Ch 11 - IMFs Flashcards
1
Q
Intermolecular Force
A
attractive forces between all molecules and atoms.
2
Q
High thermal energy
A
tend to be gasses
3
Q
low thermal energy
A
tend to be solids or liquids
4
Q
What has the strongest IMFs?
A
Solids, then liquids, then gasses
5
Q
Crystalline Solid
A
the atoms or molecules are arranged in a well-ordered three-dimensional array
6
Q
Amorphous Solid
A
the atoms or molecules have no long range order
7
Q
How do you change the state of matter?
A
Increase temperature and/or pressure.
8
Q
Higher pressure =
A
more dense state
9
Q
Higher temperature =
A
less dense state
10
Q
4 types of IMFs
A
dispersion force, dipole-dipole force, hydrogen bond, and ion-dipole force
11
Q
Dispersion Force
A
london dispersion force.
all molecules and atoms have these.
fleeting charges create an unbalance of electrons around the atoms/molecules creating temporary dipoles.
12
Q
What can dispersion forces do?
A
the temporary +/- in one can trigger a chain reaction in others creating strength temporarily.
13
Q
Major factors affecting dispersion forces?
A
the larger the electron cloud the greater the dispersion forces can be.
All other things being equal then larger molar mass = larger dispersion force.
Shape also plays an important factor: more surface area = more dispersion forces.
14
Q
Greater dispersion force =
A
higher boiling point.
more energy is required to break the bonds.
15
Q
Why does shape affect LDFs?
A
long chains have greater LDFs than clumped patterns because more surface area = chance to tangle and increase attraction.
all other factors = then longer chain = more LDFs
C6H14>C5H12
16
Q
Dipole-dipole Force
A
exist in all molecules that are polar.
Has permanent dipole - a molecule with an electron rich and electron poor regions
17
Q
Permanent Dipole
A
a molecule with an electron rich and electron poor regions
18
Q
Dipole-Dipole Forces still have _____.
A
Dispersion forces
19
Q
Dipole-Dipole Forces have higher ______
A
melting and boiling points than molecules with similar masses because there are more forces to overcome.
20
Q
A greater dipole moment typically_____.
A
has a higher boiling point when compared to non dipole-dipole force molecules of similar molar mass.
21
Q
Miscibility
A
the ability to mix without separating into two states of liquids.
polar with polar.
nonpolar with nonpolar.
22
Q
Hydrogen Bonding
A
incredibly strong type of Dipole-Dipole Force caused by very large electronegativity difference AND close proximity to H since it only has 1 electron.
H with FON
23
Q
Hydrogen bonding is the _____.
A
strongest of LDFs, Dipole-Dipole, and H Bonds because it has all 3 types.
Unique.
24
Q
Ion-dipole Force
A
an ionic compound mixed with a polar compound.
Especially important with aqueous solutions.
Very strong compared to other IMFs.
25
What is the most important manifestation of IMFs?
the very existence of liquids and solids.
26
Surface Tension
the energy required to increase the surface area by a unit amount.
27
Molecules at the surface have _____.
less molecules to interact with and are inherently less table(higher potential energy)
28
liquids tends to minimize surface area because
more surface area = higher potential energy and subsequently more work must be done to maintain.
Creates a kind of skin on the surface that resist penetration.
29
Surface tension decreases as _____.
IMFs decrease.
30
What shape reduces surface area the most?
A sphere.
31
A sphere has_____
the smallest surface area to volume ratio(minimizes potential energy)
32
Viscosity
the resistance of a liquid to flow.
33
What is viscosity measured in?
poise(P)
1g/cm*s
34
Liquid water has a viscosity of
1 centipoise(cP)
35
Major factors affecting viscosity.
IMFs, Shape, and temperature.
36
How do IMFs affect viscosity?
more surface area for IMFs to interact tends to have higher viscosity.
Long chains can tangle and IMFs interact more.
37
How does molar mall affect viscosity?
generally higher molar mass = higher viscosity as more IMFs are able to interact.
38
How does temperature affect viscosity?
Higher temperature breaks IMFs and LOWERS viscosity.
39
Capillary Action
the ability of a liquid to flow against gravity up a narrow tube.
blood sample, trees
40
Cohesive Force
the attraction between molecules of a liquid
41
adhesive force
the attraction between molecules and the surface of a tube
42
Adhesive force>cohesive force
then liquid will be drawn up the tube
43
Cohesive force shape in tube
Convex.
pulling away from tube and bubbling up in middle.
44
Adhesive force shape in tube
concave.
being drawn up the side of the tube looking depressed in center of liquid.
45
the higher the temperature the greater the average_____
energy of the collection of molecules which can be vaporized more easily.
46
Vaporization
the transition from liquid to gas.
Endothermic
47
Condensation
the transition from gas to liquid.
Exothermic
48
Condensation and vaporization both______
occur simultaneously but under normal conditions more evaporation then condensation occurs on the substance,
49
What increases the evaporation rate?
higher temperature.
more surface area.
Decreased IMFs
Will not increase temperature until all liquid is turned to gas.
50
Volatile
liquids that evaporate easily.
Acetone
51
Nonvolatile
liquids that do not evaporate easily.
Motor oil
52
if no new energy is added to a system then the temperature of the liquid will_______.
decrease as evaporation occurs.
Sweat cools the body.
Endothermic process.(takes heat away from body)
53
Vaporization is a _____ process.
endothermic.
Absorbs energy to turn to a gas.
54
Condensation is a _____ process.
Exothermic
energy must be released from the gas to turn it to a liquid.
Steam burns condense on skin releasing a significant amount of heat which hurts.
55
Heat of Vaporization
the amount of heat required to vaporize one mol of a liquid to gas.
From liquid to gas is a + number as energy is required.
Gas to liquid is a - number as energy is released.
56
Hvap of H2O at normal boiling point(100C 1atm)
+40.7kJ/mol
Endothermic.
57
Hvap of H2O at normal boiling point(100C 1atm)
To condense
-40.7kJ/mol
Exothermic
58
Dynamic Equilibrium
the point when the rate of evaporation = rate of condensation.
Both still occur.
59
Vapor Pressure
The pressure of a gas in dynamic equilibrium with its liquid.
Dependent on the IMFs and temperature.
60
Weak IMFs =
volatile liquid with high vapor pressure
61
Strong IMFs =
nonvolatile liquid with low vapor pressure.
Stronger IMFs = less gas and more liquid so less pressure.
62
What happens when dynamic equilibrium is disturbed?
the system responds so as to minimize the disturbance and return to a state of equilibrium.
Holds true as long as gas and liquid are present in system.
63
If pressure decreases in a system then ____
some liquid vaporizes to restore dynamic equilibrium.
64
If pressure increases in a system then _____
some gas condenses to liquid to restore dynamic equilibrium.
65
Temperature and vapor pressure.
Higher temperature = higher vapor pressure.
More gas in system.
66
Change in temps effect on a systems vapor pressure
Small change in temperature = large difference in the number of molecules which can vaporize
= large increase in vapor pressure.
Exponential relationship.
67
Boiling Point
the temperature at which the liquids vapor pressure equals the external pressure.
the point where the liquids internal molecules have enough energy to break free and enter the gaseous state.
68
Normal Boiling Point
the temperature at which the liquids vapor pressure equals 1 atm.
At a lower pressure liquids will boil at a lower temperature.
Takes longer to cook food at high altitudes because not as hot.
69
Adding heat ______ raise the temperature but only ______ the rate of boiling until all liquid has changed to gas.
does not
increases
70
As long as liquid is present boiling _____
will not change the temperature
71
Clausius-Clapeyron Equation
gives a linear relationship between the natural log of vapor pressure and the inverse of temperature.
a convenient way to measure the heat of vaporization in the lab.
72
760mmHg =_____=_____
760 torr
1 atm
73
Supercritical Fluid
the point in a sealed container where the meniscus between gas and liquid disappears and the to comingle due to increased temperature(forcing gas) and increased pressure(forcing liquid)
74
Critical Temperature(Tc)
the temperature at which the transition to a supercritical fluid happens.
75
Critical Pressure(Pc)
the pressure at which the transition to a supercritical fluid happens.
76
Sublimation
transition from solid to gas
77
deposition
the transition from gas to solid
78
typically sublimation happens _____than deposition
more rapidly.
ice will shrink over time even if below freezing
79
Melting Point
molecules have enough thermal energy to partially overcome IMFs and begin to move around each other.
80
Freezing point
liquid to solid.
Enough energy is lost to stop movement of molecules.
81
increasing temperature _____ while melting.
only expedites the change of states.
does not increase temperature until all solid has turned to liquid.
82
Heat of Fusion(Hfus)
endothermic
amount of heat required to melt 1 mol of a solid
83
Water Hfus
6. 02kJ/mol(fusion/melting)
| - 6.02kJ/mol(freezing)
84
Hfus is ______ energy then the Hvap
significantly less
85
Melting only ______ overcomes IMFs while vaporizing _____ overcomes IMFs
partially
completely
86
Heating curve for water
5 stages:
ice,melting,water,boiling, vapor
2 formulas:
[3 states accounting for heat]q=m(Cssolid/liquid/gas)(delta T) in j/mol
[2 transitions no temperature change]q=n(deltaHfus/Hvap) in kJ/mol
calculate, convert to proper units, and add together to get total energy.
87
Phase Diagram
a map of the state or phase of a substance as a function of pressure(torr)(y axis) and temperature(C)(x axis)
88
triple point
represents the unique set of conditions at which the three states of matter are equally stable and in equilibrium.
89
H2O triple point
0.0098C and 4.58torr.
90
Critical point
represents the temperature and pressure above which a super critical fluid exists.
91
H2O fusion curve on a phase diagram
Atypical.
Most substances have a positive slope but water is negative.
increasing pressure favors the liquid state as it is more dense than the solid state.
92
X-ray diffraction
powerful laboratory technique that enable us to determine the arrangement of atoms and measure the distance between them.
93
destructive interference
two waves interact with the crests of one aligning with the trough of another.
weakens/cancels each other out.
94
constructive interference
two waves troughs and crests align amplifying the effect.
95
interference pattern
pattern of light and dark spots made by waves to measure them.
96
Braggs law
n(Alpha)=2(d)(sin theta)
alpha = 154pm, theta = 32.6 degrees, n= 1
d=((1)(154pm))/(2sin(32.6degrees))
= 143pm
97
crystalline lattice
the regular arrangements of atoms within a crystalline solid.
natures way of aggregating the particles to minimize their energy.
98
unit cell
a small collection of atoms, ions, or molecules repeated over and over to make up the crystalline lattice.
many different ones exist.
99
Coordination number
the number of atoms with which each atom is in DIRECT contact in a unit cell.
the number of atoms each atom can strongly interact with.
100
packing efficiency
the percentage of volume of the unit cell occupied by spheres.
higher efficiency = more interaction = stronger bonds
101
simple cubit unit cell
1/8 atom at each corner of cell.
```
# atoms = 1
coordination # = 6
edge length(l) = 2r
packing efficiency = 52%
```
102
body-centered cubic unit cell
1/8 atom at each corner and 1 atom in center.
```
# atoms = 2
coordination # = 8
edge length(l) = 4r/sqrt3
packing efficiency = 68%
```
atoms do not touch along the same edge but from one corner along the diagonal through the center to opposite corner.
103
face-centered cubic unit cell
1/8 atom in each corner and 1/2 atom in each face of cell.
```
# atoms = 4
coordination # = 12
edge length(l) = 2(sqrt2)(r)
packing efficiency = 74%
```
atoms do not touch along the edge but along the diagonal of the face from one corner through the center of the face to the opposite corner.
104
Hexagonal closest packing
the first and third layers mirror each other while the second is offset to fit closer together.
packing efficiency = 74%.
ABAB set up
touch 6 atoms in its layer and 3 in the layer above and below it.
NOT cubic but a hexagonal arrangement.
105
Cubic Closest packing
the first and third layers are backwards of each other while the second is offset and the 4th and 1st are mirrors of each other.
ABCABC
identical to the face-centered cubit unit cell structure.
Packing efficiency = 74%
106
3 types of crystalline solids
molecular, ionic, and atomic
107
Molecular Solids
ice
- composite units = molecules
- low melting points.
- held together by IMFs
- strong IMFs can change melting point(H2O and hydrogen bonding)
108
Ionic Solids
table salt
- composite unit = formula units(cations and anions)
- high melting points
- held together by coulombic interactions
- coordination number represents the close cation-anion interactions
109
Ionic Solids and Potential Energy
higher coordination numbers = lower PE.
Goal is to maximize the coordination number while keeping charge neutrality(unit cell must be neutral)
Size and charge neutrality will determine the structures.
High disproportions of size can significantly change the geometric shape.
110
Atomic Solids
Composit Units = atoms
3 types:
nonbonding
metallic
network covalent
111
Atomic Solids:
nonbonding
solid xenon
- held together by relatively weak dispersion forces
- low melting points
- tend to form closest packed structures to maximize IMFs, coordination number, and minimize distance.
- ONLY exist as noble gases in solid form
112
Atomic Solids:
metallic
gold
- held together by metallic bonds
- variable melting points
- metallic bonds are not directional
- tend to form closest packed structures of varying strength
113
Atomic Solids:
Network Covalent
Quartz, silicon dioxide
- held together by covalent bonds
- high melting points
- diamond, graphite, and silicon dioxide
- more restricted by geometrical constraints of the covalent bonds
- DO NOT tend to form closest packed structures
- covalent atomic solids have strong covalent bonds and subsequently high melting points
- silicates are the most commmon
- shape plays a huge factor in these
114
Band Theory
- combination of the atomic orbits of the atoms within a solid crystal to form orbitals that are not localized on individual atoms, but delocalized over the entire crystal.
- infinitely small gaps between the orbitals creates one super band.
115
Valence Band
the occupied molecular orbitals.
116
Conduction Band
the unoccupied orbitals
117
Band Gap
a gap between the valence and conduction bands
118
The size of the band gap determines
if the molecules are a conductor, semiconductor, or insulator.
the entire basis for computers
119
Conductors have ____ gap.
No.
120
Semi conductors have _____ gap.
A small
121
Insulators have _____ gap.
a lage
122
Dopant
a molecules which allows more or less electrons to jump between the valence and conduction bands.
123
Dopants control _____.
the conductivity of semiconductors.
124
N-Type Semiconductor
semiconductor with an additional electron which is forced into the conduction band because the valence band is full
- results in a negatively charged conduction band allowing charge to be passed through.
125
P-Type semiconductor
semiconductor with a deficiency in electrons in the valence band creating a "hole" in the electrons
- effectively adds a positive charge which electrons can jump around.
126
P-N junctions
tiny spots that are p-type on one side and n-type semiconductors on the other which can serve a variety of functions be being able to control the flow of charge.
Computers.
127
Diodes
circuit elements that allow the flow of electrical current in only one direction
128
Amplifiers
elements that amplify a small electrical current into a large one
129
where do you read the meniscus?
At its lowest point in the middle.
130
Can non polar molecules be viscous?
yes.
If the chain length is long the molecules can still get tangled.
131
3 names for dispersion forces
Dispersion forces, London Dispersion Forces, and Van der Waals Force.
Interchangeable.
132
What is an important force in solutions?
Ion-dipole forces.
+ and - interact
- salts dissolving in water
133
What is boiling?
the point where vapor pressure = atmospheric pressure
134
Diethyl Ether(CH3OCH3)
What forces present?
CH3 non polar.
C-O polar.
dispersion and dipole-dipole forces.
135
Ethanol(C2H6O)
What forces present?
Dispersion, dipole-dipole, and hydrogen bonds(OH group)
high boiling point
136
Cyclohexane(C6H12)
What forces present?
Dispersion forces.
low boiling point.
