1. Chemistry 103
Lecture 8

2. Outline Electron Dot Symbols Lewis Dot Diagrams Predicting Formulas
I. Electronic Structure (CH5)
Brief Review
Orbital Diagrams
Valence Electrons
Atomic & Ionic Radii
II. Chemical Bonds - Ionic Compounds (CH6)
Electron Dot Symbols
Lewis Dot Diagrams
Predicting Formulas

3. Electron Configuration
Write the full electron configuration and the nobel gas shorthand electron configuration for tin (Sn, element 50)

4. Basic Questions about the Model
Which of the following orbitals will not be found in our quantum mechanical model?
A. 1s
B. 3p
C. 2f
D. 4p
Justify your answer based on quantum number rules. When n=2, can only have l=0 (s) and l=1(p) orbitals.

5. Basic Questions about the Model
What shell will be the first in which a “4” orbital will be found?
n = 4 , l = 0, 1 ,2, 3
n = 4 (first time l can equal “3” (f))
How many individual orbitals are in the n=2 shell?
1 (s), l = 0, ml = (1 value, 1 orbital)
3 (p), l = 1, ml = -1, 0, (3 orbitals)
TOTAL = 4 orbitals in the n=2 shell

6. Orbital (Box) Diagrams
Visual Representations of Electron Configurations
Each line (or circle) represents an orbital
Arrows represent electrons
The rules of electron configurations are followed
1s s 2p s p s

7. Orbital Occupancy
The Maximum Number of Electrons any single orbital can hold is two.
They are distinguished from each other an arrow convention - one must be up, one must be down if electrons are in the same orbital.
He ____ 1s

8. Writing Orbital Diagrams
Electron configurations tells us which energy levels the electrons for each element are located.
THREE rules:
1. Electrons fill orbitals starting with lowest energy first
2. There can be no more than 2 electrons in any orbital, and those electrons must have different “spins”

9. Writing Orbital Diagrams
Electron configurations tells us which energy levels the electrons for each element are located.
THREE rules:
1. Electrons fill orbitals starting with lowest energy first
2. There can be no more than 2 electrons in any orbital, and those electrons must have different “spins”
3. For orbitals in the same subshell, electrons fill each orbital singly before any orbital gets a second electron

10. Writing Orbital Diagrams
Nitrogen
___ ___ ___ ___ ___ 2s 2p 1s

11. Writing Orbital Diagrams
Unpaired electrons
Nitrogen
___ ___ ___ ___ ___ 2s 2p 1s
Paired electrons

12. Learning Check Write the orbital diagrams for: A. oxygen B. phosphorus
C. calcium

13. Question!
Which electrons are most likely to be involved in chemical reactions?
A) those nearest to the nucleus
B) those farthest from the nucleus
C) all are equally likely to be involved

14. Atomic Families What differentiates one family from another? Li Na K
Let’s look at the electron configurations of the alkali metals to find out. Li Na K Rb Cs Fr

15. The Properties of Atoms Are Dependent on Their Valence Electrons
Elements that similar properties have the same number of valence electrons!
Valence Electrons:
For elements in the “s” and “p” block - valence electrons are the electrons in the outer-most electron shell (the shell with the highest n value)
Example: Mg: 1s22s22p63s2
Example: Cl: 1s22s22p63s23p5

16. How Many Valence Electrons?
How many valence electrons do the following elements have?
Na: 1
Al: 3
S: 6
Cl: 7

17. Atomic Radius
Atomic radius is the distance from the nucleus to the valence electrons.
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18. Atomic Radius Within a Group
Atomic radius increases going down each group of representative elements.
Li Atom
Na Atom
K Atom
INCREASE
Radius = distance between nucleus and valence electrons

19. Atomic Radius Across a Period
Atomic radius decreases from left to right across a period because more protons increase nuclear attraction for valence electrons.
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20. Learning Check Select the element in each pair with the larger atomic
radius.
A. Li or K
B. K or Br
C. P or Cl

21. EXAM I - Summer 09 Material

22. General Course structure
Atoms ---> Compounds ---> Chemical Reactions

23. Atomic Radius
Atomic radius is the distance from the nucleus to the valence electrons.
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24. Ionic Radii
An atom (or group of atoms) that is electrically charged as a result of the loss or gain of electrons.

25. Sizes of Metal Atoms and Ions
A positive ion
Lost its valence electrons
Is smaller (about half the size) than its corresponding metal atom
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26. Size of Sodium Ion The sodium ion Na+ is smaller than Na atom because
the valence electrons have been lost from the 3rd
energy level.
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27. Sizes of Nonmetal Atoms and Ions
A negative ion
Increased its number of valence electrons.
Is larger (about twice the size) than its corresponding metal atom.
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28. Size of Fluoride Ion The fluoride ion F- is larger than F atom because
the added valence electron increase repulsions
between the outer electrons.
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29. Chapter 6 - Names & Formulas of Compounds
Compounds - Compounds result from the formation of chemical bonds between two or more different elements.

30. Bonds
Chemical bond: attractive force holding two or more atoms together.

31. Bonds
Ionic Bonds - electron transfer process. Typically between a metal and a nonmetals
Covalent Bonds - electrons shared. Typically involving nonmetals.

32. Periodic Table of Elements

33. Electron Dot Symbols
Focus on valence electrons - highest “n” quantum number for representative elements
Electron Dot Symbols consists of symbol for the element with one dot for each valence e-.
Na .

34. Rules for Electron Dot Symbols
1. Representative elements in the same group of the periodic table have the same number of valence electrons
2. The number of valence electrons is represented by “dots” around the symbol of the representative element.
3. The maximum number of valence electrons for any representative element is 8 - Octet Rule.

35. The Representative Elements
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36. Octet Rule - What is special about “8”?
An octet
Is 8 valence electrons
Is associated with the stability of the noble gases
He is stable with two valence electrons (duet).
valence electrons
He 1s
Ne 1s2 2s2 2p
Ar 1s2 2s2 2p6 3s2 3p
Kr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 8

37. Ionic and Covalent Bonds
Atoms acquire octets
To become more stable
By losing, gaining, or sharing valence electrons
By forming ionic bonds or covalent bonds
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38. Focus on Ionic Compounds
Formed between a positively charged metal ion and a negatively charged nonmetal ion

39. Metals Form Positive Ions
Octets by losing all of their valence electrons.
Positive ions have the electron configuration of the nearest noble gas.
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40. Metal Ions – Why Are They Positive?
Let’s look at Mg. Its electron configuration is 1s22s22p63s2.
If it loses two electrons to become Mg2+, it will have the electron configuration 1s22s22p6
8 Valence Electrons!
If it gains 6 electrons to become Mg6-, it will have the electron configuration 1s22s22p63s23p6
The question is which is easier? To lose TWO electrons or to gain SIX?
Losing two: Therefore, Mg tends to form Mg2+ ions

41. Nonmetals Form Negative Ions
Octets by gaining valence electrons to acquire eight.
Negative ions have the electron configuration of the nearest noble gas
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42. Nonmetal Ions – Why Are They Negative?
Let’s look at Cl. Its electron configuration is 1s22s22p63s23p5.
If it loses seven electrons to become Cl7+, it will have the electron configuration 1s22s22p6
8 Valence Electrons!
If it gains 1 electron to become Cl-, it will have the electron configuration 1s22s22p63s23p6
The question is which is easier? To lose SEVEN electrons or to gain ONE?
Gaining ONE: Therefore, Cl tends to form Cl- ions