Transcript Chapter 6 Chemical Bonding

Chapter 6
Chemical Formulas
OBJECTIVES
• 1. Distinguish between ionic and
molecular compounds.
• 2. Define cation and anion and relate them
to metal and nonmetal.
Molecules and Molecular Compounds
1. Molecules
a. the smallest electrically
neutral unit of a substance
b. Made up of two or more
atoms
c. examples:
1. Oxygen in air: O2
2. water: H2O
d. diatomic molecule: two of the
same atom
: seven of them
hydrogen (H2)
oxygen (O2)
fluorine (F2)
bromine (Br2)
iodine (I2)
nitrogen (N2)
chlorine (Cl2)
Molecular compounds
1. made up of molecules
2. have low melting and
boiling points
3. Exist as gases or liquids at room
temperature
4. composed of two nonmetals
5. examples: CO (carbon monoxide)
CO2 (carbon dioxide)
Ions and Ionic Compounds
1. Ions
a. Atoms or group of atoms that are
either + charged or – charged
b. Cation: + ion
lose electrons
c. Anion : - ion, gain electrons
: change ending to -ide
d. Examples:
1. Sodium loses 1 electron
Na becomes Na+1
2. Chlorine gains 1 electron
Cl becomes Cl-1
chloride
2. Ionic compounds
a. Consist of:
cation (+), anion (-)
metal, nonmetal
b. Represented by a formula unit
c. Solids, high melting point
OBJECTIVES
• 3. Distinguish between chemical formulas,
molecular formulas, and formula units.
• 4. Use experimental data to show that a
compound obeys the law of definite
proportions.
Chemical formulas
1. shows symbols, subscripts
2. gives the number of atoms
3. example:
Ca3N2
Ca = 3
N=2
Al2(SO4)3
Al = 2
S = 1*3 = 3
O = 4*3 = 12
Molecular Compounds
1. Shows the type and number of atoms
in a molecule
2. Does not show shape or how the
atoms are arranged
3. NaCl , PH3
Models that show arrangement of atoms
1. Structural formula
H–N–H
H
2. Space-filling
CO2
3. Ball and stick molecular model
4. Perspective drawing
Formula Units
1. represents ionic compounds
2. lowest whole-number ratio of ions
3. example:
a. MgCl2
b. AlCl3
ratio is 1:2
ratio is 1:3
 Law of Definite Proportions
1. Any sample of a chemical compound,
the elements are always combined in
the same proportion by mass
ex. H2O consists:
2 hydrogen atoms, mass of 2 g
1 oxygen atom, mass of 16 g
ratio 1:8
 Law of multiple proportions
1. When two elements form more than
one compound, the different masses
of one element that combine with the
same mass of the other element are
in the ratio of small whole numbers
2. Examples: FeS
FeS2
OBJECTIVES
• 5 Use the periodic table to determine the
charge of an ion.
• 6. Define a polyatomic ion and give the
names and formulas of the most common
polyatomic ions.
Ionic Charges
A. Monatomic ions
1. ions of one atom
2. examples:
Na+ or Cl-
B. Transition metals (Groups 3-12)
1. Have more than 1 positive charge
2. Charge is shown with a roman
numeral
3. Examples: iron (II) and iron (III)
4. Old names
a. Iron (II): ferrous ion
b. Iron (III): ferric ion
c. lower roman numeral ends in –ous
d. Higher roman numeral ends in -ic
 Polyatomic ions
1. Tightly held group of atoms acting like
one ion
2. Examples:
a. SO4 -2
b. NH4+1
OBJECTIVES:
7. Apply the rules for naming and writing
formulas for binary ionic compounds
8. Apply the rules for naming and writing
formulas for ternary ionic compounds.
 Ionic Compounds
A. Binary compounds
1. has 2 elements
2. ALWAYS metal followed by a
nonmetal
B. Formula writing
1. Write the symbol and charge of
the first element
2. Write the symbol and charge of
the second element
3. Charges must cancel and = 0
4. Examples:
a. sodium fluoride
Na+1
F-1
NaF
+1-1=0
b. Calcium iodide
Ca+2 I-1
CaI2
+2 -1 ≠ 0
c. Aluminum sulfide
Al+3
S-2
+3 -2 ≠ 0
Al2S3
d. Copper(II) oxide
gives the charge of copper
Cu+2 O-2
+2-2 =0
CuO
C. Using polyatomic ions
1. tertinary – contains three different
elements
2. must use ( ) if there is more than one
3. examples:
a. Sodium hydroxide
b. Calcium phosphate
c. Ammonium sulfide
D. Naming binary
1. write the name of the first element
2. write the name of the second element
and change the ending to “ide”
3. Transition metals
**Must include roman numerals
4. Examples:
a. CaF2
Calcium fluoride
b. KBr
potassium bromide
c. CuCl2
copper (II) chloride
Objectives
9. Apply the rules for naming and writing
formulas for binary molecular compounds.
Writing covalent compound formulas
A. Rules
1. Write the symbol and subscript of
the first element.(look at prefix)
2. Write the symbol and subscript of
the second element.(look at prefix)
3. prefixes
**1-mono
(only for 2nd element)
2- di
3-tri
4-tetra
5-penta
6-hexa
7-hepta
8-octa
9-nona
10-deca
4. Examples:
a. tricarbon difluoride
C3F2
b. heptabromine tetrachloride
Br7Cl4
B. Naming
1. Write the name of the element using
the prefixes.
*** Do not use mono for the 1st
element
2. Write the name of the second
element using prefixes, change
ending to –ide.
3. Examples:
a. P3O5
b. SO4
c. N2O
OBJECTIVES
10. Name and write formulas for common
acids.
 Naming and writing acids
A. Naming binary hydrides
1. use the prefix hydro- and the
suffix –ic
2. HF
hydrofluoric acid
3. HBr
hydrobromic acid
B. Writing binary hydrides
1. hydrochloric acid
HCl
2. hydrosulfuric acid
H2S
Naming acid containing oxygen
1. use the suffixes –ous or -ic with the
negative ion
2. -ous (lower oxygen state)
3. -ic (higher oxygen state)
4. examples:
a. HNO3
b. HNO2
nitric acid
nitrous acid
c. H3PO4
phosphoric acid
d. H3PO3
phosphorous acid
** If there is only one ion use –ic.
Writing acids with oxygen
1. acetic acid
H+1
C2H3O2-1
HC2H3O2
2. oxalic acid
H+1
C2O4-2
H2C2O4
Extra oxygens
1. ClO4-1 perchlorate
2. ClO3-1 chlorate
3. ClO2-1 chlorite
4. ClO-1 hypochlorite
5. sulfurs
a. Persulfate
SO5-2
b. Hyposulfite
SO2-2
Homework problems
1. Name the following
acids
a. HC2H3O2
• 2. Write the formulas
• a. Hydrophosphoric acid
b. HI
• b. Hydrofluoric acid
c. HNO3
• c. pernitric acid
d. HBr
• d. Carbonic acid