Chapter 11 ­ Modern Atomic Theory Notes
Electromagnetic radiation ­ energy that travels through space as waves.
Waves have three primary characteristics:
Wavelength ( lambda) ­ distance between two consecutive peaks or troughs in a wave. Unit = meter
Frequency ( nu) ­ indicates how many waves pass a given point per second. Unit = Hertz (Hz)
c = λν
Speed ­ velocity (c = speed of light = 3.0 x 108 m/sec) ­ indicates how fast a given peak moves in a unit of time
ex: If red light was a wavelength of 7.00 x 10­9 m ,what is its frequency?
So 3.0 x 108 m/sec = 7.00 x 10­9 m (ν)
ν = 3.0 x 108 m/sec = 4.3 x 1016 1/sec or 4.3 x 10 16 Hz
7.00 x 10 ­9 m
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Electromagnetic radiation (light) is divided into various classes according to wavelength.
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Wave­ Particle Theory ­ Light as waves ­ Light as photons
"In 1905 Einstein ... suggested that light itself could behave like little particles or quanta, with energy proportional to the frequency (the colour) of the light. These particles of light are what we now call photons. Einstein's suggestion went straight to the heart of Quantum Theory, and began to expose the fundamental conceptual difficulties associated with it. Indeed it was for this work that he was awarded the Nobel prize in 1921." http://www.tcd.ie/Physics/Schools/what/atoms/quantum/duality.html
Photon/quantum ­ packet of energy ­ a “particle” of electromagnetic radiation
Excited State ­ atom with excess energy
When a Photon/light wave hits an electron its transfers energy to that electron and thus the electron is moved to a higher energy level.
Ground State ­ lowest possible energy state When an electron releases that extra energy it also emits a photon/light wave.
• Wavelengths of light carry different amounts of energy per photon
• Only certain types of photons are produced (see only certain colors)
• Quantized ­ only certain energy levels (and therefore colors) are allowed
Bohr Model ­ suggested that electrons move around the nucleus in circular orbits
Incorrect but usefull for discribing excited and ground states.
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The electrons on this Bohr model atom have been hit by photons and are currently in an Excited State. In order to release the excess energy they must give off a photon of distinct light which corresponds to how far the electron drops. After giving off this energy the atom will be back in its Ground State . Title: Jul 19­9:12 AM (4 of 13)
We can only see the changes between an Excited State and the Ground State. However, this occurs throughout the entire electronmagnetic spectrum and these changes in energy levels have been studied for every atom known today. An atom always gives off the same pattern of light.
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Wave Mechanical Model ­ (Quantum Mechanical Model)
The Bohr model is severely lacking in desribing many of the atom to atom interactions which have been observed through experimentation. To account for these observations a mathematical model was created by a large team of scientists and it is called the Wave Mechanical Model (WMM). The WMM divides and subdivides the areas electrons can be found into Energy Levels, then Sublevels, and finally Orbitals. The description of an electron within an Orbital is said to be 95% accurate allowing for some transitions between orbitals as energy is exchanged. This also takes into account the Heisenberg Uncertainty Principle that states it is impossible to know both the location and velocity of an object. The more you know about where an electron occupies a certain point in space the less you know about how it moves and the converse is also true.­ Principle Energy Level (n) ­ Largest area which is divided up into smaller areas.
The energy levels are designated by numbers 1­7.
They are also called principle quantum numbers
Sublevel (l) ­ The largest subdivision within each principle energy level
The electron in a given sublevel have the same energy
If electrons are in different sublevels then the energy is slightly different
lowest energy sublevel is s, then p, then d, and then f is the most energy
Orbital ­ The further subdivision of the sublevels where 2 electrons can be found
The s sublevel contains 1 orbital
The p sublevel contains 3 orbitals
The d sublevel contains 5 orbitals
The f sublevel contains 7 orbitals
­ ** No more than two electrons can occupy an orbital**
an orbital can be empty, half­filled, filled
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Summary
Principle # of
Energy Level sublevels
# of orbitals present
s p d f
Total # of orbitals
Maximum # of electrons
1
1
1s
1
2
2
2
2s 2p 2p 2p 4
8
3
3
3s 3p 3p 3p 3d 3d 3d 3d 3d 9
18
4
4
4s 4p 4p 4p 4d 4d 4d 4d 4d
4f 4f 4f 4f 4f 4f 4f 32
16
Orginally it was believed that electrons always filled up orbitals starting with the lowest energy level and proceeded up higher and higher energy levels until all the electrons in the atom would occupy an orbital. However, this was adjusted drastically after Aufbau suggested a new order.
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Order of Electrons for Hydrogen does not work for all of the elements, which led us to Aufbau's Order.
With Hydrogen
Start Here 1
2
3
4
5
6
7
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1. Start Here
1s2
2s2 3s2 4s2 5s2 6s2 7s2 2. Then follow this diagonal
3. And so on
2p6
3p6 4p6 5p6 6p6 7p6 4. And so on
5. Until
6. all the 3d10
4d10 4f14
5d10 5f14
6d10 7. electrons
8. have orbitals
Shapes of orbitals
­
All s orbitals are spherical as the principle energy level increases the diameter increases.
­
All p orbitals are dumbbell shaped ­ all have the same size and shape within an energy level
n = 1 s­ orbital in the middle
n = 2 Each additional energy level of electrons is farther from the nucleus with more room between orbitals at the higher energy levels which leads to the introduction of the d and f orbitals.
s and p orbitals
second layer
n = 3 s, p and d(not drawn)
third layer
I know the drawing isn't that good and that my skill is limited. You will not be required to draw, I just want you to understand what the s, p, d, and f's represent when we write them out.
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Electron Spin ­ motion that resembles earth rotating on its axis­ clockwise or counterclockwise
Represented by arrows point up and down.
Pauli Exclusion Principle
­ two electrons in the same orbitals must have opposite spins
_________
1s2 correct orientation
_________
1s2 no possible
_________
1s2 not possible
Orbital Diagram ­ describes the placement of electrons in orbitals
­
use arrows to represent electrons with spin
K has 19 electrons 1s2 2s2 2p6 3s2 3p6 4s1
­
line represents orbital
________ full ________ half­full
________ empty Hund’s Rule ­ All orbitals within a sublevel must contain at least one electron before any orbital can have two
Electron Configuration ­ arrangement of the electrons among the various orbitals of the atom
Aufbau Order ­ Tool to predict the order in which sublevels will fill (On the back of your periodic table.)
Ar has 18 electrons 1s 2 2s2 2p6 3s2 3p6 the superscripts 2+2+6+ 2+6 = 18 electrons
K has 19 electrons 1s 2 2s2 2p6 3s2 3p6 4s1 the superscripts 2+2+6+ 2+6+1 = 19 electrons
Fe has 26 electrons 1s 2 2s2 2p6 3s2 3p6 4s2 3d6 the superscripts 2+2+6+ 2+6+2+6 = 26 electrons
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Noble Gas Configuration ­ Shorthand configuration that substitutes a noble gas for electrons
Lets look at our previous 3 examples. Ar is a Noble Gas and its configuration is the same as K and Fe up to the last couple of orbitals. We can write [Ar] to represent that part of the electron configuration.
Ar has 18 electrons 1s2 2s2 2p6 3s2 3p6 <­­­ will be represented by [Ar] in larger elements
K has 19 electrons 1s2 2s2 2p6 3s2 3p6 4s1 but using the Noble Gas Configuartion:
K is [Ar] 4s1
Fe has 26 electrons 1s2 2s2 2p6 3s2 3p6 4s2 3d6 but using the Noble Gas Configuartion:
Fe is [Ar] 4s2 3d6 This can be very usefull for large elements like Barium 1s2 2s2 2p6 3s2 3p6 4s2 3d6 4p6 5s2 4d6 5p6 6s2 = [Xe] 6s2
Valence Electrons ­ Electrons in the outermost (highest) principle energy level in an atom
Look at previous examples. 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Core Electrons ­ innermost electrons ­ not involved in bonding
Those electrons which are not at the HIGHEST energy level.
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Periodic Table Dimitri Mendeleev­1869­ developed the first version of the periodic table.
He expressed the regularities as a periodic function of the atomic mass.
Henry Moseley ­ revised Mendeleev’s periodic table by describing regularities in
physical and chemical properties as periodic functions of the atomic
number
Groups (family) ­ vertical column
Elements with similiar valence electrons configurations
Group 1 ­ alkali metals ­ reactive
Group 2 ­ alkaline earth metals ­ reactive
Group 3­12 ­ transition metals
Group 15 ­ nitrogen family
Group 16 ­ oxygen family ­ reactive
Group 17 ­ halogens ­ very reactive
Group 18 ­ noble gases
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Periods ­ horizontal rows
Period number corresponds to the principal quantum number of valence electrons
Metallic Character Increases ­ down a group
Decreases ­ across a period
Ionization Energy ­ energy required to remove an electron from an individual
atom in a gas phase M(g) ­­> M+(g) + e­
• Metals lose electrons to non­metals so relatively low energy is needed
• High ionization energy means an electron is hard to remove
Decreases ­ down a group
Increases ­ across a period
Atomic Size Increases ­ down a group
Decreases ­ across a period
Ionic Size depends upon how many electrons the atom has lost or gained.
Cation
Ca+2/Ca
Ca larger because Ca +2 lost 2 electrons
Anion
S­2/S
S­2 larger because S ­2 gained 2 electrons
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Attachments
h2linespectrum.mov
orbitalenergies.mov