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Chem 2 Heider > Ch 11 - IMFs > Flashcards

Ch 11 - IMFs Flashcards

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1
Q

Intermolecular Force

A

attractive forces between all molecules and atoms.

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2
Q

High thermal energy

A

tend to be gasses

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3
Q

low thermal energy

A

tend to be solids or liquids

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4
Q

What has the strongest IMFs?

A

Solids, then liquids, then gasses

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5
Q

Crystalline Solid

A

the atoms or molecules are arranged in a well-ordered three-dimensional array

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6
Q

Amorphous Solid

A

the atoms or molecules have no long range order

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7
Q

How do you change the state of matter?

A

Increase temperature and/or pressure.

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8
Q

Higher pressure =

A

more dense state

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9
Q

Higher temperature =

A

less dense state

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10
Q

4 types of IMFs

A

dispersion force, dipole-dipole force, hydrogen bond, and ion-dipole force

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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.

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12
Q

What can dispersion forces do?

A

the temporary +/- in one can trigger a chain reaction in others creating strength temporarily.

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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.

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14
Q

Greater dispersion force =

A

higher boiling point.

more energy is required to break the bonds.

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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

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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

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17
Q

Permanent Dipole

A

a molecule with an electron rich and electron poor regions

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18
Q

Dipole-Dipole Forces still have _____.

A

Dispersion forces

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19
Q

Dipole-Dipole Forces have higher ______

A

melting and boiling points than molecules with similar masses because there are more forces to overcome.

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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.

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21
Q

Miscibility

A

the ability to mix without separating into two states of liquids.

polar with polar.

nonpolar with nonpolar.

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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

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23
Q

Hydrogen bonding is the _____.

A

strongest of LDFs, Dipole-Dipole, and H Bonds because it has all 3 types.

Unique.

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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.

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25
Q

What is the most important manifestation of IMFs?

A

the very existence of liquids and solids.

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26
Q

Surface Tension

A

the energy required to increase the surface area by a unit amount.

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27
Q

Molecules at the surface have _____.

A

less molecules to interact with and are inherently less table(higher potential energy)

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28
Q

liquids tends to minimize surface area because

A

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.

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29
Q

Surface tension decreases as _____.

A

IMFs decrease.

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30
Q

What shape reduces surface area the most?

A

A sphere.

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31
Q

A sphere has_____

A

the smallest surface area to volume ratio(minimizes potential energy)

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32
Q

Viscosity

A

the resistance of a liquid to flow.

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33
Q

What is viscosity measured in?

A

poise(P)

1g/cm*s

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34
Q

Liquid water has a viscosity of

A

1 centipoise(cP)

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35
Q

Major factors affecting viscosity.

A

IMFs, Shape, and temperature.

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36
Q

How do IMFs affect viscosity?

A

more surface area for IMFs to interact tends to have higher viscosity.

Long chains can tangle and IMFs interact more.

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37
Q

How does molar mall affect viscosity?

A

generally higher molar mass = higher viscosity as more IMFs are able to interact.

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38
Q

How does temperature affect viscosity?

A

Higher temperature breaks IMFs and LOWERS viscosity.

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39
Q

Capillary Action

A

the ability of a liquid to flow against gravity up a narrow tube.

blood sample, trees

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40
Q

Cohesive Force

A

the attraction between molecules of a liquid

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41
Q

adhesive force

A

the attraction between molecules and the surface of a tube

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42
Q

Adhesive force>cohesive force

A

then liquid will be drawn up the tube

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43
Q

Cohesive force shape in tube

A

Convex.

pulling away from tube and bubbling up in middle.

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44
Q

Adhesive force shape in tube

A

concave.

being drawn up the side of the tube looking depressed in center of liquid.

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45
Q

the higher the temperature the greater the average_____

A

energy of the collection of molecules which can be vaporized more easily.

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46
Q

Vaporization

A

the transition from liquid to gas.

Endothermic

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47
Q

Condensation

A

the transition from gas to liquid.

Exothermic

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48
Q

Condensation and vaporization both______

A

occur simultaneously but under normal conditions more evaporation then condensation occurs on the substance,

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49
Q

What increases the evaporation rate?

A

higher temperature.

more surface area.

Decreased IMFs

Will not increase temperature until all liquid is turned to gas.

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50
Q

Volatile

A

liquids that evaporate easily.

Acetone

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51
Q

Nonvolatile

A

liquids that do not evaporate easily.

Motor oil

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52
Q

if no new energy is added to a system then the temperature of the liquid will_______.

A

decrease as evaporation occurs.

Sweat cools the body.

Endothermic process.(takes heat away from body)

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53
Q

Vaporization is a _____ process.

A

endothermic.

Absorbs energy to turn to a gas.

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54
Q

Condensation is a _____ process.

A

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.

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55
Q

Heat of Vaporization

A

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
Q

Hvap of H2O at normal boiling point(100C 1atm)

A

+40.7kJ/mol

Endothermic.

57
Q

Hvap of H2O at normal boiling point(100C 1atm)

To condense

A

-40.7kJ/mol

Exothermic

58
Q

Dynamic Equilibrium

A

the point when the rate of evaporation = rate of condensation.

Both still occur.

59
Q

Vapor Pressure

A

The pressure of a gas in dynamic equilibrium with its liquid.

Dependent on the IMFs and temperature.

60
Q

Weak IMFs =

A

volatile liquid with high vapor pressure

61
Q

Strong IMFs =

A

nonvolatile liquid with low vapor pressure.

Stronger IMFs = less gas and more liquid so less pressure.

62
Q

What happens when dynamic equilibrium is disturbed?

A

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
Q

If pressure decreases in a system then ____

A

some liquid vaporizes to restore dynamic equilibrium.

64
Q

If pressure increases in a system then _____

A

some gas condenses to liquid to restore dynamic equilibrium.

65
Q

Temperature and vapor pressure.

A

Higher temperature = higher vapor pressure.

More gas in system.

66
Q

Change in temps effect on a systems vapor pressure

A

Small change in temperature = large difference in the number of molecules which can vaporize

= large increase in vapor pressure.

Exponential relationship.

67
Q

Boiling Point

A

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
Q

Normal Boiling Point

A

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
Q

Adding heat ______ raise the temperature but only ______ the rate of boiling until all liquid has changed to gas.

A

does not

increases

70
Q

As long as liquid is present boiling _____

A

will not change the temperature

71
Q

Clausius-Clapeyron Equation

A

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
Q

760mmHg =_____=_____

A

760 torr

1 atm

73
Q

Supercritical Fluid

A

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
Q

Critical Temperature(Tc)

A

the temperature at which the transition to a supercritical fluid happens.

75
Q

Critical Pressure(Pc)

A

the pressure at which the transition to a supercritical fluid happens.

76
Q

Sublimation

A

transition from solid to gas

77
Q

deposition

A

the transition from gas to solid

78
Q

typically sublimation happens _____than deposition

A

more rapidly.

ice will shrink over time even if below freezing

79
Q

Melting Point

A

molecules have enough thermal energy to partially overcome IMFs and begin to move around each other.

80
Q

Freezing point

A

liquid to solid.

Enough energy is lost to stop movement of molecules.

81
Q

increasing temperature _____ while melting.

A

only expedites the change of states.

does not increase temperature until all solid has turned to liquid.

82
Q

Heat of Fusion(Hfus)

A

endothermic

amount of heat required to melt 1 mol of a solid

83
Q

Water Hfus

A
  1. 02kJ/mol(fusion/melting)

- 6.02kJ/mol(freezing)

84
Q

Hfus is ______ energy then the Hvap

A

significantly less

85
Q

Melting only ______ overcomes IMFs while vaporizing _____ overcomes IMFs

A

partially

completely

86
Q

Heating curve for water

A

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
Q

Phase Diagram

A

a map of the state or phase of a substance as a function of pressure(torr)(y axis) and temperature(C)(x axis)

88
Q

triple point

A

represents the unique set of conditions at which the three states of matter are equally stable and in equilibrium.

89
Q

H2O triple point

A

0.0098C and 4.58torr.

90
Q

Critical point

A

represents the temperature and pressure above which a super critical fluid exists.

91
Q

H2O fusion curve on a phase diagram

A

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
Q

X-ray diffraction

A

powerful laboratory technique that enable us to determine the arrangement of atoms and measure the distance between them.

93
Q

destructive interference

A

two waves interact with the crests of one aligning with the trough of another.

weakens/cancels each other out.

94
Q

constructive interference

A

two waves troughs and crests align amplifying the effect.

95
Q

interference pattern

A

pattern of light and dark spots made by waves to measure them.

96
Q

Braggs law

A

n(Alpha)=2(d)(sin theta)

alpha = 154pm, theta = 32.6 degrees, n= 1

d=((1)(154pm))/(2sin(32.6degrees))

= 143pm

97
Q

crystalline lattice

A

the regular arrangements of atoms within a crystalline solid.

natures way of aggregating the particles to minimize their energy.

98
Q

unit cell

A

a small collection of atoms, ions, or molecules repeated over and over to make up the crystalline lattice.

many different ones exist.

99
Q

Coordination number

A

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
Q

packing efficiency

A

the percentage of volume of the unit cell occupied by spheres.

higher efficiency = more interaction = stronger bonds

101
Q

simple cubit unit cell

A

1/8 atom at each corner of cell.

# atoms = 1
coordination # = 6
edge length(l) = 2r
packing efficiency = 52%
102
Q

body-centered cubic unit cell

A

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
Q

face-centered cubic unit cell

A

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
Q

Hexagonal closest packing

A

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
Q

Cubic Closest packing

A

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
Q

3 types of crystalline solids

A

molecular, ionic, and atomic

107
Q

Molecular Solids

A

ice

  • composite units = molecules
  • low melting points.
  • held together by IMFs
  • strong IMFs can change melting point(H2O and hydrogen bonding)
108
Q

Ionic Solids

A

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
Q

Ionic Solids and Potential Energy

A

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
Q

Atomic Solids

A

Composit Units = atoms

3 types:
nonbonding
metallic
network covalent

111
Q

Atomic Solids:

nonbonding

A

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
Q

Atomic Solids:

metallic

A

gold

  • held together by metallic bonds
  • variable melting points
  • metallic bonds are not directional
  • tend to form closest packed structures of varying strength
113
Q

Atomic Solids:

Network Covalent

A

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
Q

Band Theory

A
  • 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
Q

Valence Band

A

the occupied molecular orbitals.

116
Q

Conduction Band

A

the unoccupied orbitals

117
Q

Band Gap

A

a gap between the valence and conduction bands

118
Q

The size of the band gap determines

A

if the molecules are a conductor, semiconductor, or insulator.

the entire basis for computers

119
Q

Conductors have ____ gap.

A

No.

120
Q

Semi conductors have _____ gap.

A

A small

121
Q

Insulators have _____ gap.

A

a lage

122
Q

Dopant

A

a molecules which allows more or less electrons to jump between the valence and conduction bands.

123
Q

Dopants control _____.

A

the conductivity of semiconductors.

124
Q

N-Type Semiconductor

A

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
Q

P-Type semiconductor

A

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
Q

P-N junctions

A

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
Q

Diodes

A

circuit elements that allow the flow of electrical current in only one direction

128
Q

Amplifiers

A

elements that amplify a small electrical current into a large one

129
Q

where do you read the meniscus?

A

At its lowest point in the middle.

130
Q

Can non polar molecules be viscous?

A

yes.

If the chain length is long the molecules can still get tangled.

131
Q

3 names for dispersion forces

A

Dispersion forces, London Dispersion Forces, and Van der Waals Force.

Interchangeable.

132
Q

What is an important force in solutions?

A

Ion-dipole forces.

+ and - interact

  • salts dissolving in water
133
Q

What is boiling?

A

the point where vapor pressure = atmospheric pressure

134
Q

Diethyl Ether(CH3OCH3)

What forces present?

A

CH3 non polar.
C-O polar.

dispersion and dipole-dipole forces.

135
Q

Ethanol(C2H6O)

What forces present?

A

Dispersion, dipole-dipole, and hydrogen bonds(OH group)

high boiling point

136
Q

Cyclohexane(C6H12)

What forces present?

A

Dispersion forces.

low boiling point.

Chem 2 Heider (9 decks)

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