valence Electrons

• Valence electrons The electrons in the outermost
  energy level n (where n 1, 2, 3 )
• The most important (chemically)
• Always found in the outermost s or p sublevels
• Group number equals the valence electrons for
  each element in that group
• Applies only to the groups 1A-8A

58
Periodic Trends of the Elements/Valence Electrons

• Group IA elements have one valence electron ns1
• Group IIA elements have two valence electron ns2
• Group VIIA elements have seven valence electron
  ns2np6

59
Periodic Trends of the Elements/Valence Electrons

• Write the electron configuration for lithium
• Write the electron configuration for sodium
• Each group 1A element has a single electron in an
  s-sublevel. This is the (one) valence electron

Li 1s22s1
Na 1s22s22p63s1
60
Atomic Size

• For representative (main group) elements only
• Describes the volume of the electron cloud in the
  atoms
• Dependent upon the electron configuration of the
  atoms

61
Atomic Size

• Within groups The atomic radius increases from
  top to bottom
• Increase in the period number
• Principal E level (n) increases
• Valence electron is further from the nucleus

62
Atomic Size

• Across periods The atomic radius decreases from
  L to R with increasing atomic number
• Each element increases in proton and electron
  number
• Increase in nuclear charge
• Valence electrons pulled closer to the nucleus

63
Size of Atoms and Their Ions

• The formation of a positive ion requires the loss
  of one or more valence electrons
• Loss of the outermost (valence) causes a
  reduction in atomic size
• Positive ions are always smaller than their
  parent ions

64
Size of Atoms and Their Ions

• The formation of a - ion requires the addition of
  one or more electrons to the valence shell of an
  atom
• There is no increase in nuclear charge to
  offset the added electrons - charge
• Increase in size due to repulsion between
  electrons

65
Ionization Energy

• The minimum energy required to remove one
  electron from an atom of an element
• The more tightly an electron is held, the higher
  the ionization energy

66
Ionization Energy

• In the same group (top to bottom) Ionization
  Energy decreases
• Energy required to remove an electron decreases
• Due to larger principal energy level (larger n
  value)
• This puts outer electron farther from nucleus
• As n increases, ionization energy decreases
• Across same period (left to right) Ionization
  Energy increases
• Metals (left end) have lower ionization E
• Tend to lose electrons to form ions
• Nonmetals (right end) have higher ionization E
• Tend to gain electrons in chemical reactions

67

• End