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Chapter 5 1
Example
NO2 ( ) nitrogen ( ) oxide
2 O’s nitrogen dioxide
1 N Mononitrogen dioxide
often omit “mono”
Nitrogen dioxide
Chapter 5 2
Name the following:
CCl4
N2O3
SiS2
Chapter 5 3
Name the following:
CCl4 carbon tetrachloride
N2O3dinitrogen trioxide
SiS2 silicon disulfide
Chapter 5 4
Write the formula for
dihydrogen monoxide
Silicon tetrafluoride
dinitrogentetroxide
Chapter 5 5
Write the formula for
dihydrogen monoxide H2O
Silicon tetrafluoride SiF4
dinitrogen tetroxide N2O4
Chapter 5 6
Lewis Dot Structures
Helpful in determining 3-D Shape of molecule
Can use 3-D shape to predict properties of molecules
Chapter 5 7
Rules for Lewis structures for molecules
1.Put in the atoms and arrange them to show which atoms are connected to which other atoms.
H always on outside
2.Count the total number of outer shell electrons available to form bonds.
Each atom contributes its group # of e-
Chapter 5 8
3. Draw bonds between atoms. Add in the remaining available electrons in pairs, starting with outside atoms to make octets
4. Make double or triple bonds if necessary to form complete octets around each atom. (Move e- pairs)
Exceptions to full octetH needs only 2 e-
B needs only 6 e-
Period 3 or higher can have > 8
Chapter 5 9
Double check to make sure all atoms have full octets!!!
Chapter 5 10
Draw Lewis structures for:
H2
HCl
PH3
CO2
Chapter 5 11
Resonance
Some compounds can have “equivalent” resonance structures (SO2)
Only difference is “placement” of double bond
Two structures are known as resonance forms
Chapter 5 12
In actual fact, neither double bond structure exists
True situation is resonance hybrid
(midway between two resonance structures)
Chapter 5 13
Types of Covalent Bonds
Bond Type Pairs of e- Total electrons
Single 1 2
Double 2 4
Triple 3 6
Unshared pairs of electrons are known as non-bonding pairs or lone pairs
Chapter 5 14
Bond Energy
Bond energy is amount of energy required to break a bond
Bond energies measured in J or kJ
Stronger bonds have higher bond energy
Triple > Double > Single
Chapter 5 15
Bond Length
Bond length is distance between two nuclei
Shorter in multiple bonds
Length of Bond:
Triple < Double < Single
Chapter 5 16
Valence Shell Electron Pair Repulsion (VSEPR)
Electron “pair” groups in the outer shell of atoms arrange themselves as far away from each other as possible.
Electron group occupies one region of space:
bonding pair of electrons
nonbonding pair of electrons (lone pair)
double or triple bond
Chapter 5 17
Molecular Geometry
Arrangement of atoms around a central atom
Look at one center at a time
Chapter 5 18
Central atom with:
Two bonding regions
(attached to two atoms)
No lone (nonbonding) pairs
Arrangement is linear
Bond angles are 180o
Chapter 5 19
Chapter 5 20
Central atom with:
Three bonding regions
(attached to three atoms)
No lone (nonbonding) pairs
Arrangement is planar
(molecule is flat)
Bond angles are 120o
Chapter 5 21
Chapter 5 22
Central atom with:
Four bonding regions
(attached to four atoms)
No lone (nonbonding) pairs
Arrangement is tetrahedral
Bond angles are 109.5o
Chapter 5 23
Chapter 5 24
Representing 3-D Structures
Solid line: bonds in plane of paper
Dotted wedges: bonds that project behind (or beneath) the plane
Solid wedges: bonds that project in front of (or above) the plane
Chapter 5 25
Chapter 5 26
Geometry of Atoms with Lone Pairs
Two bonding regions
(attached to two atoms)
One lone (nonbonding) pair
Arrangement is angular (bent)
(molecule is flat)
Bond angles are 120o
Chapter 5 27
Chapter 5 28
Geometry of Atoms with Lone Pairs
Three bonding regions
(attached to three atoms)
One lone (nonbonding) pair
Arrangement is trigonal pyramidal
Bond angles are 107o
Chapter 5 29
Chapter 5 30
Geometry of Atoms with Lone Pairs
Two bonding regions
(attached to two atoms)
Two lone (nonbonding) pairs
Arrangement is angular (bent)
Bond angles are 104.5o
Chapter 5 31
Why are bond angles smaller than 109.5o?
Bonding electrons have an atom on both sides of the bond
Lone pairs tend to spread and force the bonding pairs closer together
Chapter 5 32
Chapter 5 33
Fig. 4.8
Chapter 5 34
Fig. 4.9
Chapter 5 35
Summary of Effects of Lone Pairs on Bond Angles
OHH
104.5O107O
NHH
HC
H
HHH109.5O
Chapter 5 36
Summary of Molecular Geometry
# Electron Lone Bond Angle Shape regions pairs
2 0 180o Linear
3 0 120o Planar 2 1 Angular (bent)
4 0 109.5o Tetrahedral 3 1 Trigonal pyramidal 2 2 Angular (bent)
Chapter 5 37
Review: Electronegativity
Measure of the relative pull of an atom on a shared pair of electrons
Arbitrary scale ranging from 0 to 4
Most electronegative element is fluorine
F has E = 4
Chapter 5 38
Electronegativity Values
Chapter 5 39
Polarity of bondsCovalent Bonds between atoms with similar electronegativity values are nonpolar
Δ E < 0.5
Covalent Bonds between atoms with different electronegativity values are polar covalent
0.5 < Δ E < 1.9
Bonds between atoms with very different electronegativity values are ionic
Δ E > 1.9
Chapter 5 40
Polar Bond has positive and negative ends to the bond (uneven distribution of charge)
Polar molecule has positive and negative ends to the molecule (even distribution of charge)
Chapter 5 41
Polar Bond has positive and negative ends to the bond
Polar molecule has positive and negative ends to the molecule
Molecule acts like a dipole
Chapter 5 42
Polar Covalent Bonds---Unequal Sharing of Electrons
Chapter 5 43
Dipoles can align in an electrical field
Nonpolar molecules do not have +/- ends and do not align in electrical field
Chapter 5 44
Nonpolar molecule
All of the bonds are C—H bonds.
C end is slightly negative compared to H end for each bond
Overall, molecule is nonpolar because it does not have positive and negative ends---Polarities cancel out
Chapter 5 45
Polar Molecule
Each O—H bond is polar
O end is more negative than H end
Negative end to molecule
Polarities do not cancel each other----instead are additive
Chapter 5 46
• Intramolecular Forces– Forces within molecules
Covalent bonds
• Intermolecular Forces– Forces between molecules
Hydrogen bonds Dipole-dipole interactions Dispersion Forces
Chapter 5 47
Intermolecular Forces
• Hold matter together
• Boiling point and melting point a good measure of how strong intermolecular forces are
Chapter 5 48
More energy is required to separate molecules held together by strong intermolecular forces than weak intermolecular forces.
Materials with strong intermolecular forces have higher boiling points and melting points
Chapter 5 49
Ionic Forces
• Interaction between + and – charges on ions
• Strongest intermolecular forces– Found in salts– NaCl melts at 801oC
Chapter 5 50
Hydrogen Bonds• H that is covalently attached to O, N or F is
attracted to a different O, N or F
• This attraction is a “hydrogen bond”
• Much weaker than ionic forces or covalent bonds (intramolecular bond) between atoms
• Important in biological systems
• Water melts at 0oC
Chapter 5 51
Hydrogen Bond
Chapter 5 52
Dipole Forces
• Result from attraction between partially positive and partially negative ends (poles) of molecules
• Works only for polar molecules
• HCl melts at -112oC
Chapter 5 53
Dipole Forces (a=solid, b=liquid)
Chapter 5 54
Dispersion Forces (London)
• Present in all molecules
• Weak temporary forces that result from movement of electrons within molecules and around atoms– Important in nonpolar materials– Each individual London force is very weak,
but together they add up, especially for large molecules
Chapter 5 55
Intermolecular Forces
Strongest
Ionic forces (much stronger)Hydrogen bondingDipole ForcesDispersion Forces
Weakest
Chapter 5 56
Properties affected by Intermolecular Forces
Melting points
NaCl > H2O > HCl > CH4
mp 801 0 -112 -182.5oC
Ionic H-bond Dipole London
Chapter 5 57
States of Matter
Solid
Liquid
Gas
Chapter 5 58
States of Matter
Solid Liquid Gas
Definite shape
Yes No No
Definite volume
Yes Yes No
Is fluid (Pours or Flows)
No Yes Yes
Chapter 5 59
Solid Liquid Gas
Interactions between Particles
Strong Moderate None
Particles touching
Yes Yes (some holes)
No
Space between particles
No Some holes
LOTS of space
Chapter 5 60
Changes of State
Melting:
Change from the solid to the liquid state
Freezing:
Change from the liquid to the solid state
(solidification or crystallization)
Chapter 5 61
Changes of State
Vaporization:
Change from the liquid to the gaseous state (Evaporation or Boiling)
Condensation:
Change from the gaseous to the liquid state
Chapter 5 62
Changes of State
Sublimation:
Change from the solid to the gaseous state directly (skips liquid state)
Deposition:
Change from the gaseous to the solid state directly (skips liquid state)
Chapter 5 63
Melting point:
Temperature at which substance goesfrom solid to liquid state
Boiling point:
Temperature at which substance goes from liquid to gaseous state
Chapter 5 64
Summary of State Changes
Chapter 5 65
Energy is needed to overcome the forces between molecules
Need to add energy for substance to melt or vaporize (or sublime)
Need to remove energy for substance to condense or freeze (or deposit)
Heat of Vaporization is amount of energy needed to change given amount of liquid into gas at its boiling point
Chapter 5 66
Gases
Lots of space between molecules
Gas molecules can be pushed closer together because there is plenty of space between them
Gases are compressible
Chapter 5 67
Molecules move in straight line until they hit the sides of their container and move in a different direction
When molecules hit container, they exert a force
Pressure: Force per unit area
Chapter 5 68
Pressure units
Atmosphere:
1 atm is pressure at sea level
mm Hg:
760 mm Hg is pressure at sea level
1 atm = 760 mm Hg
Chapter 5 69
Boyle’s Law
Pressure (P) and Volume (V) of a gas are inversely proportional
P α 1/V or PV = constant
Raise Pressure → Lower Volume
Lower Pressure → Increase Volume
Chapter 5 70
Chapter 5 71
Chapter 5 72
Charles’s Law
Temperature (T) and Volume (V) of a gas are directly proportional
P α T or P/T = constant
Raise Temperature → Raise Volume
Lower Temperature → Lower Volume
Chapter 5 73
Chapter 5 74
• Solution is homogeneous mixture (mixed uniformly)
Chapter 5 75
Properties of Liquids
Molecules are close together, so liquids are only slightly compressible
Each liquid has a unique vapor pressure
Vapor pressure is the pressure of gas molecules that have escaped the liquid state at a given temperature
(closed container)
Chapter 5 76
Boiling Point
Boiling Point:
Temperature at which vapor pressureis equal to atmospheric pressure
Normal Boiling Point:
Boiling point at 1 atm pressure
Chapter 5 77
Water: Unique Properties
• Water is liquid at room temperature– Hydrogen bonds between water molecules
give water a higher melting point
• Density of ice is less than density of water
• Water has high heat capacity.– Can absorb lots of energy with very little
change in temperatue
Chapter 5 78
Hydrogen Bond
Chapter 5 79
Water: Unique Properties
• Water has high heat of vaporization
• Water has high surface tensionSurface tension: Molecules at surface are only attracted on one side and form a “skin”
• Water is an excellent solvent
Chapter 5 80
Solutions
Solvent: Substance present in greater amount
Solute:Substance present in smaller amount
Solubility:Amount of solute that will dissolve in a given amount of solvent at a given temperature
Chapter 5 81
Saturated solution:
Contains the maximum amount ofsolute it can dissolve
Unsaturated solution:
Contains less than the maximum amount of solute possible (can dissolve more)
Chapter 5 82
Insoluble:Solute will not dissolve in the solvent
Miscible:Solute and solvent will dissolve in all proportions
Chapter 5 83
Solubility
Like dissolves like.
Polar substances dissolve in polar substancesSalt and water, sugar and water
Nonpolar substances dissolve in nonpolar substances
Grease and CCl4
Unlike substances do not dissolve each other
Oil and vinegar
Chapter 5 84
Ionic Compounds in SolutionIons dissolve in water---Charges on ions attracted to partial charges on polar water molecules
Chapter 5 85
Compounds in Solution
Electrolyte:*Compound that conducts electricity when it is melted or dissolved in water *Ions carry charges and conduct electrons*Electrolytes are ionic compounds
Non-electrolyte:*Compound that does not conduct
electricity when it is melted or dissolved in water
*Nonelectrolytes are not ionic compounds
Chapter 5 86
Strong electrolyte:
*Completely separates into ions in solution
*Conducts electricity well
Weak electrolyte:
*Partly separates into ions in solution
*Conducts electricity poorly
