1. Ch. 7/8 Notes Day 1

2. Objectives
SWBAT explain how elements and atoms become ions when forming bonds
SWBAT to identify trends in the periodic table (atomic radius, ionization energy, electronegativity)
SWBAT explain why some bonds might form over others

3. Quick Review
Electron Configuration – how electrons are arranged in energy levels and orbitals around a nucleus
Energy Level – the location (“house”) that contains orbitals (“rooms”) where electrons “live.”
Also called SHELLS
The BIGGER the energy level, the BIGGER the house

4. Think of it like this…. High Energy Low Energy 1st Energy Level
2nd Energy Level
3rd Energy Level
1st Energy Level
Orbitals
Orbital
Orbitals

5. Electron Configuration for Energy Levels
Energy Level (Shell)
# of Electrons 1 2 8 3 18 4 32 5 50 6 72
Away from Nucleus

6. How to construct electron configurations
Rules
Start on the inside and work your way out
COMPLETELY fill each shell as you work your way out
Whatever is left over in the outer-most shell, is your number of VALENCE electrons

7. Electron-Dot Diagrams
Write the chemical symbol for each element, then place the number of valence electrons around the symbol on 4 sides.
Keep placing one dot on each side, until you have the number of valence electrons
NOTE: you should not have more than 2 per side, for a total of 8.

8. New Stuff

9. Ions
Ion – an atom that has a NET positive or NET negative charge
Anion – an atom that has a NET negative charge
Nonmetals will always become ANIONS
Cation – an atom that has a NET positive charge
Metals will always become CATIONS
**GAINS one electron to fill the octet. Now has 10 e- total.
Protons +9
Electrons -9
Charge F
Protons +9
Electrons
-10
Charge -1 F
**LOSES two electron to empty the octet. Now has 18 e- total.
Protons
+20
Electrons
-20
Charge
Protons
+20
Electrons
-18
Charge +2 Ca Ca

10. Practice… What’s the charge? Anion or Cation?

11. Ch. 7/8 Notes Day 2

12. Objectives
Students will be able to use electronegativity relates to electron affinity, and how it relates to the strength of an ionic bond.
Students will be able to explain and demonstrate naming and formula writing conventions for ionic bonds given elements, or formula names

13. Review

14. Atomic Radius
Atomic Radius – the distance from the center of an atom (nucleus) to the outermost electron orbital
Atomic Radius DECREASES
Atomic Radius INCREASES

15. Ionization Energy
Ionization Energy – how easy it is to pull an electron out of an atom
The LARGER the atomic radius, the lower the ionization energy. It’s easier to pull an electron out of a large atom than a small atom.
Ionization Energy INCREASES
Ionization Energy DENCREASES

16. Electron Affinity
Electron Affinity – how much an atom will attract additional electrons
How “hungry” an atom is for more electrons.
The opposite of ionization energy
The smaller the atomic radius, the higher the electronegativity.
Electron Affinity INCREASES
Electron Affinity DECREASES

17. New Stuff

18. Electronegativity
Trick: Move right to left first, then check up and down
Electronegativity – how much an atom will attract additional electrons
How “hungry” an atom is for more electrons.
The opposite of ionization energy
The smaller the atomic radius, the higher the electronegativity.
Pauling Scale
Rates electronegativity on a scale from 0.7 to 4.0.
The bigger the difference in electronegativity, the stronger the bond
Electronegativity INCREASES
Electronegativity DENCREASES

19. Ionic Bond Review
Ionic Bond – The force of attraction that holds cations (+) and anions (-) together
This force is caused by the TRANSFER of electrons from a cation to an anion. Na F

20. Dot Bonding Review Na Cl Li O Li

21. Writing Formulas Rules
Cation (+) ALWAYS comes first
Usually a METAL
Anion (-) ALWAYS comes second
Usually a NonMetal
If there is more than 1 of an atom in the formula, write it as a subscript Li O Li
Li2O
AlCl3

22. Ch. 7/8 Notes Day 3

23. Objectives
Students will be able to check their Lewis Dot models of ionic bonding with a shortcut method of checking atomic charge and bond ratios

24. Ionic Formula Shortcut
You MUST KNOW how to construct dot models. The shortcut is just a quick check…
You DO NOT need to write a 1 in the formula.
If the numbers are the same, they reduce to 1’s, and don’t write them.
Mg +2
O -2
B +3
S -2
Mg +2
I -1
Mg2O2
B2S3
Mg1I2
MgO
B2S3
MgI2

25. Naming Ionic Bond Rules
The name of the CATION always comes FIRST
Use the FULL NAME of the cation
The name of the ANION always comes LAST
Use the ROOT NAME of the anion, then add –IDE
Anion Name Examples
Fluorine -> Fluoride
Iodine -> Iodide
Sulfur -> Sulfide
Nitrogen -> Nitride
Phosphorous -> Phosphide

26. Li3P NaCl BaS Fr2O Lithium Phosphide Sodium Chloride Barium Sulfide
Ionic Naming Practice
Li3P
NaCl
BaS
Fr2O
Lithium Phosphide
Sodium Chloride
Barium Sulfide
Francium Oxide

27. Ionic Naming & Bonding Practice
Name of the Compound
Write the Formula for the Compound
NaBr
CaO
Li2S
MgBr2
BeO2
NaCl
Cs2Se
Potassium Iodide
Magnesium Oxide
Aluminum Chloride
Beryllium Phosphide
Calcium Chloride
Aluminum Oxide
Ammonium Chloride

28. Atomic Dating Activity Review
(Disregard. We are not doing this, this year.)

29. Dating Data Symbol Electroneg. Radius (pm) H 2.20 53 B 2.04 87 Li 0.98
167 C
2.55 67 Be
1.57
112 N
3.04 56 Na
0.93
190 O
3.44 48 Mg
1.31
145 F
3.98 42 K
0.82
243 Si
1.00
111 Ca
194 P
2.19 98 Rb
265 S
2.58 88 Sr
0.95
219 Cl
3.18 79 Cs
0.79
298 As
2.18
114 Ba
0.80
253 Se
103 Fr
0.7
270 Br
2.06 94 Ra
278 I
2.68
115 Te
2.10
123 At
127

30. Ch. 7/8 Notes Day 4

31. Objectives
SWBAT explain the difference between ionic bonds and covalent bonds
SWBAT to correctly name and draw Dot- Bonds for the compounds created

32. Polyatomic Ions
Polyatomic Ions- naturally occurring combinations of elements that are “stuck” so tightly together that they act as a single atom
Naming Rules:
(-ide)- element off of the periodic table
(-ite)- polyatomic ion
(-ate)- polyatomic ion
Examples
BaS – Barium Sulfide
BaSO3 – Barium Sulfite
BaSO4 – Barium Sulfate

33. Variable Charge Elements
Elements that can have more than 1 charge
Typically Transition Metals (Groups 3-12)
The charge is given by Roman Numerals
Examples
Lead (II) = Pb+2
Lead (IV) = Pb+4
Copper (I) = Cu+1
Copper (III) = Cu+3

34. Ch. 7/8 Notes Day 5

35. Objectives
SWBAT explain the difference between ionic bonds and covalent bonds
SWBAT to correctly name and draw Dot- Bonds for the compounds created

36. Covalent Bonds
Covalent Bond – a chemical bond where two or more atoms SHARE a pair of valence electrons
Molecule – a neutral group (no charge) of atoms that are joined together by covalent bonds
Lewis Structure – a model of boding between elements, usually represented by dots.
A “Dot-Bond”
Prefix – a Latin root that comes BEFORE the name of an atom that tells how many atoms are there

37. Special Covalent Bonds
Diatomic Molecules – seven NONMETALS that ALWAYS stick to themselves
Covalent naming rules DO NOT apply. Just name from periodic table.
N2 - Nitrogen
H2 - Hydrogen
F2 - Fluorine
O2 - Oxygen
I2 - Iodine
Cl2 - Chlorine
Br2 - Bromine
Never
Have
Fear Of
Ice
Cold
Beer

38. F Cl Cl S Cl Covalent Bonds Group 16 Elements
Usually form 2 Bonds, can form up to 6 Bonds
Group 17 Elements
Usually form 1 Bond, can form up to 7 Bonds F Cl Cl S Cl
TWO electron pairs So
TWO BONDS
ONE electron pair So
ONE BOND

39. F Cl F C F Cl N Cl F Covalent Bonds Group 15 Elements
Usually form 3 Bonds, can form up to 5 Bonds
Group 14 Elements
Usually form 4 Bonds F
THREE electron pairs So
THREE BONDS Cl F C F Cl N Cl F
FOUR electron pairs So
FOUR BONDS

40. Writing Covalent Compounds
Made of two anions (2 nonmetals or 1 nonmetal & 1 polyatomic ion)
Identify the less electronegative element 1st (trend = less EN is the one farthest left & farthest down on Periodic Table)
The 1st nonmetal is just given the name as found on the Periodic Table (just like ionic)
The 2nd nonmetal ending is changed to –ide (just like ionic…don’t change polyatomics)
Difference = numerical prefixes are used to express how many of each nonmetal are present

41. Writing Covalent Compounds (cont.)
You never use “criss-cross” method
The only time you do not use a prefix is when there is only one of the first nonmetal.
Number
Prefix 1
Mono 2 Di 3
Tri 4
Tetra 5
Penta 6
Hexa 7
Hepta 8
Octa 9
Nona 10
Deca

42. Covalent Compound Practice
Sulfur Dioxide Dinitrogen Monoxide Nitrogen Dioxide Carbon Tetrachloride Dichlorine Heptaoxide Phosphorus Trichloride Sulfur Hexaflouride Trisilicon Tetranitride Nitrogen Pentabromate Water  - can use common name
SO2 N2O NO2 CCl4 Cl2O7 PCl3 SF6 Si3N4 N(BrO3)5 H2O

43. Covalent Compound Practice
Phosphorus Trioxide Dinitrogen Pentacarbide Tellurium Noniodide Carbon Monoxide Selenium Heptaflouride Tetraphosphorous Decoxide Arsenic Hexabromide Silicon Dichloride
PO3 N2C5 TeI9 CO SeF7 P4O10 AsBr6 SiCl2

47. Ch. 7/8 Notes Day 6

48. Objectives
SWBAT explain the how transition metals have different properties than Alkali and Alkali Earth Metals
SWBAT to explain ionic bonding
SWBAT identify charge of transition metals from Roman Numerals

49. Naming Acids

50. Transition Metals (Groups 3-12)
Become Cations (+ charge)
Transition metals are very large
Have lots of extra, empty orbitals.
Metallic Bond – a bond between transition metals. Electrons are “given off” into a “sea” of electrons that surrounds the metal cations.
electron
metal cation