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AP Chemistry > Unit 3 - Intermolecular Forces & Properties > Flashcards

Unit 3 - Intermolecular Forces & Properties Flashcards

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1
Q

Intramolecular forces

A

Bonds

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2
Q

Intermolecular forces

A

Forces of attraction between molecules

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3
Q

Intramolecular forces examples

A

Includes covalent, ionic, and metallic

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4
Q

Intermolecular forces

A 
Ion-Dipole
Dipole-Dipole
Hydrogen Bonds
Dipole-Induced Dipole
London Dispersion Forces
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5
Q

Ion-Dipole

A

Forces of attraction between an ion & polar molecule

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6
Q

Ion-dipole forces of attraction increase as

A

Radius of ion decreases
Charge of ion increases
Magnitude of dipole on polar molecule increases

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7
Q

Ion-Dipole between H2O and NaCl

A

Water molecules Na+

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8
Q

Dipole-dipole

A

Attractive forces between negative end of one polar molecule & positive end of another polar molecule

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9
Q

Molecules with dipole moments experience

A

Coulombic interactions when they are in close proximity to one another

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10
Q

When positive & negative dipoles line up well, attractive forces are …. and repulsive forces are …

A

Attractive forces are STRONGER

Repulsive forces are WEAKER

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11
Q

When positive & negative dipoles do not line up well, attractive forces are …. and repulsive forces are …

A

Attractive forces are WEAKER

Repulsive forces are STRONGER

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12
Q

Hydrogen Bonds

A

Occurs between a hydrogen that is covalently bonded to Fluorine, Oxygen, or Nitrogen and another F, O, or N with at least one lone pair

5 to 10 times STRONGER than other dipole-dipole attractions

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13
Q

Why are H-bonds so strong?

A

F - H, O - H, and N -H are VERY POLAR

ATOMS are very SMALL so PARTIAL CHARGES caused by differences in electronegativity are HIGHLY CONCENTRATED

LONE PAIR(s) on F, O, or N increases the already partially negative charge on atoms, creating a STRONGER ATTRACTION for slightly positive hydrogen

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14
Q

H-Bonds can occur

A

IN water molecules
Between water and methanol
In acetamide

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15
Q

Ethanal does not

A

form H-Bonds

Bonded to C not FON

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16
Q

Hydrogen bonds causes

A

Boiling points of elements in groups 5 & 6 to increase

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17
Q

PE of electrons associated with negative pole of a molecule DECREASES as it

A

Approach the POSITIVE pole of another molecule

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18
Q

Molecules with stronger dipoles have

A

Stronger attractions for one another which pull them closer together

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19
Q

What must be done in order to weaken or break these intermolecular forces?

A

Energy must be ADDED

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20
Q

Dipole-Induced Dipole

A

Attractions between a polar molecule and non-polar species or polarized molecule

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21
Q

Attractions resulting from dipole-induced dipole forces are STRONGER when

A

MAGNITUDE of the dipole in the polar molecule is LARGER

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22
Q

Molecules that have larger dipoles have a

A

GREATER ability to INDUCE a larger DIPOLE in a nonpolar molecule

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23
Q

Strength of dipole-induced dipole forces increases when

A

nonpolar molecule has a LARGER electron cloud & is MORE POLARIZABLE

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24
Q

Induced Dipole - Induced Dipole is also known as

A

London dispersion forces

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25
Q

London dispersion forces

A

Exist between ALL species: atoms, ions, non-polar molecules, and polar molecules

CONTRIBUTE to OVERALL force of attraction between all particles

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26
Q

What is the only IMF found in nonpolar molecules?

A

London Dispersion Forces

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27
Q

LDF are caused by

A

Coulombic interactions between temporarily induced dipoles of neighboring species that result from their electron distributions

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28
Q

Species with more electrons and larger electron clouds are

A

MORE POLARIZABLE

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29
Q

When moving down a group or constructing molecules with more atoms, the resulting species has

A

MORE ELECTRONS which translates to more polarizability

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30
Q

Molecular shape plays a role in

A

STRENGTH of LDF and physical state

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31
Q

Dispersion forces increase as contact area between molecules

A

INCREASES

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32
Q

Presence of pi-bonds

A

INCREASES polarizability

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33
Q

Why does pi-bonds increase polarizability?

A

Electrons are more DELOCALIZED and have more FREEDOM TO MOVE and assist with polarization

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34
Q

Rank the IMF from strongest to weakest

A 
ion-ion
ion-dipole
H-bonds
Dipole-Dipole
Ion-Induced dipole
Dipole-Induced dipole
London dispersion
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35
Q

When H-bonds are present, small molecules can have

A

very STRONG INTERMOLECULAR ATTRACTIONS

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36
Q

LDF is often the STRONGEST force of attraction between

A

LARGE molecules

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37
Q

Biomolecules

A

Long chains can be constructed through reaction

May contains hundreds of amino acids

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38
Q

H-bonds contributes to

A

Secondary structures in amino acids
a-helix
b-pleated sheet

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39
Q

IMF between R groups contributes to

A

Tertiary structures

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40
Q

Properties of Ionic Solids

A

Strong bonds
Cleave along planes
Soluble in polar solvents
Conduct electricity when dissolved or molten

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41
Q

Why do ionic solids have strong bonds?

A

Very strong Coulombic forces of attraction between cations & anions

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42
Q

Why do ionic solids cleave along planes?

A

– Brittle 3D structure
– Ions line up in a repetitive pattern that maximizes
attractive forces and minimizes repulsive forces.
– Not malleable or ductile

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43
Q

Why properties does strong bonds cause?

A
  • High melting points
  • Very hard
  • Low volatility
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44
Q

Solubility & conductivity of ionic solids

A

– Most are soluble in polar solvents.
– They conduct electricity only when molten or dissolved in a polar solvent, as the charged particles are free to
move

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45
Q

Solubility & conductivity of molecular solids

A

Most do not conduct electricity when molten or dissolved in water.
– The individual molecules have NO NET CHARGE, as they valence electrons are tightly held within covalent bonds and lone pairs.

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46
Q

Most molecular solids are held together by

A

inter-molecular forces which are much weaker than actual bonds

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47
Q

How does these properties compare between ionic and molecular?
Vapor pressures
Melting point
Boiling point

A

Molecular solids have
higher vapor pressure
lower melting point
lower boiling point

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48
Q

In a molecular solid, molecules are held close together in a regular pattern by intermolecular forces that

A

attempt to maximize attractions and minimize

repulsions.

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49
Q

Heat of Fusion (∆Hfus)

A

heat absorbed as 1 mole of a solid liquefies

energy REQUIRED to sever IMF between molecules

always POSITIVE and ENDOTHERMIC

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50
Q

∆Hfus for Ionic Compounds

A

As ionic bonds are much stronger than intermolecular forces, the ΔHfus values for ionic compounds are very large.

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51
Q

In a molecular liquid, intermolecular forces attempt to

A

maximize attractions and minimize repulsions.

molecules have more
freedom to move.

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52
Q

Heat of Vaporization (∆Hvap)

A

heat absorbed as 1 mole of a liquid becomes gaseous

energy REQUIRED to sever IMF between molecules

always POSITIVE and ENDOTHERMIC

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53
Q

When molecules leave the surface of a liquid to enter the gas phase, they

A

exert a pressure.

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54
Q

The vapor pressure exerted depends on the

A

rate of evaporation per unit area of the liquid’s

surface.

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55
Q

Rate of evaporation and vapor pressure

increase as

A

temperature increases.

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56
Q

When two substances are at the same temperature, the rate of evaporation and vapor pressure will be higher in the substance that
has

A

weaker intermolecular forces.

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57
Q

Boiling Points

A

A liquid boils when its vapor pressure equals the atmospheric pressure.

Evaporation occurs inside the liquid

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58
Q

Boiling points decrease as elevation

A

increases

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59
Q

Boiling points increase as the strength of IMF

A

increases

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60
Q

Sublimation

A

Solids can evaporate and have a vapor pressure.

Solids with high vapor pressures, have relatively
weak intermolecular forces

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61
Q

Vapor Pressures of Ionic Solids

A

Ionic compounds have very low vapor pressures and very high boiling points.

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62
Q

Covalent Network Solids

A

Composed of one or two non-metals held together by networks of covalent bonds

Very often contain carbon

High melting points
Very hard as atoms are covalently bonded

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63
Q

Examples of Covalent Network Solids

A 
Graphite
Diamond
SiO2 (quartz)
SiC (quartz)
Si (covalent network with itself)
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64
Q

Graphite

A 
• Weak π-bonds and London dispersion forces allow sheets to slide over one another (pencils).
• If hooked up to a potential 
difference, electrons will flow.
• High melting point, as
covalent bonds between
carbon in each layer are
relatively strong.
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65
Q

Water soluble proteins have

A

polar ‘R’ groups that

face out and non-polar ‘R’ groups that face in

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66
Q

Plastics

A

Non-polar

Held together by LDF

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67
Q

Properties of Synthetic Polymers

A

Generally flexible solids or viscous liquids

Heating increases flexibility/ allows molding

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68
Q

Particulate Characteristics of Solids

A

Limited motion
Close together
Held by IMF or bonds
Structure influenced by ability to pack together

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69
Q

Amorphous Solids

A
  • Random arrangement of particles
  • Particles have no orderly structure

• Macroscopic structures lack well defined faces
and shapes

• Many are mixtures of molecules that do not stack Many are mixtures of molecules that do not stack
up well together.

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70
Q

Examples amorphous solids

A

Glass

rubber

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71
Q

Crystalline Solids

A

• Atoms, ions, or molecules are arranged in an
orderly fashion that follows a pattern of
repetition in three dimensions

• Macroscopic structures usually have flat
surfaces that make definite angles to one
another.

72
Q

Examples crystalline solids

A

quartz

ionic solids

73
Q

Properties of Liquids

A

Constant motion & collisions
Close together
Motion influenced by strength of IMF

74
Q

Volume of Solid and Liquid Phases

A

• The solid and liquid phases for a particular
substance normally have similar molar
volumes.
• The density of particles is similar in both
phases.
• Ice has a slightly larger molar volume than
liquid water.
• Most solids have a slightly smaller molar
volume than their liquids.

75
Q

Pressure

A

Force / Area or N/m^2

76
Q

Gases exert pressure by

A

bouncing off surfaces

77
Q

Gas in a container exert pressures

A

evenly in all direction

78
Q

When you suck gas particles out of a can, the pressure

A

on the outside is greater than that on the inside so the can gets crushed

79
Q

Barometers

A

Measure Gas Pressure

80
Q

1 atm

A

760 mm Hg or torr

81
Q

Boyle’s Law

A

Relationship between
Pressure and Volume of Gases

V1P1 = V2P2

Volume is inversely proportional to pressure.

82
Q

Temperature

A

A measure of the average kinetic energy of atoms

and molecules in a system

83
Q

The Kelvin (K) temperature scale is proportional to

A

temperature

84
Q

Kinetic Energy of Gas Molecules

A

Translational, Rotational, Vibrational

Most of a gas particle’s KE is related to its translational velocity

85
Q

Charles’ Law

A

Relationship between
Temperature and Volume of Gases

V1/T1 = V2/T2

Volume is proportional to temperature.

86
Q

Properties of Gases

A 
Constant motion
Expand to fill volume
Low density
Highly compressible
Exert pressure
Form homogeneous mixtures
No definite shape or volume
87
Q

The volume of a gas doubles when the pressure

A

is halved

88
Q

Avogadro’s Principle

A

Equal volumes of different gases at the same

temperature and pressure contain equal numbers of particles

89
Q

Ideal Gas Equation

A

PV = nRT

P = pressure (atm)
V = volume (L)
n = number of moles
R = 0.0821 L•atm/K•mol
T = temperature (K)
90
Q

The Combined Gas Law

A

PiVi / Ti = PfVf / Tf

n remains constant while P,T, and T changes

91
Q

Dalton’s Law of Partial Pressures

A

For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone

Ptotal = P1 + P2 + P3 + …

92
Q

Mole Fraction

A

percent composition by moles of a single component in a mixture, represented in its decimal form

XA = nA / nA + nB + nC +nD + … +nZ

93
Q

You can find the partial pressure of any component

in a gas mixture by

A

multiplying the total pressure

by its mole fraction.

94
Q

When measuring
the volume of gas
collected, one must
first

A

line up the water levels inside

and outside the graduated cylinder.

95
Q

Why do you line up the water levels inside and outside the graduated cylinder?

A 
This ensures that
the pressure inside
the cylinder is
equal to the
atmospheric
pressure.
96
Q

KMT I states

A

Gases consist of particles (molecules and/or
individual atoms) that are in continuous random
movement.

97
Q

KMT II states

A

The total volume of all of the gas particles in a system is negligible when
compared with the total
volume of the system.

98
Q

KMT III states

A

Coulombic forces of attraction or repulsion do not exist between gas particles in a system.

99
Q

KMT IV states

A

Collisions experienced by gas particles are elastic.

- Kinetic Energy is conserved.

100
Q

KMT V

A

The average KE of the gas particles in a system is
proportional to the absolute temperature.

The gas particles in any system that is kept at the
same temperature will have the same average KE.

101
Q

The average KE of the particles in a system

increases as the temperature

A

increase.

102
Q

The total pressure exerted by the gas particles in a system is an average of

A

the particles with less KE and the particles with more KE

103
Q

Average velocity increases as mass

A

decreases

104
Q

All Real Gases Do Not Behave Ideally When…

A
  • Under high pressures (P > 5 atm) or low volume

* At low temperatures

105
Q

Volume adjustment for gases under high pressure

A

At high pressures, the volume occupied by particles is significant

Because in an ideal gas, volume occupied is assumed to be negligible and thus, zero, the ideal gas equation fails once volume is significant

106
Q

Ideal Gas Law assumes there are no

A

forces of attraction between gaseous particles

107
Q

When gas particles are very close together, the pressure they exert may be

A

less than what the Ideal
Gas equation would predict.

.

108
Q

Why do gas particle at low volume not exhibit `ideal gas behavior?

A

Neighboring molecules exert forces of attraction on one another when they are very close together.

Such forces pull a gas molecule in the direction
opposite to its motion.

This reduces the pressure resulting from impacts
with the walls of the container

109
Q

Gases do not behave ideally at low temperatures

A

• The Ideal Gas law assumes that gases experience no intermolecular forces of attraction.

• At low temperatures, gas particles move slower
and are closer together. Attractions between
molecules exist under these conditions.

110
Q

Condensation of gases may occur at

A

sufficiently low

temperatures and/or exceptionally high pressures.

111
Q

Why does condensation occur?

A

At low temp or high pressure, collisions between the particles may result in the particles sticking together due to IMF

112
Q

How does condensation affect ideal gas behavior?

A

• When a gaseous system is approaching the point
where condensation will occur, the forces of
attraction between gas particles are at a maximum.

• This results in the largest possible decrease in
measured pressure, and thus, a large deviation
from ideal behavior

113
Q

Suspension is also know as a

A

mechanical mixture

114
Q

Suspension

A

A heterogeneous mixture of two or more substances.

• Macroscopic properties are different at different Macroscopic properties are different at different
locations within the sample.
– The sizes, shapes, and concentrations of
particles can vary.

• In some cases, components can be separated through filtration.

115
Q

Example suspension

A

Sand & water

116
Q

Solution is also known as

A

Homogeneous mixture

117
Q

Solution

A

A homogeneous mixture of two of more substances

• Macroscopic properties do not vary within the
sample.
• Components cannot be separated by filtration.
• Components can be separated by methods that
alter intermolecular forces.
• No components are large enough to scatter visible
light.

118
Q

Solvent

A

The substance that is more plentiful in a

solution.

119
Q

Solute

A

The substance that is less plentiful in a solution.

120
Q

Saturated Solution

A

• When the solvent has dissolved the maximum
amount of solute possible at a certain temperature, and some solid particles remain
undissolved.

• This is an equilibrium system where solid
particles continually dissolve in the solvent and
dissolved particles fall out of solution.

121
Q

Miscible

A

Soluble in all proportions

Never become saturated

122
Q

Many ionic compounds dissolve in

A

polar

solvents. (ion-dipole)

123
Q

Polar solids, such as glucose, dissolve in

A

polar solvents. (dipole-dipole or H-bonds)

124
Q

Non-polar solids, such as mothballs,

A

dissolve in non-polar solvents. (dispersion)

125
Q

gas-liquid solutions includes

A

carbonated drinks

oxygen gas dissolves in water

126
Q

Gases are always

A

infinitely soluble in

one another.

127
Q

Gas-solid solutions

A

H2 gas can occupy the spaces between some

metal atoms such as iron, and palladium.

128
Q

Solid-solid solutions

A

Alloys

129
Q

Two methods for expressing concentration:

A

Molarity

Mole fraction

130
Q

Molarity

A

moles solute / liters solution

Change with temperature

131
Q

Mole fraction

A

moles A/moles A moles B + … + moles Z

Does not change with temperature

132
Q

Factors Affecting Solubility

A

Structure
• “Like dissolves Like”
– Polar dissolves polar; non-polar dissolves non-polar
Temperature
• Different rules for different types of solutions
Pressure
• Applies to Gas-Liquid solutions

133
Q

Like Dissolves Like

A

Substances that share similar intermolecular
interactions tend to be soluble or miscible in
one another.

134
Q

Smaller ions have stronger electric fields so they

A

drag more water molecules around with them

135
Q

If cation-anion attractions are stronger than ion-dipole attraction, the compound will .

A

not be soluble

136
Q

Chromatography paper

A

Non-polar carbon chains with -OH groups that can form H-bonds

137
Q

Max height on paper traveled by

A

non-polar solvent

138
Q

Solution contains

A

non-polar solvent, solute A, and solute B

139
Q

Stationary phase

A

Chromatogrpahy paper

140
Q

Mobile phase

A

solvent used

141
Q

The distances that the different solute particles travel up the paper depend on

A

their relative attractions for the moving solvent and the stationary paper.

142
Q

Polar solute will

A

not travel very far up the paper because the solute will form H bonds with the paper

143
Q

Non-polar solute will

A

travel further up the paper because the solute will have weak attractions for the paper and relatively strong attractions for the mainly non-polar solute.

144
Q

As temperature increases, solubility generally

A

increases

145
Q

Fractional Distillation

A

The separation of volatile liquids in a liquid-liquid

solution on the basis of boiling points

146
Q

The solubility of most gases decreases as

temperature

A

increases.

147
Q

Henry’s Law

A

The solubility of a gas is directly proportional to the partial pressure of that
gas above the solution

Pressure only affects the solubility of gases

148
Q

Hybrid Orbital Theory

A

• Atomic orbitals on the same atom combine in
order to form hybrids.
• Atomic orbitals on different atoms overlap in
order to form covalent bonds.
• Each atom in the compound retains its
associated orbitals and electrons.
• This theory correlates with observed bond
angles in molecules.

149
Q

Molecular Orbital (MO) Theory

A

Views a molecule as a whole instead of a
collection of individual atoms.
• MOs are similar to atomic orbitals.
– They both have specific energy levels.
– They both have specific sizes and shapes.
– They can both hold a maximum of two
electrons that spin in opposite directions.

150
Q

Bonding Orbital

A

Bonding Orbital

151
Q

Anti-Bonding Orbital

A
  • A MO that is higher in energy than any atomic orbitals from which it was derived
    – Electrons that occupy these orbitals cause instability
152
Q

• Non-Bonding Orbital

A

– A MO that is at the same energy level as the one
atomic orbital that it was derived from
– Electrons that occupy these orbitals do not cause
stability or instability.
– Orbitals that contain lone pairs

153
Q

Electromagnetic Spectrum

A 
radio waves
microwaves
infrared light
visisble light
ultraviolet rays
x-rays
gamma rays
154
Q

Spectroscopy

A

– A method of analysis which is based upon the
absorbance of electromagnetic radiation by matter.
– Used to acquire data pertaining to the structure of a
molecule or the concentration of a species.

155
Q

The intensity of light striking sample is equal to

A

intensity of light exiting sample and absorbed by the sample

156
Q

Ultraviolet/Visual Spectroscopy

A

• Examines transitions in electronic energy levels
– Is used to probe the electronic structure of
certain compounds
• Is used to determine concentrations of solutions
that contain certain compounds
• An absorption spectrometer is used to measure the
absorbance of a sample at wavelengths between
about 200 nm and 800 nm.
• The peaks represent wavelengths that correspond to the energy associated with possible electronic
transitions within the molecule.

157
Q

Colorless species can only absorb

A

UV light between

about 200 nm and 400 nm

158
Q

Colored species will always absorb light from

A

visual spectrum, but could also absorb UV light

159
Q

Beer-Lambert Law

A 
A = Ɛbc
A = absorbance
Ɛ = molar absorptivity (M-1cm-1)
b = path length of sample (cm)
c = concentration (M)
160
Q

Ɛ describes how

A

intensely a sample of ions or molecules absorbs light at a specific wavelength

161
Q

Infrared (IR) spectroscopy

A

• Examines transitions in molecular vibrations
– Is used to detect the presence of different
types of bonds and to identify molecules

162
Q

Covalent bonds have a

A

vibrational frequency within the IR region of electromagnetic spectrum

163
Q

Vibrational frequencies depend on the

A

mass of the atoms and the strength of the bonds

164
Q

IR Spectra can be used to identify

A

bond types, functional groups, and compounds

165
Q

Microwave spectroscopy

A

• Microwaves cause polar molecules to rotate.
• Each type of polar molecule has specific rotational frequencies that it can exhibit.
• The peaks in the microwave spectra below
correlate with the different rotational frequencies
for a specific polar molecule

166
Q

frequency

A

the number of times a wave repeats itself per second

167
Q

Quantum theory

A

Energy only increases in discrete units - by a full quantum or not at all

168
Q

Photoelectric Effect I

A

Highly intense low frequency light does
not eject any electrons, even if it shines on the
surface for several days.

169
Q

Photoelectric Effect II

A

When the threshold frequency is reached,

electrons are ejected immediately

170
Q

Photoelectric Effect III

A

Increasing the intensity of the light at a
frequency that will cause electrons to eject results
in a higher ejection rate. However, all ejected
electrons share the same velocity.

171
Q

Photoelectric Effect IV

A

Increasing the frequency of the light increases the velocity of the ejected electrons.
However, all ejected electrons share the same
velocity.

172
Q

Duality of light

A
  • it behaves like a wave, and

* it behaves like a particle

173
Q

When a photon is absorbed by an atom or

molecule, an electron

A

moves up one or more

energy levels

174
Q

The increase in energy is equal to

A

the energy of the photon that was absorbed
AND
the difference in energy between the two energy
levels

175
Q

When a photon is emitted from an atom or

molecule, an electron

A

moves down one or more

energy levels.

176
Q

The decrease in energy is equal to

A

the energy of the photon that was released
AND
the difference in energy between the two energy
levels

AP Chemistry (9 decks)

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