Chapters 7-9: Chemical Composition

I. “Intra-Chemical” Forces

A. Intra = within

B. Atoms (elements) held together by an __________________________

C. Three Types of Bonding

1) Metallic:

a) atoms of a metal _____________________ the valence electrons because they

________________________________

2) Ionic:

a) valence electrons are ______________________ between two elements.

b) strongest

3) Covalent:

a) Definition: valence electrons are ______________________ between two elements.

b) Weaker than ionic

II. Ionic Bonding

A. Definition: electrons are _________________________from one element to another

(opposite charges is attractive force)

B. Ionic compounds commonly referred to as: _____________________

C. Atom that donates electron: (+) ____________________

D. Atom that accepts electron: (–) _____________________

E. Oxidation State: refers to the ____________________ of an atom

F. Lewis Dot Formulas are used to show how outer shell electron transfer

1) Octet Rule: every element wants ________________ in its outer shell

a) potassium + chlorine potassium chloride

b) magnesium + fluorine magnesium fluoride

G. Types of Ions

1) Monatomic Cation:

2) Monatomic Anion:

3) Polyatomic Ions:

H. Writing Ionic Chemical Formulas

1)

2) ____________________of elements

3) Writing chemical formulas (from the name)

a) Recognize the (+) and (–) ions

b) Write the symbols of the elements with their charge

c) A Roman numeral will tell you what the charge is on the cation if there is more than one

possibility

d) Adjust the number of each ion (with subscripts) as needed so the positive charge is equal

and opposite the negative charge.

e) If the ions are polyatomic and there is more than one, the ion is enclosed with parentheses

with a subscript on the outside.

Examples

1. sodium chloride _______________________

2. calcium sulfide _______________________

3. calcium sulfate _______________________

4. barium phosphate _______________________

f) Naming Ionic Compounds

a. Consists of __________________words:

b. Name the cation

c. Name the anion

d. If the cation has more than one possible charge, a Roman numeral is used to

show the charge.

e. All transition metals need Roman numerals except:

i. _____________ that always has a charge of __________ and

ii. _____________ that always has a charge of __________

Examples

1. FeCl3 Fe3+ iron (III) chloride

2. FeCl2 Fe2+ iron (II) chloride

3. NH4Cl _______________________________

4. Cu2SO4 _______________________________

5. NaC2H3O2 _______________________________

6. Ca(NO3)2 _______________________________

7. Zn(ClO)2 _______________________________

8. Cu2O _______________________________

9. CuO _______________________________

III. Covalent Bonding A. Definition: valence electrons are ______________________ between two elements

1) F2

2) Types:

a) Polar Covalent (______________________): _______________________ sharing of

electrons one pulls more than the other – the more electronegative element)

b) Non-Polar Covalent (_____________________): ____________________ sharing of

electrons

c) Writing Formulas for Covalent Compounds

Examples

1. carbon dioxide _______________________

2. carbon monoxide _______________________

3. dinitrogen monoxide _______________________

4. carbon tetrafluoride _______________________

5. triphosphorus pentachloride _____________________

d) Naming Formulas for Covalent Compounds

a. _________________________covalent compounds (2 elements)

b. Formulas with two ________________________

c. Rules

i. First word:

1. prefix indicating the number of atoms for the first element (if there is

more than one)

2. name of first element

ii. Second word:

1. prefix for the number of atoms of the second element

2. name of second element

3. prefixes on page 228

Examples

1. NO _______________________________

2. NO2 _______________________________

3. CBr4 _______________________________

4. P4O10 _______________________________

5. BF3 _______________________________

6. SiI5 _______________________________

7. H2O _______________________________

8. S6Cl8 _______________________________

9. Se7O9 _______________________________

e) Lewis Structures

a. The number of covalent bonds formed by an atom equals the number of

___________________electrons in the Lewis Dot Formula.

b. Examples

i. water (H2O)

ii. hydrogen gas (H2)

iii. hydrochloric acid (HCl)

iv. ammonia (NH3)

v. methane (CH4)

c. Multiple Bonds

i. Double bonds: _________________pairs of electrons are shared

O2

ii. Triple Bonds: __________________pairs of electrons are shared

N2

f) Hybridization: combining of __________________________ orbitals of nearly the same

energy into new orbitals of equal energy (occurs most often in groups IIA, IIIA, IVA)

a. Beryllium: [He]2s2 sp hybrid

b. Boron: [He]2s22p1 sp2 hybrid

c. Carbon: [He]2s22p2 sp3 hybrid

g) Molecules with more than one element (polar vs. non-polar)

a. Depends On:

i. Electronegativity difference (2 elements)

ii. Non-bonded electron pairs (2+ elements)

iii. Structure (symmetry) (2+ elements)

h) Shapes of Molecules (VSEPR handout)

IV. “Inter-Chemical” Forces

A. Inter = between

B. Definition: whole ____________or_________________attract and “bond” with one another.

C. Four Types

1) Ion-Dipole (strongest): occurs between an ion and a polar molecule

2) Hydrogen Bonding (medium strength):

3) Dipole-Dipole (weaker than hydrogen bonding): occurs between

____________________molecules

4) ________________________________Forces (weakest): the intermolecular attractions

resulting from the constant motion of electrons and the creation of instantaneous dipoles.