1. Orbital Diagrams, Valence Electrons, Lewis Electron Dot Structures and the Periodic Table

2. Orbital Notation (Diagrams)
Since we already know how to do electron configurations of different elements, we can draw what they look like in an orbital diagram
All we need to know is the element, how many orbitals it contains and how many electrons can fit within each orbital level!
Before we can start, we must follow some rules…

3. Pauli Exclusion Principle
No two electrons in the same atom can have the same set of four quantum numbers
This is due to the opposite spins of the electrons within the orbitals
Aufbau Principle
Orbitals of lowest energy are filled first
Hund’s Rule
Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron

4. 1s1 Let’s look at Hydrogen’s electron configuration for example:
The “s” sublevel of our notation means that we will have one orbital (represented by a circle) that can hold a maximum of 2 electrons.
Since we know that the “s” orbital can hold two electrons and we only have one electron to put inside of the circle (Hydrogen is atomic number 1), we just put one arrow (representing an electron) inside of the circle.
The arrows that are placed within the circles represent electrons, which MUST have different spins (directions)! Let’s look at another example…
1s1

5. What about the orbital diagram for Nitrogen?
Electron Configuration: 1s2 2s2 2p3
1s s p3
Make sure that you always fill the lower energy level orbitals first. They must be full before you can move to the next energy level.

6. Orbital Diagrams Using our Nitrogen example: 1s2 2s2 2p3
We can write orbital diagrams as a “step” progression as well. Instead of circles, draw lines representing the orbitals in the energy levels and fill in the electrons just as you would using circles.
Using our Nitrogen example: 1s2 2s2 2p3
2p ______ ______ ______
2s ______
1s ______

7. Practice: Orbital Diagrams
Draw the orbital diagrams for: Carbon Sodium Phosphorus Argon

8. Valence Electrons

9. Periods: Energy Levels (n)
Each row in the periodic table is called a “period”
The period corresponds to a specific energy level of the atom
The top row, Period 1, is closest to the nucleus, the next one down is Period 2, etc…until you end with Period 7.
Level 1: s
Levels 2 and 3: s,p
Levels 4 and 5: s,p,d
Levels 6 and 7: s,p,d,f

10. Groups: Valence Electrons
Each column in the Periodic Table is called a “group”
Each element in a group has the same number of electrons in their outer energy level (the valence level).
The electrons in the outer shell are called “Valence Electrons”
Red: Group Orange: Group Yellow: Group 13
Green: Group Sky Blue: Group Baby Blue: Group 16
Dark Blue: Group Purple: Group 18

11. Valence Electrons
Valence electrons are the electrons in the highest occupied energy level of the atom.
Valence electrons are the only electrons generally involved in bond formation (which we will talk about in the next unit!)

12. Bohr Atomic Structures
Electron Configuration of Na: 1s22s22p63s1
The first energy level contains 2 electrons. (s orbital…1s2)
The second level containw 8 electrons. (2s and 2p orbitals…2s22p6)
How many electrons do you see in the outermost level? 3s1… 1electron!
This is the Valence number. Sodium has 1 Valence electron.

13. Electron Dot Structure: Lewis Dot Diagrams
A notation showing the valence electrons surrounding the atomic symbol. How many valence electrons in Cl? C?

14. Lewis Dot Structures Find out which group (column) your element is in.
This will tell you the number of valence electrons your element has.
You will only draw the valence electrons.

15. C Lewis Structures 1)Write the element symbol.
2) Carbon is in the 4th group, so it has 4 valence electrons.
3) Starting at either the right or left of the element symbol, draw 4 electrons, (dots), around the element symbol. C

17. For those of you who will be gone on the band field trip, you will find some practice on the next slide. Write the following Lewis Dot structures on a separate sheet of paper and turn this in when you come back to class.

19. The History of the Periodic Table and Trends

20. In the olden days…
Many elements were known in the ancient world- copper, gold, silver, lead, etc.
For several hundred years, elements were discovered by alchemists
Alchemy was the ultimate search for wisdom and immortality.

21. By 1860, more than 60 elements had been discovered….HOWEVER,
There was no consistent organization of the elements.
No one was using the same method to determine mass of atoms, or the ratios of atoms in compounds
In 1860, Stanislao Cannizzaro of Italy presented a convincing method to measure the mass of atoms, thus creating standard values for atomic mass.
Now that there are some common standards…

22. The Matter of Mendeleev
In 1869, Dmitri Mendeleev began to try to arrange the elements. Inspired by solitaire, he started to find patterns in the properties of elements, and arranged the known elements by atomic mass in a Periodic (repeating) Table.

23. Mendeleev’s Genius
Mendeleev recognized there were undiscovered elements.
By using his periodic table, he could predict the chemical properties of the undiscovered elements.
Years later, Scandium, Gallium, Technetium and Germanium were discovered and were characteristic of the properties Mendeleev predicted! *swoon*

24. Henry Moseley
1911: Henry Moseley (working under the direction of Rutherford) rearranged the Periodic Table to go horizontally, and put the elements in order by atomic number.

25. Variations on the Periodic Table
The Mayan Periodic Table of Elements by Mitch Fincher
A Spiral Periodic Table by Prof. Thoedor Benfey

26. The future of the periodic table?

27. Periodic Trends
Now that the periodic table is organized, what patterns can we find?
What does it even mean to be “periodic”?

28. Thanks, to Moseley, we learn that patterns arise because of PROTONS!!!
This led to the development of the Periodic Law: the physical & chemical properties of the elements are periodic functions of their atomic numbers.
** In other words, when the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

29. Pattern: Families
Elements in column share similar traits, and are called families:
These columns are also called groups.

30. The Alkali Metals 1 valence electron Highly reactive with water
Form ionic compounds
Do not occur in nature as pure elements (always in compounds)

31. Alkali-Earth Metals Have 2 valence electrons
Reactive, but less reactive than alkali metals
Are ductile, malleable and have a silvery luster

32. Transition metals… and inner transition metals
Are less reactive than groups 1 and 2.
Tend not to react in water.
Are malleable and ductile, but still harder than group 1 & 2.
Tend to be solids at room temperature.
Have variable chemical properties
Are good conductors of electricity and heat.
**Inner transition metals tend to be radioactive

33. Nonmetals Poor conductors of heat and electricity
Often are found as gases or liquids, sometimes solids.

34. Halogens Are nonmetals
highly reactive with metals- most reactive is fluorine, lease reactive is astatine
Mostly exist as gases or liquids (except At -solid)
Have 7 valence electrons

35. Noble gases At room temperature, exist as gases.
Are completely unreactive
Have full s and p orbitals
Are odorless, colorless, nonflammable

36. Metalloids Tend to be solids
Have properties similar to both metals and nonmetals
Tend to be semiconductors (which means they are useful for technological uses)

37. Patterns and trends tend to follow the Representative Elements.

38. Valence Electrons:
The outermost s & p electrons

39. Ions Ionization energy
Charged atoms that become charged by losing or gaining electrons
Ionization energy
Energy necessary to make an ion by removing an electron from a neutral atom

40. Rule #1 to remember!
When an element loses an electron, we can think of it as being given away, which is a good thing or POSITIVE thing to do.

41. Group 1 Elements H, Li, Na, K, Rb, Cs, Fr
Achieve a stable octet (full outer shell) by losing 1 electron, which forms a +1 ion
H+ , Li+ , Na+ , etc…

42. Group 2 Elements Be, Mg, Ca, Sr, Ba, Ra
Achieve a stable octet by losing 2 electrons, which forms a +2 ion
Be2+, Mg2+, Ca2+, etc…

43. Rule #2 to remember!
When an element gains an electron, we can think of it as it is being stolen from another ion, which is a bad or NEGATIVE thing to do.

44. Group 7 Elements F, Cl, Br, I, At
Achieve a stable octet by stealing (gaining) 1 electron, which forms a -1 ion
F-, Cl-, Br- , etc…

45. Group 6 Elements O, S, Se, Te, Po
Achieve a stable octet by stealing (gaining) 2 electrons, which forms a 2- ion
O2- , S2- , Se2-, etc…