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11th Grade > Intermolecular Forces and Properties > Flashcards

Intermolecular Forces and Properties Flashcards

1
Q

Types of Intermolecular Forces

A

London Dispersion Factors/Induced Dipole-Induced Dipole, Dipole-Dipole Attractions, Hydrogen Bonding, Ion-Dipole Attractions

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2
Q

Types of Intramolecular Forces

A

ionic, metallic, and covalent bonds

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3
Q

Which forces are strong? Intermolecular or Intramolecular

A

Intramolecular

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4
Q

London Dispersion Forces

A

Motion of Electrons create momentary dipoles
Occurs between all molecules and sometimes stronger than dipole-dipole forces

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5
Q

Polarizability

A

the ease with which the electron distribution in a molecule can be distorted (‘squashiness’ of the electron cloud)

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6
Q

+ Polarizability =

A

+ Dispersion Forces

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7
Q

What molecules have more polarizability?

A

Molecules with a greater molar mass have a greater # of electrons and, therefore, greater polarizability and more dispersion forces
Molecular shape also influences dispersion forces → if molecules can pack more tightly together in a long/cylindrical shape then they have greater dispersion forces than a spherical molecule

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8
Q

Dipole-Dipole Interactions

A

Attraction between the partially charged ends of polar molecules
+ polarity = stronger forces

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9
Q

Dipole-Induced Dipole Forces

A

A polar molecule induces or creates a momentary dipole in a neighboring nonpolar molecule

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10
Q

Hydrogen Bonds

A

Attraction between a hydrogen atom that is covalently bonded to a highly EN atom (N, O, F) and a nearby highly EN atom (N, O, F) in another molecule
Stronger than dipole-dipole but weaker than ion-dipole

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11
Q

Ion-Dipole Interactions

A

Attraction between an ion and a polar molecule
Increase in charge or polarity = increase in force
Smaller ions have stronger attractions with polar molecule

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12
Q

Comparing Relative Strengths of IMFs in two substances

A

If they have similar molar masses and shapes then dispersion forces are ~equal in the substances-more polar molecules have greater attraction

If they have very different molar masses, then dispersion forces are the most important attractive forces-the bigger the molecule, the greater the number of electrons, the greater the polarizability, the greater the dispersion forces

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13
Q

Ranking Strengths of Intermolecular Forces

A

Dispersion Forces, Dipole-Dipole, Hydrogen Bond, Ion Dipole Interactions

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14
Q

Vapor Pressure

A

The pressure exerted by a liquid’s vapor phase when the liquid and vapor states are in equilibrium
Increase in temp = Increase in vapor pressure

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15
Q

Volatile

A

Liquids that evaporate easily due to low IMFs

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16
Q

Increase in IMFs = increase…

A

Melting point, boiling point, surface tension, viscosity (resistance to flow), heat of vaporization (energy required to evaporate)

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17
Q

Increase in IMFs = decrease…

A

Vapor Pressure, volatility

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18
Q

Types of Solids

A

Metallic Solids, Ionic Solids, Covalent-network solids, Molecular Solids

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19
Q

Ionic Solids

A

Held together by electrostatic attraction between cations and anions held in a 3D lattice structure
High Melting and Boiling points
Brittle
Poor conductors of electricity unless they’re aqeuous

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20
Q

Molecular Solids

A

Held together by IMFs
Relatively low melting and boiling point due to the weak IMFs holding the molecules together
Poor conductors of electricity because the electrons are held tightly in covalent bonds

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21
Q

Covalent-Network Solids

A

Formed by atoms all held together in large networks by covalent bonds
Hard solids, high melting points due to being held in solid state by covalent bonds

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22
Q

Types of Covalent-Network Solids (8)

A

diamond, graphite, silicon, germanium, quartz (SiO­2), silicon carbide (SiC), boron nitride (BN), and boron carbide (BC)

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23
Q

Metallic Solids

A

Held together by delocalized ‘sea’ of collectively shared valence electrons
Great conductors of heat and electricity
Melting points vary depending on the element

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24
Q

Polymers

A

Contain long chains of atoms (usually carbon), where atoms within a chain are held together by covalent bonds but adjacent chains are held together by IMFs

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25
State of matter depends on...
balance between the kinetic energies of the particles the attractive forces between the particles
26
Ideal Gas Law
PV=nRT
27
STP
0 degrees Celsius, 1 atm, and 1 mole of gas occupies 22.4L
28
1atm =
760 mmHg=760 torr=101.3 kPa
29
Molar mass equation
M = dRT/p
30
Dalton's Law of Partial Pressure
P total = P1 + P2 +...
31
Calculating Vapor Pressure over water
P total = P H2O + P gas
32
Mole Fraction
Xa (mole fraction) = na (mole of substance)/ ntotal (total moles)
33
Calculating partial pressure
Pa (partial pressure of gas) =Xa (mole fraction)/ Ptotal (total pressure)
34
Kinetic-Molecular Theory of Gas
Gas particles are in constant, random motion Gas particles do not take up space (volume of particles is negligible) Gas particles do not attract or repel each other Average kinetic energy of gas particles remains constant (energy is conserved when collisions occur) Average kinetic energy of gas particles is directly proportional to temperature in Kelvin
35
Effusion
Diffusion of gas through a small hole
36
Diffusion
Diffusion of gas from a high concentration to a low concentration
37
The speed of gas is dependent on what?
Molar mass and temperature High temp = faster particle Small molar mass = faster particle
38
Graham's Law of Effusion
r1/r2 = square root (molar mass of 2/molar mass of 1)
39
Deviation from ideal gas law
High pressure and low temperature At low pressures the volume is small relative to the space between the particles so it is negligible. But at high pressures the idea gas law predicts a volume to low because it doesn't account for the volume of the particles. At high temperatures, the particles move to faster for there to be any attraction. But at low temperatures, the attraction impacts the pressure so the ideal gas law predicts a too large pressure.
40
PARTICLES will deviate the most if...
Particles have the potential to have greater IMFs (more important) Large particles
41
Molarity =
mols solute/L
41
Homogeneous vs. Heterogeneous mixtures
Homogeneous mixtures are uniform in physical appearance while Heterogeneous mixtures are not
42
When can dispersion forces be greater than dipole-dipole forces?
When the molecule has a big molar mass which leads to more electrons which leads to more polarizability which leads to more Dispersion forces
43
Increase in temperature =
Increase in vapor pressure
44
Using mole fraction to calculate partial pressure
Pa=XaPtotal or Xa=Pa/ptotal
45
When creating or analyzing particulate models, pay attention to...
Size of ions, Orientation of solute and solvent particles, and relative number of particles
46
Chromatography
Used to separate components in a solution due to relative attraction forces among the components in the mobile (solvent) and stationary (paper) phases In this, the solvent is usually polar and the paper is usually nonpolar
47
What are the three types of chromatography?
paper, thin-layer, and column chromatography
48
Polar components in a solution in chromatography will...
interact less with the stationary phase and will travel further by interacting more readily with the polar mobile phase
49
Less polar components of the solution in chromatography will.
interact more with the stationary phase and will not travel as far (by interacting more readily with the nonpolar stationary phase
50
Distillation
Used to separating components in a solution due to different strengths of intermolecular forces and therefore different boiling points
51
Distillate
The substance that boils at a lower temperature that vaporizes and condenses into another container
52
Substances with similar IMFs will tend...
To be more soluble (or miscible)
53
What dissolves in what? (Ionic compounds, polar molecules, Nonpolar molecules)
Ionic compounds dissolves better in polar solvents (ion-dipole forces) Polar molecules dissolves better in polar solvents (dipole-dipole forces) Nonpolar molecules will dissolve better in Nonpolar molecules (LDFs)
54
Interactions in a solution
Solute-solute Interactions, solvent-solvent Interactions, and solvent-solute interactions
55
Solutions form when...
solvent-solute interactions are greater than the SUM of the solute-solute and solvent-solvent interactions
56
Hydration
When a solvent is dissolving in water
57
Spectroscopy
Matter can absorb or emit radiation in different regions of the electromagnetic radiation spectrum, and those regions are associated with molecular motion and electronic transitions
58
Types of Spectroscopy
X-ray: Highest; Removes core electrons; tells us binding energy (helps identify the element) due to measuring how tightly e- are held; known as photoelectron spectroscopy Ultraviolet: High; Excites valence electrons and trasitions in energy levels; tells us the identity of a molecule or element; Visible Spectroscopy Visible Light: Medium; Excites valence electrons; identifies the concentration of a molecule; Visible spectroscopy Infrared: Low; Changes/Transitions the vibrations in the bonds; tells us the types of bonds/atoms/functional groups; vibrational spectroscopy Microwaves: Quite low; changes/transitions the rotations of the atoms in the bonds; tells us the location of hydrogen atoms within a molecules; rotational spectroscopy
59
Electromagnetic Radiation
Light
60
Speed of Light
3.00 * 10^8 m/s
61
Frequency
Number of cycles/waves that pass a given point in 1 second Symbols: f or nu (v thing) Units: s^-1 or Hz
62
Wavelength
Distance from 1 peak to the next Symbol: lambda (tent thing) Units: meters
63
C=
frequency times wavelength
64
Planck's Idea
Energy can be released or absorbed in specific amounts called quanta
65
A quantum of energy
Smallest amount of energy that can be emitted/absorbed
66
Planck's Constant (h(
6.626 * 10^-34 Js
67
E=
hv
68
Spectrometer
Measures the amount of light absorbed by a sample by comparing the intensity of light emitted from a light source with the intensity of the light emerging from the example
69
Beer's Law
Relates amount of light absorbed to the concentration of the substance absorbing the light
70
Dynamic Equilibrium
Condition in which two opposing processes occurring simultaneously at equal rates
71
A=
Ebc A= measured absorbance E (thing) = molar absorptivity constant (how intense a sample absorbs light of a particular wavelength) b= path length (distance light is traveling when it goes through the sample) c= concentration
72
Normal Boiling Points
Boiling point at 1 atm
73
Mean free path
Average distance traveled by a molecule between collisions

11th Grade (10 decks)

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