new Chapter 6 notes.notebook
April 27, 2015
Chapter 6 ­ Structure and Properties of Substances
In chapter 5 you learned about the different types of bonds and their individual properties:
1. Ionic
2. Covalent/Molecular
3. Metallic
In this chapter you will be looking at the structure of molecules and molecular shape is linked to the structures.
Dec 4­8:43 PM
Chap 6.1 ­ Covalent Bonds and Structures
‐ Molecular compounds come in a great variety of shapes
‐ These shapes are determined by the covalent bonds which form the molecules
‐ Lewis structures can be used to predict the structures and properties of molecules
Oct 16­7:42 PM
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METHOD 1
1.Find the valence number for each element.Put the element with the lowest number in the middle. Place the other elements around it.
2.Fill in the electrons for the center element.Add outside element's electrons so that all valence electrons are used. Check for the octet rule.You may have to move pairs to make double or triple bonds.
Apr 7­7:44 AM
Example: Draw the Lewis structure for CH2O.
Step 1: Total # of valence electrons:
C ‐ H ‐
O ‐
Total = Step 2: Skeleton Structure:
Step 3: Put lone pairs around the outer electrons:
Step 4: Put lone pairs around the central atoms ‐ make double or triple bonds if necessary.
May 16­3:29 PM
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Example 1: H2O
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Example 2: CO2
Dec 4­9:40 PM
Example 3: NH3
Dec 4­9:41 PM
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VSEPR ‐ Valence Shell Electron Pair Repulsion Theory
‐ The VSEPR theory states that the bonding pairs and lone pairs of electrons in the valence level of an atom repel each other due to their negative charges. This helps us to predict the shapes of molecules.
‐ To determine the shape of a molecule, first look at the number of electron groups that the molecule has. Electron groups can be either bonded electrons or lone pairs of electrons.
Oct 16­8:16 PM
VSEPR has 2 basic rules.
1.Bonding electron pairs repel each other, therefore adjust to be as far apart as possible.
Example: Methane
Apr 7­7:56 AM
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2. Unshared pairs of electrons are held closer to the atom than the bonding pairs.The unshared pairs of electrons strngly repel the bonding electron pairs,pushing them closer together. EXAMPLE AMMONIA
WATER
CARBON DIOXIDE
Apr 7­7:58 AM
Electron pair repulsions are not always equal.The repulsions can be ranked as :
STRONGEST: 2 unshared pairs
MEDIUM:
One unshared and one shared pair (bond)
WEAKEST:2 shared pairs(bonds)
Apr 7­8:02 AM
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Number of Electron Groups
Name of Molecular Shape
2
Shape
Angle
Example
Linear
180
0
CO
3
Trigonal Planar
120
0
CH 2 O
4
Tetrahedral
109.5
4
Pyramidal
107
0
NH
4
Bent
105
0
H 2O
0
2
CH 4
3
May 17­3:43 PM
How to predict molecular shape using VSEPR:
1. Draw a Lewis structure for the molecule
2.
Determine the total number of electron groups around the central atom . (**double and triple bonds count as 1 group).
3.
Look at where the bonds and lone pairs are, and determine which of the 5 shapes best accommodates the combination of electron groups.
May 17­3:46 PM
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Example: Draw a Lewis structure for CH2Cl2 and use VSEPR to determine its shape and bond angle.
May 17­3:50 PM
Example: Use VSEPR to determine the shape and bond angle of PH3.
May 17­5:04 PM
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BUILD ASSIGNMENT
skeleton
shape
bond angle
SH
GeSe
SiH
PH
Apr 7­8:21 AM
Question: Draw Lewis structures for each of the following. a﴿ CBr4
b﴿ NCl3
c﴿ NCl4+
d﴿ PS2+
e﴿ NS2­
Oct 16­8:02 PM
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Drawing Lewis Dot Diagrams (with a central atom): method 2
Step 1 ­ Determine the total number of valence electrons of all the atoms in the molecule.
Step 2 ­ Draw a skeleton structure.
­ Put the atom with the lowest group number in the middle
­ Join the atoms with a pair of bonding electrons (subtract 2 electrons from the total for each bond you make)
Step 3 ­ Put lone pairs around all atoms except the central atom (follow the octet rule). Note: The most electronegative atoms get the electrons first!
Step 4 ­ Put the remaining electrons around the central atom (octet rule!)
­ If all the valence electrons are used up but the central atom does not have an octet of electrons, move one or more lone pairs from the outer atoms to form double or triple bonds.
Oct 16­7:44 PM
Apr 7­9:50 AM
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Resonance Structures
­> Occur when there is more than one possible Lewis structure
Example: O3 (ozone) Dec 4­9:47 PM
Example: SO2
Dec 4­10:12 PM
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More Challenging Lewis Structures
How do we determine the central atom?
The atom with the most unpaired valence electrons will be the central atom of a molecule.
Example: CH3NH2
Example: CH3OH
Dec 4­10:13 PM
Drawing Lewis Structures for Polyatomic Ions (charges)
Example : PO43­
Dec 5­1:07 PM
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Coordinate Covalent Bonds: ­ Bonds formed when one atom contributes both electrons to make a shared pair to satisfy an octet.
Example : The ammonium ion, NH4+, is formed when ammonia joins with H+. Example : CO
May 16­3:32 PM
Example : Show the coordinate covalent bond formed when water joins with H+. May 16­7:11 PM
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Polarity of Molecules
A non­polar molecule has no net dipole.
This is when molecules have:
­only non­polar bonds
­polar bonds arranged symmetrically so that dipoles cancel out
A polar molecule has a net dipole.
­This can be found in molecules made up of polar bonds arranged so that the dipoles do not cancel out.
Dec 11­3:10 PM
Polarity of Bond Types:
• Linear ­ polar or non­polar
• Trigonal Planar ­ polar or non­polar
• Tetrahedral ­ polar or non­polar
• Pyramidal ­ always polar
• Bent ­ Always polar
Oct 16­8:38 PM
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Oct 24­1:27 PM
How to tell if a molecule is polar or non­polar? Are there electrons on the central atom?
YES
POLAR
Check change in EN:
Is there a diff. greater than 0.5? If yes then, POLAR. 1) Draw arrows in the direction of the most electronegative atom. Are there different atoms bonded to the central atom?
YES
2) If the arrows cancel our, then the molecule is non­
polar. If they do not cancel out, the molecule is polar and has dipoles. NO
NON­POLAR
Dec 11­3:13 PM
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Determine the polarity of the following molecules:
1) SO3
2) SCl2
3) CS2
4) CH2O
5) CH2Cl2
Oct 16­8:49 PM
Use VSPER to predict the shape of the following molecules, then determine it's polarity.
1) COCl2
2) BeF2
3) [NO2]+
4) NOF Oct 16­8:51 PM
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sa
Chap 6.2 : Intermolecular Forces
Ever Wonder? .....
-How Salt Dissolves?
-Why Water only freezes on the Top of the lake?
-What holds our DNA together?
ANSWER = Intermolecular Forces!
Oct 16­8:55 PM
When a substance melts or boils, intermolecular forces are broken (NOT intramolecular bonds)
* a high boiling point indicates strong attractive forces
Dec 12­6:33 PM
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-Intramolecular Forces = Forces that bond the
ATOMS to each other within a molecule.
Example = Covalent Bonds, ionic bonds, metallic
bonds.
-Intermoleccular Forces = Forces that bond the
MOLECULES to each other.
Weak relative to covalent and ionic bonds.
Oct 16­8:58 PM
Types of Intermolecular Forces
1 : Dipole-Dipole Forces
2 : Dispersion (London) Forces
3 : Hydrogen Bonding
Oct 16­9:01 PM
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1.Dipole Interactions(Dipole­dipole)
­occurs when the slightly positive end of one dipole attracts to the slighty negative end of another dipole.
­They help explain why polar molecules have higher boiling points than non­polar molecules.
Apr 7­1:14 PM
2: Dispersion (London) Forces
h
Also called Londen Forces.
In a non­polar molecule,the motion of the electrons causes a momentary uneven dstribution of charge. A non­polar moleule becomes slightly polar for an instant.The result is an attraction similar to a dipole interaction.
These are the main forcec between non­polar molecules.
The larger the molecule ,the stonger the Londen forces
F and CL are gases at room temp.
Br is a liquid
I is a solid
Oct 16­9:21 PM
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3: Hydrogen Bonding
­Electrostatic Attraction between the nucleus of a
hydrogen atom (bonded to a highly electronegative
atom such as O, F, N) and the negative end of a
nearby dipole.
~ only 5% as strong as a covalent bond.
Oct 16­9:31 PM
WATER
­ hydrogen bonding accounts for the many unique properties of water
­ in liquid water hydrogen is bonded to at least four other water molecules ­ is the only pure substance that exists in nature in all three states of matter at the same time.
Oct 16­9:53 PM
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Hydrogen bonds explain why:
1. Ice floats in water
Water is not like most liquids. ~Normally a liquid contracts when cooled so it becomes more dense and sinks in its own liquid, but water expands when frozen. ~Water expands when frozen because as the temperature of the water is getting lower the H­bonds pull the water into an orderly pattern called a HONEYCOMB FRAMEWORK. ~At low temperatures the KE of the water molecules is too low to break out of this shape. This honeycomb framework gives the ice a bigger volume for its mass which in turn decreases the density allowing ice to float in water.
INSULATES OCEAN LIFE
Oct 16­9:55 PM
2. Water has a high Boiling Point
Water has an unusually high boiling point. As a general rule, the lower the molar mass the lower the boiling point, but the H­bonds in water disrupt this rule. This allows the earth to be cooled. Water acts as a heat sink! Moderates th seasons.
3. Water has a high surface tension INSECTS FLOAT ON WATER
SURFACE TENSION ­ an inward force or pull that tends to minimize the surface area of a liquid.
H­bonds cause the inward pull to create water’s high surface tension
(Activities to accompany)
SURFACTANT – a substance that decreases the surface tension of water. Eg detergent.
4.ADHESION­Water molecules attach to objects.Water sticks to the xylem.
5.COHESION_Water molecules attach to other molecules. This is how water goes up a tree after being started by transpiration. Oct 16­9:57 PM
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Intermolecular Forces ­ YouTube
Oct 17­10:17 AM
Chap 6.3 ­ Structure Determines Properties
TEMPERATURE is defined as a measure of the average kinetic energy of the particles. To increase the KE of particles you increase the temperature.
~A substance will change states from a solid to a liquid or a liquid to a gas when the average kinetic energy of the particles is great enough to overcome the force of attraction holding the particles together. When ‘no’ attractive forces exist with other particles the substance will exist as a gas.
Oct 16­10:00 PM
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Melting Point and Boiling Point
­ Depends on the attractive forces holding the particles together. ­ In order for metals and ionic compounds to melt or boil, they must break the ionic bonds or metallic bonds. ­ BUT when molecular compounds melt or boil, the intermolecular attractions between the molecules must break (NOT the covalent bonds within)
Ionic and metallic bonds are very strong while intermolecular attractions are weak. So metals and ionic compounds melt or boil at much higher temperatures than molecular compounds. Polar molecules boil at higher temperatures than nonpolar substances because of the extra intermolecular attractions (and the bond energy of these attractions)
Metals > ionic > polar > non­polar
Oct 16­10:03 PM
Ionic Compounds Metallic Compounds
Molecular Compounds
Eg. NaCl
Eg. Mg Eg. C6H12O6 ­ must break ionic bonds ­ intermolecular ­ must break metallic bonds between bonds molecules break, not the covalent bonds May 23­3:14 PM
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Strength of Intramolecular and Intermolecular Forces
STRONGEST
1. Metallic Ex: Cu, Fe, CoI
2. Ionic Ex: LiCl (NOT aqueous)
3. Covalent Ex: O­H bond in water
4. hydrogen bonding
5. dipole­dipole
6. London dispersion
INTRA
INTER
WEAKEST
Dec 16­12:31 PM
How would you determine the higher boiling point between two ionic compounds?
• Check charges! Higher set of charges wins!
Examples:
1. NaCl and BaO
2. Na3N and K2O
Dec 16­12:36 PM
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How would you determine the higher boiling point between two London dispersion?
• Size!
Example:
1. Only type I will ever give!
a) C2H6 and C3H8
b) C4H10 and C5H12
Dec 16­12:45 PM
Circle the one that has the highest boiling point:
1. Cu or H2O 2. CF4 or LiBr
3. Sr3N2 or NaCl
4. NH3 or PH3
5. LiCl(aq) or C2H6
6. HF or NaCl(aq)
7. C2H6 or C3H8 Dec 16­12:50 PM
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Electrical Conductivity
To conduct an electrical current, electrons or ions must be able to move independently of oppositely charged ions.
~Metals are good conductors of electricity as solids. ~Ionic compounds are good conductors of electricity when melted or dissolved in water.
~Molecular compounds are poor conductors of electricity.
Oct 16­10:10 PM
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