Chapter 7 Electron Configurations & the Periodic Table

1
Chapter 7Electron Configurations the Periodic
Table

• General Chemistry I
• S. Imbriglio

2
Part A Electron Configurations

• The arrangement of all of the electrons in an
  atom is called the electron configuration
• Electron configurations can be used to explain
• Reactivity properties of the elements
• Trends in reactivity properties (periodic
  table!)
• The electron configuration of an atom is best
  investigated using electromagnetic radiation

3
A. Electromagnetic (EM) Radiation

• Electromagnetic (EM) Waves oscillating
  perpendicular magnetic electric fields that
  travel through space at the same rate (the speed
  of light c 3.00x108 m/s)
• - Unlike sound waves, electromagnetic waves
  require no medium for propagation
• eg. This allows the suns electromagnetic
  radiation to reach the earth as sunlight.

http//ww2.unime.it/dipart/i_fismed/wbt/mirror/ntn
ujava/emWave/emWave.html
4
a) Wavelength Frequency

• All EM waves can be described in terms of
  wavelength frequency.
• Wavelength
• (? - lambda)
• distance between
• adjacent crests
• (or troughs) in a
• wave

5
a) Wavelength Frequency

• - Frequency (? - nu) the number of complete
  waves passing a point in a given period of time
  (remember c 3.00x108 m/s)
• Unit of frequency is the Hertz (Hz)
• 1 Hz 1 s-1
• 1 per second

6
a) Wavelength Frequency

• For EM radiation, frequency is related to
  wavelength by ?? c.
• If you know one, you know the other.
• Calculate the frequency of an X-ray that has a
  wavelength of 8.21 nm.
• ?? c

7
i. Electromagnetic Spectrum

• The type of electromagnetic radiation is defined
  by its frequency wavelength
• Remember, as ? increases, ? decreases, and vice
  versa.

8
b) Amplitude

• The intensity of radiation is related to its
  amplitude.
• Amplitude height
• of the wave crest
• In the visible portion
• of the spectrum,
• brighter light is light
• with a greater
• amplitude.

9
c) Refraction

• Classifying EM radiation (light) as a wave
  explains many fundamental properties of light.
• Refraction
• When white light passes through a narrow slit
  then through a glass prism, the light separates
  in a continuous spectrum.
• The spectrum is continuous because each color
  merges into the next without a break all
  wavelengths (or frequencies) of visible light are
  observed.

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d) Diffraction

• Diffraction Waves can add constructively or
  destructively to amplify or cancel each other.

11
e) Black-Body Radiation

• At high temperatures, matter emits
  electromagnetic radiation.
• As the temperature increases, the maximum
  intensity of the emitted radiation increases in
  frequency.
• The observed
• spectrum depends
• only on temperature
• not on the particular
• elements present.

12
b) Black-Body Radiation

• THE EXPLANATION According to classical physics,
  as the temperature of a solid increases, the
  atoms vibrate more vigorously some of the
  vibrational energy is released as EM radiation.
• THE PROBLEM Using
• the classical picture of
• light as a wave,
• scientists were unable
• to explain the shape of
• the observed black-body
• radiation spectra.

13
3. Plancks Quantum Hypothesis THE ANSWER?

• According to classical physics, the energy scale
  is continuous there are no limitations on the
  amount of energy a system can gain or lose.
• Planck proposed that variations in energy are
  discontinuous energy changes occur only by
  discrete amounts.

eg. The Quantization of Elevation
Classical (continuous)
Quantized (1 step 1 quantum)
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3. Plancks Quantum Theory

• For electromagnetic radiation of a certain
  frequency, the smallest amount of energy, called
  a quantum, is defined by the relationship
• E h?
• (h Plancks constant 6.626x10-34 Js)
• Energy can be absorbed or emitted only as a
  quantum, or some whole-number multiple of a
  quantum.

15
3. Plancks Quantum Theory

• According to Plancks theory, the energy of one
  quantum of EM radiation is dependent on the
  frequency (and wavelength) of the radiation
• E h? hc/?
• The energy per quantum increases as the frequency
  gets higher the wavelength gets shorter.

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3. Plancks Quantum Theory

• The energy per quantum increases as the frequency
  gets higher the wavelength gets shorter.
• E h? hc/?
• Which has more energy a quantum of microwave
  radiation (? 1x10-2 m) or a quantum of infrared
  radiation (? 1x10-6 m)?

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3. Plancks Quantum Theory

• Planck proposed that vibrating atoms in a heated
  solid can absorb and emit EM radiation only in
  certain discrete amounts.
• Plancks quantum theory allowed him to
  successfully explain black-body radiation
  spectra, but his radical assertion that energy
  is quantized was difficult for the scientific
  community to accept.
• Fortunately, five years after its inception,
  Einstein used Plancks Quantum Theory to explain
  another well-known phenomenon called the
  photoelectric effect.

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4. The Photoelectric Effect

• Certain metals exhibit a photoelectric effect
  when illuminated by light of certain wavelengths
  (photo-), they emit electrons (-electric).
• In order for the photoelectric effect to occur,
  the frequency of the light must be higher than a
  certain minimum value called the threshold
  frequency.
• Each photosensitive metal has a different
  threshold frequency.

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4. The Photoelectric Effect

• When light of a high enough energy (frequency) is
  used, the number of electrons ejected is
  proportional to the intensity of the light.
• Light below the threshold frequency will not
  cause an electric current to flow no matter how
  bright (intense) the light is.
• eg. Light meters use the
• photoelectric effect to
• measure the intensity
• (brightness) of light.

http//jchemed.chem.wisc.edu/JCEDLib/WebWare/colle
ction/open/JCEWWOR006/peeffect5.html
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a) Photons

• Classical physics could not explain the existence
  of a threshold frequency, so Einstein turned to
  Plancks Quantum Theory.
• Einstein defined a quantum of electromagnetic
  radiation as a photon.
• Einstein proposed that light could be thought of
  as a stream of photons with particle-like
  properties as well as wave properties.
• For light of frequency ?
• Ephoton h? hc/l

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a) Photons

• The photoelectric effect can be explained by
  assuming that light has particle-like properties
• Removing one electron from a photosensitive metal
  requires a certain minimum energy (Emin).
• Each photon has an energy given by E h?.
• Only photons with E gt Emin have enough energy to
  knock an electron loose.
• Photons of lower frequency (lower energy) do not
  have enough energy to knock an electron loose.

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a) Photons

• If the intensity of light is proportional to the
  number of photons, then more intense light means
  more photons.
• If each photon ejects an electron, then more
  photons means more electrons ejected.
• The number of electrons ejected is proportional
  to the intensity of light.

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The Photoelectric EffectExplained

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b) Wave-Particle Duality of Light

• Depending on the circumstances, light (all EM
  radiation) can appear to have either wave-like or
  particle-like characteristics.
• Both ideas are needed to fully explain lights
  behavior in different phenomena.
• It not only prohibits the killing of two birds
  with one stone, but also the killing of one bird
  with two stones.
• - James Jeans, commenting on Einsteins
  explanation of the photoelectric effect

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Nobel Prize Winners

• Max Planck won the Nobel Prize for Physics in
  1918 for his quantum theory.
• Blackbody radiation spectra explained
• Albert Einstein won the Nobel Prize for Physics
  in 1921 for his theory on the quantized nature of
  light and how it relates to lights interaction
  with matter (not for his theory of relativity!).
• Photoelectric effect explained

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5. Line Emission Spectra

• In the 1920s, another phenomenon was left
  unexplained by classical physics the observance
  of atomic line emission spectra.
• When a voltage is applied to a gaseous element at
  low pressure, the atoms absorb energy become
  excited.
• The excited atoms then emit the extra energy as
  EM radiation.

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5. Line Emission Spectra

• When this radiation is passed through a prism, a
  limited number of discrete colored lines are seen
  a discontinuous spectrum.
• This discontinuous spectrum is called a line
  spectrum, or a line emission spectrum.
• Unlike black-body radiation, each element has a
  unique line emission spectrum.

Why dont these atoms emit continuous spectra?
28
B. Bohrs Hydrogen Atom A Planetary Model

• Classical physics could not explain the presence
  of line emission spectra.
• Not long after Einstein used quantum theory to
  explain the photoelectric effect, Niels Bohr used
  quantum theory to explain the behavior of the
  electron in a hydrogen atom.
• Bohrs model provided the first explanation of
  the discontinuous line emission spectrum of
  hydrogen.

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B. Bohrs Hydrogen Atom A Planetary Model

• Bohr assumed that the single electron in a
  hydrogen atom moves around the nucleus in a
  circular orbit.
• Bohr applied quantum theory to his model by
  proposing that the electron is restricted to
  circling the nucleus in orbits of certain radii,
  each of which corresponds to a specific energy.
• Thus, the energy of the electron is quantized,
  and the electron is restricted to certain energy
  levels orbits.

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B. Bohrs Hydrogen Atom A Planetary Model
31
B. Bohrs Hydrogen Atom

• 1. Energy Levels (Orbits)
• Each allowed orbit is assigned a principal
  quantum number (n 1,2,3,?).
• The energy of the electron and the radius of its
  orbit increase as the value of n increases.
• An atom with its electron in the lowest energy
  level is said to be in the ground state.

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1. Energy Levels (Orbits)
En _ 2.179x10-18 J (n 1, 2, 3, )
n2

• The allowed energies of an electron (orbit) in a
  hydrogen atom are restricted by the principal
  quantum number (n), according to the equation
  above.
• The negative sign is a result of Bohrs choice to
  define En 0 when n ?.

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a) Excited State vs. Ground State

• Transitions Between Levels Electrons can move
  from one energy level to another
• An electron must absorb energy to transition from
  a lower energy level to a higher energy level
• Energy is emitted when an electron transitions
  from a higher energy level to a lower energy
  level
• When an electron absorbs energy and moves to a
  higher energy level, that atom is said to be in
  an excited state.

http//www.upscale.utoronto.ca/GeneralInterest/Har
rison/BohrModel/Flash/BohrModel.html
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a) Excited State vs. Ground State

• Absorb energy to move to a higher energy orbit.
• Emit energy to move to a lower energy orbit.

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a) Excited State vs. Ground State

• When an excited
• electron returns to the
• ground state, energy
• is emitted as a photon
• with an energy
• corresponding to the
• difference in energy between the two levels.
• In the Bohr model, n? is the excited state in
  which enough energy has been added to completely
  separate the electron from the proton Bohr
  arbitrarily assigned this state as having E 0
  (hence the negative energy values).

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2. Explanation of Line Spectra

• Bohrs model of the hydrogen atom can be used to
  explain the line emission spectrum of hydrogen
• ?E Efinal - Einitial
• ?E h?

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2. Explanation of Line Spectra

• Using Bohrs equation for allowed energies in a
  hydrogen atom
• ?E Ef - Ei
• ?E _ 2.179x10-18 J _ _ 2.179x10-18 J
  
• nf2 ni2
• ?E 2.179x10-18 J x 1 _ 1
• ni2 nf2
• Only certain energies of light (?E ) can be
  absorbed or emitted by electrons in a hydrogen
  atom.

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2. Explanation of Line Spectra

• Now, coupling that equation with E h? allows us
  to describe the frequencies of light that can be
  absorbed or emitted by an electron in a hydrogen
  atom.
• E h? where E ?E
• The frequencies (?) determined by this equation
  correlate with the frequencies of light observed
  in the line emission spectrum of hydrogen.

39

• The discrete lines in the line emission spectra
  correspond to photons of specific frequencies
  that are emitted when electrons relax from higher
  energy levels to lower energy levels

40
Using Bohrs model, calculate the frequency of
the radiation released by the transition of an
electron in a hydrogen atom from the n 5 level
to the n 3 level.
41
Using Bohrs model, calculate the wavelength of
the radiation absorbed by a hydrogen atom when
the electron undergoes a transition from the n
4 to n 5 level.

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C. Quantum Mechanical Model of the Atom

• By the early 1920s, the theory of the
  Wave-Particle Duality of light had been accepted,
  but a young scientist named Louis De Broglie was
  ready to shock the scientific community with
  another hypothesis.
• De Broglie proposed that matter can exhibit
  wave-like properties.
• eg. Electrons exhibit diffraction
• similar to that observed with light.

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1. De Broglie Matter as Waves

• De Broglie proposed that a particle of mass m
  moving at speed v will have a wave nature
  consistent with a wavelength given by the
  equation
• h/mv
• Large (macroscale) objects have wavelengths too
  short to observe.
• Small (nanoscale) objects have longer more
  readily observable wavelengths.

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a) Quantum Mechanics

• Current ideas about atomic structure are based on
  De Broglies theory.
• The treatment of atomic structure using the
  wave-like properties of the electron is called
  quantum mechanics (or wave mechanics)
• In contrast to Bohrs precise atomic orbits,
  quantum mechanics provides a less certain
  picture of the hydrogen atom.

45
b) Wave Equation Wave Functions

• In 1926, Erwin Schrödinger used De Broglies
  theory to develop an equation (Schrödingers wave
  equation) describing the locations energies of
  the electron in a hydrogen atom.
• Acceptable solutions to Schrödingers wave
  equation are called wave functions (?).
• Unlike Bohrs model, these wave functions do not
  describe the exact location of an electron.

46
b) Wave Equation Wave Functions

• The square of a wave function (?2) gives the
  probability of finding an electron in a
  particular infinitesimally small volume of space
  in an atom.
• Because we are treating electrons as waves (not
  particles) we cannot pinpoint the specific
  location of an electron.
• Instead, mathematical solutions to the wave
  functions give 3-dimensional shapes (orbitals)
  within which electrons can usually be found.

47
b) Wave Equation Wave Functions

• These 3-D orbitals (probability clouds) take the
  place of Bohrs simple well-defined orbits in the
  modern model of the atom. We dont know exactly
  where the electrons are.
• This less certain model is justified by an
  important principle of science established in
  1927.

48
2. Heisenbergs Uncertainty Principle

• It is impossible to determine the exact location
  and the exact momentum of a tiny particle like an
  electron.
• The very act of measurement would affect the
  position and momentum of the electron because of
  its very small size and mass.
• The collision of an electron with a high-energy
  photon (required to locate the electron) would
  change the momentum of the electron.
• The collision of an electron with a low-energy
  photon would not provide much information about
  the location of the electron.

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2. Heisenbergs Uncertainty Principle

• A macroscale analogy
• High Shutter Speed Low Shutter Speed
• Can judge location, Can judge speed,
• but not speed. But not location

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D. Quantum Numbers Atomic Orbitals

• According to quantum mechanics, each electron in
  an atom can be described using four quantum
  numbers
• n Principal Quantum Number
• l Angular Momentum Quantum Number
• ml Magnetic Quantum Number
• ms Electron Spin Quantum Number
• The first three numbers describe the atomic
  orbital in which the electron resides the
  fourth differentiates electrons that are in the
  same atomic orbital.

51
1. Principal Quantum Number (n)

• The principal quantum number (n) has only integer
  values, starting with 1
• n 1, 2, 3, 4, . . .
• a) The value of n corresponds to the Principal
  Electron Shell that the orbital is in.
• b) The principal electron shell is the major
  factor in determining the energy of the
  electron(s) in that orbital a higher n value
  means a higher energy.

52
2. Angular Momentum Quantum Number (l)

• The angular momentum quantum number (l ) is an
  integer that ranges from zero to a maximum of n
  1
• l 0, 1, 2, 3, . . . (n 1)
• a) The value of l indicates the subshell that the
  orbital is in (within the larger energy shell).
• n 1 l 0 (1 subshell)
• n 2 l 0 or 1 (2 subshells)
• n 3 l 0, 1 or 2 (3 subshells)
• n 4 l 0, 1, 2 or 3 (4 subshells)

53
2. Angular Momentum Quantum Number (l)

• Each subshell (l) is designated with a letter
• b) Each letter (s, p, d, f) symbolizes a subshell
  containing one specific type of orbital with a
  unique shape.
• eg. All s orbitals are spherical (l 0) all p
  orbitals are shaped like dumbbells (l 1) more
  on this in a minute.

s orbital
p orbital
54
2. Angular Momentum Quantum Number
In the third principle shell, there is one s
subshell containing one s orbital, one p subshell
containing three p orbitals one d subshell
containing five d orbitals..
In the second principle shell, there is one s
subshell containing one s orbital one p
subshell containing three p orbitals.
In the first principle shell, there is one s
subshell containing one s orbital.
Within a p or d subshell, how do you distinguish
between the individual orbitals?
55
3. Magnetic Quantum Number (ml )

• The magnetic quantum number (ml) can have any
  integer value between l and - l, including zero
• ml l, . . . , 1, 0, -1, . . . , - l
• a) The magnetic quantum number (ml) is related to
  the directional orientation of the orbital.
• eg. There are three possible p orbitals each
  pointing along a different axis in space.

56
3. Magnetic Quantum Number (ml )
(s) l o ml 0 (1 s orbital) (p) l 1
ml -1,0,1 (3 p orbitals) (d) l 2 ml
-2,-1,0,1,2 (5 d orbitals) (f) l 3 ml
-3,-2,-1,0,1,2,3 (7 f orbitals) eg. There is
only one type of directional orientation for
any given s orbital in an l 0 subshell
because ml must equal 0.
57
3. Magnetic Quantum Number (ml )

• There are three different p orbitals in every l
  1 subshell because ml -1,0,1.
• Each of the three
• p orbitals is
• pointed along a
• different axis
• (x,y,z).

58
3. Magnetic Quantum Number (ml )

• There are five different d orbitals in every l
  2 subshell because ml -2,-1,0,1,2.
• Four of the five
• d orbitals are
• pointed along a
• different axis.
• The fifth has a
• slightly different
• shape.

59
4. Shells (n), Subshells (l ) Orbitals (ml ) A
Summary

60
4. Shells (n), Subshells (l ) Orbitals (ml ) A
Summary

This picture shows all of the orbitals in the
first three electron shells (n 1,2,3).
61
State whether an electron can be described by
each of the following sets of quantum number. If
a set is not possible, state why not.

• n 2, l 1, ml -1
• n 1, l 1, ml 1
• n 4, l 3, ml 3
• n 3, l 1, ml -3

62
Replace the question marks by suitable responses
in the following quantum number assignments.

• n 3, l 1, ml ?
• n 4, l ?, ml -2
• n ?, l 3, ml ?

63
Provide the three quantum numbers describing each
of the three p orbitals in the 2p subshell.
n l ml 2px 2py 2pz
64
5. Electron Spin Quantum Number (ms)

• The first three quantum numbers (n, l, ml) fully
  characterize all of the orbitals in an atom.
• But, one more quantum number is necessary to
  describe all of the electrons in an atom.
• This is because every orbital can hold two
  electrons.

65
5. Electron Spin Quantum Number (ms)

• The spin quantum number (ms) can have just one of
  two values (1/2 -1/2).
• Each electron exists in one of two possible spin
  states.
• - The spinning electron induces
  an external magnetic field. Opposite
  spins induce opposing magnetic fields.

66
5. Electron Spin Quantum Number (ms)

• When two electrons have the same ms quantum
  number, those spins are said to be parallel.
• When two electrons in the same orbital have
  different ms quantum numbers, those electrons are
  said to be paired.

?? Paired spins
? ? Parallel spins
67
a) Pauli Exclusion Principle

• The Pauli Exclusion Principle states that no more
  than two electrons can be assigned to the same
  orbital in an atom those two electrons must
  have opposite spins.
• In other words
• No two electrons in the same atom can have the
  same set of four quantum numbers (n, l, ml, ms).
• If two electrons occupy the same orbital, their
  spins must be paired (1/2 -1/2).

68
Quantum Numbers A Macroscale Analogy

• n - indicates which train (shell)
• l - indicates which car (subshell)
• ml - indicates which row (orbital)
• ms - indicates which seat (spin)
• No two people can have exactly the same ticket
  (sit in the same seat).

69
For n 1, determine the possible values of l.
For each value of l, assign the appropriate
letter designation determine the possible
values of ml.
n 1 How many orbitals in shell n
1? How many electrons possible?
70
For n 2, determine the possible values of l.
For each value of l, assign the appropriate
letter designation determine the possible
values of ml.
n 2 How many
orbitals in shell n 2? How many electrons
possible?
71
For n 3, determine the possible values of l.
For each value of l, assign the appropriate
letter designation determine the possible
values of ml.
n 3 of
Orbitals? of Electrons?
72
For n 4, determine the possible values of l.
For each value of l, assign the appropriate
letter designation determine the possible
values of ml.
73
Provide the four quantum numbers describing each
of the two electrons in the 3s orbital.
n l ml ms
74
E. Electron Configurations

• The electron configuration of an atom is the
  complete description of the orbitals occupied by
  all of its electrons
• eg. The electron in a ground state hydrogen atom
  occupies the 1s orbital
• There are several ways to represent electron
  configurations. . .

75
1. Representations of Electron Configuration

• In most cases, it is sufficient to write a list
  of all of the occupied subshells and indicate the
  number of electrons in each subshell with a
  superscript.
• H 1s1
• C 1s2 2s2 2p2
• Ar 1s2 2s2 2p6 3s2 3p6

76
1. Representations of Electron Configuration

• a) Expanded Electron Configuration In some
  cases, it is more informative to write a list of
  each occupied orbital and indicate the number of
  electrons in each orbital.
• N 1s2 2s2 2p3 versus N 1s2 2s2 2p1 2p1 2p1
• The expanded configuration indicates that there
  is one electron in each of the three 2p orbitals
  the original configuration doesnt.

77
1. Representations of Electron Configuration

• b) An orbital box diagram goes one step further
  by also illustrating the spins of the elctrons.
• P 1s2 2s2 2p2 2p2 2p2 3s2 3p1 3p1 3p1
• P
• The orbital box diagram indicates that the three
  electrons in the 3p subshell all have parallel
  (unpaired) spins.

1s
2s
2p
3s
3p
??
??
??
??
??
??
?
?
?
78
i) Hunds Rule

• In the last example, we saw that
• Atoms can have half-filled orbitals
• the electrons in the half-filled orbitals tend to
  have parallel spins
• Hunds Rule The most stable arrangement of
  electrons in the same subshell has the maximum
  number of unpaired electrons, all with the same
  spin
• In other words, electrons pair only after each
  orbital in a subshell is occupied.

79
Write the expanded electron configuration and the
box orbital diagram for oxygen (1s2 2s2 2p4).
O 1s2 2s2 2p4 O O
80
Write the expanded electron configuration and the
box orbital diagram for boron (1s2 2s2 2p1).
B 1s2 2s2 2p1 B B
81
1. Representations of Electron Configuration

• c) When you get deeper into the periodic table,
  electron configurations can be abbreviated by
  using noble gas notation.
• The noble gases are the elements in group 8A (He,
  Ne, Ar, Kr, Xe, Rn)
• Each noble gas has a filled outer subshell
  (enough electrons to fill its highest energy
  subshell)

82
c) Noble Gas Notation

• Electron Configurations of Noble Gases
• He 1s2
• Ne 1s2 2s2 2p6
• Ar 1s2 2s2 2p6 3s2 3p6
• Kr 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
• To use noble gas notation, write the symbol for
  the preceding noble gas in brackets to
  represent all of the electrons in its electron
  configuration.
• Add the rest of the electrons at the end.

83
c) Noble Gas Notation

• Write the following electron configurations using
  noble gas notation
• O 1s2 2s2 2p4
• Si 1s2 2s2 2p6 3s2 3p2
• Now we know how to write electron configurations.
  How do we know what the ground state electron
  configuration for an element is???

84
2. Ground State Configuration

• Afbau Principle Every atom has an infinite
  number of possible electron configurations.
• For an atom in its ground state, electrons are
  found in the energy shells, subshells orbitals
  that produce the lowest energy for the atom.
• Other configurations correspond to excited states.

85
2. Ground State Configuration

• In other words, when deciding where to put the
  electrons in the ground state, always start
  filling the lowest energy orbitals first.
• In general
• Orbital energy increases as n increases
• Within the same shell (n), orbital energy
  increases as l increases (E sltpltdltf)

86
a) Order of Subshell Filling

• The electron configurations of the first ten
  elements illustrate this point.

87
a) Order of Subshell Filling

• In general, subshells are filled in order of
  increasing n l value
• If two orbitals have the same value for n l,
  fill the subshell with lowest n value first

88
a) Order of Subshell Fillingi) Using the
Periodic Table

• You dont have to memorize the order of the
  subshells, just use the periodic table!
• Start at H move through the table in order
  until the desired element is reached.

Notice (n 1)d orbitals are filled after ns and
before np orbitals.
89
a) Order of Subshell Fillingi) Using the
Periodic Table

• Write the electron configuration for Al.
• Al
• Al

Ne
Al
90
a) Order of Subshell Fillingi) Using the
Periodic Table

• Write the electron configuration for As.
• As
• As

Ar
As
91
a) Order of Subshell Fillingi) Using the
Periodic Table
Write the electron configuration for Sn. Sn Sn
Kr
Sn
92
ii) Transition Metals

• Remember, (n 1)d orbitals are generally filled
  after ns orbitals and before np orbitals.
• There are some exceptions
• When it is possible to half-fill or fill the
  (n-1)d shell, the ns subshell can be left
  half-filled
• This is an example of Hunds Rule. The ns and
  (n-1)d orbitals are very close in energy, so the
  more parallel spins, the better.

93
ii) Transition Metals
4s
3d


?


??

• Sc Ar3d14s2
• - 4s filled before 3d
• Ti Ar3d24s2
• V Ar3d34s2

4s
3d


?

??
?
4s
3d
?

??
?
?

94

ii) Transition Metals

Cr might expect Ar3d44s2 Physical
properties indicate that this is not the electron
configuration. It is actually Cr
Ar3d54s1 Notice the 3d subshell is
half-filled. This configuration maximizes
unpaired electrons - Hunds Rule.
4s
3d

?
?
??
?
?
4s
3d
?
?
?
?
?
?
95

ii) Transition Metals

Having a filled subshell is also energetically
favorable, so copper has an unexpected
configuration Cu Ar3d104s1 The energetic
stability gained from having either a filled or a
half-filled subshell has an effect on the
reactivity of different elements.
4s
3d
?
??
??
??
??
??
96
iii) Magnetic Properties

• The electron configuration of an atom determines
  its magnetic properties.
• In atoms (or ions) with completely filled shells,
  all of the electron spins are paired, so their
  individual magnetic fields effectively cancel
  each other out.
• Such substances are called diamagnetic.

97
iii) Magnetic Properties

• Atoms (or ions) with unpaired electrons (parallel
  spins) are attracted to a magnetic field.
• More unpaired electrons, stronger attraction.
• Such substances are called paramagnetic.
• eg. Metallic nickel is paramagnetic
• Ni

4s
3d
?
??
??
?
??
??
98
iii) Magnetic Properties

• Ferromagnetic substances are permanent magnets.
• Spins of electrons in a cluster of atoms are
  aligned in same direction, regardless of external
  magnetic field
• Metals in the iron, cobalt nickel groups
  exhibit ferromagnetism

99
iv) Valence Electrons

• The atomic electron configuration of an element
  determines the chemical reactivity of that
  element, but it is not the total number of
  electrons that is important.
• If that were the case, each element would have
  unique reactivity we would not observe
  periodicity in atomic trends and reactivity.
• How do we explain the trends in the periodic
  table?
• Valence Electrons!

100
iv) Valence Electrons

• When considering the principal electron shells (n
  1,2,3,), there are two types of electrons
• Core Electrons electrons in the filled inner
  shell(s) of an atom
• Valence Electrons electrons in the unfilled
  outer shell of an atom
• All elements in the same group have similar
  chemical properties because they have the same
  number of valence electrons in their outer shell!

101
iv) Valence Electrons

• For elements in the first three periods
• The core electrons are those in the preceding
  noble gas configuration.
• The additional electrons in the outer shell are
  the valence electrons.
• eg. B 1s2 2s2 2p1
• B He2s2 2p1
• Core 1s2 Valence 2s2 2p1
• (Shell with n 1) (Shell with n 2)

102
iv) Valence Electrons
Cl Cl Core Valence
103
iv) Valence Electrons

• For elements in the fourth period and below in
  groups 3A 7A, the filled d subshells are also
  part of the core, even though they are not
  included in the noble gas configuration.
• Se
• Se
• Core Valence

104
iv) Valence Electrons

• In each A group, the number of valence electrons
  is equal to the group number.

1
8
3 4 5 6 7
2
of valence electrons
p-block
s-block
105
v) Lewis Dot Symbols

• The number of valence electrons in an atom is
  directly related to its reactivity.
• Gilbert Lewis came up with a way to represent an
  element its valence electrons.
• One dot equals one valence electron.

106
3. Ion Electron Configurations

• The number of valence electrons determines the
  type of cation () or anion (-) an atom will
  form.
• When s- and p-block elements form ions, electrons
  are removed or added such that a noble gas
  configuration is achieved.
• The ions are said to be isoelectronic with the
  noble gas more on this later.
• In general
• Metals lose electrons form cations
• Non-metals gain electron form anions

107
a) Cations

• Li 1s22s1 (loses an electron)
• Li 1s2 He
• - Group 1 metals form cations with 1 charge.
• Mg 1s22s22p63s2 (loses two electrons)
• Mg2 1s22s22p6 Ne
• - Group 2 metals form cations with 2 charges.

108
b) Anions

O 1s22s22p4 (gains two electrons) O2- 1s22s22p6
Ne - Group 6 elements form anions with -2
charge. F 1s22s22p5 (gains one electron) F-
1s22s22p6 Ne - Group 7 elements (halogens)
form anions with -1 charge.
109
Periodic Trends

• When Mendelev created the first periodic table,
  he organized the elements based on similarities
  in chemical properties reactivity.
• Now, we can use the electron configurations of
  the elements to explain the trends in the
  periodic table.
• Atomic Ionic Radii
• Ionization Energy
• Electron Affinity

110
F. Atomic Radii

• The atomic radius of an atom is defined as
  one-half of the internuclear distance between two
  of the same atoms in a simple diatomic molecule.
• - In this simplified picture, we assume
  that each atom is spherical the radius
  is the distance from the center to the edge
  of the sphere.

111
1. Trends Atomic Radii

• The size (radius) of an atom is determined by two
  main factors
• a) Principal Quantum Number (n) the larger the
  principal quantum number (n), the larger the
  orbitals
• - As you move down a group in the periodic table,
  the atomic radii of the atoms increase because n
  increases.

112
1. Trends Atomic Radii

• b) Effective Nuclear Charge (Z) the nuclear
  positive charge experienced by outer-shell
  electrons in a many-electron atom
• Outer-shell electrons are shielded from the full
  nuclear positive charge (Z) by the inner-shell
  electrons (electron-electron repulsion)
• The effective nuclear charge (Z) felt by an
  outer-shell electron is less than the actual
  charge of the nucleus (Z).

113
1. Trends Atomic Radii

• b) Effective Nuclear Charge
• As Z increases, the outer electrons are pulled
  closer to the nucleus the atomic radius
  decreases.
• Z increases across a period in the periodic
  table (additional attraction to nucleus stronger
  than electron-electron repulsion/shielding).
• Atomic Radius decreases across a period in the
  periodic table.

114
1. Trends Atomic Radii

• Main Group Elements

Increasing Atomic Radius
Decreasing Atomic Radius
115
G. Ionic Radii

• Periodic trends in ionic radii parallel the
  trends in atomic radii within the same group.

Ionic radii increase as you move down a group in
the periodic table (n increases).
116
G. Ionic Radii

• 1. Cations The radius of a cation is always
  smaller than that of the atom from which it is
  derived.
• Nuclear charge (Z) remains the same.
• Electron-electron repulsion (shielding)
  decreases.
• Effective nuclear charge (Z) increases.

117
G. Ionic Radii

• 2. Anions The radius of an anion is always
  larger than that of the atom from which it is
  derived.
• Nuclear charge (Z) remains the same.
• Electron-electron repulsion (shielding)
  increases.
• Effective nuclear charge (Z) decreases.

118
a) Isoelectronic Ions

• Atoms or ions with identical electron
  configurations are said to be isoelectronic.
• In general
• Anions in a given period are isoelectronic with
  the noble gas in the same period.
• Cations in a given period are isoelectronic with
  the noble gas in the preceding period.

119
a) Isoelectronic Ions

120
a) Isoelectronic Ions

• When comparing isoelectronic ions, the radii
  depend on the number of protons in the nucleus
  (Z)
• The more protons in the nucleus, the smaller the
  radius.

121
a) Isoelectronic Ions
122
Rank the following series in order of increasing
atomic or ionic radii. (1 smallest, 3 largest)
123
H. Ionization Energies

• The first ionization energy of an atom is the
  minimum energy required to remove the highest
  energy (outermost) electron from the neutral atom
  in the gas phase.
• The larger the I.E., the harder it is to remove
  the electron.
• eg. The first ionization energy of lithium is
  illustrated by the following equation
• Li (g) ? Li (g) e- ?E 520 kJ/mol
• 1s22s1 1s2

124
1. Trends Ionization Energies

• Ionization energies tend to decrease as you move
  down a group in the periodic table.
• Size - it is easier to remove an electron that is
  further from the nucleus.
• Ionization energies tend to increase as you move
  across a period in the periodic table.
• Effective Nuclear Charge As Z increases, it
  becomes harder to remove an electron.

125
1. Trends Ionization Energies

Generally Increasing
Generally Decreasing
126
1. Trends Ionization Energies

• Ionization energies do not increase smoothly
  across the periods in the periodic table.
• In general - It is easier to remove
  an electron if it results in the
  formation of a filled or half-filled
  subshell.
• - It is harder to remove an
  electron from a filled or
  half-filled subshell.

127
1. Trends Ionization Energies
eg. In contrast to periodic trends, the
ionization energy of nitrogen is higher than the
ionization energy of oxygen. N N O O
128
1. Trends Ionization Energies
In contrast to periodic trends, the ionization
energy of beryllium is higher than that of boron.
Why? Be Be B B
129
2. Subsequent Ionization Energies

• The first, second third ionization energies are
  the energies associated with removing the first,
  second third highest energy electrons in an
  atom.
• Ionization energies increase with each successive
  electron removed because Z increases (same
  number of protons, less electron repulsion).

130
2. Subsequent Ionization Energies

• Ionization energies can be used to explain why Li
  forms Li cations and Be forms Be2 cations.

131
I. Electron Affinities

• The electron affinity of an element is the energy
  change resulting from an electron being added to
  an atom to form a 1- anion.
• Electron affinity is the measure of the
  attraction an atom has for an additional
  electron.
• eg. The electron affinity of fluorine is
  illustrated by
• F (g) e- ? F- (g) ?E EA -328
  kJ/mol
• 1s22s22p5 1s22s22p5
• Fluorine readily accepts an electron to gain a
  stable noble gas configuration (filled n 2
  shell).

132
I. Electron Affinities

• Most electron affinities are lt0 (favorable).
• Electron affinities are generally ? 0
  (unfavorable) for atoms with filled subshells
  (Groups 2A 8A).

133
Knowing the Trends

• You should be able to explain the trends use
  the periodic table to predict relative magnitudes
  for the following properties
• Atomic Radius
• Ionic Radius
• Ionization Energy
• The trends in electron affinities are less
  regular, but you should be able to explain
  differences in EAs based on filled/unfilled
  subshells.