Transcript Document 7446379

CHAPTER 9
COVALENT BONDS
• OCTET RULE = atoms lose, gain or
share electrons in order to acquire
a full set of 8 valence electrons
– (form stable configuration)
• COVALENT BOND- 2 or more
elements combine by sharing
electrons; generally occurs with
elements close together on the
table
• MOLECULE- formed when 2 or more
atoms bond covalently
– Examples: DNA; fats; cotton; polymers
– video
• ELECTRONEGATIVITY = tendency for
an atom to attract electrons to
itself when bonded
– Covalent bonded atoms have
electronegativities that are close
• The difference between the two values is
less than 1.70
Dog Bonds
•
http://ithacasciencezone.com/chemzone/lessons/03bonding
/dogbonds.htm
CHARACTERISTICS OF IONIC VS
COVALENT BONDS
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CHARACTERISTIC
Type of Particle
IONIC
formula unit
•
Bond formed
by transferring e-
by sharing e-
•
Types of elements
metal to nonmetal
nonmetal to
nonmetal
•
Physical state
•
Melting point
high
•
Solubility in water
high
•
ConductivityElectrical
solid
good
COVALENT
molecule
solid, liq., gas
low
low (some
will dissolve)
poor to nonconductive
FORMATION OF COVALENT BOND
• Covalent bond forms when the two
atoms come close together and the
distance is just right for the
attraction between one atom’s
protons and the other atom’s
electrons
Each covalent bond
represents one pair
of shared electrons
DIATOMIC MOLECULE
• Occur in nature as a molecule of two
atoms because they are more stable than
the individual atoms alone
• Examples
–
–
–
–
–
–
–
H2
O2
F2
Cl2
Br2
I2
N2
Except for Hydrogen, the
diatomic molecules take
the shape of the number
7 on the periodic table
NAMING MOLECULES
• NON-ORGANIC OR INORGANIC
– Uses prefixes that stand for the
number of atoms present
– Prefixes =
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1 = mono
2= di
3=tri
4=tetra
5= penta
6= hexa
7= hepta
8= octa
9= nona
10= deca
Rules for inorganic mol.
• 1st element– name as written on
periodic table, use a prefix if the #
of atoms is greater than one
• 2nd element- always add a prefix
indicating # of atoms present
– Change ending to – ide
– Example:
• SO2 – sulfur dioxide
• N2O3 – dinitrogen trioxide
INORGANIC PRAC. PROB.
• Write the name for the following
compounds:
– SiF4
– S5O6
– PCl3
• Write the formula for the following
compounds:
– Oxygen difluoride
– Disulfur trioxide
ACID NAMING
• ACID = a compound that produces
H+ ions in solution
• 2 types of acids:
– 1. Binary acid-contains hydrogen and
one other element (or pair without
Oxygen)
– 2. Oxyacid- has a polyatomic ion
containing oxygen with hydrogen
attached
Naming Binary Acids
• Use the prefix (hydro-)
• Add the root of the second element
and attach a (-ic) ending
• Examples:
– Hydrochloric acid
– Hydrofluoric acid
HCl
HF
Naming oxyacids
• If the ion ends in (-ate) change to
(-ic) and add the word “acid”
Example: HNO3
NO3 is nitrate
so change the ending and the
compound is then called nitric acid
• If the ion ends in (-ite) change to (ous) and add the word “acid”
– Example: H2SO2 (SO2 is sulfite)
change to sulfurous acid
Naming organic molecules
• HYDROCARBON = a compound
composed of hydrogen and carbon
• Organic names- consist of a prefix
that indicates the # of carbon
atoms followed by an ending that
indicates the # of bonds between
carbons
Organic Prefixes
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1 carbon atom = meth2 carbon atoms = eth3 carbon atoms = prop4 carbon atoms = but5 carbon atoms = pent6 carbon atoms = hex7 carbon atoms = hept8 carbon atoms = oct9 carbon atoms = non10 carbon atoms = dec-
Types of organic bonds
• ALKANE- all carbon-carbon bonds
are single
• Have “-ane” ending
• General formula= CnH2n+2 where
n= number of carbons
H
H
• C2H6 = ethane
H- C - C -H
H H
• ALKENE= have at least one carbonto-carbon bond that is double
C = C
• Have “-ene” ending
• General formula is CnH2n
• Example:
C3H6
propene
• ALKYNE – at least one carbon-tocarbon bond is a triple bond
C
C
• Have “-yne” ending
• General formula CnH2n-2
• Example:
C4H6 butyne
Drawing hydrocarbons
• Use sticks to represent bonds
– Single bond
-Double bond
– Triple bond
share 1 pair of e-= share 2 pairs of eshare 3 pairs of e-
• Use the symbol to represent the
element
• C= carbon H = hydrogen F= Fluorine
• Examples:
H H H H
H-C-C-C-C-H C4 H10
H H H H butane
Remember:
Carbon
Needs 4
electrons to
be shared
4 lines of
attachment
Hydrogen
Fluorine,
chlorine,
bromine,
iodine
Needs 1
electron
1 line of
attachment
Oxygen,
sulfur
Needs 2
electrons
2 line of
attachment
Nitrogen
3 electrons
3 lines
Practice Problems:
• Draw the following structures and write
their formula:
– Methane
octene
heptyne
– Hexyne
propane
butene
LEWIS STRUCTURE
• Model that uses electron dot
structures to show how electrons are
arranged in molecules
• Pairs of dots or lines represent
bonding pairs
• Hydrogen needs 2 dots, all other
elements need 8 dots around them
• Example:
: :
H: H
H:O:
H
or H - H
or H-O
H
SHAPES- MOLECULAR GEOMETRY
:
• For covalent bonds shape is very
important, it determines some of
its properties ( ex. Smells)
• Ex: NH3 has shape
N
H
H
H
• VSEPR THEORY- (valence shell
electron pair repulsion theory)
– In a small molecule, the electrons are
arranged as far apart as possible
SHAPES
• 1. LINEAR- straight line
– Bond angle = 180 degrees
– Includes molecule with just 2 atoms
(O = O, H-Cl) or more O=C=O
• 2. TRIGONAL PLANAR
– Flat triangle
– Central atom with 3 atoms attached
120o
Cl
B
Cl
Cl
• 3. TETRAHEDRAL
– Four-surfaced shape
– Triangular base
109.5o
H
C
H
H
H
• 4. PYRAMIDAL
:
– Has central atom bonded to three
others with an unshared pair of
electrons
N
H
H
H
107o
• 5. BENT
– Has one angle
O
H
H
105o
Shape simulation
Animated molecules
WEAK FORCES
• Intramolecular forces- forces within
a molecule that holds atoms
together in a covalent or ionic bond
– Van der Waals- weak forces
involving the attraction of the
electrons of one atom for the
protons of another (covalent)
• Intermolecular forces- forces
of attraction between
molecules (like in solid or
liquid state)
– Dipole-dipole forces- forces of
attraction between two polar
molecules
– Simulation
• Bond length- distance between
two bonded atoms (also called
bond axis)
Bond axis (length)
O
H
H
Bond Angle
• Bond angle- distance between axis
• Bond energy- energy required to
break a chemical bond
• COVALENT RADIUS-radius of an
atom when bonded to another
Cl
Cl
Covalent radius
MOLECULAR POLARITY
• Not all covalent bonds are the same
• The bonding shared pairs of
electrons are pulled between the
nuclei of the atoms sharing them
• Look at electronegativity
differences to see what type of
covalent bond
• Types
– Nonpolar covalent
– Polar covalent
NONPOLAR COVALENT
• When the atoms are chemically
similar
• The bonding e- are shared
EQUALLY (ex. H2, O2)
• OR bonds are symmetrically
arranged (everything cancels)
– Ex. O = C= O or CH4
• Electronegativity difference is < .4
generally
POLAR COVALENT
• The 2 atoms are joined by sharing
electrons UNEQUALLY (E.N.
difference is usually .4-1.7)
• The atom with stronger eattraction (higher E.N. #) acquires
a slightly negative charge
• The atoms with the lower E.N. #
gets a slightly positive charge
• Set up poles
Points to side w/higher EN #
Electronegativity
Difference
Type of bond
0.0 - 0.4
0.4 – 1.0
covalent- nonpolar
covalent- moderately
polar
covalent- very polar
ionic
1.0 – 1.7
> than 1.7
Non- Polar: Symmetrical
(POLAR bonds that aren’t POLAR molecules)
• A polar molecule always contains
polar bonds but some molecules
with polar bonds are non-polar
molecules (symmetrical)
– CF4 = non polar
• Polar molecules develop partially
charged ends, and ARE NOT
symmetrical, like water
DIPOLE
• Created by equal but opposite
charges that are separated by a
short distance (use to show
direction of (-) charge)
DISSOLVING RULE
• RULE = “LIKE DISSOLVES LIKE”
• Polar molecules dissolve in polar
water or polar solvent
– Washable markers
– Ionic dissolves in polar
• Nonpolar molecules dissolve in
nonpolar solvents
– Perm. Markers- alcohol
– Dry cleaning
CONDUCTIVITY
• Transmission of electric current by
ions
• Ionic bonds- have electric charge
(e-) so they conduct
• Covalent bonds- no charges, so no
charged particles- no conduction
• ELECTROLYTE- substance that
dissolves in water to give a
conducting solution (salts, acids,
bases)
• In order to have conductivity:
– 1. you need charged particles
– 2. the charged particles (ions) must
be free to move
– 3.the higher the concentration, the
more ions, the greater the conductivity
CHROMATOGRAPHY
• Used to separate the components
of a solution to identify them
(method of fractionationseparating parts from a whole)
• Separates by polarity
Parts of chromatography:
• 1. Solvent- used to move the
mixture (needs to be same polarity
as mobile phase)
• 2. mobile phase- consists of the
mixture to be separated
• 3. stationary phase- thing the
mobile phase is going to travel
through- can be a solid or a liquid
adhered to the surface of the solid
– Has an attraction for polar molecules
– Most polar found –bottom of paper
– Least polar (non polar) found- top of
paper
Chromatography process:
• The substances in the mobile phase
will travel at different rates thru
the stationary phase depending on
their polarity
– Slowest moving substance- has
greatest attraction for the stationary
phase (most polar) found at the
bottom of the paper
– Fastest moving substance- least
attraction (least polar); top of paper;
closer to the solvent front
TYPES OF CHROMATOGRAPHY
• Paper
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Column
TLC (thin layer)
Gas chromatography
Ion exchange