Unit 4 Test Flashcards
(129 cards)
1
Q
intermolecular forces
A
the attractive forces that exist among the particles that compose matter
2
Q
dispersion force (London dispersion forces or Van der Waals forces)
A
An intermolecular force (also referred to as London force) exhibited by all atoms and molecules that results from fluctuations in the electron distribution
3
Q
dipole–dipole force
A
An intermolecular force exhibited by polar molecules that results from the uneven charge distribution
4
Q
How do you determine if a molecule has dipole-dipole forces?
A
(1) determine if the molecule contains polar bonds; and (2) determine if the polar bonds add together to form a net dipole moment
5
Q
hydrogen bond
A
A strong attraction between the H atom on one molecule and the F, O, or N on its neighbor
6
Q
ion–dipole force
A
An intermolecular force that occurs when an ionic compound is mixed with a polar compound
7
Q
Rank the weakest to strongest intermolecular forces:
A
Dispersion, dipole-dipole, hydrogen bonding, ion-dipole bonding
8
Q
How do you determine if a covalent bond has hydrogen bonding?
A
(1) draw the molecule (2) check to see if hydrogen is directly bonded with oxygen, fluorine, or nitrogen
9
Q
How do you determine if a covalent bond has ion-dipole bonding?
A
(1) draw molecule (2) observe formal charges of atoms (3) check for polar charges (4) if an ionic compounds and polar compounds are present then there is a ion-dipole bonding
10
Q
How do you predict boiling point trends and solubility trends based on intermolecular forces for given covalent compounds?
A
(1) compare intermolecular forces, the stronger the force, the higher the boiling point (2) if the intermolecular forces are the same, compare the molecular mass; the higher the molecular mass, the higher the boiling point
11
Q
What type of intermolecular force is responsible for the solubility of ionic compounds?
A
Ion-dipole
12
Q
surface tension
A
The energy required to increase the surface area of a liquid by a unit amount; responsible for the tendency of liquids to minimize their surface area, giving rise to a membrane-like surface
13
Q
Surface tension decreases as intermolecular forces …
A
decreases
14
Q
viscosity
A
A measure of the resistance of a liquid to flow
15
Q
Viscosity increases as intermolecular forces …
A
increase
16
Q
As temperature rises, viscosity…
A
decreases
17
Q
As temperature rises, surface tension …
A
decreases
18
Q
Rank the intermolecular forces from highest surface tension and viscosity to lowest:
A
ion-dipole
hydrogen bonding
dipole-dipole
dispersion forces
19
Q
vaporization
A
the phase transition from liquid to gas
20
Q
condensation
A
the phase transition from gas to liquid
21
Q
sublimation
A
the transition phase from solid to gas
22
Q
deposition
A
the transition phase from gas to solid
23
Q
melting or fusion
A
the transition phase from solid to liquid
24
Q
freezing
A
the transition phase from liquid to solid
25
The rate of vaporization ... with increasing temperature.
increases
26
The rate of vaporization ... with increasing surface area.
increases
27
The rate of vaporization ... with decreasing strength of intermolecular forces.
increases
28
Endothermic or exothermic: vaporization
endothermic
29
Endothermic or exothermic: condensation
exothermic
30
Endothermic or exothermic: sublimation
endothermic
31
Endothermic or exothermic: deposition
exothermic
32
Endothermic or exothermic: fusion
endothermic
33
Endothermic or exothermic: freezing
exothermic
34
heat (or enthalpy) of vaporization (ΔHvap)
The amount of heat required to vaporize 1 mol of a liquid to a gas
35
heat (or enthalpy) of sublimation (ΔHsub)
The amount of heat required to sublime 1 mol of a solid to a gas
36
heat of fusion (ΔHfus)
The amount of heat required to melt 1 mol of a solid
37
heat (or enthalpy) of sublimation (ΔHsub)
The amount of heat required to sublime 1 mol of a solid to a gas
38
The sign of heat of vaporization is ...
positive because it is an endothermic reaction
39
vapor pressure
The partial pressure of a vapor in dynamic equilibrium with its liquid
40
The stronger the intermolecular forces, the ... the vapor pressure
lower
41
Formula for heat of sublimation:
heat of sublimination = heat of fusion + heat of vaporization
42
Formula for energy of ice warming:
q = m * C(ice) * delta(T)
43
Formula for energy of melting point:
q = n * delta(H_fusion)
n = number of moles (mol)
delta(H_fusion) = kJ/mol
44
Formula for energy of liquid water warming:
q = m * C_water * delta(T)
45
Formula for energy of boiling point:
q = n * delta(H_vaporization)
46
Formula for energy of steam warming:
q = m * C_steam * delta(T)
47
boiling point of a liquid
the temperature at which the pressure of the liquid's vapor equals the pressure of its surroundings
48
normal boiling point of a liquid
the temperature at which its vapor pressure equals the standard atmospheric pressure at sea level, which is 760 millimeters of mercury or one atmosphere
49
draw and label the typical phase diagram using the following: solid, liquid, gas, sublimation curve, vaporization curve, fusion curve, triple point, critical point, supercritical fluid
compare with notes
50
triple point
The unique set of conditions at which all three phases of a substance are equally stable and in equilibrium
51
critical point
The temperature and pressure above which a supercritical fluid exists
52
critical pressure
The pressure required to bring about a transition to a liquid at the critical temperature
53
supercritical fluid
a point in a phase diagram where the highly compressed gases which combine properties of gases and liquids in an intriguing manner
54
critical temperature
The temperature above which a liquid cannot exist, regardless of pressure
55
Formula for pressure
P = Force / Area
56
What is the average air pressure at sea level for pascal?
101,325 Pa
57
What is the average air pressure at sea level for psi?
14.7 psi
58
What is the average air pressure at sea level for Torr?
760 Torr
59
What is the average air pressure at sea level for inches of mercury?
29.92 in Hg
60
What is the average air pressure at sea level for atm?
1 atm
61
barometer
an instrument used to measure pressure (mmgHg)
62
manometer
An instrument used to determine the pressure of a gaseous sample, consisting of a liquid-filled U-shaped tube with one end exposed to the ambient pressure and the other end connected to the sample
63
As temperature rises, pressure ...
increases
64
Boyle’s law
The law that states that volume of a gas is inversely proportional to its pressure (V ∞ 1/P)
65
As pressure increases, volume ...
decreases
66
Charles’s law
The law that states that the volume of a gas is directly proportional to its temperature (V ∞ T)
67
Avogadro’s law
The law that states that the volume of a gas is directly proportional to its amount in moles (V ∞ n).
68
Formula for ideal gas law:
PV = nRT
P = pressure (atm or Kpa)
V = volume (l)
n = number of moles (mol)
R = universal gas constant (8.314 J / mol)
T = temperature (K)
69
Standard temperature and pressure
T = 0 degrees celsius and 273 K and P = 1.00 atm
70
molar volume
The volume occupied by 1 mol of a substance; the molar volume of an ideal gas at STP is 22.4 L
71
formula for density of gas:
molar mass/molar volume = g/L or PM/RT
72
Dalton’s law of partial pressures
The law stating that the sum of the partial pressures of the components in a gas mixture must equal the total pressure
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Formula for Dalton's law of partial pressures:
Ptotal=Pa+Pb+Pc---
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Formula for total pressure given mass or moles:
Ptotal=(ntotal)RT/V
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Formula of mole fraction:
Xa= na/ntotal
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Formula of partial pressure using mole fraction:
Pa=Xa x Ptotal
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Collecting gas over water, also known as water displacement, can help you calculate the ... of gas collected and the .... of the gas:
volume; pressure
78
Why is subtracting the water vapor pressure from the total pressure collected important?
because the collected gas pressure includes both the pressure from the gas itself and the pressure from the water vapor
79
What are the three basic ideas of the kinetic molecular theory?
(1) The size of a particle is negligibly small.
(2) The average kinetic energy of a particle is proportional to the temperature in kelvins.
(3) The collision of one particle with another (or with the walls of its container) is completely elastic.
80
How does kinetic molecular theory relate to Newton's 2nd law?
According to kinetic molecular theory, a gas is a collection of particles in constant motion. The motion results in collisions between the particles and the surfaces around them. As each particle collides with a surface, it exerts a force upon that surface. The result of many particles in a gas sample exerting forces on the surfaces around them is a constant pressure.
81
How does kinetic molecular theory relate to Boyle's law?
According to kinetic molecular theory, if we decrease the volume of a gas, we force the gas particles to occupy a smaller space. As long as the temperature remains the same, the number of collisions with the surrounding surfaces (per unit surface area) must necessarily increase, resulting in a greater pressure.
82
How does kinetic molecular theory relate to Charles's law?
According to kinetic molecular theory, when we increase the temperature of a gas, the average speed, and thus the average kinetic energy, of the particles increases. Since this greater kinetic energy results in more frequent collisions and more force per collision, the pressure of the gas increases if its volume is held constant (Gay-Lussac’s law).
83
How does kinetic molecular theory relate to Dalton's law?
According to kinetic molecular theory, the particles have negligible size and they do not interact. Consequently, the only property that distinguishes one type of particle from another is its mass. However, even particles of different masses have the same average kinetic energy at a given temperature, so they exert the same force upon collision with a surface.
84
How does kinetic molecular theory relate to Avogadro's law?
According to kinetic molecular theory, when we increase the number of particles in a gas sample, the number of collisions with the surrounding surfaces increases. The greater number of collisions results in a greater overall force on surrounding surfaces; the only way for the pressure to remain constant is for the volume to increase so that the number of particles per unit volume (and thus the number of collisions) remains constant.
85
diffusion
The process by which gas molecules spread out in response to a concentration gradient
86
effusion
The process by which a gas escapes from a container into a vacuum through a small hole
87
What are the factors that affect diffusion and effusion?
temperature, molar mass, concentration gradient, surface area, distance, density
88
Higher temperature increases the kinetic energy of gas particles, causing them to move ____ and diffuse or effuse more ____
faster; quickly
89
Lighter gases diffuse and effuse ... than heavier gases
faster
90
Diffusion occurs from areas of ... . concentration to areas of ... concentration. The larger the ... ..., the faster the diffusion
high, low; concentration difference
91
A larger surface area allows for ... diffusion
faster
92
The molecular speed of a gas is ... proportional to its speed and ... proportional to its molar mass
directly, inversely
93
The van der Waals equation is a correction to the ideal gas law that accounts for the ... of gas molecules and the ... forces between them.
volume, attractive
94
How are solutions formed?
A solution forms when a substance dissolves, or breaks apart, into another substance. The substance that dissolves to form a solution is called a solute.
95
solute
The minority component of a solution
96
solution
A homogeneous mixture of two substances
97
solvent
The majority component of a solution
98
How do you identify the solute and solvent if the components of solutions are given?
The compound being dissolved is the solute, while the compound (usually a liquid) that is dissolving is the solvent
99
What molecule is bigger? solute or solvent?
solvent
100
Does a solution form? Solvent-solute interactions > solvent-solvent and solute-solute interactions
solution forms
101
Does a solution form? Solvent-solute interactions = solvent-solvent and solute-solute interactions
solute forms
102
Does a solution form? Solvent-solute interactions < solvent-solvent and solute-solute interactions
solution may or may not form. depending on relative disparity
103
saturated solution
a solution in which the dissolved solute is in dynamic equilibrium with the solid (undissolved) solute
104
unsaturated solution
a solution containing less than the equilibrium amount of solute
105
saturated or unsaturated? If you add additional solute to a ... solution, it will not dissolve
saturated
106
saturated or unsaturated: if you add additional solute to a ... solution it will dissolve
unsaturated
107
supersaturated
a solution containing more than the equilibrium amount of solute that causes it to be unstable, however, it can exist for an extended period of time
108
What's the driving force for solution formation?
tendency toward greater entropy (a state function that measures how spread out the energy of atoms and molecules is in a system) (or greater dispersal)
109
solubility
the amount of a substance is the amount of the substance that dissolves in a given amount of solvent
110
What determines the solubility of one substance in another substance?
types of intermolecular forces that exist between the substances as well as within each substance
111
How do you determine the overall enthalpy changes for solution formation?
(1) separation of the solute particles, (2) separation of the solvent particles, and (3) mixing the solute and solvent particles
112
dynamic equilibrium
a solution that occurs when the rates of dissolution and recrystallization in a solution are equal
113
With increasing temperature, the solubility of most solids in water ...
increases
114
The solubility of gases in water generally ... with increasing temperature, but it ... with increasing pressure
decreases, increases
115
formula for mass percentage
g solute / g solution
116
formula for molarity (M)
mol solute/L solution
117
formula for molality (m)
mol solute / kg of solvent
118
mole fraction
mole of solute / (mol solute + mole solvent)
119
formula for finding the solubility of gas
Henry's Law: S = k *p
S = solubility of gas
k = henry's constant
p = partial pressure
120
colligative property
a property that depends on the number of particles dissolved in solution, not on the type of particle
121
what are the four colligative properties?
vapor pressure, freezing point depression, boiling point elevation, and osmotic pressure
122
freezing point depression
a colligative property observed in solutions that results from the introduction of solute molecules to a solvent
123
boiling point elevation
the boiling point of a liquid will be higher when another compound is added, meaning that a solution has a higher boiling point than a pure solvent
124
osmotic pressure
the minimum pressure which needs to be applied to a solution to prevent the inward flow of its pure solvent across a semipermeable membrane
125
formula for boiling point elevation
delta_T_b = K_b *Cm, delta_T_b = T_solution -T_b solvent
change in boiling point elevation, boiling point elevation constant, Cm is boiling point elevation
126
formula for electrolytes
[(deltaTb = Kb x iCm), [(deltaTb = Tb solution – Tb solvent)]
127
formula for melting point depression
n (deltaTf = Kf x Cm), [(deltaTf = Tf solvent – Tf solution)]
128
formula for electrolytes
(Tf = Kf x iCm), [(Tf = Tf solvent – Tf solution)]
129
formula for osmotic pressure
=MRT
