Unit 3 - Intermolecular Forces & Properties Flashcards
1
Q
Intramolecular forces
A
Bonds
2
Q
Intermolecular forces
A
Forces of attraction between molecules
3
Q
Intramolecular forces examples
A
Includes covalent, ionic, and metallic
4
Q
Intermolecular forces
A
Ion-Dipole Dipole-Dipole Hydrogen Bonds Dipole-Induced Dipole London Dispersion Forces
5
Q
Ion-Dipole
A
Forces of attraction between an ion & polar molecule
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
7
Q
Ion-Dipole between H2O and NaCl
A
Water molecules Na+
8
Q
Dipole-dipole
A
Attractive forces between negative end of one polar molecule & positive end of another polar molecule
9
Q
Molecules with dipole moments experience
A
Coulombic interactions when they are in close proximity to one another
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
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
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
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
14
Q
H-Bonds can occur
A
IN water molecules
Between water and methanol
In acetamide
15
Q
Ethanal does not
A
form H-Bonds
Bonded to C not FON
16
Q
Hydrogen bonds causes
A
Boiling points of elements in groups 5 & 6 to increase
17
Q
PE of electrons associated with negative pole of a molecule DECREASES as it
A
Approach the POSITIVE pole of another molecule
18
Q
Molecules with stronger dipoles have
A
Stronger attractions for one another which pull them closer together
19
Q
What must be done in order to weaken or break these intermolecular forces?
A
Energy must be ADDED
20
Q
Dipole-Induced Dipole
A
Attractions between a polar molecule and non-polar species or polarized molecule
21
Q
Attractions resulting from dipole-induced dipole forces are STRONGER when
A
MAGNITUDE of the dipole in the polar molecule is LARGER
22
Q
Molecules that have larger dipoles have a
A
GREATER ability to INDUCE a larger DIPOLE in a nonpolar molecule
23
Q
Strength of dipole-induced dipole forces increases when
A
nonpolar molecule has a LARGER electron cloud & is MORE POLARIZABLE
24
Q
Induced Dipole - Induced Dipole is also known as
A
London dispersion forces
25
London dispersion forces
Exist between ALL species: atoms, ions, non-polar molecules, and polar molecules
CONTRIBUTE to OVERALL force of attraction between all particles
26
What is the only IMF found in nonpolar molecules?
London Dispersion Forces
27
LDF are caused by
Coulombic interactions between temporarily induced dipoles of neighboring species that result from their electron distributions
28
Species with more electrons and larger electron clouds are
MORE POLARIZABLE
29
When moving down a group or constructing molecules with more atoms, the resulting species has
MORE ELECTRONS which translates to more polarizability
30
Molecular shape plays a role in
STRENGTH of LDF and physical state
31
Dispersion forces increase as contact area between molecules
INCREASES
32
Presence of pi-bonds
INCREASES polarizability
33
Why does pi-bonds increase polarizability?
Electrons are more DELOCALIZED and have more FREEDOM TO MOVE and assist with polarization
34
Rank the IMF from strongest to weakest
```
ion-ion
ion-dipole
H-bonds
Dipole-Dipole
Ion-Induced dipole
Dipole-Induced dipole
London dispersion
```
35
When H-bonds are present, small molecules can have
very STRONG INTERMOLECULAR ATTRACTIONS
36
LDF is often the STRONGEST force of attraction between
LARGE molecules
37
Biomolecules
Long chains can be constructed through reaction
May contains hundreds of amino acids
38
H-bonds contributes to
Secondary structures in amino acids
a-helix
b-pleated sheet
39
IMF between R groups contributes to
Tertiary structures
40
Properties of Ionic Solids
Strong bonds
Cleave along planes
Soluble in polar solvents
Conduct electricity when dissolved or molten
41
Why do ionic solids have strong bonds?
Very strong Coulombic forces of attraction between cations & anions
42
Why do ionic solids cleave along planes?
– Brittle 3D structure
– Ions line up in a repetitive pattern that maximizes
attractive forces and minimizes repulsive forces.
– Not malleable or ductile
43
Why properties does strong bonds cause?
* High melting points
* Very hard
* Low volatility
44
Solubility & conductivity of ionic solids
– 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
45
Solubility & conductivity of molecular solids
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.
46
Most molecular solids are held together by
inter-molecular forces which are much weaker than actual bonds
47
How does these properties compare between ionic and molecular?
Vapor pressures
Melting point
Boiling point
Molecular solids have
higher vapor pressure
lower melting point
lower boiling point
48
In a molecular solid, molecules are held close together in a regular pattern by intermolecular forces that
attempt to maximize attractions and minimize
| repulsions.
49
Heat of Fusion (∆Hfus)
heat absorbed as 1 mole of a solid liquefies
energy REQUIRED to sever IMF between molecules
always POSITIVE and ENDOTHERMIC
50
∆Hfus for Ionic Compounds
As ionic bonds are much stronger than intermolecular forces, the ΔHfus values for ionic compounds are very large.
51
In a molecular liquid, intermolecular forces attempt to
maximize attractions and minimize repulsions.
molecules have more
freedom to move.
52
Heat of Vaporization (∆Hvap)
heat absorbed as 1 mole of a liquid becomes gaseous
energy REQUIRED to sever IMF between molecules
always POSITIVE and ENDOTHERMIC
53
When molecules leave the surface of a liquid to enter the gas phase, they
exert a pressure.
54
The vapor pressure exerted depends on the
rate of evaporation per unit area of the liquid’s
| surface.
55
Rate of evaporation and vapor pressure
| increase as
temperature increases.
56
When two substances are at the same temperature, the rate of evaporation and vapor pressure will be higher in the substance that
has
weaker intermolecular forces.
57
Boiling Points
A liquid boils when its vapor pressure equals the atmospheric pressure.
Evaporation occurs inside the liquid
58
Boiling points decrease as elevation
increases
59
Boiling points increase as the strength of IMF
increases
60
Sublimation
Solids can evaporate and have a vapor pressure.
Solids with high vapor pressures, have relatively
weak intermolecular forces
61
Vapor Pressures of Ionic Solids
Ionic compounds have very low vapor pressures and very high boiling points.
62
Covalent Network Solids
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
63
Examples of Covalent Network Solids
```
Graphite
Diamond
SiO2 (quartz)
SiC (quartz)
Si (covalent network with itself)
```
64
Graphite
```
• 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.
```
65
Water soluble proteins have
polar ‘R’ groups that
| face out and non-polar ‘R’ groups that face in
66
Plastics
Non-polar
| Held together by LDF
67
Properties of Synthetic Polymers
Generally flexible solids or viscous liquids
| Heating increases flexibility/ allows molding
68
Particulate Characteristics of Solids
Limited motion
Close together
Held by IMF or bonds
Structure influenced by ability to pack together
69
Amorphous Solids
* 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.
70
Examples amorphous solids
Glass
| rubber
71
Crystalline Solids
• 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
Examples crystalline solids
quartz
| ionic solids
73
Properties of Liquids
Constant motion & collisions
Close together
Motion influenced by strength of IMF
74
Volume of Solid and Liquid Phases
• 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
Pressure
Force / Area or N/m^2
76
Gases exert pressure by
bouncing off surfaces
77
Gas in a container exert pressures
evenly in all direction
78
When you suck gas particles out of a can, the pressure
on the outside is greater than that on the inside so the can gets crushed
79
Barometers
Measure Gas Pressure
80
1 atm
760 mm Hg or torr
81
Boyle’s Law
Relationship between
Pressure and Volume of Gases
V1P1 = V2P2
Volume is inversely proportional to pressure.
82
Temperature
A measure of the average kinetic energy of atoms
| and molecules in a system
83
The Kelvin (K) temperature scale is proportional to
temperature
84
Kinetic Energy of Gas Molecules
Translational, Rotational, Vibrational
Most of a gas particle’s KE is related to its translational velocity
85
Charles’ Law
Relationship between
Temperature and Volume of Gases
V1/T1 = V2/T2
Volume is proportional to temperature.
86
Properties of Gases
```
Constant motion
Expand to fill volume
Low density
Highly compressible
Exert pressure
Form homogeneous mixtures
No definite shape or volume
```
87
The volume of a gas doubles when the pressure
is halved
88
Avogadro’s Principle
Equal volumes of different gases at the same
| temperature and pressure contain equal numbers of particles
89
Ideal Gas Equation
PV = nRT
```
P = pressure (atm)
V = volume (L)
n = number of moles
R = 0.0821 L•atm/K•mol
T = temperature (K)
```
90
The Combined Gas Law
PiVi / Ti = PfVf / Tf
n remains constant while P,T, and T changes
91
Dalton’s Law of Partial Pressures
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
Mole Fraction
percent composition by moles of a single component in a mixture, represented in its decimal form
XA = nA / nA + nB + nC +nD + … +nZ
93
You can find the partial pressure of any component
| in a gas mixture by
multiplying the total pressure
| by its mole fraction.
94
When measuring
the volume of gas
collected, one must
first
line up the water levels inside
| and outside the graduated cylinder.
95
Why do you line up the water levels inside and outside the graduated cylinder?
```
This ensures that
the pressure inside
the cylinder is
equal to the
atmospheric
pressure.
```
96
KMT I states
Gases consist of particles (molecules and/or
individual atoms) that are in continuous random
movement.
97
KMT II states
The total volume of all of the gas particles in a system is negligible when
compared with the total
volume of the system.
98
KMT III states
Coulombic forces of attraction or repulsion do not exist between gas particles in a system.
99
KMT IV states
Collisions experienced by gas particles are elastic.
| - Kinetic Energy is conserved.
100
KMT V
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
The average KE of the particles in a system
| increases as the temperature
increase.
102
The total pressure exerted by the gas particles in a system is an average of
the particles with less KE and the particles with more KE
103
Average velocity increases as mass
decreases
104
All Real Gases Do Not Behave Ideally When…
* Under high pressures (P > 5 atm) or low volume
| * At low temperatures
105
Volume adjustment for gases under high pressure
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
Ideal Gas Law assumes there are no
forces of attraction between gaseous particles
107
When gas particles are very close together, the pressure they exert may be
less than what the Ideal
Gas equation would predict.
.
108
Why do gas particle at low volume not exhibit `ideal gas behavior?
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
Gases do not behave ideally at low temperatures
• 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
Condensation of gases may occur at
sufficiently low
| temperatures and/or exceptionally high pressures.
111
Why does condensation occur?
At low temp or high pressure, collisions between the particles may result in the particles sticking together due to IMF
112
How does condensation affect ideal gas behavior?
• 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
Suspension is also know as a
mechanical mixture
114
Suspension
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
Example suspension
Sand & water
116
Solution is also known as
Homogeneous mixture
117
Solution
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
Solvent
The substance that is more plentiful in a
| solution.
119
Solute
The substance that is less plentiful in a solution.
120
Saturated Solution
• 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
Miscible
Soluble in all proportions
Never become saturated
122
Many ionic compounds dissolve in
polar
| solvents. (ion-dipole)
123
Polar solids, such as glucose, dissolve in
polar solvents. (dipole-dipole or H-bonds)
124
Non-polar solids, such as mothballs,
dissolve in non-polar solvents. (dispersion)
125
gas-liquid solutions includes
carbonated drinks
| oxygen gas dissolves in water
126
Gases are always
infinitely soluble in
| one another.
127
Gas-solid solutions
H2 gas can occupy the spaces between some
| metal atoms such as iron, and palladium.
128
Solid-solid solutions
Alloys
129
Two methods for expressing concentration:
Molarity
| Mole fraction
130
Molarity
moles solute / liters solution
Change with temperature
131
Mole fraction
moles A/moles A moles B + ... + moles Z
Does not change with temperature
132
Factors Affecting Solubility
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
Like Dissolves Like
Substances that share similar intermolecular
interactions tend to be soluble or miscible in
one another.
134
Smaller ions have stronger electric fields so they
drag more water molecules around with them
135
If cation-anion attractions are stronger than ion-dipole attraction, the compound will .
not be soluble
136
Chromatography paper
Non-polar carbon chains with -OH groups that can form H-bonds
137
Max height on paper traveled by
non-polar solvent
138
Solution contains
non-polar solvent, solute A, and solute B
139
Stationary phase
Chromatogrpahy paper
140
Mobile phase
solvent used
141
The distances that the different solute particles travel up the paper depend on
their relative attractions for the moving solvent and the stationary paper.
142
Polar solute will
not travel very far up the paper because the solute will form H bonds with the paper
143
Non-polar solute will
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
As temperature increases, solubility generally
increases
145
Fractional Distillation
The separation of volatile liquids in a liquid-liquid
| solution on the basis of boiling points
146
The solubility of most gases decreases as
| temperature
increases.
147
Henry's Law
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
Hybrid Orbital Theory
• 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
Molecular Orbital (MO) Theory
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
Bonding Orbital
Bonding Orbital
151
Anti-Bonding Orbital
- A MO that is higher in energy than any atomic orbitals from which it was derived
– Electrons that occupy these orbitals cause instability
152
• Non-Bonding Orbital
– 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
Electromagnetic Spectrum
```
radio waves
microwaves
infrared light
visisble light
ultraviolet rays
x-rays
gamma rays
```
154
Spectroscopy
– 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
The intensity of light striking sample is equal to
intensity of light exiting sample and absorbed by the sample
156
Ultraviolet/Visual Spectroscopy
• 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
Colorless species can only absorb
UV light between
| about 200 nm and 400 nm
158
Colored species will always absorb light from
visual spectrum, but could also absorb UV light
159
Beer-Lambert Law
```
A = Ɛbc
A = absorbance
Ɛ = molar absorptivity (M-1cm-1)
b = path length of sample (cm)
c = concentration (M)
```
160
Ɛ describes how
intensely a sample of ions or molecules absorbs light at a specific wavelength
161
Infrared (IR) spectroscopy
• Examines transitions in molecular vibrations
– Is used to detect the presence of different
types of bonds and to identify molecules
162
Covalent bonds have a
vibrational frequency within the IR region of electromagnetic spectrum
163
Vibrational frequencies depend on the
mass of the atoms and the strength of the bonds
164
IR Spectra can be used to identify
bond types, functional groups, and compounds
165
Microwave spectroscopy
• 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
frequency
the number of times a wave repeats itself per second
167
Quantum theory
Energy only increases in discrete units - by a full quantum or not at all
168
Photoelectric Effect I
Highly intense low frequency light does
not eject any electrons, even if it shines on the
surface for several days.
169
Photoelectric Effect II
When the threshold frequency is reached,
| electrons are ejected immediately
170
Photoelectric Effect III
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
Photoelectric Effect IV
Increasing the frequency of the light increases the velocity of the ejected electrons.
However, all ejected electrons share the same
velocity.
172
Duality of light
* it behaves like a wave, and
| * it behaves like a particle
173
When a photon is absorbed by an atom or
| molecule, an electron
moves up one or more
| energy levels
174
The increase in energy is equal to
the energy of the photon that was absorbed
AND
the difference in energy between the two energy
levels
175
When a photon is emitted from an atom or
| molecule, an electron
moves down one or more
| energy levels.
176
The decrease in energy is equal to
the energy of the photon that was released
AND
the difference in energy between the two energy
levels
