1. Bonding
Unit III

2. I. Bond N A. What is a bond? Attraction of an electron by two nuclei
B. What electrons are involved in bonding?
Valence electrons
Electrons in the outermost energy level
Represented by an electron-dot diagram
example N
5 valence electrons N

3. C. Why do atoms bond?
Atoms will form bonds to achieve 8 valence electrons
Octet Rule
“8 is GREAT”
Stable configuration
Noble gas configuration (Group 18)
Compound has lower potential energy than the un-bonded atoms

4. D. Energy Terms Ionization Energy Electronegativity
Amount of energy needed to remove an electron
Compare Na Cl Ar
Ionization Energy kJ kJ kJ
Why?
Electronegativity
Attraction for electrons
Compare Na Cl Ar
Electronegativity

5. E. Types of Bonds Atomic Bonds Molecular Bonds
Holds atoms together to make a molecule
Molecular Bonds
Holds molecules together to make a solid or liquid

6. II. Atomic Bonds Ionic Bond
Transfer of electrons from metals to nonmetals
Metals lose electrons
Become positive
Decreases in size
Nonmetals gains electrons
Become negative
Increases in size
Properties
High melting points
Hard
Poor conductor as solid
Good conductor in liquid or aqueous states
Nonmetals
Metals

7. Bond Diagram Write electron dot diagram for each element
Transfer electrons from metal to nonmetal until all single electrons are paired up

9. Sharing of electrons between two nonmetals Types Nonpolar Covalent
Covalent Bond
Sharing of electrons between two nonmetals
Types
Nonpolar Covalent
equal sharing of electrons
Atoms have the same electronegativity values
diatomic elements (H2 O2 F2 Br2 I2 N2 Cl2)
Polar Covalent
unequal sharing of electrons
Atoms have the different electronegativity values
higher electronegative element is the negative end of the molecule

10. Properties
Low melting points
Soft
Poor conductor
Bond Diagram
Write electron dot diagram for each element
Share electrons of any all single electrons
Each “loop” is a bond

11. example of a nonpolar covalent bonds: Diatomic molecules H O F Br I N Cl

12. example of a polar covalent bond:

13. Network Solid Giant covalent bonds
sharing of electrons occurs in the entire crystal
examples
C (diamond)
C (graphite),
SiO2 (silicon dioxide)
SiC (silicon carbide)
asbestos
Properties
Hardest solids
Highest melting points
Poor conductor

14. Metallic Bonding
Metals do not have enough valence electrons to stabilize the structure
Delocalize electrons
Allow them to move anywhere in the structure
Properties
Hard
High melting points
Good conductor in the solid state

15. III. Nomenclature Naming Compounds 1. Binary Compounds
Compounds that contain two elements
Name the first element
Name the second element using an “ide” ending
Check the oxidation number of the first element
If it has more than one positive value, use a Roman numeral to indicate the number that was used

16. CaCl2
Calcium chloride
Ca is +2 only
No Roman numeral needed

17. CCl4 Carbon chloride C is -4, +2 and +4
Since there are two positive choices, use a Roman numeral
C Cl4
Carbon IV chloride +4 -4 =0 +4 -1

18. B. Ternary Compounds
Compounds that contain more than two elements
Polyatomic ion is present
Found on Table E
Identify the polyatomic used
Name the first substance
Name the second substance
Check the oxidation number of the first element
If it has more than one positive value, use a Roman numeral to indicate the number that was used

19. K2SO4 Potassium sulfate Potassium is +1 only
No Roman numeral is needed

20. Cu(NO3) 2 Copper nitrate Copper can be +1 or +2 Copper II nitrate
Needs a Roman numeral
Cu (NO3)2 +2 -2 =0 +2 -1
Copper II nitrate

21. FeSO4 Iron sulfate Iron can be +2 or +3 Iron II sulfate
Needs a Roman numeral
Fe SO4 +2 -2 =0 +2 -2
Iron II sulfate

22. NH4Cl Ammonium Ammonium chloride Ammonium ion is +1
No choice means no Roman numeral needed

23. B. Writing Formulas
Use ending to tell if compound is binary or ternary
Binary usually ends in “ide”
[except hydroxide and cyanide]
Ternary usually ends in “ate” or “ite”
Write the symbols
Assign oxidation numbers
Positive atom on left, negative on right
Reduce if possible and “criss-cross”
These numbers become the subscripts of the formula

24. Strontium phosphide “ide” means binary Strontium (Sr)
Phosphide is phosphorus (P)
Sr+2 P-3
Sr3P2

25. Iron II Oxide “ide” means binary Iron (Fe) Oxide is oxygen (O)
reduce
FeO

26. Manganese IV Carbonate
“ate” means ternary
Manganese (Mn)
Carbonate is a polyatomic ion
(CO3-2)
Mn+4 CO3-2
Mn2(CO3 )4
reduce
Mn (CO3 )2

27. IV Intermolecular Bonds
Definition
Holds completed molecules together to form a liquid or solid
Strength is reflected by melting point or boiling point
Types
Ionic bonds, network solids and metallic bonds can be considered both atomic or intermolecular
Covalent bonds will join by
Hydrogen bonding
Dipole bonding
Van der Waals forces

28. Hydrogen Bonding Dipole Van der Waals forces
Attraction between small highly electronegative elements
H to F, O or N
Dipole
Attraction between asymmetrical molecules
Polar molecule
Greater polarity, stronger bond
(electronegativity difference)
Van der Waals forces
Attraction between symmetrical molecules
Nonpolar molecule
Larger the molecule, stronger the bond
“bigger is better”

29. Asymmetrical molecule
Nonpolar molecule
Asymmetrical molecule
Polar molecule
Dipole

30. H2O vs H2S vs H2Se vs H2Te Highest Boiling Point
H2O vs H2S vs H2Se vs H2Te Highest Boiling Point? Strongest bond H2O Why? Hydrogen bonded Most polar

31. F2 vs Cl2 vs Br2 vs I2 Smallest F2 and Cl2 Br2 Largest I2 Gases Liquid
Solid