1. Write the noble gas notation for the following elements:
Germanium (Ge): #32
Niobium (Nb): #41
Osmium (Os): #76

2. Find the noble gas notation of the following elements:
Germanium (Ge): [Ar] 4s2 3d10 4p2
Niobium (Nb): [Kr] 5s2 4d3
Osmium (Os): [Xe] 6s2 4f14 5d6

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

4. 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…

5. 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

6. 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

7. What about the orbital diagram for Nitrogen?
Electron Configuration: 1s2 2s2 2p3
1s s p
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.

8. 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 ______

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

10. Valence Electrons

11. 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

12. 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

13. 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!)

14. Bohr Atomic Structures
Electron Configuration of Na: 1s22s22p63s1
The first energy level contains 2 electrons. (s orbital…1s2)
The second level contains 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.

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

16. 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.

17. N Lewis Structures 1)Write the element symbol.
2) Nitrogen is in the 5th group, so it has 5 valence electrons.
3) Starting at either the right or left of the element symbol, draw 5 electrons, (dots), around the element symbol. N

19. The History of the Periodic Table and Trends

20. In the old 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, 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. Alkaline-Earth Metals
Have 2 valence electrons
Reactive, but less reactive than alkali metals
Are ductile, malleable and have a silvery luster
Form alkaline (basic) solutions) when put in water

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.
Are good conductors of electricity and heat.
**Inner transition metals tend to be radioactive

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

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

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

37. 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)

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…

46. Periodic Trends

47. What we will investigate:
Atomic size
How big the atoms are
Ionization energy
How much energy to remove an electron
Ionic size
How big ions are
Electronegativity
The attraction for the electron in a compound

48. What we know about the atom:
All elements have a unique electron configuration, which determines what their atomic and orbital structures will look like
A positive nucleus is present and pulls on electrons
The more electrons present in the outermost energy level, the stronger the attraction is towards the nucleus of the atom

49. Keep in mind… + Group trends- as you go down a group:
Increase in energy levels (n)
Outermost electrons not as attracted by the nucleus because the outermost energy level is further away +

50. Atomic Size
The electron cloud doesn’t have a definite edge, so where do I start to measure?

51. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of molecule

52. Trends in Atomic Size Influenced by two factors: Energy Levels (n)
Charge on nucleus- More charge pulls electrons in closer

53. Group Trends- Atomic SizeH
Group Trends- Atomic Size Li Na
As we go down a group, each atom increases by an energy level of n+1
Increased shielding
Atoms become larger as you go down group K Rb

54. Periodic Trends- Atomic Size
As we go across a period, the radius gets smaller
Same shielding and energy level within a period
Increased nuclear charge pulls outermost electrons closer, creating a smaller, more dense atom Na Mg Al Si P S Cl Ar

56. For each of the following pairs which atom is larger?
Mg, Sr
Sr, Sn
Ge, Sn
Ge, Br,
Cr, W Sr Sn Ge W

57. Driving Force for Ionization Energy
Full Energy Levels are very low energy
Noble Gases have full orbitals (all outer shells end in p6, with the exception of Helium at s2)
Atoms behave in ways to achieve noble gas configuration (stable octet)

58. Ionization Energy
Energy necessary to make an ion by removing an electron from a neutral atom
Removing one electron makes a +1 ion
The energy required is called the first ionization energy

59. Ionization Energy
The second ionization energy is the energy required to remove a second electron, which is always greater than first IE
The 3rd IE is the energy required to remove a third electron, which is greater than 1st or 2nd IE

60. What determines ionization energy?
The greater the nuclear charge, the greater IE
Increased shielding (energy levels) decreases IE
Filled and half filled orbitals have lower energy, so achieving them is easier. This results in lower IE.

61. Group Trends- Ionization Energy
As you go down a group, IE decreases because of more shielding
Outer electron is less attracted to the nucleus, so less energy is required to remove that electron

62. Periodic trends- Ionization Energy
All the atoms in the same period have:
Same shielding (energy level)
Increasing nuclear charge
Ionization energy generally increases from left to right.
Exceptions: full and half full orbitals hardly require any energy to obtain electrons

63. Which of the following pairs has a higher ionization energy?
Mg, Na
S, O
Ca, Ba
Cl, I
Na, Al
Se, Br Mg O Ca Cl Al Br

64. Ionic Size Cations Metals form cations
are positive ions
form by losing electrons (giving electrons away)
Metals form cations
Cations of representative elements have noble gas configuration. (8 electrons in outermost shell)

65. Ionic size Anions are negative ions
form by gaining electrons (stealing electrons from other ions)
Nonmetals form anions
Anions of representative elements have noble gas configuration. (8 electrons in outermost shell)

66. Configuration of Ions
Ions of representative elements have noble gas configuration
Example: Na is 1s22s22p63s1
Forms a 1+ ion: 1s22s22p6 (Same configuration as Neon)
Metals form ions with the configuration of the noble gas before them on the periodic table- they lose (give away) electrons

67. Configuration of Ions
Non-metals form ions by gaining (stealing) electrons to achieve noble gas configuration.
Example: Fluorine is 1s22s22p5
Forms a 1- ion: 1s22s22p6(Same configuration as neon)
Non-metals end up with the configuration of the noble gas after them in the periodic table.

68. Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically bonded to another element.

69. Group Trend- Electronegativity
The further down a group:
More shielding (energy levels)
more electrons an atom has
Less attraction for electrons (Low electronegativity)

70. Periodic Trend- Electronegativity
Metals (Left side of Periodic Table)
Low nuclear charge
Low attraction for electrons
Low electronegativity
Non-metals (Right side of Periodic Table)
High nuclear charge
Large attraction for electrons
High electronegativity
Only exclusion is Group 8: Noble gases… they do not bond with other elements at all

71. Electronegativity Trend
In other words, electronegativity increases across a period and decreases down a group

72. Which of the following pair has a higher electronegativity?
Mg, Na
Na, Al
Cl, I
Ca, Ba
S, O
Se, Br Mg Al Cl Ca O Br