Transcript Slide 1

Chapter :
Chemical
Bonding
Cartoon courtesy of NearingZero.net
Chemical Bonds Lect1
• Forces that hold groups of atoms
together and make them function as a
unit.
• Ionic bonds – transfer of electrons
Metal + Nonmetal
Ex) NaCl Li2O
• Covalent bonds – sharing of electrons.
2 nonmetals
Ex) CH4 CO2
• Metallic bonds- electrons are free to
move throughout the material. Metals
Covalent Bonding
• Molecule- is the smallest unit quanitity of
matter which can exist by itself and retain
all the properties of the original
substance.
• When a compound is formed by sharing
electrons, the compound is called a
• ____Molecule__________________
• Examples H2O & O2
• Diatomic molecule- is a molecule containing
2 identical atoms. (H2 N2 O2 F2 Cl2 Br2 I2)
• H NO F
Covalent Bonding
• Diatomic molecule- is a molecule containing
2 identical atoms.
•
• H NO F
(H2 N2 O2 F2 Cl2 Br2 I2)
Naming Covalent Compounds
•
•
•
•
Two words, with prefixes
Prefixes tell you how many.
1-mono,
6- hexa
2-di,
7- hepta (greek)
septa (latin)
• 3-tri,
8-octa
• 4-tetra,
9-nona
• 5-penta,
10- deca
Naming Covalent
Compounds
• First element whole name with the
appropriate prefix, except mono
• Second element, -ide ending with
appropriate prefix
• Practice
Naming Covalent Compounds
• CO2
• CO
• CCl4
• N2O4
• XeF6
• P2O9
• H2O
• Carbon Dioxide
• Carbon Monoxide
• Carbon tetrachloride
• Dinitrogen tetroxide
• Xenon hexaflouride
• Diphosphorus Nonaoxide
• Dihydrogen Monoxide
Water!
Covalent compounds
• The name tells you how to write the
formula
• duh
• Sulfur dioxide
SO2
• diflourine monoxide
F 2O
• nitrogen trichloride
NCl3
• diphosphorus pentoxide P2O5
Empirical formula
• Shows the lowest, simplified ratio of
elements in a compound
Molecule
Molecular
formula
Empirical
formula
1
2
3
C2H4
C4H8
C3H8
CH2
CH2
?
• Chemical Formula- represents the
relative # of atoms of each kind in a
chemical compound by using atomic
symbols and numerical subscripts.
Example: H2O H=2 O=1
• Molecular compound (Covalent
Compounds) - simplest formula unit are
molecules. Have low melting & boiling pts.
• Molecular formula- shows the types and
numbers of atoms combined in a single
molecule.
• Bond Length- is the average
distance between 2 bonded atoms.
• Bond Energy- is the energy required
to break a bond.
• It gives us information about the
strength of a bonding interaction.
I. Lewis Diagrams
(p. 202 – 213)
Lecture 2
The Octet Rule
• Chemical
compounds tend to
form so that each
atom, by gaining,
losing, or sharing
electrons, has an
(8) octet of
electrons in its
valence shell.
H
He
8 is
Great!
Lewis Dot
• Shows how valence
electrons are
arranged among
atoms in a molecule.
• Reflects central idea
that stability of a
compound relates to
noble gas electron
configuration.
I
started
it for
you
CH4
Try one yourself NH3
H2O
A. Octet Rule
• Remember…
– Most atoms form bonds in order to have
8 valence electrons.
•
B. 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.
B. Drawing Lewis
Diagrams
• CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
F
F C F
F
B. Drawing Lewis
Diagrams
• CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e-
O C O
B. Drawing Lewis
Diagrams
• BeCl2
1 Be × 2e- = 2e2 Cl × 7e- = 14e16e- 4e12e-
Cl Be Cl
A. Octet Rule
• Exceptions:
F
F
 Hydrogen  2 valence e
F
B
F
 Groups F
1,2,3 get
2,4,6
valence e
S
F
H
N
O
O
H
 Expanded octet

more
than
8
F
Very
unstable!!
valence
e
(e.g.
S,
P,
Xe)
F
F
-
-
-
 Radicals  odd # of valence e-
C. 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
• ClO41 Cl × 7e- = 7e4 O × 6e- = 24e31e+ 1e32e- 8e24e-
O
O Cl O
O
C. Polyatomic Ions
• NH4+
1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 8e0e-
H
H N H
H
Electornegativity
Lecture 3
Is it Covalent or ionic
1. Nonpolar-Covalent bonds (H2)
• Electrons are equally shared
• Electronegativity difference of 0 to
0.5
2. Polar-Covalent bonds (HCl)
• Electrons are unequally shared
• Electronegativity difference between .5
and 2.1
3. Ionic Bonds 2.1- 3.3
Using Electronegativity
differences
• CO
• C=2.5
• O= 3.5
3.5-2.5= 1.0 polar covalent
look on table pg 198
2.1 thru 3.3= Ionic Bond
Electronegativity
Nonpolar Covalent
Polar covalent
Ionic
0-------------0.5 --------------------------------------- >2.1 ---------3.3
practice
• NaCl
•
• CO
• CH4
Na=0.9 Cl=3.2 3.2-0.9=2.3
NaCl is ionic
3.4- 2.6= 0.8 Polar covalent
2.6 -2.2= 0.4 Non polar covalent
D. 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.
D. Resonance Structures
 SO3
O
O S O
O
O S O
O
O S O
D. Resonance-Occurs when more
than one valid Lewis structure can
be written for a particular
molecule.
Structural Formula
• Shows shared pair
of electrons by a
dashed line.
•
•
•
•
•
•
Single bond- 1 pair of electrons
Double bond- 2 pair of electrons
Triple bond- 3 pair of electrons
Try a couple:
O2
N2
Ionic Bonding and
Compounds
• Ionic Compound- is composed of
positive and negative ions combined
so that the positive & negative
charges are equal. (Metal + nonmetal)
• Formula Unit- is the simplest
collection of atoms from which a
compounds formula can be
established.
1.
2.
3.
Writing Formulas
Write the symbols of each element
Put their charge in their upper right corner
Crisscross the numbers down (Not the charges).
Example:
Write the formula for Magnesium Chloride
Mg Cl
Mg+2 Cl-1
Mg+2 Cl-1
MgCl
MgCl22
Writing Formulas Practice
+1
•
1.
2.
3.
4.
+2
Write the formula for:
Aluminum Bromide
Calcium Oxide
Calcium Nitride
Sodium Chloride
+3
4 -3 -2 -1
• Lattice Energy- is the energy
released when 1 mole of an ionic
crystalline compound is formed from
gaseous ions.
• Ionic Compounds have high melting
points & boiling points, are hard and
brittle, have crystalline structure.
Polyatomic Ions
• Many atoms with a
charge.
• Example SO4-2
Metallic Bonding
• Metals- conduct heat, have low
ionization energy & electronegativity,
give up e• Metallic Bond- is a chemical bond
resulting from the attraction
between positive ions and surrounding
mobile e-.
• Malleability and ductility
Molecular Geometry
• VSEPR Theory- “Valence- shell,
electron-pair repulsion”
• states that repulsion between the
sets of valence-level electrons
surrounding an atom cause these sets
to be oriented as far apart as
possible.
Determining VSEPR
•
1.
2.
3.
4.
5.
Determine the VSEPR
for H2O
Draw the Lewis Dot
Draw the Structural
Formula
Label the central atom
as A
Label any atoms
attached to the center
atom as B
Label any paired
electrons on the central
atom that are not used
in the bond as E
H-O-H
B
E2
A B
VSEPR AB2E2
Shape Bent (look on chart)
VSEPR Chart
VSEPR
AB or AB2
AB2E
AB3
AB4
AB3E
AB2E2
AB5
AB6
SHAPE
Linear
Bent
Trigonal-Planar
Tetrahedral
Trigonal-Pyramidal
Bent
Trigonal-Bipyramidal
Octahedral
• Hybridization-The Blending of Orbitals.
• Dipole- is created by equal but opposite charges
that are separated by a short distance.
• Dipole-Dipole Attractions-Attraction between
oppositely charged regions of neighboring
molecules.
• Hydrogen Bonding- Bonding between hydrogen
and more electronegative neighboring atoms such
as oxygen and nitrogen. Hydrogen bonding in
Kevlar, a strong polymer used in bullet-proof
vests.
• London Dispersion Forces- The temporary
separations of charge that lead to the London
force attractions are what attract one nonpolar
molecule to its neighbors. London forces increase
with the size of the molecules.
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“Electronegativity chart”. Table. Aug. 9, 2006.
http://www.chemistry210.com/notes/u01s06f.htm
“Lewis Structures”. Drawings. Aug. 9, 2006. http://www.avonchemistry.com/chem_bond_explain.html
“Oscar”. Photo. Aug. 9, 2006. http://www.musicmerchant.com/22061.htm
“Water Structural Formula”. Drawing. Aug. 10, 2006.
http://www.accs.net/users/kriel/ch4notes/water_structura
l_formula.gif
“Periodic Table of Elements”. Chart. Aug. 9, 2006.
http://users.erols.com/kdennis/periodictable.jpg
“Information”. Aug 11, 2006.
http://www.sciencegeek.net/Chemistry/Powerpoin
t/Unit3/Unit3_files/frame.htm
Holt, Rinehart and Winston. Modern Chemistry. Harcourt Brace & Company.
1999.