Periodicity and Atomic Structure

1
Chapter 5

• Periodicity and Atomic Structure

2
Development of the Periodic Table

• The periodic table is the most important
  organizing principle in chemistry.
• Periodic table powerpoint elements of a group
  have similar properties
• Chapter 2 elements in a group form similar
  formulas
• Predict the properties of an element by knowing
  the properties of other elements in the group

3
Light and the Electromagnetic Spectrum

• Radiation (light) composed waves of energy
• Waves were continuous and spanned the
  electromagnetic spectrum

4
Light and the Electromagnetic Spectrum
5
Light and the Electromagnetic Spectrum
6
Light and the Electromagnetic Spectrum

• Speed of a wave is the wavelength (in meters)
  multiplied by its frequency in reciprocal
  seconds.
• Wavelength x Frequency Speed
• ? (m) x ? (s1) c (m/s1)
• C speed of light - 2.9979 x 108 m/s1

7
Electromagnetic Radiation and Atomic Spectra

• Classical Physics does not explain
• Black-body radiation
• Photoelectric effect
• Atomic Line Spectra

8
Particlelike Properties of Electromagnetic
Radiation The Plank Equation

• Blackbody radiation is the visible glow that
  solid objects emit when heated.
• Max Planck (18581947) Developed a formula to
  fit the observations. He proposed that energy is
  only emitted in discrete packets called quanta.
• The amount of energy depends on the frequency

9
Particlelike Properties of Electromagnetic
Radiation The Plank Equation

• A photons energy must exceed a minimum threshold
  for electrons to be ejected.
• Energy of a photon depends only on the frequency.

10
Electromagnetic Radiation and Atomic Spectra

• Atomic spectra Result from excited atoms
  emitting light.
• Line spectra Result from electron transitions
  between specific energy levels.

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Electromagnetic Radiation and Atomic Spectra
1/? R 1/m2 1/n2
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Quantum Mechanics and the Heisenburg Uncertainty
Principle

• Niels Bohr (18851962) Described atom as
  electrons circling around a nucleus and concluded
  that electrons have specific energy levels.
• Erwin Schrödinger (18871961) Proposed quantum
  mechanical model of atom, which focuses on
  wavelike properties of electrons.

13
Quantum Mechanics and the Heisenburg Uncertainty
Principle

• Werner Heisenberg (19011976) Showed that it is
  impossible to know (or measure) precisely both
  the position and velocity (or the momentum) at
  the same time.
• The simple act of seeing an electron would
  change its energy and therefore its position.

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Wave Functions and Quantum Mechanics

• Erwin Schrödinger (18871961) Developed a
  compromise which calculates both the energy of an
  electron and the probability of finding an
  electron at any point in the molecule.
• This is accomplished by solving the Schrödinger
  equation, resulting in the wave function, ?.

15
Wave Functions and Quantum Mechanics

• Wave functions describe the behavior of
  electrons.
• Each wave function contains three variables
  called quantum numbers
• Principal Quantum Number (n)
• Angular-Momentum Quantum Number (l)
• Magnetic Quantum Number (ml)

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Wave Functions and Quantum Mechanics

• Principal Quantum Number (n) Defines the size
  and energy level of the orbital. n 1, 2, 3,
  ???
• As n increases, the electrons get farther from
  the nucleus.
• As n increases, the electrons energy increases.
• Each value of n is generally called a shell.

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Wave Functions and Quantum Mechanics

• Angular-Momentum Quantum Number (l) Defines the
  three-dimensional shape of the orbital.
• For an orbital of principal quantum number n, the
  value of l can have an integer value from 0 to n
  1.
• This gives the subshell notation l 0 s
  orbital l 1 p orbital l 2 d
  orbital l 3 f orbital l 4 g orbital

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Wave Functions and Quantum Mechanics

• Magnetic Quantum Number (ml) Defines the spatial
  orientation of the orbital.
• For orbital of angular-momentum quantum number,
  l, the value of ml has integer values from l to
  l.
• This gives a spatial orientation ofl 0 giving
  ml 0 l 1 giving ml 1, 0, 1l 2 giving
  ml 2, 1, 0, 1, 2, and so on...

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Wave Functions and Quantum Mechanics
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Problem

• Why cant an electron have the following quantum
  numbers?
• (a) n 2, l 2, ml 1 (b) n 3, l 0, ml
  3
• (c) n 5, l 2, ml 1
• Give orbital notations for electrons with the
  following quantum numbers
• (a) n 2, l 1, ml 1 (b) n 4, l 3, ml
  2
• (c) n 3, l 2, ml 1

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The Shapes of Orbitals

• s Orbital Shapes

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The Shapes of Orbitals

• p Orbital Shapes

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The Shape of Orbitals

• d and f Orbital Shapes

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Orbital Energy Levels in Multielectron Atoms
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Orbital Energy Levels in Multielectron Atoms

• Zeff is lower than actual nuclear charge.
• Zeff increases toward nucleus ns gt np gt nd gt
  nf
• This explains certain periodic changes observed.

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Orbital Energy Levels in Multielectron Atoms

• Electron shielding leads to energy differences
  among orbitals within a shell.
• Net nuclear charge felt by an electron is called
  the effective nuclear charge (Zeff).

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Wave Functions and Quantum Mechanics

• Spin Quantum Number
• The Pauli Exclusion Principle states that no two
  electrons can have the same four quantum
  numbers.x

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Electron Configurations of Multielectron Atoms

• Pauli Exclusion Principle No two electrons in an
  atom can have the same quantum numbers (n, l, ml,
  ms).
• Hunds Rule When filling orbitals in the same
  subshell, maximize the number of parallel spins.

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Electron Configurations of Multielectron Atoms

• Rules of Aufbau Principle
• Lower n orbitals fill first.
• Each orbital holds two electrons each with
  different ms.
• Half-fill degenerate orbitals before
  pairingelectrons.

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Electron Configurations and Multielectron Atoms
Li ?? ? 1s2 2s1 1s 2s Be
?? ?? 1s2 2s2 1s 2s B
?? ?? ? 1s2 2s2 2p1
1s 2s 2px 2py 2pz C ?? ??
? ? 1s2 2s2 2p2 1s 2s
2px 2py 2pz
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Electron Configurations and Multielectron Atoms
N ?? ?? ? ? ? 1s2 2s2 2p3
1s 2s 2px 2py 2pz O ?? ?? ?? ?
? 1s2 2s2 2p4 1s 2s 2px 2py
2pz Ne ?? ?? ?? ?? ?? 1s2 2s2 2p5
1s 2s 2px 2py 2pz S Ne ?? ?? ?
? Ne 3s2 3p4 3s 3px
3py 3pz
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Problems

• Give the ground-state electron configurations
  for
• Ne (Z 10) Mn (Z 25) Zn (Z 30)
• Eu (Z 63) W (Z 74) Lw (Z 103)
• Identify elements with ground-state
  configurations
• 1s2 2s2 2p4 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2
  4d6
• 1s2 2s2 2p6 Ar 4s2 3d1 Xe 6s2 4f14 5d10 6p5

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Electron Configurations and the Periodic Table
34
Some Anomalous Electron Configurations

• Anomalous Electron Configurations Result from
  unusual stability of half-filled full-filled
  subshells.
• Chromium should be Ar 4s2 3d4, but is Ar 4s1
  3d5
• Copper should be Ar 4s2 3d9, but is Ar 4s1
  3d10
• In the second transition series this is even more
  pronounced, with Nb, Mo, Ru, Rh, Pd, and Ag
  having anomalous configurations (Figure 5.20).

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Electron Configurations and Periodic Properties
Atomic Radii
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Optional Homework

• Text 5.24, 5.26, 5.28, 5.30, 5.32, 5.34, 5.44,
  5.56, 5.58, 5.66, 5.68, 5.70, 5.72, 5.76, 5.78,
  5.82, 5.84, 5.94, 5.98, 5.108
• Chapter 5 Homework online

37
Required Homework

• Assignment 5