1. Chapter 5 CH1g,1h,2g,2i Electrons in Atoms
Actually, the Chemical History powerpoint talked about a lot of the stuff from Chapter 5 too, specifically part of section 1 and most of section 3.
This powerpoint will talk about what is still left over in Chapter 5.

2. Quantum Mechanics…where we left off from Chem History http://www
Better than any previous model, quantum mechanics does explain how the atom behaves.
Quantum mechanics treats electrons not as particles (which they are), but as waves (like light) which can gain or lose energy.
But they can’t gain or lose just any amount of energy. They can only gain or lose a “quantum” of energy.
A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level. Max Planck, another German Nobel Prize winning scientist first came up with this idea.

3. What the heck is a Quantum. http://www. blogcdn. com/www. slashfood
Think of a quantum as a “packet” of energy, much like a sugar packet at a restaurant. A sugar packet contains a teaspoonful of sugar.
If the electron absorbs energy, it moves to a higher energy level. If it emits (loses) energy, it moves to a lower energy level.
But like Bohr suggested in his model, the electron has to gain or lose exactly the right amount. That amount is a quantum of energy.
C12H19O8Cl3 is the formula for sucralose, which is the chemical name for Splenda. That molecule (in yellow) is sucralose. It’s an Organic compound.

4. Neils Bohr: The Planetary Model & Energy Levels (http://www. usd
You can’t just step anywhere.
You have to step on the rungs of a ladder.
An electron has to jump from one level to another.
The steps on a ladder are all the same distance apart.
But in Bohr’s model, the energy levels get closer and closer the further away you get from the nucleus.
An energy level is like a step or rung on a ladder. In the Solar System, as you get further away from the Sun, the planets get further apart.

5. Quantum Mechanics “borrowed” the concept of energy level.
The electron really doesn’t orbit (like a little planet) around the nucleus.
Quantum mechanics describes “electron clouds” and where they are in relation to the nucleus.
Again, the electron can move from one energy level to another, IF it absorbs a quantum of energy.
I told you Schroedinger just “borrowed” the stuff he wanted from everyone else’s work. But he did explain what they all had in common and how it all “fit” together.

6. Energy Levels & Quantum Numbers http://www. chem4kids
Quantum mechanics has a principal quantum number.
It is represented by a little n. It represents the “energy level” similar to Bohr’s model.
n=1 describes the first energy level
n=2 describes the second energy level. Etc.
Each energy level represents a period or row on the periodic table.
Isn’t it amazing how all this stuff just “fits” together?
Red n = 1
Orange n = 2
Yellow n = 3
Green n = 4
Blue n = 5
Indigo n = 6
Violet n = 7

7. Atomic Orbitals http://milesmathis.com/bohr2.jpg
The energy levels in quantum mechanics describe locations where you are likely to find an electron cloud.
Schroedinger used calculus to calculate the PROBABILITY of finding an electron in a particular location.
These locations are called ORBITALS.
Orbitals are “geometric shapes” around the nucleus where electrons are found.
There must be at least a 90% probability of finding an electron there.
The 4 different types of orbitals are s, p, d, and f.

8. Atomic Orbitals http://courses. chem. psu. edu/chem210/quantum/quantum
Think of orbitals  as sort of a "border” for spaces around the nucleus inside which electrons are allowed.
No more than 2 electrons can ever be in 1 orbital.
The orbital just defines an “area” where you can find 1 or 2 electrons.
No more than 2 can fit into any one orbital.
What is the chance of finding an electron in the nucleus?
Yes, of course, it’s zero.
There aren’t any electrons in the nucleus.
A node = a location where the probability of finding an electron there = 0.

9. Atomic Orbitals define an area where electrons are moving
3s s s
Quantum mechanics doesn’t predict a SPECIFIC orbit, like the Bohr model does.
We don’t really know how the electron is moving, or if it follows any particular path as it moves.

10. Energy Sub-level = Specific Atomic Orbital
Each energy level has 1 or more “sub- levels” which describe the specific “atomic orbitals” for that level.
n = 1 has 1 sub-level (the “s” orbital)
n = 2 has 2 sub-levels (“s” and “p”)
n = 3 has 3 sub-levels (“s”, “p” and “d”)
n = 4 has 4 sub-levels (“s”, “p”, “d” and “f”)
s, p, d, f refer to specific areas on the Periodic Table where those orbitals are being filled with electrons.
A second quantum number identifies the specific orbital.
Blue = s block (0)
Yellow = p block (1)
Red = d block (2)
Green = f block (3)

11. Shapes of These Orbitals (the nucleus is ALWAYS at the center of the orbital)
The s orbital looks like a ball or sphere.
The p orbital looks like a dumb-bell.
These orbitals are all perpendicular to each other.
The d orbitals have two shapes.
4 of the 5 look like “4-leaf clovers.”
The 5th one looks like a “big dumb-bell” with a “hula-hoop” around the middle.
The shapes of the f orbitals are complex.
We have a slide showing them, but you don’t need to remember them, nor will they be on the test. But s, p and d will be.

12. Shapes of s, p, and d Orbitals http://media-2. web. britannica
In the s block, electrons are going into s orbitals.
In the p block, the s orbitals are full.
New electrons are going into the p orbitals.
In the d block, the s is full but the p orbitals are not full.
New electrons are going into the d orbitals, because we are in the transition metals. THIS is characteristic of the d block.

13. f orbitals http://antoine. frostburg

14. g orbitals = Science Fiction. 2,8,18,32…50. http://antoine. frostburg
Dr. Seaborg predicted the g orbitals would start with element number 121, which has not been invented yet. The g block will have 18 elements.
Will his hypothesis be proven true?

15. To Summarize Complete the chart in your notes as we discuss this.
Energy Level
Sub-levels
Total Orbitals
Total Electrons
Total Electrons per Level
n = 1 s
1 (1s orbital) 2
n = 2 p
1 (2s orbital)
3 (2p orbitals) 6 8
n = 3 d
1 (3s orbital)
3 (3p orbitals)
5 (3d orbitals) 10 18
n = 4 f
1 (4s orbital)
3 (4p orbitals)
5 (4d orbitals)
7 (4f orbitals) 14 32
Complete the chart in your notes as we discuss this.
The first level (n=1) has an s orbital. It has only 1. There are no other orbitals in the first energy level.
We call this orbital the 1s orbital.

16. Island of Stability http://www. nytimes
This is another hypothesis from Dr. Seaborg. His thought was that element 114 would be an “island of stability,” especially if it also had 184 neutrons. It would aehv a mass number of 298.
However, other “islands” might be 120 or Detailed and complicated math calculations are necessary to figure out these numbers.
Most synthesized elements only last for fractions of seconds. However, in researchers synthesized element 114 and it lasted for 30 seconds. Perhaps this is the “shore” of the Island of Stability that Dr. Seaborg hypothesized.
The element 114 was made using some of the original Pu-244 that Dr. Seaborg himself made in the early 1940s. They bombarded plutonium with Ca-48 atoms to form some of the new element 114.
Element 114 is now know as Flerovium (symbol Fl); it was named in 2012.
It took 14 years to agree on the name.
All of the atoms so far have had mass numbers of Therefore, the “island” still remains undiscovered.

17. Island of Stability http://www. sciencecodex
Famous picture of the “Island of Stability” showing the island off in the distance (top right) with 114 protons and 184 neutrons. An element with Z = 184 is also predicted to be another “island of stability.”

18. Electron Configurations Section 2
What do I mean by “electron configuration?”
The electron configuration is the specific way in which the atomic orbitals are filled.
Think of it as being similar to your address. The electron configuration tells me where all the electrons “live.”

19. Rules for Electon Configurations https://teach. lanecc
In order to write an electron configuration, we need to know the RULES.
3 rules govern electron configurations.
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Using the orbital filling diagram at the right will help you figure out HOW to write them
Start with the 1s orbital.
Fill each orbital completely and then go to the next one, until all of the electrons have been accounted for.
FOLLOW the arrows!!!

20. Fill Lowest Energy Orbitals FIRST http://www. meta-synthesis
Each line represents ONE orbital.
1 (s), 3 (p), 5 (d), 7 (f)
The Aufbau Principle states that electrons enter the lowest energy orbitals first.
The lower the principal quantum number (n) the lower the energy.
Within an energy level
s orbitals have the lowest energy
followed by p then d and then f.
f orbitals are the highest energy for that level.
High Energy
Low Energy

21. No more than 2 Electrons in Any Orbital…ever. http://www. fnal
The next rule is the Pauli Exclusion Principal.
The Pauli Exclusion Principle states that an atomic orbital may only have 1 or 2 electrons and then it is full.
The spins have to be paired.
We usually represent this with an up arrow and a down arrow.
Wolfgang Pauli, yet another German Nobel Prize winner
Quantum numbers describe an electrons position in the atom, and no 2 electrons can have the exact same quantum numbers. Because of that, electrons must have opposite spins from each other in order to “share” the same orbital.

22. Hund’s Rule (Dog’s Rule. ) http://intro. chem. okstate
Hund’s Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons.
Degenerate means they have the same energy.
p orbitals are degenerate because there are 3 of them on EACH level.
d and f orbitals are also degenerate.
Don’t pair up the 2p electrons until all 3 orbitals are half full.

23. Let’s Try Some…
NOW that we know the rules, we can try to write some electron configurations.
Remember to use your orbital filling guide to determine WHICH orbital comes next in the sequence (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc).
Follow the arrows!!
Lets write some electron configurations for the first few elements, and let’s start with hydrogen.
There are also shorthand electron configurations, but we will look at those after Chapter 6.

24. Electron Configurations
Element
Configuration
H Z=1
1s1
He Z=2
1s2 (1s is now full)
Li Z=3
1s22s1
Be Z=4
1s22s2 (2s is now full)
B Z=5
1s22s22p1
C Z=6
1s22s22p2
N Z=7
1s22s22p3
O Z=8
1s22s22p4
F Z=9
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), = 10

25. Electron Configurations of Alkali Metals (and H)
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
This similar configuration causes them to behave the same chemically.
It’s for that reason they are in the same family or group on the periodic table.
Each group will have the same ending configuration, in this case something that ends in s1.

26. Exceptions to the Rules for Electron Configurations
Element
Configuration
Cr should be
1s22s22p63s23p64s23d4 (Z=24)
BUT Cr is
1s22s22p63s23p64s13d5 (d half full)
Cu should be
1s22s22p63s23p64s23d9
BUT Cu is
1s22s22p63s23p64s13d10 (d is full)
Exceptions in the d block (transition metals) occur because a half full OR totally full set of d orbitals is energy favorble.
This is illustrated for Chromium and Copper.
These are the only 2 exceptions in the first row of the d block.
LOTS of other exceptions occur in the d block, but these are the only 2 you need to know…for now.

27. More HW…OMG!
Chemistry: write full electron configurations for elements 1-36.
Also write orbital diagrams for 3-10
Advanced Chemistry: write full electron configurations for Rb, Sr, Y, Ag, I, Kr, Cs, Ba, La, Ce, Hf, Pb.
Also write orbital diagrams for 11-18

28. Emission Spectra = Fingerprint of the Elements (Section 3) http://www
Atomic emission spectrum is sometimes called a line spectrum, to distinguish it from the continuous spectrum.

29. Emission Spectra = Fingerprint of the Elements (Section 3) http://www
The top 3 (H, Hg, Ne) are emission spectra.
The bottom one is an absorption spectrum of H.

30. Emission Spectra = Fingerprint of the Elements
Atomic emission spectra are “unique.” You can use the spectrum to identify the element (like a fingerprint).
Bohr’s model predicted and explained emission spectra by pointing out how electrons can move from one energy level to another.
His model also explained why metals glow red when they are heated.
Scientists can look at light from a distant star and analyze it and determine what types of elements make up that star.
Just by looking at the light!
No element (except H) has those same 4 lines in its spectrum.

31. All the Rest of Section 3….
…was covered in the Chemical History power point.
Photoelectric effect
A photon is a quantum of light. It is light behaving as a particle. A photon has a certain wavelength, frequency and energy.
De Broglie equation
Showed that particles could also act as waves.
Heisenberg’s uncertainty principle
Principal = Dr. Gordon
Principle = a statement that explains how or why something works scientifically

32. Organic Flavor Carbon can bond and form long chains
Like in soap molecules
Carbon can form rings
Like in sugar molecule
Carbon can form huge networks of carbon atoms
Like in diamonds
Carbon is always bonded to 4 things. Carbons unique size (atomic radius) and electronegativity (we’ll get to that in a minute) means that it can form very strong COVALENT bonds between itself and H and O and N and other atoms (that means electrons are SHARED).

33. The End