Intermolecular Forces - Research at OSU Chemistry

Report
Chapter 10:
Structure of Solids and Liquids
Chem 1110
Figures: Basic Chemistry 3rd Ed., Timberlake and Timberlake
Electron-Dot Symbols
Electron-dot symbols:
• Show the valence electrons of an atom
• Electrons are arranged in s and p orbitals
around the element symbol
• Can mix orbitals and create “hybrid” orbitals
with equal energy (like in Chapter 6)
o sp, sp2, and sp3
Covalent Bond Revisited
Covalent Bond is a chemical bond that
results from two nuclei attracting the same
shared electron pair
• Shared electrons in the bond
For Example: F2
Bonding and Lone Pair Electrons
Bonding Pair are shared to create the
covalent bond
Non-Bonding Pair or lone pair make up
the octet
Bonding and Lone Pair Electrons
Bonding Pair are shared to create the
covalent bond
Non-Bonding Pair or lone pair make up
the octet
H2
HCl
H 2O
Rules for Drawing
Electron Dot Formulas
1. Find the total number of valence electrons
contributed by all atoms
2. Write chemical symbols in the order that they
are bonded
3. Place one covalent bond between each atom
•
Determining the central atom is important
•
First atom is usually the central atom
o EXCEPT when it is hydrogen (H)
Rules for Drawing
Electron Dot Formulas
4. Add lone pairs to the bonded atoms first, then
to central atom
5. Check the Octet Rule for ALL atoms
• May need to make multiple bonds
For Example: SF2
Rules for Drawing Electron Dot Formulas
1. Find the total number of valence electrons
contributed by all atoms
2. Write chemical symbols in the order that they
are bonded
3. Place one covalent bond between each atom
4. Add lone pairs to the bonded atoms first, then
to central atom
5. Check the Octet Rule for ALL atoms
• May need to make multiple bonds
Electron Dot Formulas
EXAMPLES:
CO2
NH3
SiCl4
Exceptions to the Octet Rule
Boron is stable with a total of 6 electrons
• Boron will make compounds with only
three bonds to B
BCl3
Period 3 (and above) non-metals can
form compounds with expanded octets
(> 8 valence electrons): P, S, Se
Multiple Bonds
Multiple bonds involve the sharing of more
than one pair of electrons:
Bond Order
Bond Order denotes the level of bonding
in a molecule:
• Single bond: Share one electron pair
Bond Order of 1
• Double bond: Share two electron pairs
Bond Order of 2
• Triple bond: Share three electron pairs
Bond Order of 3
Bond Length and Bond Energy
As we add additional bonds between atoms
the atoms are drawn closer together:
• Bond length decreases with more bonds:
Single > Double > Triple
As we add additional bonds between atoms
the bonds get stronger:
• Bond energy increases with more bonds:
Triple > Double > Single
Bond Length and Bond Energy
Practicing organic compounds:
• Single bond (alkane): Ethane, C2H6
• Double bond (alkene): Ethene, C2H4
• Triple bond (alkyne): Ethyne (Acetylene), C2H2
Electron Dot Formulas
EXAMPLES:
O2
HCN
CS2
CO
Resonance Structures
Resonance structures are equivalent structures
that show where multiple bonds may occur
• Electrons are delocalized over entire molecule
• Resonance changes bond order, bond length,
and energy
Examples:
SO2
SO3
Learning Check
Draw the three resonance structures of
carbonate, CO32-
Electron Dot Formulas
How can we draw electron dot structures for
polyatomic ions?
NH4Br or
Na2SO3
These compounds have both an ionic and
covalent:
[NH4]+ [Br]2 [Na]+ [SO3]2-
Electron Dot Formulas
[NH4]+
[SO3]2-
Learning Check
Draw the electron dot formula for ClO3-
Learning Check
Write two resonance structures for nitrite.
Orbital Geometry and
Molecular Shape
• Electrons in bonds and lone pairs want to get
as far apart from each other as possible
• We have seen the tetrahedral shape of carbon
• Causes the molecule to conform to a set of
shapes
Determining Orbital Geometry
• Electron pairs do NOT like to share space
with other electron pairs
• Electrons will move away as far as possible
to obtain “personal space”
o An electron pair can be in a bond
o An electron pair can be a “lone pair”
(non-bonding)
VSEPR Theory: Valence-Shell Electron
Pair Repulsion Theory
Areas of electron density around an atom
want to maximize their “personal space”
• A lone pair of electrons is ONE area
• Two electrons in a single bond are ONE area
• Four electrons in a double bond are ONE area
• Six electrons in a triple bond are ONE area
VSEPR Theory
Determining Orbital Geometry:
• Count ALL areas of electron density to
determine the base geometry:
• Two is linear (sp)
• Three is trigonal planar (sp2)
• Four is tetrahedral (sp3)
Orbital Geometry
Linear (sp)
Orbital Geometry
Trigonal Planar (sp2)
Orbital Geometry
Tetrahedral (sp3)
Orbital Geometry
Linear and Trigonal Planar Molecules
Tetrahedral Molecules
Molecular Shape is Altered by
Lone Pairs
Two different “bent” structures which differ in
their bond angle: 120° vs. 109.5°
• Lone pairs take up more room and squeeze
bond angles together:
o Trigonal planar: 120°
o Trigonal planar with lone pair: < 120°
o Tetrahedral: 109.5°
o Tetrahedral with lone pair: < 109.5°
Orbital Geometry and
Molecular Shape
REMEMBER:
1) Electrons repel each other
2) Electrons in bonds have a fixed location
3) Electrons in lone pair take up more space
4) Two atoms always define a line.
Molecular Shape
Let’s look at two seemingly similar molecules:
CO2 and H2O
We first have to draw the molecule to determine
their shapes:
Molecular Shape
We can base molecular shape on orbital
geometry:
• Linear: CO2
• Trigonal Planar: BF3, CH2O
• Tetrahedral: CH4 , NH3, H2O
Learning Check
Determine the shape of the N2O molecule:
Electron Dot Formulas
Practice drawing electron dot structures and
identify the proper molecular shape for:
AlH3
NF3
SiCl4
HCN
CO
Electronegativity
In the preceding structures we had sharing of
electrons (covalent bonds) and charged species
(transfer of electrons, ionic bonds):
How do we know what bond type we have?
• General Rule of Thumb:
o A metal with a non-metal forms an ionic bond
o A non-metal binding with a non-metal forms a
covalent bond
WHY????
Electronegativity
The type of bond interaction is determined
by differences in electronegativity:
Electronegativity is the ability of an atom
to attract shared electrons in a bond
towards itself
• Time to revisit Periodic Trends…
Electronegativity
Increasing Electronegativity
Increasing Electronegativity
Fluorine is the MOST electronegative
Cs/Fr are the LEAST electronegative
Electronegativity Values
• We can assign values for electronegativity:
Learning Check
Place the following in order of increasing
electronegativity: O, K, and C
Determining Bond Type
Difference in Electronegativity (ΔEN) allows
us to determine the type of bond formed:
1. Difference > 1.8 → ionic bond
2. Difference = 0.0 → pure covalent bond
3. Difference between 0.0 and 0.4
•
Non-polar covalent
4. Difference between 0.4 and 1.8
•
Polar covalent bond
Common Pure Covalent Bond
Species
O2
Cl2
Br2
I2
H2
N2
S8
P4
Bond Polarity
Bonds which have an unequal sharing of
electrons are said to be Polar:
• The electron cloud is pulled toward the
more electronegative atom
• This sets up a separation of charges: dipole
• The more electronegative element attracts
the cloud more strongly - partially negative
• The less electronegative element attracts
the cloud less - partially positive
Bond Polarity
Bond Polarity
Blue: low electron density
Red: high electron density
Green: non-polar
Bond Polarity
A different way of looking at polarity:
EXAMPLES
Molecule
NH3
O2
Δ EN and Bond Type
0.9 → polar covalent
0 → pure covalent
NaCl
2.3 → ionic
SO2
1.0 → polar covalent
Electronegativity and Bond Types
Predicting Bond Types
Learning Check
Use electronegativity differences to classify each of
the following bonds as nonpolar covalent (NP), polar
covalent (P), or ionic (I):
A bond between:
1) K and N
2) N and O
3) Cl and Cl
4) H and Cl
Drawing Bond Dipoles
K–N
N– O
Cl – Cl
Molecular Shape and Polarity
Dipoles can add across a molecule
• Polar molecules – have an overall dipole
• Non-polar molecules – no overall dipole
To determine whether a molecule is polar or
non-polar, we must determine the 3-D shape
of the molecule
Molecular Polarity
• Additive dipoles across a molecule make
it polar:
• If dipoles cancel, no net dipole and the
molecule is non-polar:
Molecular Polarity
EXAMPLES:
CO2
NH3
CF4
Learning Check
Identify each of the following molecules as
(P) polar or (NP) nonpolar:
A. PBr3
B. HBr
C. Br2
D. SiBr4
Learning Check
Draw the Lewis Structure for SeCl2 and
determine its shape and molecular polarity:
Forces in Matter
Intermolecular Forces: are attractive forces
between molecules:
• Weaker than bonds (covalent or ionic)
• Help to determine the state of matter:
solid, liquid or gas
Intermolecular forces organize matter and
are opposed by motion of molecules which
disorganize matter
States of Matter
Solid
Attractive Forces >> Disruptive Forces
Liquid
Attractive Forces ≈ Disruptive Forces
Gas
Attractive Forces << Disruptive Forces
Intermolecular Forces
Three primary types of forces that help to hold
covalent molecules together:
1. London Dispersion Forces
2. Dipole – Dipole Interactions
3. Hydrogen Bonding
Intermolecular Forces
London Dispersion Forces occur in non-polar
molecules by forming a temporary dipole:
• Temporary dipoles induce dipoles in nearby
molecules
• Lasts for milliseconds!
• ALL matter has London Dispersion Forces
• Let’s consider H2 (non-polar molecule, H–H)
Intermolecular Forces
London Dispersion Forces
• Increase with increasing size
• Increase with increased number of electrons
Examples: CH4, CF4, CCl4, CBr4
• CBr4 will have the strongest London dispersion
forces as it is the largest
• CH4 will have the weakest London dispersion
forces
Intermolecular Forces
London Dispersion Forces
• Increase with increasing size
• Increase with increased number of electrons
o Electrons on larger atoms are further from the
nucleus therefore easier to induce temporary
dipole
Examples: CH4, CF4, CCl4, CBr4
• CBr4 will have the strongest London dispersion
forces as electrons on Br are further from nucleus
Intermolecular Forces
Dipole-Dipole Interactions occur between
polar molecules where δ+ end attracts the δend of another molecule
• There is a permanent dipole in polar
molecules: attraction of molecules
• Occur only between polar molecules
• Consider FCl:
Fluorine is more electroneagtive (δ-)
Intermolecular Forces
Dipole-Dipole Interactions
• Stronger than London dispersion forces
• Strength of interaction increases with
increasing size of the permanent dipole
For Example:
FCl (ΔEN= 1.0), FBr (ΔEN= 1.2), FI (ΔEN= 1.5)
• FI will have the strongest dipole–dipole interaction
• FCl will have the weakest dipole–dipole forces
Intermolecular Forces
Hydrogen bonds are a special type of dipoledipole force when we have a VERY large
difference in electronegativity
• To have a Hydrogen Bond:
1.
2.
3.
4.
Need to have H
Need to have F, O, or N
H must be bonded to F, O, or N
At least one lone pair on the F, O, or N
Intermolecular Forces
Hydrogen bond: The hydrogen on one
molecule is attracted to the lone pair on
another molecule
• The Hydrogen bond is the STRONGEST
intermolecular force!
• Let’s consider H2O!
Hydrogen Bonds in Water
Hydrogen Bonds
Hydrogen Bonds
Predicting strength of H-bonds:
1) CH3OCH3 vs. C2H5OH
2) CH3NH2 vs. CH3NHCH3 vs. (CH3)3N
Learning Check
Identify the major type of attractive force in
each of the following substances:
A. NCl3
B. H2O
C. Br2
D. KCl
E. NH3
Equilibrium
• When the forces acting on matter are in
balance, we have Equilibrium
• Or, Equilibrium occurs when two or more
opposing forces balance each other
Equilibrium: Balance of Forces
Equilibrium occurs when two or more
opposing forces balance each other, and is
a dynamic process:
• Represented by the symbol:
⇄
• For Example (Changes of State):
Solid ⇄ Liquid
Liquid ⇄ Gas
Melting point
Boiling point
Changes of State
Intermolecular Forces
Strength of intermolecular forces increase as
previously described:
London < Dipole-Dipole < Hydrogen Bond
• As the strength of intermolecular forces increases,
the properties of molecules change:
• Want to “stick together” more!
• Polar molecules tend to be liquids or solids at
room temperature
• Requires increased kinetic energy (disruptive
forces) to separate them!
Intermolecular Forces
Strength of intermolecular forces increase as
previously described:
London < Dipole-Dipole < Hydrogen Bond
• As the strength of intermolecular forces increases,
they change the properties of molecules:
o Melting Point Increases
o Boiling Point Increases
o Vapor Pressure Decreases
Learning Check
Identify the compound in each pair that has the
higher melting point. Explain your choice:
A. NCl3 or NH3
B. HBr or Br2
C. KCl or HCl
Balance of Forces
Let’s consider H2O:
• Hydrogen bonds are the strongest force
• Requires lots of energy to break apart
intermolecular interactions
• Consequently, water has an increased
melting point, decreased vapor
pressure, and increased boiling point
Liquid Water

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