Chapters 4 & 5
Chemical Bonding
Valence Electrons
• Outermost electrons
• s and p electrons for main group elements
• Responsible for chemical properties of
atoms
• Participate in chemical reactions
Core Electrons
Valence Electron
Problems
•
1)
2)
3)
4)
Write out the electron configurations for
the following elements and identify how
many core and valence electrons each has.
Mg
S
Br
Kr
Lewis Dot Structures
• LDS: a representation of an atom using its
chemical symbol surrounded by dots that
signify valence electrons
Problems
• Write the Lewis Dot Structures for the following
atoms
• Li
• Be
• Br
• C
• N
• Ne
Li: [He]1s1
Na: [Ne]2s1
K: [Ar]3s1
Octet Rule
• Octet Rule: the tendency for atoms to seek
8 electrons in their outer shells
– Natural electron configuration of the Noble
Gases
– Done by gaining, losing, or sharing electrons
– Increases stability
– H and He seek a “Duet”
Ionic Bonding
• Ions: atoms that have a charge due to gain or loss of
electrons
– Anion: (-) charged atom
– Cation: (+) charged atom
• Ionic Bond: a bond formed through the transfer of one
or more electrons from one atom or group of atoms to
another atom or group of atoms
Formula Unit
Ionic Compounds: compounds composed of oppositely
charged ions that are held together by their attraction to each
other
• Metal + Non-metal
– NaCl
• Metal + Polyatomic Ion
– NaNO3
• Polyatomic Ion + Non-metal
– NH4Cl
• Polyatomic Ion + Polyatomic Ion
– NH4NO3
• Net charge on compound equal to zero
Oxyanions
SO42-
Sulfate
SO32-
Sulfite
PO43-
Phosphate
PO33-
Phosphite
-
Nitrate
NO2-
Nitrite
NO3
ClO4-
Perchlorate
ClO3-
Chlorate
ClO2-
Chlorite
ClO-
Hypochlorite
Rules For Naming Ionic Compounds
1) Name the cation by its elemental/polyatomic
name
2) If the metal is a transition metal with a variable
charge, indicate its charge with a Roman
Numeral in parentheses
3) Next, name the anion and change its ending to
“-ide”
4) If the anion is polyatomic, do not change the
ending to “-ide”
5) Do NOT use prefixes (mono, di, tri etc.) to
indicate how many of each atom are present
Problems
Write the name for the following compounds:
1) KI
2) MgBr2
3) Al2O3
4) FeCl2
5) CaSO4
6) Ba(NO2)2
7) Cu(NO3)2
Write the Formula for the following ionic compounds:
8) Sodium Fluoride
9)
Calcium Sulfite
10) Calcium Chloride
11) Iron (III) Oxide
12) Cobalt (II) Hydroxide
13) Ammonium Bromide
14) Ammonium Carbonate
15) Aluminum Carbonate
Iron (II) Chloride
Iron (III) Chloride
Covalent Compounds
• Covalent Compounds: compounds composed of atoms
bonded to each other through the sharing of electrons
• Electrons NOT transferred
• No + or – charges on atoms
• Non-metal + Non-metal
• Also called “molecules”
• Examples:
– H2O
– CO2
– Cl2
– CH4
or
Duet
or
H-H
Naming Covalent Compounds
1) Name the first non-metal by its elemental name
2) Add a prefix to indicate how many
1
2
3
4
5
6
7
8
9
10
3) Name the 2nd non-metal and change its ending to “-ide”
4) Add a prefix to indicate how many
Problems
Write the name of the following compounds:
1) CO
2) NI3
3) N2O
4) SF6
5) B2O3
Write the formula for the following compounds:
6) Phosphorous Pentachloride
7) Nitrogen Monoxide
8) Dinitrogen Tetroxide
9) Tetraphosphorous Decoxide
Problems
1) KCl
2) Na2S
3) H2O
4) SO2
5) K3PO4
6) FeCl3
7) (NH4)2SO4
8) SCl2
9) Cu(OH)2
10) P2O5
8) Sodium Iodide
9) Aluminum Sulfate
10) Phosphorous Pentabromide
11) Magnesium Nitride
Naming Acids
•
Acids that do not contain oxygen
1) Begin the name with “hydro”
2) Name the anion, but change the ending to “-ic”
3) Add “acid” on the end
•
•
HCl
HF
•
Acids that contain oxygen
1) Do not put “hydro” at the beginning
2) Begin the name with the anion
3) If the anion has the ending “-ate,” change this
to “-ic acid”
4) If the anion has the ending “-ite,” change this
to “-ous acid”
•
•
•
•
HClO4
HClO3
HClO2
HClO
Problems
•
Name the following
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
HBr(g)
HBr(aq)
HNO2(aq)
HNO3(aq)
HI (aq)
HI (g)
H2CO3 (aq)
H3PO4 (aq)
H3PO3 (aq)
HCN (aq)
Molecular Structures
Ball & Stick Models
Water
Methane
Space-Filling
Models
Ethanol
Lewis Dot Structures
1) Count the total number of valence electrons in the
molecule. Ex: PCl3
2) Use atomic symbols to draw a proposed structure with
shared pairs of electrons.
• Atoms don’t tend to bond to other atoms of the same
element when they can avoid it
• Exception: Carbon
3) Place lone pair electrons around each (except H) to
satisfy the octet rule, beginning with the terminal atoms
4) Place any leftover electrons on the central atom
5) If the number of electrons around the central atom is less
than 8, change single bonds to the central atom to
multiple bonds (double or triple).
•
Ex: CH2O
Problems
Draw the LDS’s for the following molecules:
1)Cl2O
2)C2H4
3)C2H6O
What Things Like To Do
• Halogens
– Like to be terminal
– Like to have one single bond and 3 lone pairs
(non-bonding electrons)
• Carbon
– Likes to have 4 single bonds and no lone pairs
• A double bond counts as two singles
• A triple bond counts as three singles
– Likes to be central
– Likes to bond to other carbons
• Silicon
– Likes to do what carbon does
• Oxygen
– Likes to have two single bonds and 2 lone pairs
• Sulfur
– Likes to do what oxygen does
• Nitrogen
– Likes to have 3 single bonds and one lone pair
• Phosphorous
– Likes to do what nitrogen does
• Hydrogen
– Likes to be terminal with only one single bond
– No lone pairs!
Problems
1)
SH2
2)
C3H8
3)
Si2H6
4)
PI3
5)
CH3OH
6)
C2H2
7)
CCl2O
8)
N2H4
9)
CH2OS
10) C2H6O
11) CO
12) BrHO
Electronegativity
• The measure of the ability of an atom to
attract electrons to itself
–
–
–
–
Increases across period (left to right) and
Decreases down group (top to bottom)
fluorine is the most electronegative element
francium is the least electronegative element
Electronegativity Scale
Types of Bonding
1) Non-Polar Covalent Bond:
•
•
•
Difference in electronegativity
values of atoms is 0.0 – 0.4
Electrons in molecule are
equally shared
Examples: Cl2, H2, CH4
ENCl = 3.0
3.0 - 3.0 = 0
Pure Covalent
2) Polar Covalent Bond:
•
•
Difference in
electronegativity values of
atoms is 0.4 – 2.0
Electrons in the molecule
are not equally shared
•
•
•
The atom with the higher
EN value pulls the electron
cloud towards itself
Partial charges
Examples: HCl, ClF, NO
ENCl = 3.0
ENH = 2.1
3.0 – 2.1 = 0.9
Polar Covalent
3) Ionic Bond:
•
•
•
Difference in EN
above 2.0
Complete transfer of
electron(s)
Whole charges
ENCl = 3.0
ENNa = 1.0
3.0 – 0.9 = 2.1
Ionic
Problems
•
1)
2)
3)
4)
Predict the type of bonding in the following
compounds using differences in EN values of
the atoms. Indicate the direction of the dipole
moment if applicable
KBr
HF
BrI
FI
Valence Shell Electron Pair
Repulsion Theory
• VSEPR theory:
– Electrons repel each other
– Electrons arrange in a
molecule themselves so as
to be as far apart as
possible
• Minimize repulsion
• Determines molecular
geometry
Defining Molecular Shape
• Electron pair geometry: the geometrical
arrangement of electron groups around a
central atom
– Look at all bonding and non-bonding e-’s
• Molecular Geometry: the geometrical
arrangement of atoms around a central atom
– Ignore lone pair electrons
• 2 e- groups surrounding
the central atom
– e- pair geometry: linear
– MG: linear
– AXE designation: AX2E0
• A: Central Atom
• X: Bonding pairs
• E: Non-bonding pairs
– Example: BeCl2
3
e groups
• 3 Bonds, 0 Lone Pairs
– e- PG: Trigonal Planar
(Triangular planar)
– MG: Trigonal Planar
– AX3E0
– BF3
• 2 Bonds, 1 Lone Pair
– e- PG: Trigonal Planar
(Triangular planar)
– MG: Bent/angular
– AX2E1
– GeCl2
4 e- groups
• 4 bonds, 0 Lone Pairs
–
–
–
–
e- PG: Tetrahedral
MG: Tetrahedral
AX4E0
CH4
• 3 bonds, 1 Lone Pair
–
–
–
–
e- PG: Tetrahedral
MG: Triangular Pyramidal
AX3E1
NH3
• 2 bonds, 2 Lone Pairs
–
–
–
–
e- PG: Tetrahedral
MG: Bent/Angular
AX2E2
H2O
Drawing LDS With
Correct Geometry
Molecular Polarity
Problems
Draw the Lewis Dot Structures for the following molecules
and then identify the direction of polarity, if any.
1)
2)
3)
4)
5)
NF3
CH2O
CBr4
CHCl3
CH2Cl2