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CHEM 112 > CHEM 112 - MIDTERM I > Flashcards

CHEM 112 - MIDTERM I Flashcards

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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
Q

What is Hydrogen bonding?

A

When the electrostatic attraction between a partial (+) dipole on an H donor and a partial (-) dipole result in a hydrogen bond between two atoms

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26
Q

Hydrogen bond donor is _____

A

the atom that is covalently bonded to H (gives the H to another electronegative atom)

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27
Q

The N-H molecule bonds to an outside N atom, where N in the N-H bond is the Hydrogen bond _____.

A

Donor

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28
Q

Hydrogen bond acceptor is ____

A

the atom that is not covalently bonded to H (the electronegative atom that accepts the H bonding)

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29
Q

Hydrogen bonds are the _____

A

strongest dipole-dipole interactions

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30
Q

Hydrogen bonds occur between ____

A

molecules that have a hydrogen bond donor and a molecule with a hydrogen bond acceptor

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31
Q

When H bonds directly to F, O, or N - the bonding atoms acquire ____ partial charges

A

large (promotes dipole-dipole attractions between neighboring molecules)

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32
Q

How many hydrogen bonds can water form?

A

4

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33
Q

The ability to hydrogen bond ____ attraction between molecules

A

increases

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34
Q

Hydrogen bonds have ____ boiling points than those who do not hydrogen bond

A

higher

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35
Q

When a molecule can form hydrogen bonds, it will have stronger IM forces than a molecule of the ___ mass without an O-H present

A

same

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36
Q

What intermolecular forces does CH2Cl2 have?

A

dispersion; dipole-dipole

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37
Q

What IM forces does CH4 have?

A

dispersion

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38
Q

What IM forces does water have?

A

dispersion; dipole-dipole; H-bonding

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39
Q

Can CHO2H bond with itself and water?

A

Yes it can bond to both because it contains an Hydrogen acceptor and a Hydrogen donor

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40
Q

What are the two types of structures that solids can form?

A

Crystalline or amorphous structures

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41
Q

Amorphous solid

A

Lacks a regular arrangement of atoms

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42
Q

Examples of amorphous solids

A

glass, rubber, and plastics

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43
Q

Crystalline solid

A

Ordered internal structure

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44
Q

Examples of Crystalline solids

A

Table salt, diamonds, ice

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45
Q

Forces responsible for the stability of the crystal solid are _____

A

intermolecular forces

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46
Q

Lattice structure

A

the internal structure of a crystalline solid

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47
Q

Unit cells

A

basic repeating structural units in a lattice

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48
Q

Most of the unit cell’s atoms are shared by ____

A

neighboring unit cells

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49
Q

How many atoms are found on the corners of the unit cells?

A

1/8 of an atom ( #of atoms on the corners x 1/8)

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50
Q

How many atoms are found on the edges of the unit cells?

A

1/4 of an atom (# of atoms on the edges x 1/4)

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51
Q

How many atoms are found on the faces of the unit cells?

A

1/2 of an atom (# of atoms on the face x 1/2)

52
Q

How many atoms are found directly in the middle of the unit cells?

A

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
Q

Why do we care about unit cells?

A

The # of atoms in them can tell us the empirical formula of a chemical compound

54
Q

Phase changes are defined as _____

A

The transformation of one state of matter (or phase) to another

55
Q

When do phase changes occur?

A

When energy (heat) is added or removed from a substance

56
Q

Melting

A

solid -> liquid

57
Q

Vaporization

A

liquid -> gas

58
Q

Condensation

A

gas -> liquid

59
Q

Freezing

A

liquid -> solid

60
Q

Sublimation

A

solid -> gas

61
Q

Deposition

A

gas -> solid

62
Q

Endothermic phase changes are ___

A

melting (fusion), vaporization, sublimation (closer atom arrangement -> released)

63
Q

Exothermic phase changes are ____

A

condensation, freezing, deposition (looser atom arrangement -> closer)

64
Q

Phase diagrams

A

Demonstrates how the states of matter for a given molecule change as a function of its temperature and pressure

65
Q

When a molecule is changing from one phase to another ____

A

both phases coexist

66
Q

Sublimation curve

A

Where solid and gas are in equilibrium (where the solid and gas lines intersect)

67
Q

Fusion curve

A

Where solid and liquid are in equilibrium (where solid and liquid lines intersect) - establishes freezing point

68
Q

Vaporization curve

A

Where liquid and gas are in equilibrium (where liquid and gas lines intersect) - establishes boiling point

69
Q

Triple point

A

Where all three phases exist

70
Q

Critical point

A

Where the vaporization curve ends (in some phase diagrams)

71
Q

Single phase in the critical point

A

Supercritical fluid

72
Q

Solution

A

A homogenous mixture of 2 or more substances

73
Q

Solvent

A

Medium that the compound is dissolved in (present in the greatest quantity)

74
Q

Common example of a solvent

A

Water

75
Q

Solute

A

Compound that is dissolved in solvent

76
Q

Common example of solute

A

Salt/salt water

77
Q

Solubility

A

The amount of substance that will dissolve in a given amount of solvent at a specific temperature

78
Q

If the solute-solvent is stronger than solvent-solvent & solute-solute attraction, the solution process is ______.

A

Exothermic

79
Q

If the solute-solvent attraction is weaker than the solvent-solvent & solute-solute attraction, the solution process is ______.

A

Endothermic (but the solutes can still dissolve in a solvent because it is entropically favorable and increases disorder)

80
Q

Number #1 rule with the solution process?

A

Like dissolves like!

81
Q

Solutes with only dispersion forces will dissolve in a solvent with _____.

A

only dispersion forces

82
Q

The solution may or may not form if the solvent-solute interaction is ________ than solvent-solvent interactions.

A

weaker

83
Q

Ionic compounds usually dissolve only in _____ solvents.

A

polar

84
Q

Molecular compounds made of polar molecules usually dissolve best in ____ solvents.

A

polar

85
Q

Molecular compounds made of nonpolar molecules usually dissolve best in _____ solvents.

A

nonpolar

86
Q

Hydrogen bonding compounds dissolve best in ____ solvents.

A

hydrogen-bonding solvents

87
Q

Molarity (M)

A

amount of solute (mol)/volume of solution (L)

88
Q

Molality (m)

A

amount of solute (mol)/mass of solvent (kg)

89
Q

Mole fraction (x)

A

amount of solute (mol)/total amount of solute + solvent (mol)

90
Q

Mole percent (%)

A

amount of solute (mol)/total amount of solute + solvent (mol) x 100%

91
Q

Mass percent

A

mass solute/mass of total solution x 100%

92
Q

Unlike Molarity, Molality is _______ independent.

A

Temperature!!

93
Q

Solubility of solids increases with ______ in temperature

A

Increases

94
Q

Solubility of gases decreases with ____ in temperature

A

Increases

95
Q

Solubility of gases increases with ____ in pressure

A

Increases

96
Q

Solubility of liquids and solids is ______ by pressure

A

not impacted

97
Q

Solubility of gases decreases as you _____ the temperature

A

increase

98
Q

Henry’s Law

A

the quantitative relationship between gas solubility and pressure (cgas = kH * Pgas)

99
Q

cgas

A

concentration of the gas in molarity (M)

100
Q

kH

A

Henry’s law constant (unique to solute and associated with temp)

101
Q

Solubility of a gas in a liquid is ____ to the pressure of the gas over the solution

A

proportional

102
Q

Colligative properties

A

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
Q

4 colligative properties***

A

vapor pressure, freezing point, osmotic pressure, boiling point

104
Q

Raoult’s Law

A

Psolution = Xsolvent x Ppure solvent

Decrease vapor pressure of pure solvent or decrease in moles= decrease in pressure of the solution

105
Q

For a pure solvent the mole fraction will always be ____.

A

1

106
Q

The mole fraction of a solution will be ______

A

less than 1

107
Q

The vapor pressure of the solution will always be _____ than it would be if no solute was present***

A

less

108
Q

Psolution

A

Pressure a vapor reaches in a sealed container once dynamic equilibrium is established

109
Q

remember we will always be interested in the vapor pressure of the _____.

A

solvent

110
Q

Vapor pressure lowering impacts the ______ and ____ points

A

boiling; freezing

111
Q

Vapor pressure _____ the boiling point

A

increases

112
Q

Vapor pressure ____ the freezing point

A

decreases

113
Q

Decrease in vapor pressure = _____ in energy to change phases

A

increase

114
Q

The freezing point ____ when solute is added to pure solvent

A

decreases

115
Q

Freezing point depression

A

Delta T = Tf pure - Tf sol = i Kf m

116
Q

For a molecular compound, i will always be _____.

A

1 (non-electrolyte solution)

117
Q

For an ionic compound, i will be the amount of _____ in a compound.

A

ions - electrolyte solution (ie. NaCl will produce an i number of 2)

118
Q

Boiling point elevation

A

Delta T = Tb solution- Tb pure solvent = i Kb m

119
Q

Boiling point of water

A

100 Celsius

120
Q

Freezing point of water

A

0 Celsius

121
Q

Pure water will generally have the _____ boiling point

A

lowest

122
Q

Pure water will generally have the ____ freezing point

A

highest

123
Q

Van’t Hoff factor

A

the number of particles produced per formula unit (i)

124
Q

NaCl’s Van’t Hoff factor would be ___

A

i = 2 ; Na+ Cl- (2 particles produced)

125
Q

Measured Van’t Hoff factors are often ____ than you might expect due to ion pairing in solution

A

lower

126
Q

Pure water will generally have the _____ freezing point

A

highest

127
Q

Pure water will have the _____ boiling point

A

lowest

CHEM 112 (3 decks)

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