Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

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Unit 6 Chapters 11-12. Pages 295-366 ATOMIC
ELECTRON CONFIGURATIONS AND PERIODICITY
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Bohr Model

• First model of the electron behavior
• Vital to understanding the atom
• Does not work for atoms with
• more than 1 electron

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Collision of Ideas
Matter
Dalton
Thompson
Rutherford
Bohr
?
De Broglie
Einstein
Plank
Maxwell
Newton
Light
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The Photoelectric Effect

• Duality of Light
• Wave behavior
• Particle behavior

1905
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de Broglies Novel Notion
1923

• Light was known (thought) to be a wave, but
• Einstein showed that it also acts particle-like
• Electrons were particles with known mass charge
• What if

electrons behaved as waves also
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Evidence for de Broglies Notion

• Diffraction pattern obtained with firing a beam
  of electrons through a crystal.
• This can only be explained if the electron
  behaves as a wave!
• Nobel Prize for de Broglie in 1929

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Electron Characteristics

• Extremely small mass
• Located outside the nucleus
• Moving at very high speeds
• Have specific energy levels
• Standing wave behavior

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Baseball vs Electron
A baseball behaves as a particle and follows a
predictable path. BUT An electron behaves as a
wave, and its path cannot be predicted. All we
can do is to calculate the probability of the
electron following a specific path.
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What if a baseball behaved like an electron?

• Characteristic wavelength
• baseball ? 10-34 m
• electron ? 0.1 nm

All we can predict is..
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Werner Heisenberg(1901-1976)
The Uncertainty Principle
speed
position

• Proposed that the dual nature of the electron
  places limitation on how precisely we can know
  both the exact location and speed of the electron
• Instead, we can only describe electron behavior
  in terms of probability.

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Erwin Schrodinger (1887-1961)
Wave Equation Wave Mechanics

• In 1926, Austrian physicist, proposed an equation
  that incorporates both the wave and particle
  behavior of the electron
• When applied to hydrogens 1 electron atom,
  solutions provide the most probable location of
  finding the electron in the first energy level
• Can be applied to more complex atoms too!

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Solutions to Schrodingers Wave Equation

• Gives the most probable location of electron in
  3-D space around nucleus (probability map)
• -
  most probable
• location called an
• orbital
• - orbitals can hold a
• maximum of 2 e-

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Most Successful Theory of 20th Century
Matter
Dalton
Thompson
Rutherford
Quantum Mechanics
Bohr
De Broglie
Heisenberg
Einstein
Schrödinger
Plank
Maxwell
Wave Mechanics
Newton
Light
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Quantum Mechanics ModelDescribes the arrangement
and space occupied by electrons in atoms
Electrons energy is quantized
Quantum Mechanics
Mathematics of waves to define orbitals (wave
mechanics)
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Bohr Model v. Quantum Mechanics
Bohr Q. Mech.
Energy Electron Position/Path
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Dartboard Analogy
Suppose the size of the probability distribution
is defined as where there is a chance
of all hits being confined.
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Quantum Mechanics Model
The electron's movement cannot be known
precisely. We can only map the probability of
finding the electron at various locations outside
the nucleus. The probability map is called an
orbital. The orbital is calculated to confine 90
of electrons range.
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Arrangement of Electrons in Atoms

• Electrons in atoms are arranged as
• SHELLS (n) distance from nucleus
• 1, 2, 3,
• SUBSHELLS (l) shape of region of probability
• s, p, d, f
• ORBITALS (ml) orientation in space

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Arrangement of Electrons in Atoms

• There is a relationship between the quantum
  number (n) and its the number of subshells.

Principal quantum number (n) number of
subshells
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Representing s Orbitals
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Comparison of 1s and 2s Orbitals
The 2s orbital is similar to the 1s orbital, but
larger in size. Larger means that the highest
probability for finding the electron lies farther
out from the nucleus. Each can hold a maximum of
electrons.
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Probability Maps of the Three 2p Orbitals
The 2p orbital is in the n energy
level. There are 2p orbitals oriented in
three directions. Each orbital can hold a
maximum of electrons. The maximum number of
electrons in the 2p sublevel is . Adding
all 2p orbitals would result in a sphere.
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Probability Maps of the Five 3d Orbitals
The five 3d orbitals are generally oriented in
different directions. Adding all five orbitals,
would result in a sphere. The five orbitals,
taken together, make up the d subshell of the n
3 shell. Each orbital can hold a maximum of two
electrons. This sublevel has a maximum of
electrons.
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Probability Maps of 7 f Orbitals
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Arrangement of Electrons in AtomsElectron Spin
Quantum Number- ms

• Each orbital can be assigned no more than 2
  electrons! And each electron spins in opposite
  directions.

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Electron Spin Quantum Number
Diamagnetic NOT attracted to a magnetic
field Paramagnetic substance is attracted to a
magnetic field. Substance has unpaired electrons.
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Summary
4 QUANTUM NUMBERS

• n ---gt shell 1, 2, 3, 4, ...
• l ---gt sublevel s, p, d, f
• ml ---gt orbital -l ... 0 ... l
• ms ---gt electron spin 1/2 and -1/2

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Pauli Exclusion Principle- No two electrons in
the same atom can have the same set of 4 quantum
numbers.

• Determine the quantum numbers for the outer two
  valence electrons in the lithium atom.

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Aufbau Principle-Electrons fill open lower energy
levels sequentially? lower energy to higher
energy
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Writing Electron Configurations

• Two ways of writing configs. One is called the
  spdf notation.

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Broad Periodic Table Classifications

• Representative Elements (main group) filling s
  and p orbitals (Na, Al, Ne, O)
• Transition Elements filling d orbitals (Fe, Co,
  Ni)
• Lanthanide and Actinide Series (inner transition
  elements) filling 4f and 5f orbitals (Eu, Am,
  Es)

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Writing Orbital Notations

• Two ways of writing configs. Other is called the
  orbital box notation.

One electron has n 1, l 0, ml 0, ms
1/2 Other electron has n 1, l 0, ml 0, ms
- 1/2
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Energy ordering of orbitals for multi-electron
atoms
Different subshells within the same principal
shell have different energies. The more complex
the subshell, the higher its energy. This
explains why the 3d subshell is higher in energy
than the 4s subshell.
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Rules for Filling Orbitals Bottom-up (Aufbaus
principle) Fill orbitals singly before doubling
up (Hunds Rule) Paired electrons have opposite
spin (Pauli exclusion principle)
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Cobalt Symbol Atomic Number Full
Configuration Valence Configuration Shorthand
Configuration
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Orbital diagram and electron configuration for a
ground state lithium atom
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Orbital diagram and electron configuration for a
ground state carbon atom
Hunds Rule- electrons in the same sublevel will
spread out into their own orbital before doubling
up.
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Silicon's valence electrons
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Selenium's valence electrons
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Core electrons and valence electrons in germanium
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Outer electron configuration for the elements
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The periodic table gives the electron
configuration for As
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Valence Electrons by Group
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Ion charges by group
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Periodic Law

• All the elements in a group have the same
  electron configuration in their outermost shells
• Example Group 2
• Be 2, 2
• Mg 2, 8, 2
• Ca 2, 2, 8, 2

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Question

• Specify if each pair has chemical properties
  that are similar (1) or not similar (2)
• A. Cl and Br
• B. P and S
• C. O and S

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General Periodic Trends

• 1. Atomic and ionic size 2. Electron affinity
• 3. Ionization energy 4. Metallic Character
  

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Effective Nuclear Charge, Z

• Z is the nuclear charge experienced by the
  outermost electrons. Screen 8.6.
• Explains why E(2s) lt E(2p)
• Z increases across a period owing to incomplete
  shielding by inner electrons.
• Estimate Z by --gt Z - (no. inner electrons)
• Z number of electrons
• Charge felt by 2s e- in Li Z 3 - 2 1
• Be Z 4 - 2 2
• B Z 5 - 2 3 and so on!

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Effective Nuclear Charge
Figure 8.6
Electron cloud for 1s electrons
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Effective Nuclear Charge, Z

• Atom Z Experienced by Electrons in Valence
  Orbitals
• Li 1.28
• Be -------
• B 2.58
• C 3.22
• N 3.85
• O 4.49
• F 5.13

Increase in Z across a period
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Beryllium
Lithium
Sodium
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Atomic Size

• Size goes UP on going down a group. See Figure
  8.9.
• Because electrons are added further from the
  nucleus, there is less attraction.
• Size goes DOWN on going across a period.

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Atomic Radii
Figure 8.9
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Trends in Atomic SizeSee Figures 8.9 8.10
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Ion Sizes
Does the size go up or down when losing an
electron to form a cation?

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Ion Sizes
Forming a cation.
Li,152 pm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they come.

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Ion Sizes

• Does the size go up or down when gaining an
  electron to form an anion?

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Ion Sizes
Forming an anion.

• ANIONS are LARGER than the atoms from which they
  come.

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Trends in Ion Sizes
Figure 8.13
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Ionization EnergySee Screen 8.12

• IE energy required to remove an electron from
  an atom in the gas phase.

Mg (g) 738 kJ ---gt Mg (g) e-
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Ionization EnergySee Screen 8.12

• Mg (g) 735 kJ ---gt Mg (g) e-
• Mg (g) 1451 kJ ---gt Mg2 (g) e-

Mg2 (g) 7733 kJ ---gt Mg3 (g) e-
Energy cost is very high to dip into a shell of
lower n. This is why ox. no. Group no.
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Trends in Ionization Energy
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Trends in Ionization Energy

• IE increases across a period because Z
  increases.
• Metals lose electrons more easily than nonmetals.
• Metals are good reducing agents.
• Nonmetals lose electrons with difficulty.

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Trends in Ionization Energy

• IE decreases down a group
• Because size increases.
• Reducing ability generally increases down the
  periodic table.
• See reactions of Li, Na, K

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Electronegativity

• A measure of the ability of an atom that is
  bonded to another atom to attract electrons to
  itself.

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Electron Affinity

• A few elements GAIN electrons to form anions.
• Electron affinity is the energy involved when an
  atom gains an electron to form an anion.
• A(g) e- ---gt A-(g)
• E.A. ?E

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Electron Affinity of Oxygen

• ?E is EXOthermic because O has an affinity for an
  e-.

EA - 141 kJ
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Trends in Electron Affinity

• See Figure 8.12 and Appendix F
• Affinity for electron increases across a period
  (EA becomes more negative).
• Affinity decreases down a group (EA becomes less
  negative).

Atom EA F -328 kJ Cl -349 kJ Br -325 kJ I -295
kJ
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Trends in Electron Affinity
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Metallic character trends in the periodic table
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Metallic Character
The text links metallic character to the tendency
to lose electrons in chemical reactions, and
nonmetallic character to the tendency to gain
electrons in chemical reactions. The metallic
character trends therefore follow the ionization
energy trends
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The metallic character trends explain the
location of metals, metalloids, and nonmetals
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Which is the more metallic element, Sn or Te?
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Which is the more metallic element, Si or Sn?