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CHEM 112 - MIDTERM I Flashcards

1
Q

Between dispersion forces, dipole-dipole, and hydrogen bonding, which one is the strongest intermolecular force?

A

Hydrogen bonding

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2
Q

What are intermolecular forces?

A

Noncovalent interactions

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3
Q

What are the intermolecular forces?

A

Dispersion forces
Dipole-dipole force
Hydrogen bonding
Ion-dipole
Ionic bonding

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4
Q

Why are intermolecular forces critical to enzyme function?

A

Hydrogen bonding and ion-dipole interactions stabilize the enzyme structure

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5
Q

Intermolecular forces are ____ than intramolecular forces

A

weaker

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6
Q

Intermolecular forces are interactions ____ molecules

A

between

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7
Q

What is the weakest of all intermolecular forces?

A

Dispersion forces

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8
Q

What are dispersion forces?

A

Result from the attraction of instantaneous partial positive and negative charges across molecules; they arise from temporary dipoles induced in atoms or molecules

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9
Q

What are dispersion forces?

A

Result from the attraction of instantaneous partial positive and negative charges across molecules; they arise from temporary dipoles induced in atoms or molecules

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10
Q

How are dispersion forces present between two Helium atoms?

A

They promote attractive interactions between induced partial positive charge in one He atom with the induced negative charge with the second He atom to form an overall neutral charge

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11
Q

How are the temporary partial positive and negative charges between electrons produced?

A

Asymmetric electron placement in the electron cloud via instantaneous dipole moments

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12
Q

How do dispersion forces present in heavier molecules?

A

Dispersion forces are stronger in heavier molecules because they have more electrons that can asymmetrical distributed to produce temporary dipoles

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13
Q

Between He and Xe, which atom has stronger dispersion forces?

A

Xe - because it has more electrons

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14
Q

Molecules experiencing stronger dispersion forces have _____ boiling points

A

Higher

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15
Q

Which halogen will have the highest boiling point: Cl2 or Br2 or I2?

A

I2 - it is the largest halogen in the group and has the most electrons (requires more energy to break the bond)

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16
Q

Place the following compounds in order of increasing strength of intermolecular forces: CH4, CH3CH2CH3, CH3CH3

A

CH4 < CH3CH3 < CH3CH2CH3

CH3CH2CH3 has the most atoms with the most electrons - will have the strongest dispersion force and highest boiling point

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17
Q

What are dipole-dipole forces?

A

Attractive interactions between polar molecules

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18
Q

Why are dipole-dipole forces stronger than dispersion forces?

A

The partial charges in molecules are greater than dispersion forces (and achieve dipole moments)

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19
Q

Molecules with ____ polarity have stronger dipole-dipole forces

A

greater (because they have more significant partial positive and partial negative charges)

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20
Q

Which molecule would have stronger dipole interactions: CH3Cl or CH2Cl2?

A

CH2CL2 will have stronger dipole interactions because Cl-C-Cl bonding will promote a stronger dipole moment than just C-Cl bonding

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21
Q

A molecule that has dipole-dipole forces must also have what?

A

Dispersion forces

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22
Q

Dispersion forces are ____ than dipole-dipole forces

A

Weaker

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23
Q

Do dipole-dipole forces occur in nonpolar molecules?

A

NO - must have at least one polar molecule

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24
Q

Do dipole-dipole forces occur in polar molecules?

A

YES - duh

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25
What is Hydrogen bonding?
When the electrostatic attraction between a partial (+) dipole on an H donor and a partial (-) dipole result in a hydrogen bond between two atoms
26
Hydrogen bond donor is _____
the atom that is covalently bonded to H (gives the H to another electronegative atom)
27
The N-H molecule bonds to an outside N atom, where N in the N-H bond is the Hydrogen bond _____.
Donor
28
Hydrogen bond acceptor is ____
the atom that is not covalently bonded to H (the electronegative atom that accepts the H bonding)
29
Hydrogen bonds are the _____
strongest dipole-dipole interactions
30
Hydrogen bonds occur between ____
molecules that have a hydrogen bond donor and a molecule with a hydrogen bond acceptor
31
When H bonds directly to F, O, or N - the bonding atoms acquire ____ partial charges
large (promotes dipole-dipole attractions between neighboring molecules)
32
How many hydrogen bonds can water form?
4
33
The ability to hydrogen bond ____ attraction between molecules
increases
34
Hydrogen bonds have ____ boiling points than those who do not hydrogen bond
higher
35
When a molecule can form hydrogen bonds, it will have stronger IM forces than a molecule of the ___ mass without an O-H present
same
36
What intermolecular forces does CH2Cl2 have?
dispersion; dipole-dipole
37
What IM forces does CH4 have?
dispersion
38
What IM forces does water have?
dispersion; dipole-dipole; H-bonding
39
Can CHO2H bond with itself and water?
Yes it can bond to both because it contains an Hydrogen acceptor and a Hydrogen donor
40
What are the two types of structures that solids can form?
Crystalline or amorphous structures
41
Amorphous solid
Lacks a regular arrangement of atoms
42
Examples of amorphous solids
glass, rubber, and plastics
43
Crystalline solid
Ordered internal structure
44
Examples of Crystalline solids
Table salt, diamonds, ice
45
Forces responsible for the stability of the crystal solid are _____
intermolecular forces
46
Lattice structure
the internal structure of a crystalline solid
47
Unit cells
basic repeating structural units in a lattice
48
Most of the unit cell's atoms are shared by ____
neighboring unit cells
49
How many atoms are found on the corners of the unit cells?
1/8 of an atom ( #of atoms on the corners x 1/8)
50
How many atoms are found on the edges of the unit cells?
1/4 of an atom (# of atoms on the edges x 1/4)
51
How many atoms are found on the faces of the unit cells?
1/2 of an atom (# of atoms on the face x 1/2)
52
How many atoms are found directly in the middle of the unit cells?
whatever # of atoms are in the middle = # of atoms found (ie. 1 atom in the middle = 1 atom overall because they are not shared by other unit cells)
53
Why do we care about unit cells?
The # of atoms in them can tell us the empirical formula of a chemical compound
54
Phase changes are defined as _____
The transformation of one state of matter (or phase) to another
55
When do phase changes occur?
When energy (heat) is added or removed from a substance
56
Melting
solid -> liquid
57
Vaporization
liquid -> gas
58
Condensation
gas -> liquid
59
Freezing
liquid -> solid
60
Sublimation
solid -> gas
61
Deposition
gas -> solid
62
Endothermic phase changes are ___
melting (fusion), vaporization, sublimation (closer atom arrangement -> released)
63
Exothermic phase changes are ____
condensation, freezing, deposition (looser atom arrangement -> closer)
64
Phase diagrams
Demonstrates how the states of matter for a given molecule change as a function of its temperature and pressure
65
When a molecule is changing from one phase to another ____
both phases coexist
66
Sublimation curve
Where solid and gas are in equilibrium (where the solid and gas lines intersect)
67
Fusion curve
Where solid and liquid are in equilibrium (where solid and liquid lines intersect) - establishes freezing point
68
Vaporization curve
Where liquid and gas are in equilibrium (where liquid and gas lines intersect) - establishes boiling point
69
Triple point
Where all three phases exist
70
Critical point
Where the vaporization curve ends (in some phase diagrams)
71
Single phase in the critical point
Supercritical fluid
72
Solution
A homogenous mixture of 2 or more substances
73
Solvent
Medium that the compound is dissolved in (present in the greatest quantity)
74
Common example of a solvent
Water
75
Solute
Compound that is dissolved in solvent
76
Common example of solute
Salt/salt water
77
Solubility
The amount of substance that will dissolve in a given amount of solvent at a specific temperature
78
If the solute-solvent is stronger than solvent-solvent & solute-solute attraction, the solution process is ______.
Exothermic
79
If the solute-solvent attraction is weaker than the solvent-solvent & solute-solute attraction, the solution process is ______.
Endothermic (but the solutes can still dissolve in a solvent because it is entropically favorable and increases disorder)
80
Number #1 rule with the solution process?
Like dissolves like!
81
Solutes with only dispersion forces will dissolve in a solvent with _____.
only dispersion forces
82
The solution may or may not form if the solvent-solute interaction is ________ than solvent-solvent interactions.
weaker
83
Ionic compounds usually dissolve only in _____ solvents.
polar
84
Molecular compounds made of polar molecules usually dissolve best in ____ solvents.
polar
85
Molecular compounds made of nonpolar molecules usually dissolve best in _____ solvents.
nonpolar
86
Hydrogen bonding compounds dissolve best in ____ solvents.
hydrogen-bonding solvents
87
Molarity (M)
amount of solute (mol)/volume of solution (L)
88
Molality (m)
amount of solute (mol)/mass of solvent (kg)
89
Mole fraction (x)
amount of solute (mol)/total amount of solute + solvent (mol)
90
Mole percent (%)
amount of solute (mol)/total amount of solute + solvent (mol) x 100%
91
Mass percent
mass solute/mass of total solution x 100%
92
Unlike Molarity, Molality is _______ independent.
Temperature!!
93
Solubility of solids increases with ______ in temperature
Increases
94
Solubility of gases decreases with ____ in temperature
Increases
95
Solubility of gases increases with ____ in pressure
Increases
96
Solubility of liquids and solids is ______ by pressure
not impacted
97
Solubility of gases decreases as you _____ the temperature
increase
98
Henry's Law
the quantitative relationship between gas solubility and pressure (cgas = kH * Pgas)
99
cgas
concentration of the gas in molarity (M)
100
kH
Henry's law constant (unique to solute and associated with temp)
101
Solubility of a gas in a liquid is ____ to the pressure of the gas over the solution
proportional
102
Colligative properties
properties of solutions which are modified by the presence of solute molecules (which disrupt (but not react with) the intermolecular forces of the solvent)
103
4 colligative properties***
vapor pressure, freezing point, osmotic pressure, boiling point
104
Raoult's Law
Psolution = Xsolvent x Ppure solvent Decrease vapor pressure of pure solvent or decrease in moles= decrease in pressure of the solution
105
For a pure solvent the mole fraction will always be ____.
1
106
The mole fraction of a solution will be ______
less than 1
107
The vapor pressure of the solution will always be _____ than it would be if no solute was present***
less
108
Psolution
Pressure a vapor reaches in a sealed container once dynamic equilibrium is established
109
remember we will always be interested in the vapor pressure of the _____.
solvent
110
Vapor pressure lowering impacts the ______ and ____ points
boiling; freezing
111
Vapor pressure _____ the boiling point
increases
112
Vapor pressure ____ the freezing point
decreases
113
Decrease in vapor pressure = _____ in energy to change phases
increase
114
The freezing point ____ when solute is added to pure solvent
decreases
115
Freezing point depression
Delta T = Tf pure - Tf sol = i Kf m
116
For a molecular compound, i will always be _____.
1 (non-electrolyte solution)
117
For an ionic compound, i will be the amount of _____ in a compound.
ions - electrolyte solution (ie. NaCl will produce an i number of 2)
118
Boiling point elevation
Delta T = Tb solution- Tb pure solvent = i Kb m
119
Boiling point of water
100 Celsius
120
Freezing point of water
0 Celsius
121
Pure water will generally have the _____ boiling point
lowest
122
Pure water will generally have the ____ freezing point
highest
123
Van't Hoff factor
the number of particles produced per formula unit (i)
124
NaCl's Van't Hoff factor would be ___
i = 2 ; Na+ Cl- (2 particles produced)
125
Measured Van't Hoff factors are often ____ than you might expect due to ion pairing in solution
lower
126
Pure water will generally have the _____ freezing point
highest
127
Pure water will have the _____ boiling point
lowest

CHEM 112 (3 decks)

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