CHEM 1405

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CHEM 1405

• Class Meeting 4

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Assignments and Reminders

• Reading Assignment
• Chapter 3 by Tuesday, Jan 31st
• Homework Problems due next Thursday, Feb 2nd
• Chapter 2 problems 2, 10, 12, 14, 16, 18, 20, 24,
  27, 28, 29, 36, 38, 44, 46, 48
• Class website
• http//iws.ccccd.edu/jstankus/

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Electronic Structure of Atoms

• 5. Describe the Bohr model of the atom.
• 6. What is meant by the electronic configuration
  of an atom? What is meant by principal shell and
  subshell?
• 7. How are electronic configurations related to
  the periodic table?
• 8. What is the valence electronic configuration
  of (a) noble gas atoms, (b) alkali metal atoms,
  (c) alkaline earth metal atoms, and (d) halogen
  atoms?
• 9. Describe periodic trends in (a) atomic size,
  (b) ionization energy, and (c) electron
  affinities.

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Observation of Line Spectra of Elements
Flame Tests
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Bohrs Explanation

• Light can have only discrete amounts of energy
• Energy is quantized
• Electron can have only these energy values and no
  others

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Bohr Model

• Postulated That electrons occupy specific orbits
• Similar to planets orbiting sun

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Bohr Atom details

• Each Shell can hold a certain number of electrons
• Given by of electrons in shell 2n2

First shell (orbit) of e- 2(1)2 2 e-
Second shell (orbit) of e- 2(2)2 8 e-
-
Third shell (orbit) of e- 2(3)2 18 e-
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Bohr Atom details

• Each Orbit has a specific energy level

First shell (orbit) Lowest energy level
Second shell (orbit) Higher energy level
-
Third shell (orbit) Higher energy level
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Nature of Light

• Light of a particular color has a particular
  energy
• Ehn
• Where E is energy of a photon
• h is Plancks Constant
• n is frequency of light

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Transitions Between Orbits

• Each Orbit has a specific energy level

Must add the right amount of energy to transition
between energy levels
-
-
Ehv
A particular color of light has the right amount
of energy
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Transitions Between Orbits

• Going from higher energy to lower energy
  orbits a Photon is
  emitted

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-
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Ground and Excited States

• Electrons prefer to be in the lowest energy
  level
• levels closest to the nucleus
• Ground state
• Excited state
• electron goes from the lowest energy level to a
  higher energy level

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Ground State/Excited State
Ground State All electrons in lowest possible
energy levels

Excited State at least one electron promoted to a
higher energy levels
n3 energy level
n2 energy level
-
-
Hydrogen example
n1 energy level
-
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Ground State/Excited State

Lithium example Z 3
-
-
n3 energy level
n2 energy level
-
-
-
Excited State
Ground State
n1 energy level
-
-
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Bohr Atom Usefullness

• Bohr atom model worked very well for Hydrogen
  but did not agree with experiment for heavier
  elements

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-
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Other Elements Bohr Atom Model did not accurately
predict energy levels
-
Hydrogen Bohr Atom Model worked well
Bohr was on the right track but hypothesis needed
more work
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Quantum Mechanics

• Quantum Mechanics Provided a better model
• Calculated the probability of finding an electron
  in a given amount of space

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First Shell

• First Shell
• 1s orbital

Up to 2 electrons in each orbital
Maximum 2 electrons in first shell
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Second Shell
The three different 2p orbitals have same energy
level
2p orbitals
2s orbital
2px 2py 2pz
2s
Maximum 8 electrons in Second shell
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Third Shell

• 3s 3px 3py 3pz and 3d orbitals

Maximum 18 electrons in Third shell
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Electrons in Orbitals
Orbital Number of
Electrons per Total Number Type
Orientation Orbitals
Orbital of Electrons
s Spherical 1 2 2 p Perpendicular 3 2 6 d
Perpendicular 5 2 10
There are also f orbitals, but we wont deal with
them in this class
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Energy Levels of Orbitals
Single Electron
Multiple Electrons
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Aufbau Principle

• How do we fill up the orbitals with electrons?
• Build up the electronic configuration from the
  lowest energy level on up

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Electronic Configuration of Nitrogen
Z 7 --gt need 7 electrons
2 electrons per orbital 1 spin up, 1 spin down
4p
4d
4s
3d
3p
3s
Start filling from lowest energy level
2p
2s
1s
Energy
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This gives the order of filling the orbitals
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Example Magnesium
Mg Z12
Electrons remaining 12
Fill 1s orbital
Electrons remaining 10
Remember s orbitals 2 electrons p orbitals 6
electrons d orbitals 10 electrons
Fill 2s orbital
Electrons remaining 8
Fill 2p orbital
Electrons remaining 2
Fill 3s orbital
1s2
2s2
2p6
3s2
Mg
Electrons remaining 0
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Full shells are particularly stable electronic
configurations
Noble gas core Shorthand Notation - Use noble gas
to denote completely filled shells
1s22s22p6
Ne
Mg example
1s22s22p6
3s2
Ne
3s2
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Valence Electrons and Core Electrons
The valence shell is the outermost shell of
electrons of an atom.
Mg example
The electrons in the valence shell determine the
chemical properties
1s22s22p6
3s2
Ne
3s2
The electrons in the inner shells are the core
electrons and dont affect the chemical
properties strongly
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Periodic Table

• Dmitri Mendeleev gave us a functional scheme with
  which to classify elements.
• The periodicity of the elements was demonstrated
  by Mendeleev when he used the table to predict to
  occurrence and chemical properties of elements
  which had not yet been discovered.

The Periodic Law Similar physical and chemical
properties recur periodically when the elements
are listed in order of increasing atomic number.
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Periodicity
Periodic behavior of the chemical properties is
due to the periodic pattern of electronic
configuration
Similar chemical Properties within group
Noble gases Alkali Metals
Halogens
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Periodic Table Based on Electronic Configuration
Alkali Metals
Alkaline Earth Metals
Noble Gases
Similar Properties in Vertical Column called a
group
Halogens
Row is called a period
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Groups

• IA are called alkali metals because the react
  with water to from an alkaline solution
• 1 valence electron
• Group IIA are called the alkali earth metals
  because they are reactive, but not as reactive as
  Group IA.
• 2 valence electrons
• Group VIIA are the halogens
• These need only one electron to fill their outer
  shell (7 Valence Electrons)
• They are very reactive.
• Group VIIIA are the noble gases as they have
  completely filled outer shells
• They are almost non reactive.

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Metals, Non-Metals, and Metalloids

• Metal Elements that are usually solids at room
  temperature. Most elements are metals.
• Non-Metal Elements in the upper right corner of
  the periodic Table. Their chemical and physical
  properties are different from metals.
• Metalloid Elements that lie on a diagonal line
  between the Metals and non-metals. Their
  chemical and physical properties are intermediate
  between the two.

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The Modern Periodic Table.
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Periodic Atomic Properties

• Atomic radius is a measure of the size of an atom
  based on the measurement of internuclear
  distances
• Ionization energy is the energy required to
  remove the least tightly bound electron from a
  ground-state atom (or ion) in the gaseous state.
• Electron affinity is the energy change that
  occurs when an electron is added to an atom in
  the gaseous state.

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Atomic Radius

• We cannot measure where the outer electrons of an
  isolated atom are located
• Use half the distance between nuclei in a
  molecule of like elements

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Periodic Trends in Atomic Radii
Electrons in the inner shells shield the positive
charge on the nucleus from the outer shell
electrons holding them less tightly
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Ions

• When an atom gains or loses electrons, it becomes
  charged and is called an ion.
• Tend to go to a noble gas configuration
• Ionization energy is the energy required to
  remove the least tightly bound electron from a
  ground-state atom (or ion) in the gaseous state.
• Creates a Positive Ion
• Electron affinity is the energy change that
  occurs when an electron is added to an atom in
  the gaseous state.
• Creates a Negative Ion

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Ionization energy Trends
(difficulty of losing electron to make positive
ion)
Highest Ionization energies are noble gases -
most stable electronic configuration
Electron Affinity Trends
(ease of gaining electron to make negative ion)
Highest Electron Affinities are halogens - Going
to noble gas configuration
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Atomic Property Trends