Transcript Unit 3 Review Notes - Brinkmann chapter7_and_8_review1

• Draw Lewis Structures for the
following compounds. Give the
shape, polarity, and bond angle for
each compound.
• CH3OH
• NH3
• N2H2
Chapters 7 and 8 Review
Chemical Bonds
• electrical attraction between nuclei and
valence e- of neighboring atoms that
binds the atoms together
• bonds form in order to…
– decrease potential energy
– increase stability
• Three types:
– Ionic
– Covalent
– metallic
Ionic Bonds
 Electrons are transferred
 Electronegativity differences are
generally greater than 1.7
 The formation of ionic bonds is
always exothermic!
Sodium Chloride Crystal Lattice
– Ionic compounds form solids at ordinary
temperatures. High mp and bp.
– Good conductors of heat and electricity
– Ionic compounds
organize in a
characteristic crystal
lattice of alternating
positive and negative
ions.
– generally soluble as a
liquid
Covalent Bonds
True Molecules
–
–
–
–
Non-metals – share electrons
Liquids or gases at room temp (low mp and bp)
Poor conductors of heat and electricity
Low solubility
Diatomic
Molecule
Metallic Bonds
Metal elements sharing a “Electron Sea”
– good conductors of electricity
– malleable, ductile, lustrous
Valence electrons - electrons in the outer
energy level. These electrons determine
1 the formation of chemical bonds.
8
2
3 4 5 6 7
• Electronegativity
– a measure of an atom’s ability to
attract electrons.
– higher e- neg atom  – lower e- neg atom +
Electronegativity Trend
• Increases up and to the right.
Bond Polarity
• Most bonds are
a blend of ionic
and covalent
characteristics
• Nonpolar Covalent Bond
– e- are shared equally
– symmetrical e- density
– usually identical atoms
• Polar Covalent Bond
– e- are shared unequally
– asymmetrical e- density
– results in partial charges (dipole)
+


Covalent Compounds
• Molecules are neutral groups of
atoms that are held together by
covalent bonds.
• Diatomic molecules – H2, N2, O2,
F2, Cl2, Br2, and I2. Allotrophs
include P4 and S8.
Covalent or Molecular Compounds
• - Compounds between two nonmetals
• - Use prefixes
• - Only use mono on second element P2O5 = diphosphorus pentoxide
CO2 =
CO =
N2O =
carbon dioxide
carbon monoxide
dinitrogen monoxide
Octet Rule
• Remember…
– Most atoms form bonds in order
to have 8 valence electrons.
Drawing Lewis Diagrams
• Find total # of valence e-.
• Arrange atoms - singular atom is usually
in the middle.
• Form bonds between atoms (2 e-).
• Distribute remaining e- to give each atom
an octet (recall exceptions).
• If there aren’t enough e- to go around,
form double or triple bonds.
Octet Rule
• Exceptions:
F
F
– Hydrogen  2 valence e
F
B
F
– Groups 1,2,3 get 2,4,6 valence e
F
S
F
H
N
O
O
H
– Expanded octet  more than 8
F
valenceVery
e (e.g.
S,
P,
Xe)
unstable!!
F
F
-
-
-
– Radicals  odd # of valence e-
Drawing Lewis Diagrams
• CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
F
F C F
F
Drawing Lewis Diagrams
• CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e-
- 4e12e-
O C O
Polyatomic Ions
• To find total # of valence e-:
– Add 1e- for each negative
charge.
– Subtract 1e- for each positive
charge.
• Place brackets around the ion and
label the charge.
C. Polyatomic Ions
• NH4+
1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 8e0e-
H
H N H
H
Resonance Structures
• Molecules that can’t be correctly
represented by a single Lewis
diagram.
• Actual structure is an average of
all the possibilities.
• Show possible structures separated
by a double-headed arrow.
VSEPR Geometry
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 389
Formula
Bond Type
Name
Structure
Cl
CCl4
Covalent
Carbon tetrachloride
Cl
PbF2
NI3
Ionic
Lead(II) fluoride
Covalent
Nitrogen
tri-iodide
C
Cl
Cl
F- Pb+2 FI N
I
I
Ionic Bonding: Force of attraction
between oppositely charged ions.
Ions
• Cation: A positive ion
• Mg2+, NH4+
• Anion: A negative ion
• Cl-, SO42-
+1
+2
+3
-3 -2 -1
Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
(
Ba NO3 ) 2
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Not balanced!
Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
( NH4+) SO42-
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
2
Not balanced!
Writing Ionic Compound Formulas
Example: Iron(III) chloride
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Fe3+ Cl-
3
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum sulfide
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
3+
Al
2
2S
3
Not balanced!
Writing Ionic Compound Formulas
Example: Magnesium carbonate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
Mg2+ CO32They are balanced!
Writing Ionic Compound Formulas
Example: Zinc hydroxide
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
Zn
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
( OH- )2
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum phosphate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
3+
Al
PO4
3-
They ARE balanced!
Naming Ionic Compounds
• 1. Cation first, then anion
• 2. Monatomic cation = name of the
element
• Ca2+ = calcium ion
• 3. Monatomic anion = root + -ide
• Cl- = chloride
• CaCl2 = calcium chloride
Naming Ionic Compounds
(continued)
Metals with multiple oxidation states
some metal forms more than one cation
• - use Roman numeral in name
• -
• PbCl2
• Pb2+ is cation
• PbCl2 = lead(II) chloride
Calculating Formula Mass
Calculate the formula mass of magnesium carbonate,
MgCO3.
24.31 g + 12.01 g + 3(16.00 g) = 84.32 g
Calculating Percentage Composition
Calculate the percentage composition of magnesium
carbonate, MgCO3.
From previous slide:
24.31 g + 12.01 g + 3(16.00 g) = 84.32 g
 24.31 
Mg  
  100  28.83%
 84.32 
 12.01 
C 
  100  14.24%
 84.32 
 48.00 
O
  100  56.93%
 84.32 
100.00
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2
COMPOUND
2 elements
Binary
Compound
NaCl
more than 2
elements
Ternary
Compound
NaNO3
ION
1 atom
Monatomic
Ion
Na+
2 or more atoms
Polyatomic
Ion
NO3-
Formulas
Empirical formula: the lowest whole number
ratio of atoms in a compound.
Molecular formula: the true number of
atoms of each element in the formula of a
compound.
 molecular formula = (empirical
formula)n [n = integer]
 molecular formula = C6H6 = (CH)6
 empirical formula = CH
Formulas
(continued)
Formulas for ionic compounds are ALWAYS
empirical (lowest whole number ratio).
Examples:
NaCl
MgCl2
Al2(SO4)3
K2CO3
Formulas
(continued)
Formulas for molecular compounds MIGHT
be empirical (lowest whole number ratio).
Molecular:
H2O
C6H12O6
C12H22O11
Empirical:
H2O
CH2O
C12H22O11
Empirical Formula Determination
1. Base calculation on 100 grams of compound.
2. Determine moles of each element in 100
grams of compound.
3. Divide each value of moles by the smallest of
the values.
4. Multiply each number by an integer to obtain
all whole numbers.
Empirical Formula Determination
Adipic acid contains 49.32% C, 43.84% O, and
6.85% H by mass. What is the empirical formula
of adipic acid?
 49.32 g C 1 mol C   4.107 mol C
12.01 g C 
 6.85g H 1 mol H   6.78 mol H
1.01 g H 
 43.84 g O 1 mol O   2.74 mol O
16.00 g O 
Empirical Formula Determination
(part 2)
Divide each value of moles by the smallest of the
values.
4.107
mol
C
Carbon:
 1.50
2.74 mol O
6.78 mol H
Hydrogen:
 2.47
2.74 mol O
2.74 mol O
Oxygen:
 1.00
2.74 mol O
Empirical Formula Determination
(part 3)
Multiply each number by an integer to obtain all
whole numbers.
Carbon: 1.50
x 2
3
Hydrogen: 2.50
x 2
5
Oxygen: 1.00
x 2
2
Empirical formula: C3H5O2
Finding the Molecular Formula
The empirical formula for adipic acid is
C3H5O2. The molecular mass of adipic acid
is 146 g/mol. What is the molecular
formula of adipic acid?
1. Find the formula mass of C3H5O2
3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g
Finding the Molecular Formula
The empirical formula for adipic acid is
C3H5O2. The molecular mass of adipic acid
is 146 g/mol. What is the molecular
formula of adipic acid?
2. Divide the molecular mass by the
mass given by the emipirical formula.
3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g
146
2
73
Finding the Molecular Formula
The empirical formula for adipic acid is
C3H5O2. The molecular mass of adipic acid
is 146 g/mol. What is the molecular
formula of adipic acid?
3. Multiply the empirical formula by this
number to get the molecular formula.
3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g
146
2
73
(C3H5O2) x 2
=
C6H10O4