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Chemistry Unit 3

Chapter 4: Arrangement of Electrons in Atoms
Section 4-1: The Development of a New Atomic Model
1. Explain the mathematical relationship between the speed, wavelength, and frequency of electromagnetic radiation.
2. Discuss the dual wave-particle nature of light.
3. Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model.
4. Describe the Bohr model of the hydrogen atom.
Bohr Emission Spectrum Animation
   Electronic Structure of Atoms
Section 4-2: The Quantum Model of the Atom
1. Discuss Louis de Broglie's role in the development of the quantum model of the atom.
2. Compare and contrast the Bohr model and the quantum model of the atom.
3. Explain how the Heisenberg uncertainty principle and the Schrödinger wave equation led to the idea of atomic orbitals.
4. List the four quantum numbers, and describe their significance.
5. Relate the number of sublevels corresponding to each of an atom's main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.
Section 4-3: Electron Configurations
1. List the total number of electrons needed to fully occupy each main energy level.
2. State the Aufbau principle, the Pauli exclusion principle, and Hund's rule.
3. Describe the electron configurations for the atoms of any element using orbital notation, electron-configuration notation, and, when appropriate, noble-gas notation.
Electron Arrangement Lecture outline
   The quantum theory was used to show how the wavelike behavior of electrons leads to quantized energy states when the electrons are bound or trapped. In this section, we'll use the quantum theory to explain the origin of spectral lines and to describe the electronic structure of atoms.
Emission Spectra
   experimental key to atomic structure: analyze light emitted by high temperature gaseous elements
experimental setup: spectroscopy:  atoms emit a characteristic set of discrete wavelengths- not a continuous spectrum!
atomic spectrum can be used as a "fingerprint" for an element
hypothesis: if atoms emit only discrete wavelengths, maybe atoms can have only discrete energies an analogy  A turtle sitting on a ramp can have any height above the ground- and so, any potential energy  
A turtle sitting on a staircase can take on only certain discrete energies

Energy is required / absorbed to move the turtle up the steps (absorption)
Energy is released when the turtle moves down the steps (emission)
only discrete amounts of energy are absorbed or released (energy is said to be quantized)
Energy staircase diagram for atomic hydrogen  bottom step is called the ground state higher steps are called excited states
The quantum mechanical atom
Electrons in atoms have quantized energies
Electrons in atoms are bound to the nucleus by electrostatic attraction
Electron waves are standing matter waves
standing matter waves have quantized energies, as with the "electron on a wire" model
Electron standing matter waves are 3 dimensional
The electron on a wire model was one dimensional; one quantum number was required to describe the state of the electron
A 3D model requires three quantum numbers
A three-dimensional standing matter wave that describes the state of an electron in an atom is called an atomic orbital
The energies and mathematical forms of the orbitals can be computed using the Schrödinger equation
Every electron adds 3 variables (x, y, z) to the equation; it's very hard to solve equations with lots of variables.
Energy-level separations computed with the Schrödinger equation agree very closely with those computed from atomic spectral lines
Quantum numbers
Think of the quantum numbers as addresses for electrons
Principal quantum number, n
determines the size of an orbital (bigger n = bigger orbitals)
largely determines the energy of the orbital (bigger n = higher energy)
can take on integer values n = 1, 2, 3, ...,  
all electrons in an atom with the same value of n are said to belong to the same energy level.
Bohr used the following letters for the energy levels.
1 - K  5 - O
2 - L  6 -  P
3 - M  7 - Q
4 - N  

The azimuthal quantum number,
designates the overall shape of the orbital within a shell
affects orbital energies (bigger  = higher energy)
all electrons in an atom with the same value of  are said to belong to the same subshell
only integer values between 0 and n-1 are allowed
sometimes called the orbital angular momentum quantum number
spectroscopists use the following notation for subshells

s  spherical crosses all axies at the same time - only one orientation in space.
p peanut or dumbbell - 2 lobes
d clover - 4 lobes
f  flower - 8 lobes

The magnetic quantum number, m
determines the orientation of orbitals within a subshell
does not affect orbital energy (except in magnetic fields!)
The number of possible m values determines the number of orbitals in a subshell.  possible values of m number of orbitals in this subshell

The spin quantum number, ms
Several experimental observations can be explained by treating the electron as though it were spinning
spin makes the electron behave like a tiny magnet
spin can be clockwise or counterclockwise
spin quantum number can have values of +1/2 or -1/2  or ? or ?
Electron configurations of atoms
A list showing how many electrons are in each orbital or subshell in an atom or ion
subshell notation: list subshells of increasing energy, with number of electrons in each subshell as a superscript
1s2 2s2 2p5 means "2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 5 electrons in the 2p subshell"
1s2 2s2 2p6 3s2 3p3 is an electron configuration with 15 electrons total; 2 electrons have n=1 (in the 1s subshell); 8 electrons have n=2 (2 in the 2s subshell, and 6 in the 2p subshell); and 5 electrons have n=3 (2 in the 3s subshell, and 3 in the 3p subshell).
ground state configurations fill the lowest energy orbitals first

Electron configurations of the first 11 elements, in subshell notation. Notice how configurations can be built by adding one electron at a time.
atom Z ground state electronic configuration
H 1 1s1
He 2 1s2
Li 3 1s2 2s1
Be 4 1s2 2s2
B 5 1s2 2s2 2p1
C 6 1s2 2s2 2p2
N 7 1s2 2s2 2p3
O 8 1s2 2s2 2p4
F 9 1s2 2s2 2p5
Ne 10 1s2 2s2 2p6
Na 11 1s2 2s2 2p6 3s1

Writing electron configurations
Strategy: start with hydrogen, and build the configuration one electron at a time (the Aufbau principle)
Fill subshells in order by counting across periods, from hydrogen up to the element of interest:

Rearrange subshells (if necessary) in order of increasing n & l
Examples: Give the ground state electronic configurations for:

Electron configurations including spin
unpaired electrons give atoms (and molecules) special magnetic and chemical properties
when spin is of interest, count unpaired electrons using orbital box diagrams

Examples of ground state electron configurations in the orbital box notation that shows electron spins.

B 1s 2s  2p
C 1s 2s  2p
N 1s 2s 2p
O 1s 2s 2p
F 1s 2s 2p
Cl 1s 2s 2p 3s 3p
Mn 1s 2s 2p 3s 3p  … 3d 4s
Orbital notation
orbital box diagrams
write the electron configuration in subshell notation
draw a box for each orbital.
Remember that s, p, d, and f subshells contain 1, 3, 5, and 7 degenerate orbitals, respectively.
Remember that an orbital can hold 0, 1, or 2 electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins (Pauli principle)
within a subshell (depicted as a group of boxes), spread the electrons out and line up their spins as much as possible (Hund's rule)

Configurations with unpaired electrons are attracted to magnetic fields (paramagnetism)
Configurations with only paired electrons are weakly repelled by magnetic fields (diamagnetism)
Chemistry involves mostly the shell with the highest value of principal quantum number, n, called the valence shell
Noble gas configuration
Simplify the electron configuration notation by replacing the core with a noble gas symbol in square brackets.  Showing only the valence electrons.
Electron configuration and noble gass configuration
O 1s2 2s2 2p4 He 2s2 2p4 [He] 2s2 2p4
Cl 1s2 2s2 2p6 3s2 3p5 Ne 3s2 3p5 [Ne] 3s2 3p5
Al 1s2 2s2 2p6 3s2 3p1 Ne 3s2 3p1 [Ne] 3s2 3p1

Isoelectric Noble gas configuration
Fe [Ar] 3d6 4s2                  [Ar] 3d6 4s2
Sn [Kr] 4d10 5s2 5p2       [Kr] 4d10 5s2 5p2
Hg [Xe] 4f14 5d10 6s2     [Xe] 4f14 5d10 6s2
Pu [Rn] 5f6  7s2                 [Rn] 5f6 7s2

Quantum Mechanics
Each electron in an atom is described by four different quantum numbers. Three of these quantum numbers (n, l, and m) represent the three dimensions to space in which an electron could be found. A wave function for an electron gives the probability of finding the electron at various points in space. A wave function for an electron in an atom is called an atomic orbital. The fourth quantum number (ms) refers to a certain magnetic quality called spin.
n-The Principal Quantum Number
The n quantum number relates to the size of the atomic orbital. n can have any positive integer value from 1 to 7. The smaller the n, the lower the energy, the higher the value of n, the higher the energy. In the case of any single-electron atom, or hydrogen atom, n is the only quantum number which determines the energy. The size of an orbital depends on n. The larger the orbital, the larger the value of n. Orbitals of the same quantum state belong the same shell. To use an analogy for n, why not relate it to the size of a computer, where larger values would represent larger houses.

l-The angular momentum quantum number
l can have any integer value from 0 to 3. This quantum number distinguishes orbitals of a given n value which have different states. Or, the secondary quantum number gives the shape of the orbital so the analogy can be made to the shape of the computer with larger values associated with computers with more components.

M-magnetic quantum number
The third quantum number has to do with the orientation of an orbital in a magnetic field. Because of this, we can relate its values to different directions the computer might be facing.

The final quantum number is the spin quantum number, it describes the spin orientation of an electron.
The electron configuration of an atom is the particular distribution of electrons among available shells. It is described by a notation that lists the subshell symbols, one after another. Each symbol has a subscript on the right giving the number of electrons in that subshell. For example, a configuration of the lithium atom (atomic number 3) with two electrons in the 1s subshell and one electron in the 2s subshell is written 1s22s1.
sublevel orbital maximum # of electrons
s 1 2
p 3 6
d 5 10
f 7 14

The notation for electron configuration gives the number of electrons in each subshell. The number of electrons in an atom of an element is given by the atomic number of that element.
On the left we have a diagram to show how the orbitals of a subshell are occupied by electrons. On the right there is a diagram for the filling order of electrons in a subshell.

Here are some examples that show how to use the filling order diagram to complete the electron configuration for a certain substance.
Element # of Electrons in Element Electron Configuration
He 2 1s2
Li 3 1s22s1
Be 4 1s22s2
O 8 1s22s22p4
Cl 17 1s22s22p63s23p5
K 19 1s22s22p63s23p64s1

Often times you will be asked to find the electron configuration for something that looks like this:
The 53 denotes the number of electrons in an atom of iodine. You would now proceed to do the electron configuration by looking at the filling order chart.
1s22s22p63s23p64s2 3d104p65s24d105p5
Chapter 5: The Periodic Law
Section 5-1: History of the Periodic Table
1. Explain the roles of Mendeleev and Moseley in the development of the periodic table.
2. Describe the modern periodic table.
3. Explain how the periodic law can be used to predict the physical and chemical properties of elements.
4. Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.
Section 5-2: Electron Configuration and the Periodic Table
1. Describe the relationship between electrons in sublevels and the length of each period of the periodic table.
2. Locate and name the four blocks of the periodic table. Explain the reasons for these names.
3. Discuss the relationship between group configurations and group numbers.
4. Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases.
Section 5-3: Electron Configuration and Periodic Properties
1. Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.
2. Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.
3. Define valence electrons, and state how many are present in atoms of each main-group element.
4. Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements.
Periodic Law
Mendeleev: Russian scientist arranged the periodic table according to atomic mass concerned.  Gaps for missing elements. Given credit foe discovering the periodic law
Periodic Law: The physical and chemical properties of the element are periodic functions of their atomic number.
Moseley: Arranged the periodic table according to atomic number. Used Bohr's model to count electrons and set equal to protons.
Relationships of the chemical properties of the elements to position on a table
Location and identification of the following:Periods: seven periods each represent a filled or filling energy level.

Groups / families
Alkali metals: most reactive metals, 1 valence electron, silvery and shiny,  Forms Cations, bonds with nonmetals to form ionic bonds, not found free in nature.
Alkaline earth metals : second most reactive metals, 2 valence electrons, silver, shinny, forms ionic bonds with nonmetals, forms cations.
Transitional metals : Group B metals, true metal qualities, multiple oxidation numbers, produce color.
   Hydrogen: exception to the alkali family no metallic properties.
   Helium: exception to the octet rule for noble gases.
   Rare earth metals: Lanthanide and actinide series.
   Metalloids: contains elements with both metallic and nonmetallic  
   Halogens: most reactive nonmetals and forms ionic bonds with metals.
   Inert Noble gases Group 18, 8A,0
   A, main group, s & p block elements
   Sublevel Blocks
   s block Group 1a and 2a
   p block :groups 3A to 8A or 12=18
   d block : Group 3B- 2V , 3-12
   f block : lanthanoid and actinoides (inner              
                  Transitional metals)
Periodic properties
With increasing atomic number, the electron configuration of the atoms display a periodic variation. Because of this the elements show periodic variations of both physical and chemical behavior. The periodic law is a law stating that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. We are going to be looking at three physical properties of an atom: atomic radius, ionization energy, and electron affinity.
Atomic Radius
The size of the electron cloud increases as the principal quantum number increases. Therefore, as you look down the periodic table, the size of atoms in each group is going to increase. When you look across the periodic table, you see that all the atoms in each group have the same principal quantum number. However, for each element, the positive charge on the nucleus increases by one proton. This means that the outer electron cloud is pulled in a little tighter. One periodic property of atoms is that they tend to decrease in size from left to right across a period of the table. So finally we have a good definition for how the atomic radii increases: the atomic radii increases top to bottom and right to left in the periodic table.

Ionization Energy
The energy needed to remove the most loosely held electron from an atom is known as ionization energy. Ionization energies are periodic. The ionization energy tends to increase as atomic number increases in any horizontal row or period. In any column or group, there is a gradual decrease in ionization energy as the atomic number increases. Metals typically have a low ionization energy. Nonmetals typically have a high ionization energy.

Electron Affinity
The attraction of an atom for an electron is called electron affinity. Metals have low electron affinities while nonmetals have high electron affinities. The general trend as you go down a column is a decreasing tendancy to gain electrons. As you go across a row there is also a trend for a greater attraction for electrons.