Name ____________________________________________ Date _______________ Physical Science Period ________
Chapter 6: Chemical Bonding Guided Notes
I.
Why do Atoms Combine?
A. Chemical Formula
1. The chemical formula shows:
a. The ___________________________ in the compound.
b. The ratio of their ______________________________.
_________
H2O
____________
B. Chemical Bond
1. Strong attractive force between _____________ or ions in a molecule or compound.
2. Formed by:
a. _______________________ electrons, (e-) by losing or gaining electrons.
b. _______________________ electrons (e-)
C. Stability
1. Octet Rule
a. Most atoms form bonds in order to have ______ valence electrons (e-).
b. _______________ outer energy level.
c. Want to be like the Noble Gases (group 18).
2. Stability is the driving force behind bond formation. All atoms bond to become stable.
3. Transferring e-
Ne
4. Sharing e-
II.
Kinds of Chemical Bonds
A. Ionic Bond
1. Attraction between 2 __________________________ charged ions.
2. Ions are _____________________ atoms.
3. A _____________________ is a ______ charged ion. Forms when an atom loses an electron.
Formed by ___________________.
4. A _____________________ is a ______ charged ion. Forms when an atom gains an electron.
Formed by ___________________.
5. Formed by _______________________ electrons (e-) from a _____________ to a ______________.
6. Ions form a 3-D crystal lattice.
B. Covalent Bond
1. Attraction between neutral atoms.
a. Formed by ____________________________ electrons between two nonmetals.
b. Covalent bonds result in molecules.
Cl2
NH3
H2O
2. ________________________________ Covalent Bond
a. Electrons (e-) are _____________________________ equally.
b. Usually identical atoms.
3. ________________________________ Covalent Bond
a. Electrons (e-) are shared ______________________ between 2 different atoms.
b. Results in partial opposite charges.
C. Comparison Chart
Ionic
Covalent
Electrons
Melting Point
Soluble in Water
Conduct Electricity
Other Properties
Types of Elements
D. ____________________ types of bonds
 Some compounds have a mixture of ______________ and __________________ bonds.
 These generally contain POLYATOMIC IONS.
 _____________________ ____________ are groups of non-metal elements bonded.
_____________ together that have an overall ________________.
 Generally polyatomic ions will form an ionic bond with metal.
 They are easy to recognize because they have ______________ or more elements in the
compound.
 Examples: ______________________________, __________________________
E. Practice: What type of compound is shown or described below?
1. NaCl ________________________________________
2. CO2
________________________________________
3. H2O
________________________________________
4. Fe2O3 ________________________________________
5. Ga(C2H3O2)3
________________________________________
6. High melting point
________________________________________
7. Liquid or Gas ________________________________________
8. Doesn’t dissolve in water
________________________________________
III.
Naming Molecular (Covalent) Compounds
A. Molecular Names
1. ____________________________________________________________________
2. ____________________________________________________________________
3. ____________________________________________________________________
4. ____________________________________________________________________
Prefix
Examples:
CCl4
____________________________
Subscript
1
2
3
4
5
Prefix
Subscript
6
7
8
9
10
N2O
_____________________________
SF6
______________________________
B. Molecular Formulas
1. Write the _____________________ metallic element first.
2. Add subscripts according to prefixes.
Examples: phosphorus trichloride
dinitrogenpentoxide
___________________
___________________
dihydrogen monoxide
___________________
3. The seven Diatomic elements are:
Br2
I2
N2
Cl2
H2
O2
***In nature these elements are never alone!!***
IV.
F2
Naming Ionic Compounds
A. Oxidation Number
1. The charge on an ion.
2. Indicates the # of electrons (e-) gained/lost to become ________________________.
Oxidation Chart
Group #
Valence eOxidation #
1
1
2
2
13
3
14
4
15
5
16
6
17
7
18
8
B. Ionic Names
1. __________________________________________________________________________________
2. __________________________________________________________________________________
3. __________________________________________________________________________________
Examples:
NaBr
___________________________
Na2CO3
SnCl4
______________________________
_____________________________________
C. Ionic Formulas
1. Write each ion. Put the _______________ first.
2. Overall charge must equal zero.
a. If charges _______________, just write the symbol.
b. If not, crisscross the _______________ to find subscripts.
3. Use parentheses when more than one ______________________________ is needed.
4. Roman _______________ indicate the oxidation number (oxidation #).
D. Writing Ionic Formulas
1. Write each ion. Put the cation first.
2. Overall charge must equal zero.
3. If charges cancel, just write the symbols.
4. If not, crisscross the charges to find subscripts.
5. Use parentheses when more than one polyatomic ion is needed.
E. Writing Ionic Formulas – Crisscross method
1. IF the charges don’t balance the number of the charge becomes the subscript for the opposite
element or polyatomic ions.
2. Example 1: Mg +2 N-3
_______________________
The charges BALANCE thus disappear
Example 2: Al+3 SO4-2
_______________________
Remember to use parentheses when more than one polyatomic ion is needed.
Example 3: Na+1 C2H3O2-1 _______________________
Remember If charges cancel, just write the symbols
F.
Writing Ionic Formulas Examples:
potassium chloride
magnesium nitrate
_______________________ ______________________
aluminum oxide
__________________________