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CHM 1040 > Exam 4 > Flashcards

Exam 4 Flashcards

1
Q

why study gases?

A

1) To gain an understanding of real-world phenomena

2) To gain an understanding of how science “works”

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2
Q

gas

A
  • Uniformly fills any container
  • Mixes completely with any other gas
  • Exerts pressure on its surroundings
  • has small forces between them
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3
Q

barometer

A
  • Device used to measure atmospheric pressure
  • Mercury flows out of the tube until the pressure of the column of mercury standing on the surface of the mercury in the dish is equal to the pressure of the air on the rest of the surface of the mercury in the dish
  • mm Hg
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4
Q

units of pressure

A
  • Unit millimeters of mercury (mm Hg) is often called the torr
    a) The units “torr” and “mm Hg” are used interchangeably
  • Standard atmosphere (atm): A related unit for pressure
  • Pascal: SI unit for pressure (Pa)
  • PRESSURE = FORCE / AREA
  • SI units = Newton/meter2= 1 Pascal (Pa)
  • 1 standard atmosphere = 101,325 Pa
  • 1 standard atmosphere = 1 atm= 760 mm Hg = 760 torr
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5
Q

Unit millimeters of mercury (mm Hg) is often called the

A

torr

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6
Q

what 2 units can be used interchangeably?

A

torr and mm Hg

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7
Q

Standard atmosphere (atm)

A

A related unit for pressure

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8
Q

formula for pressure

A

P = force/area

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9
Q

1 atm = __ mm Hg

A

760

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10
Q

1 atm = __ torr

A

760

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11
Q

1 atm = __ Pa

A

101,325

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12
Q

The pressure of the air in a tire is measured to be 28 psi. Represent this pressure in both torr and pascals.

A
  1. 9 atm x 760 torr / 1 atm = 1.4 x 10^3 torr

1. 9 atm x 101325 Pa/ atm = 1.9 x 10^5 Pa

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13
Q

The vapor pressure over a beaker of hot water is measured as 656 torr. What is this pressure in atmospheres?

A

656 torr x 1 atm / 760 torr = 0.863 atm

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14
Q

Robert Boyle’s Experiment

A

Study of the relationship between the pressure of the trapped gas and its volume

  • when u increase P, you decrease V
  • ex: pillow
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15
Q

Boyle’s Law Graph

A

L shaped curve

  • pressure and volume
  • inversely proportional
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16
Q

Boyle’s Law

A
  • P1V1 = P2V2
  • Pressure and volume are inversely related or inversely proportional
  • PV = k
  • k is a constant for a given amnt of gas at a specific temp
  • V= k/p
  • T and moles are constant
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17
Q

A sample of helium gas occupies 12.4 L at 23 °C and 0.956 atm. What volume will it occupy at 1.20 atm assuming that the temperature stays constant?

A

9.88 L

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18
Q

Graphing Data for Several Gases

A
  • It is easier to write an equation for the relationship between volume and temperature if the lines are extrapolated to the origin of the graph
  • Use absolute zero for the temperature
  • Absolute zero
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19
Q

absolute zero

A

temperature of −273 °C beyond which matter cannot be cooled

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20
Q

Charles’s Law Graph

A

V on y axis and T on x axis

  • direct relationship
  • straight up
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21
Q

Charles’s Law

A
  • V1/T1 = V2/T2
  • Direct proportionality between volume and temperature(in Kelvin) represented by the equation known as Charles’s law
  • V = bT (b is a proportionality constant)
  • K = °C + 273
  • 0 K is called absolute zero
  • constant pressure and moles
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22
Q

__ K is called absolute zero

A

0

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23
Q

Suppose a balloon containing 1.30 L of air at 24.7 °C is placed into a beaker containing liquid nitrogen at –78.5 °C. What will the volume of the sample of air become (at constant pressure)?

A

0.849 L

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24
Q

Avogadro’s Law

A
  • n1/V1 = n2/V2
  • Volume and numbers of moles are directly related (constant T and P)
  • V = an (a is a proportionality constant)
  • constant temp and pressure
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25
Q

Avogadro’s Law Graph

A
  • direct relationship
  • straight up
  • volume on y axis and moles on x axis
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26
Q

If 2.45 mol of argon gas occupies a volume of 89.0 L, what volume will 2.10 mol of argon occupy under the same conditions of temperature and pressure?

A

76.3 L

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27
Q

Ideal Gas Law

A
  • PV = nRT
  • R = 0.08206 L atm / mol K
  • We can bring all the laws together under one comprehensive law:
    a) charles’s law: V =bT (constant P and n)
    b) Avogadro’s law: V =an (constant Tand P)
    c) Boyle’s law: V= k / P (constant T and n)
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28
Q

what is the universal gas constant?

A

0.08206 L atm / mol K

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29
Q

An automobile tire at 23 °C with an internal volume of 25.0 L is filled with air to a total pressure of 3.18 atm. Determine the number of moles of air in the tire.

A

3.27 mol

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30
Q

What is the pressure in a 304.0-L tank that contains 5.670 kg of helium at 25 °C?

A

114 atm

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31
Q

At what temperature (in °C) does 121 mL of CO2 at 27 °C and 1.05 atm occupy a volume of 293 mL at a pressure of 1.40 atm?

A

696 °C

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32
Q

Dalton’s Law of Partial Pressures

A
  • For a mixture of gases in a container
    a) P total = P1 + P2 + P3 + . . .
  • the total pressure exerted is the sum of the pressures that each gas would exert if it were alone
  • The pressure of the gas is affected by the number of moles of particles present
  • The pressure is independent of the nature of the particles
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33
Q

Two Crucial Things to be learned from the Dalton’s Law

A

1) The volume of the individual particles must not be very important
2) The forces among the particles must not be very important

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34
Q

Collecting a Gas over Water

A

Total pressure is the pressure of the gas plus the vapor pressure of the water

  • However, we need to determine the pressure of the gas and not the mixture collected
  • to JUST get the gas, SUBTRACT the water vapor
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35
Q

Consider the following apparatus containing helium in both sides at 45 °C. Initially, the valve is closed.
- After the valve is opened, what is the pressure of the helium gas?

A

2.25 atm

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36
Q
  1. 4 L of oxygen gas at 25.0 °C and 1.30 atm and 8.50 L of helium gas at 25.0 °C and 2.00 atm were pumped into a tank with a volume of 5.81 L at 25 °C.
    - Calculate the new partial pressure of oxygen
    - Calculate the new partial pressure of helium
    - Calculate the new total pressure of both gases
A

new partial pressure of oxygen = 6.13 atm
new partial pressure of helium = 2.93 atm
new total pressure of both gases = 9.06 atm

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37
Q

Scientific Method

A
  • A law is a generalization of observed behavior
  • Laws are useful as they help us predict behavior of similar systems
  • A model can never be proved absolutely true
  • A model is an approximation and is destined to be modified
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38
Q

The Kinetic Molecular Theory of Gases: Basic Postulates of the Kinetic Molecular Theory

A

in science, “what” always comes before why

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39
Q

implications of the Kinetic Molecular Theory: Meaning of temperature

A

Kelvin temperature of a gas is directly proportional to the average kinetic energy of the gas particles

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40
Q

implications of the Kinetic Molecular Theory: Relationship between pressure and temperature

A

Gas pressure increases as the temperature increases because the gas particles move faster and undergo collision more often

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41
Q

implications of the Kinetic Molecular Theory: Relationship between volume and temperature

A

Volume of a gas increases with temperature because the gas particles speed up

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42
Q

You are holding two balloons of the same volume. One contains helium, and the other contains hydrogen. The pressures of the gas in the two balloons are ___

A

the same

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43
Q

You are holding two balloons of the same volume. One contains helium, and the other contains hydrogen. The temperatures of the gas in the two balloons are ___

A

the same

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44
Q

You are holding two balloons of the same volume. One contains helium, and the other contains hydrogen. The numbers of moles of the gas in the two balloons are ___

A

the same

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45
Q

You are holding two balloons of the same volume. One contains helium, and the other contains hydrogen. The densities of the gas in the two balloons are ___

A

different

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46
Q

V Ne = 2V Ar

what is the mass ratio of Ne:Ar in the balloons?

A

1:1

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47
Q

The gas sample is then cooled to a temperature of 15 °C. Solve for the new condition. (Hint: Consider that a moveable piston keeps the pressure constant overall.)

A

5.43 L

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48
Q

Molar Volume of an Ideal Gas

A
  • For 1 mole of an ideal gas at 0 °C and 1 atm, the volume of the gas is 22.42 L
  • STP
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49
Q

standard temperature and pressure (STP)

A
  • 0 °C and 1 atm

- any gas with 0 °C and 1 atm will have a volume of 22.4 L at STP

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50
Q

A sample of oxygen gas has a volume of 2.50 L at STP. How many grams of O2 are present?

A

3.57 g

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51
Q

Consider the following reaction:
Zn(s) + 2HCl(aq ) -> ZnCl(aq) + H(g)
If 15.00 g of solid zinc reacts with 100.0 mL of 4.00 M hydrochloric acid, what volume of hydrogen gas will be produced at 25 °C and 1.00 atm?

A

4.89 L

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52
Q

review of liquids and solids

A
  • Gases have low density, are highly compressible, and can fill a container
  • solids have high density, are slightly compressible, and are rigid
  • Some properties of liquids lie between those of solids and of gases
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53
Q

heating/cooling curve

A
  • normal boiling pt at 1 atm is 100°C
  • normal freezing pt at 1 atm is 0 °C
  • density
    a) liquid water = 1.00g /mL
    b) ice = 0.917 g/mL
  • liquid to gas is MORE energy than solid to liquid because it is more disordered
54
Q

During the process of melting ice by adding heat, what would happen to the temperature of the ice?

A

it stays constant

55
Q

Physical changes

A
  • No chemical bonds are broken
  • When water is boiled to form steam, water molecules are separated from each other, but the individual molecules remain intact
56
Q

phase changes

A
  • When a substance undergoes changes in state from solid to liquid and from liquid to gas, the molecules remain intact
  • The changes in state are due to the changes in forces among the molecules rather than those within the molecules
57
Q

Phase changes: Solid to Liquid

A

As energy is added, the motion of the molecules increases, and they eventually achieve greater movement and the disorder characteristic of a liquid

58
Q

phase changes: liquid to gas

A

As more energy is added, the gaseous state is eventually reached, in which the individual molecules are far apart and interact relatively little

59
Q

Intramolecular Forces

A
  • WITHIN
  • Take place within the molecule
  • Molecules are formed by sharing electrons between the atoms
  • Hold the atoms of a molecule together
  • 2 types: ionic and covalent *ionic is stronger
60
Q

intermolecular forces

A
  • BETWEEN
  • Forces that occur between molecules
  • 3 types: D-D, H-bonding, L-D
61
Q

__ forces are stronger than __ forces

A

Intramolecular forces are stronger than intermolecular forces

62
Q

what are the 2 types of intramolecular forces?

A

ionic and covalent

- ionic is stronger

63
Q

what are the 3 types of intermolecular forces?

A

dipole-dipole, Hydrogen bonding, and London dispersion

  • H-bonding is strongest
  • London dispersion is weakest
64
Q

dipole-dipole forces

A
  • dipole moment
  • Molecules with dipole moments can attract each other electrostatically
  • They line up so that the positive and negative ends are close to each other
  • Only about 1% as strong as covalent or ionic bonds
65
Q

hydrogen bonding

A
  • Hydrogens have FON
  • strongest
  • Strong dipole–dipole forces occur between molecules when hydrogen is bound to a highly electronegative atom - Nitrogen, oxygen, or fluorine (FON)
  • Affects physical properties (Boiling point)
    a) stronger the force the higher the boiling pt
66
Q

London Dispersion Forces

A
  • weakest
  • occurs in ALL molecules
  • e- constantly moving
  • Forces that exist among noble gas atoms and non-polar molecules
  • Instantaneous dipole that occurs temporarily in a given atom induces a similar dipole in a neighboring atom
  • Significant in large atoms/molecules
  • Occurs in all molecules, including non-polar ones
67
Q

London Dispersion Forces: Non-polar molecules

A
  • also interact by developing instantaneous dipoles

- Become stronger as the size of atoms or molecules increases

68
Q

melting and boiling points

A

Stronger the intermolecular forces, higher the melting and boiling points
- H-bonding has highest melting and boiling points

69
Q

Which molecule is capable of forming stronger intermolecular forces? H2O or N2?

70
Q

Which gas would behave more ideally at the same conditions of P and T? CO or N2?

71
Q

Consider the following compounds: NH3, CH4, H2

How many of the compounds above exhibit London dispersion forces?

72
Q

Vaporization or Evaporation

A
  • Molecules of a liquid can escape the liquid’s surface and form a gas
  • Endothermic process: Requires energy to overcome the relatively strong intermolecular forces in the liquid
73
Q

Vapor Pressure

A
  • Amount of liquid decreases initially and then becomes constant
  • Condensation
  • When no further change is visible, the opposing processes balance each other = Equilibrium
  • Equilibrium vapor pressure
  • The system is at equilibrium when no net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other
  • Liquids in which the intermolecular forces are strong have relatively low vapor pressures
74
Q

Condensation

A

Process by which vapor molecules convert to a liquid

75
Q

equilibrium

A

opposing processes balance each other

- no net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other

76
Q

Equilibrium vapor pressure

A

Pressure of the vapor present at equilibrium

77
Q

the stronger the force the __ the vapor pressure

78
Q

London dispersion forces have __ vapor pressures

79
Q

Which of the following would be expected to have the highest vapor pressure at room temperature?

a) CH3CH2CH2OH
b) CH3CH2CH2NH2
c) CH3CH2CH2CH3
d) CH3CH2CH3

A

d) CH3CH2CH3

80
Q

types of solids

A
  • crystalline solids
  • molecular solids
  • atomic solids
81
Q

Crystalline Solids

A

Substances with a regular arrangement of their components form crystalline solids

82
Q

types of crystalline solids

A
  • ionic solids
  • molecular solids
  • atomic solids
83
Q

ionic solids

A
  • components are ions

- Ions at the points of the lattice that describes the structure of the solid

84
Q

molecular solids

A
  • components are molecules

- Discrete covalently bonded molecules at each of its lattice points

85
Q

atomic solids

A
  • components are atoms

- Atoms at the lattice points that describe the structure of the solid

86
Q

examples of 3 types of crystalline solids

A

1) diamond -atomic
2) NaCl -ionic
3) H2O (ice) -molecular

87
Q

bonding in metals

A
  • Metals are held together by strong but nondirectional covalent bonds (called the electron sea model) among the closely packed atoms
88
Q

solution

A

a homogeneous mixture

89
Q

Solvent

A

a substance present in the largest amount in a solution

90
Q

Solute

A

a substance that dissolves in a solvent to form a solution

91
Q

Aqueous solution

A

a solution with water as the solvent

92
Q

example of solution

A

solution: hot chocolate
solvent: water or milk
solute: chocolate powder

93
Q

various types of solutions

A

gas and metal can also be solutions, not just liquids

94
Q

solubility of ionic substances

A
  • Ionic substances break up into individual cations and anions when dissolved in water
  • Polar water molecules interact with the positive and negative ions of a salt
  • Ethanol is soluble in water because of its polar O—H bond
  • ion breaks into cation and anion
95
Q

solubility of polar substances

A

Why is solid sugar soluble in water?

because of the O-H polar bond

96
Q

Why is solid sugar soluble in water?

A

because of the O-H polar bond

97
Q

Substances Insoluble in Water

A
  • Non-polar oil does not interact with polar water

- Water–water hydrogen bonds keep the water from mixing with the non-polar molecules

98
Q

How Substances Dissolve

A
  • A “hole”must be made in the water structure for each solute particle
  • The lost water–water interactions must be replaced by water–solute interactions
  • “Like dissolves like”
99
Q

like dissolves __

100
Q

polar molecules dissolve in __

A

polar solvents

101
Q

a non-polar molecule dissolves in __

A

a non-polar solvent

102
Q 
Which of the following solutes will generally not dissolve in the specified solvent? Choose the bestanswer. Assume all of the compounds are in the liquid state.
a) CCl4 mixed with water or H2O
b )NH3 mixed with water or H2O
c) CH3OH mixed with water or H2O
d) N2mixed with methane or CH4
A

a) CCl4 mixed with water or H2O

103
Q

Terms Associated with the Concentration of a Solution

A
  • The solubility of a solute is limited
    a) Saturated solution
    b) Unsaturated solution
  • Solutions are mixtures
  • Amounts of substances can vary in different solutions
    a) Specify the amounts of solvent and solute
    b) Use qualitative measures of concentration
  • Relatively large amount of solute is dissolved in a concentrated solution
  • Relatively small amount of solute is dissolved in a dilute solution
104
Q

Saturated solution

A

contains as much solute as will dissolve at that temperature

105
Q

Unsaturated solution

A

is a solution that has not reached the limit of solute that will dissolve in it

106
Q

Relatively large amount of solute is dissolved in a __ solution

A

concentrated

107
Q

Relatively small amount of solute is dissolved in a __ solution

108
Q

mass percent

A

mass percent = mass of solute ÷ mass of solution x 100%

mass percent = grams of solute ÷ (grams of solute + grams of solvent) x 100%

109
Q

What is the percent-by-mass concentration of glucose in a solution made by dissolving 5.5 g of glucose in 78.2 g of water?

110
Q

calculating molarity

A
  • Molarity, M, equals the number of moles of solute per volume of solution in liters
    M = moles of solute / liters of solution
111
Q

you have 1.00 mol of sugar in 125.0 mL of solution. Calculate the concentration in units of molarity.

112
Q

A 500.0-g sample of potassium phosphate is dissolved in enough water to make 1.50 L of solution. What is the molarity of the solution?

113
Q

You have a 10.0-M sugar solution. What volume of this solution do you need to have 2.00 mol of sugar?

114
Q

You have two HCl solutions, labeled Solution A and Solution B. Solution A has a greater concentration than Solution B. Which of the following statements is true?

A

If you have equal moles of HCl in both solutions, Solution B must have a greater volume

115
Q

concentration of ions

A
  • concentration of ions stays consistent
    For a 0.25 M CaCl2 solution:CaCl2→ Ca2+ + 2Cl-
  • Ca2+: 1 × 0.25 M = 0.25 M Ca2+
  • Cl- : 2 × 0.25 M = 0.50 M Cl–
116
Q

Standard Solution

A

A solution whose concentration is accurately known

- curve is straight up

117
Q

Steps to Make a Standard Solution

A
  1. Weigh out a sample of solute
  2. Transfer it to a volumetric flask
  3. Add enough solvent to bring the volume up to the mark on the neck of the flask
118
Q

dilution

A

M1 x V1 = M2 x V2

  • The process of adding water to a concentrated or stock solution to achieve a solution of desired concentration
  • Dilution with water does not alter the numbers of moles of solute present
  • Moles of solute before dilution equals moles of solute after dilution
  • changing denominator of molarity
119
Q

Steps to Dilute a Solution

A
  1. Transfer a measured amount of original solution to a flask containing some water
  2. Add water to the flask by swirling to bring the volume up to the calibration mark
  3. Mix by inverting the flask
120
Q

What is the minimum volume of a 2.00-M NaOH solution needed to make 150.0 mL of a 0.800-M NaOH solution?

121
Q

What is the minimum volume of a 5.68-M NaOH solution needed to make 323.0 mL of a 2.96-M NaOH solution?

A

168 mL

* 155 mL water added (323 - 168 = 155)

122
Q

example of concentrated vs diluted solution

A

food coloring in a bowl of water is concentrated, food coloring in a pool is dilute

123
Q

Steps for Solving Stoichiometric Problems Involving Solutions

A
  1. Write the balanced equation for the reaction
    - For reactions involving ions, it is best to write the net ionic equation
  2. Calculate the moles of reactants
  3. Determine which reactant is limiting
  4. Calculate the moles of other reactants or products, as required
  5. Convert to grams or other units, if required
124
Q
  1. 0 mL of a 0.30-M sodium phosphate solution reacts with 20.0 mL of a 0.20-M lead(II) nitrate solution (assume no volume change)
    - What precipitate will form?
    - What mass of precipitate will form?
A
  • lead(II) phosphate, Pb3(PO4)2

- 1.1 g Pb3(PO4)

125
Q
  1. 0 mL of a 0.30-M sodium phosphate solution reacts with 20.0 mL of a 0.20-M lead(II) nitrate solution (assume no volume change)
    - What is the concentration of nitrate ions left in the solution after the reaction is complete?
126
Q
  1. 0 mL of a 0.30-M sodium phosphate solution reacts with 20.0 mL of a 0.20-M lead(II) nitrate solution (assume no volume change)
    - What is the concentration of phosphate ions left in the solution after the reaction is complete?
127
Q

dipole moment

A

molecules w polar bonds often behave in an electric field as if they had a center of positive charge
(+) end attracts (-) end

128
Q

non polar

A

equal sharing of e- (diatomics)

129
Q

polar

A

non-equal sharing of e- (HCl)

130
Q

Standard atmosphere (atm)

A

A related unit for pressure

CHM 1040 (6 decks)

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