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General chemistry II > Unit 1 (ch 6,10,11) > Flashcards

Unit 1 (ch 6,10,11) Flashcards

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1
Q

Characteristics of gas phase

A
  • no defined shape or volume
  • compressible
  • change volume with temp
  • miscible
  • less dense
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2
Q

Atmospheric pressure

1 ATM=

A

P ATM = force/area

1 ATM= 760 mmhg

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3
Q

1 ATM = how many pascals

A

101,325

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4
Q

What happens if area decreases?

A

Pressure increases

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5
Q

Boyles law

A

Related to pressure and volume
P = 1/V
PV= constant
P1V1=P2V2

pressure increases, volume decreases

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6
Q

Charles law

A

Related to volume and temperature
V=T
V1/T1= V2/T2

*volume increases, temperature increases

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7
Q

Avogadro’s law

A

Related to volume and number of moles
V1/n1= V2/n2

*as volume increases, number of moles increases

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8
Q

Amontons law

A

Related to pressure and temperature
P1/T1 = P2/T2

as pressure increases, temperature increases

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9
Q

Ideal gas law

A

PV=nRT

P1V1/n1T1 = P2V2/n2T2

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10
Q

Standard temp

A

0 degrees C

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11
Q

Standard pressure

A

1 bar/1atm

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12
Q

Standard volume

A

22.4 L/mol

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13
Q

Does gas have a greater density or less density than other forms of matter?

A

less

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14
Q

at STP density =

A

molar mass/molar volume

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15
Q

density in relation to PV=nRT

A

PM/RT = m/v = d (M= molar mass, m = mass)

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16
Q

Given that density of gases depends on their molar mass, which is the least dense?
Cl2, He, H2, N2

A

H2- least molar mass

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17
Q

highest molecular weight = most or least dense?

A

most

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18
Q

increase temperature, density (blank)

A

decreases

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19
Q

decrease temperature, density (blank)

A

increases

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20
Q

Dalton’s law of partial pressures

A

Ptot = P1 + P2 + P3…

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21
Q

mole fraction

A

moles of one element/total moles of all elements

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22
Q

Partial pressure of a gas =

A

mole fraction x Ptotal

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23
Q

Kinetic molecular theory of gases list

A

1) gas molecules are tiny compared to the volume they occupy
2) move constantly and randomly
3) motion associated with average KE (same temp= same KE)
4) collisions with other molecules and walls of container are elastic (T and V constant)
5) assume no attraction or repulsion between gas molecules

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24
Q

with gases, smaller V means (blank) collisions and (blank) pressure

A

increase, increase

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25
Q

More molecules means (blank) collisions and (blank) pressure

A

increase, increase

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26
Q

Higher temperature means (blank) collisions and (blank) pressure

A

increase, increase

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27
Q

Molecular speed and kinetic energy

A

KE = 1/2 mu^2

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28
Q

Three important values

  • Um
  • Uavg
  • Urms
A

(most common speed)
(average speed)
(root mean square speed)

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29
Q

larger gas moves faster or slower?

A

slower

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30
Q

KE avg is the same for different gases at the same (blank)

A

temperature

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31
Q

Graham’s law is related to

A

effusion and diffusion

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32
Q

Van der Waal’s equation

A

(P + n^2a/v^2) x (v-nb) = nRT

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33
Q 
solids have (blank) interparticle forces 
KE (blank) attraction
A

strong,

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34
Q 
liquids have (blank) interparticle forces
KE (blank) attraction
A

medium, >

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35
Q 
gases have (blank) interparticle forces
KE (blank) attraction
A

weak,&raquo_space;

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36
Q

strong forces = solid with (blank) boiling point

A

high

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37
Q

weak forces = gas with (blank) boiling point

A

low

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38
Q

ion-ion interactions

A

ionic compounds

strongest

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39
Q

E = (blank)/(blank)

A

E = (Q1Q2)/d (smaller distance = stronger)

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40
Q

stronger interactions with (blank) distance

A

smaller

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41
Q

Interactions between polar molecules include and they relate to differences in (blank)

A

ELECTRONEGATIVITY

  • ion-diple
  • diploe-diple
  • —hydrogen
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42
Q

ion-dipole interactions

A

ions and polarized molecules

ex) water and salt

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43
Q

Dipole-Dipole

A

weaker, less organized

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44
Q

Hydrogen

A

special type of dipole-dipole
O-H, N-H, or F-H
H bonded to highly electronegative element
increase boiling point

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45
Q

Dispersion forces

A

london dispersion, induced-dipole induced dipole
occur between all molecules
weakest
change in electron distribution
larger molecules = more force (more polarized)

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46
Q

Which is the most polarizable? I2, Br2, Cl2, F2

A

I2

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47
Q

Why do strenghts of London forces increase with size?

A

larger molecules have more electrons and therefore, more distortions and a bigger force
-nonpolar liquid (oil)= high boiling

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48
Q

Van der Waals forces

A

non-ideal gases

(P + n^2a/v^2)(v-nb)=nRT

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49
Q

Thinking about the factors that influence Van der Waals constants, which do you expect to have the larges a constant? He, H2O, N2, NO

A

H2O

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50
Q

key point with polarity and solubility

A

LIKE DISSOLVES LIKE

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51
Q

O2 has (blank) solubility in H2O

A

limited

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52
Q

increase temp, (blank) gas solubility

A

decrease

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53
Q

fish example

A

O2 less soluble with heat, fish die in warm water

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54
Q

low pressure (high elevation) = (blank) gas solubility

A

lower

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55
Q

Henry’s law

A

Cgas= KH-Pgas

concentration, constant, partial pressure

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56
Q

hydrophobic vs hydrophillic

A

fearing vs loving

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57
Q

what two things influence the phase of matter?

A

temperature and pressure

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58
Q

solid CO2 =

A

dry ice

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59
Q

qualities of water

A

higher melting and boiling point
forms strong H-bonds
surface tension

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60
Q

cohesive forces

A

molecules sticking together

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61
Q

adhesive forces

A

molecules sticking to other things

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62
Q

stronger forces = (blank) viscous

A

more

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63
Q

pentane vs car oil- which is more viscous

A

car oil

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64
Q

What is more viscous, freshwater or seawater?

A

seawater

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65
Q

cold water = (blank) dense

A

most

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66
Q

solvent dissolves (blank)

A

solute

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67
Q

breaking bonds example

A

NaCl (ionic)

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68
Q

making bonds example

A

NaCl in H2O (ion-dipole)

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69
Q

delta H solution =

A

delta H ion-ion + delta H dipole-dipole + delta H ion-dipole

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70
Q

delta H solution = break + make

A

delta H ion-ion (break) + delta H hydration (make)

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71
Q

lattice energy

A

U = K (Q1Q2)/d

-more negative U is harder to dissolve

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72
Q

lattice energy trend

A

smaller ion (top right)= larger negative energy

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73
Q

which has a higher lattice energy RbF or LiF?

A

LiF

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74
Q

Born-Haber cycle

A

series of chemical steps to make ionic solid from elements

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75
Q

delta H solution = delta H hydration (+ or -? U)

A

-

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76
Q

rate of vaporization depends on

A

pressure, temp, surface area, strength of intermolecular forces

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77
Q

temperature increases, vapor pressure (blank)

A

increases

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78
Q

delta H hydration for H2O is positive or negative?

A

negative

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79
Q

in general, if solute is nonvolatile, vapor pressure (blank)

A

decreases

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80
Q

nonvolatile examples

A

sugar and salt

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81
Q

Which of the following would not lower the vapor pressure of H2O?

  • Acetic acid (118.1 C)
  • Ethylene glycol (197)
  • Glycerol (290)
  • Methanol (64.5)
A

methanol

82
Q

Raoult’s law

A

Psolution= mole fraction solvent x P solvent

83
Q

clausius-clapeyron equation

A

ln (Pvap T1/Pvap T2) = delta H vap/ R (1/T2 -1/T1)

84
Q

vapor pressures of mixtures equation

A

Ptotal=mole fraction x vapor pressure + (same)

85
Q

Colligative properties of solutions

A
  • freezing point depression
  • boiling point elevation
  • density increase
  • osmosis and osmotic pressure
86
Q

molality

A

nsolute/kg solvent

87
Q

boiling point elevation

A

delta Tb=kb x m

88
Q

freezing point elevation

A

delta Tf = kf x m

89
Q

van’t hoff factor (i)

A

nonelectrolytes vs electrolytes

90
Q

solvent moves from (blank) solute to (blank)

A

low to high

91
Q

Osmotic pressure equation

A

pi= i x M (moles/L) x R x T (kelvin)

92
Q

ethylene glycol in water is

A

antifreeze

93
Q

What differentiates gases from the other phases?

A

no shape/volume, compressible, miscible, less dense

94
Q

Patm=

A

F/A

95
Q

What are the gas laws?

A

Boyles (P and V)
Charles (V and T)
Avogadros (V and n)
Amontons (P and T)

96
Q

Ideal gas law

A

PV = nRT

97
Q

standard conditions for the ideal gas law

A

0 degrees C and 1 atm/bar

98
Q

1 mol ideal gas = blank L

A

22.4

99
Q

gases in reactions compare (blank)

A

V and P

100
Q

gas density

A

less dense than other faces

d= molar mass/molar volume

101
Q

density =

A

P x molar mass/ RT

102
Q

As temperature increases, density (blank)

A

decreases

103
Q

how can we use density to determine molar mass?

A

molar mass = dRT/P

104
Q

Dalton’s law

A

each gas in a mixture has a partial pressure

Ptotal= sum of partial pressures (P1 + P2 + P3)

105
Q

moles of gas

A

ntotal= n1+n2…

106
Q

mole fraction Xx

A

nx/ntotal (moles of the component/ total moles)

107
Q

Partial pressure =

A

mole fraction x total pressure (Px = (Xx)(Ptotal)

108
Q

For any given mixture, the sum of the mole fractions =

A

1

109
Q

Kinetic molecular theory of gases

A
  • gas molecules tiny compared to volume they occupy
  • move constantly and randomly
  • motion associated with average kinetic energy
  • collisions are elastic
  • no attraction between molecules
110
Q

collisions are what generates (blank)

A

pressure

111
Q

smaller volume = (blank) collisions

A

more

112
Q

smaller volume = (blank) pressure

A

increase

113
Q

more molecules = (blank) collisions = (blank) pressure

A

more, increased

114
Q

KE equation

A

1/2 m u^2 (m= mass, u = speed)
um= most common speed
uavg= average speed
urms= root mean squared speed

115
Q

urms

A

speed of particle with KEavg

sqrt(3RT/molar mass)

116
Q

KEavg

A

1/2 m (urms)^2

117
Q

heavier molar mass = (blank) speed

A

slower

118
Q

Graham’s law

A

gases move from high pressure to low pressure

119
Q

rate is inversely proportional to the (blank)

A

square root of molar masses

r1/r2 = sqrt (molar mass2/molar mass1)

120
Q

Van der waal’s equation

A

(P + n^2a/v^2)(v-nb) = nRT

121
Q

solid- KE (blank) intermolecular attraction

A
122
Q

liquid- KE (blank) intermolecular attraction

A

>

123
Q

gas- KE (blank) intermolecular attraction

A

> >

124
Q

strong forces and high melting = solid, liquid, or gas?

A

solid

125
Q

strength of interparticle/intermolecular intereactions influences (blank)

A

phase, melting, and boiling points

126
Q

weak forces and low boiling = solid, liquid, or gas

A

gas

127
Q

types of interactions

A 
ion-ion (E Q1Q2/d)
ion-dipole
dipole-dipole
hydrogen
dispersion forces
128
Q

dispersion forces explained

A

nonpolar molecules form temporary or induced dipoles

129
Q

induced dipole is the (blank)

A

temporary shift in the electron distribution around the atoms/bonds

130
Q

(blank) can influence the strength of dispersion forces

A

shape

131
Q

what molecules stack well and what don’t?

A

linear, spherical

132
Q

How can dispersion forces become strong?

A

added together

133
Q

Van der Waal’s forces

A

nonideal gases interact

134
Q

Intermolecular forces also explain

A

solubility

135
Q

If we mix H2O and MeOH, they both have what type of force? What does this mean?

A

dipole-dipole/hydrogen bonds

they can form new dipole-dipole interactions with each other, increase solubility

136
Q

If we mix H2O with octane, what forces are present? What does this mean?

A

dispersion, you have to break the hydrogen bonds and dispersion forces but only left with dispersion in end= unfavorable

137
Q

Gas solubility in H2O - increased T =

A

decreased solubility (gases gain more KE, escape solvent)

138
Q

Solubility in H2O depends on

A

gas partial pressure- Henry’s law

139
Q

What molecules form cell membranes?

A

molecules that interact through multiple types of forces

- polar and non polar ends

140
Q

cell membranes are also known as

A

phosopholipids

141
Q

What do phase diagrams do?

A

tell us what phases should be present at a given T and P and where transitions between the phases should occur

142
Q

supercritical region

A

a region where a compound behaves as sort of a gas-liquid hybrid

143
Q

Why is water amazing?

A

strong hydrogen bonds

144
Q

Solutes affect solution properties such as

A

melting point, boiling point, and vapor pressure

145
Q

When we dissolve a solute in a solvent, we are doing what?

A

breaking some bonds and forming new ones

146
Q

Delta H solution =

A

delta H hydration - U (lattice energy)

147
Q

lattice energy =

A

k(Q1Q2)/d

148
Q

Born Haber cycle

A

breaking a complicated chemical reaction into its basic steps and analyzing the energy for each of those steps

149
Q

delta H hydration (NaCl) can be broken into

A

delta H hydration Na = delta H hydration Cl

150
Q

Vapor pressure

A

the pressure exerted by a gas above a liquid in a closed system when equilibrium between evaporation and condensation is reached

151
Q

Vapor pressure can tell us (blank) and increases with (blank)

A

fast; temp

152
Q

If we dissolve a nonvolatile solute, vapor pressure

A

drops

153
Q

vapor pressure of pure water is (blank) seawater

A

greater than

154
Q

Raoult’s law

A

Psolution = Xsolvent x Psolvent (mole fraction x vapor pressure of solvent)

155
Q

mole fraction is the

A

ratio of moles of one component of a mixture to the entire mixture (nx/ntotal)

156
Q

Colligative properties

A

depend on concentration of solute

157
Q

Mixtures of volatile solutes can be found using what equation?

A

clausius-clapeyron

158
Q

Clausius- Clapeyron

A

ln(PvapT1/PvapT2) = delta Hvap/R(1/T2-1/T1)

159
Q

In a volatile mixture, Ptotal =

A

X1P1 + X2P2

160
Q

Colligative properties

A

molality, boiling point elevation, freezing point depression, osmosis

161
Q

molality

A

nsolute/kg solvent

162
Q

boiling point elevation

A

delta Tb = kb x m (i)

163
Q

freezing point depression

A

delta Tf = kf x m (i)

164
Q

Van’t Hoff factor is needed to account for

A

when dissolved, not all 1 m solutions have the same number of “particles”

165
Q

Van’t Hoff factor = i =

A

number of “particles” we get when we dissolve a compound

166
Q

NaCl i =

A

2

167
Q

CaCl2 i =

A

3

168
Q

glucose i =

A

1

169
Q

at higher concentrations of solute we see (blank)

A

ion pairing, vh decreases

170
Q

Osmosis

A

movement of solvent through a semipermeable membrane from an area of low solute concentration to high solute concentration

171
Q

osmotic pressure

A

pi = iMRT

172
Q

Osmotic pressure is super important in

A

IV administration and reverse osmosis

173
Q

Explain why the boiling point of Br2 (59C) is lower than that of iodine monochloride, ICl (97C), even though they have nearly the same molar mass.

A

ICl is polar while Br2 is no polar.

174
Q

Polar liquids have higher or lower boiling point?

A

Higher, dipole dipole

175
Q

In an aqueous solution containing Sn (II) and Sn (IV) salts, which cation would you expect to be the more strongly hydrated? Why?

A

Sn4+ because it is a smaller ion with a larger charge.

176
Q

Which of the following compounds is expected to have the weakest interactions between its molecules? CO2, NO2, SO2, H2S

A

CO2

177
Q

Which of the following compounds is expected to have the strongest interactions between its molecules? CO2, NO2, SO2, H2S

A

SO2

178
Q

Why do the strengths of London dispersion forces generally increase with increasing molecular size?

A

Dispersion forces arise from dipoles caused by the electron distribution being distorted. Larger molecules have more electrons and therefore, more distractions and a bigger force.

179
Q

What substances are insoluble in water? KBr (s), Benzene, C6H6 (l), Br2(l)

A

Benzene and Br2

180
Q

Smaller ions of the same charge form compounds that are (blank) soluble than compounds from larger ions.

A

Less

181
Q

Ions with higher charges form compounds that are (blank) soluble than ions of a lower charge.

A

Less

182
Q

Rank the following compounds in order of increasing solubility in water.

A

CaO, BaO, KCl, KI

183
Q

The most soluble compounds in water is th compound with the (blank) hydrophobic hydrocarbon chain.

A

Shortest

184
Q

Solid helium cannot be converted directly into the vapor phase. Does the phase diagram of He have a triple point?

A

No, without a sublimation curve there can be no triple point.

185
Q

What kind of intermolecular forces must be overcome during a the phase change in which solid CO2 sublimes?

A

Dispersion

186
Q

What kind of intermolecular forces must be overcome during a the phase change in which CHCl3 boils?

A

Dispersion

Dipole-dipole

187
Q

What kind of intermolecular forces must be overcome during a the phase change in which ice melts?

A

Dispersion, dipole-dipole, hydrogen

188
Q

At 1 ATM and 0 C, water is

A

Solid and liquid

189
Q

A hot needle sinks when out on the surface of cold water. Will a cold needle float or sink in warm water?

A

The hot needle breaks the surface tension of the water, causing it to sink. A cold needle cannot break up the surface tension of the water, so it floats.

190
Q

Does hot water have a higher or lower surface tension than cold water?

A

Lower

191
Q

Explain why different liquids do not reach the same height in capillary tubes of the same diameter.

A

Cohesive forces stronger than adhesive forces reduce the height of the liquid in the capillary tube, whereas adhesive forces stronger than cohesive forces increase the height of the liquid.

192
Q

The smaller the diameter of a tube, the (blank) the liquid rises up the tube by capillary action.

A

Higher

193
Q

Water has higher surface tension and viscosity than

A

Methanol

194
Q

What makes a substance soluble in water?

A

Polar and can form hydrogen bonds with water molecules.

195
Q

Of the following compounds, LiF- NaCl- or KI, which should have the highest melting point?

A

LiF

196
Q

The smallest distance between the ions and highest charges =

A

Higher lattice energy and melting point

197
Q

When the vapor pressure of a liquid is equal or greater than atmospheric pressure, the liquid will (blank)

A

Boil

198
Q

The higher the vapor pressure of a given liquid, the (blank) the boiling point.

A

Lower

199
Q

As the size of atomic number increases, the Latrice energy and melting point (blank)

A

Decrease

200
Q

Vapor pressure depends on (blank)

A

Temperature

201
Q

The higher the temperature, the (blank) the vapor pressure.

A

Higher

202
Q

The higher the molarity of the, the (blank) the osmotic pressure.

A

Higher

General chemistry II (4 decks)

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