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Fundamentals of Chemistry I > Final Exam > Flashcards

Final Exam Flashcards

1
Q

Tera

A

10^12

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2
Q

Giga

A

10^9

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3
Q

Mega

A

10^6

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4
Q

Kilo

A

10^3

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5
Q

Deci

A

10^-1

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6
Q

Centi

A

10^-2

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7
Q

Milli

A

10^-3

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8
Q

Micro

A

10^-6

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9
Q

Nano

A

10^-9

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10
Q

Pico

A

10^-12

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11
Q

Scientific Notation

A

Coefficient * (10^exponent)

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12
Q

Kelvin

A

Celcius + 273.15

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13
Q

Fahrenheit

A

(9/5) + 32

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14
Q

Nonzeroes

A

Always count

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15
Q

Leading Zeroes

A

Never count

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16
Q

Interior Zeroes

A

Always count

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17
Q

Trailing Zeroes

A

Only count if…
* After a decimal
* Before a decimal, after a nonzero

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18
Q

Exact Numbers

A

Always count

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19
Q

Accuracy

A

Agreement of measured to true value

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20
Q

Precision

A

Reproducibility

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21
Q

Density

A

mass / volume

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22
Q

Intensive Properties

A

Independent of amount

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23
Q

Extensive Properties

A

Dependent of amount

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24
Q

Dimensional Analysis

A

given * (desired/given) = given

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25
Q

Element

A

Same atoms

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26
Q

Compound

A

Different atoms

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27
Q

Homogenous Mixture

A

Uniform throughout

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28
Q

Heterogenous Mixture

A

Separable mixture

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29
Q

Scientific Theory

A
  1. Observation
  2. Hypothesis
  3. Experiment
  4. Theory (model)
  5. Scientific Law
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30
Q

Scientific Law

A

What? (summarizes observation)

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31
Q

Theory

A

Why? (interpretation)

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32
Q

Conservation of Mass

A

“Matter is neither created nor destroyed”

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33
Q

Definite Proportions

A

Constant composition

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34
Q

Multiple Proportions

A
  • When two elements form two different compounds, the masses of element B combine with 1 gram of element A
  • Can be expressed as a ratio of small whole numbers
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35
Q

Dalton’s Atomic Theory

A
  1. Each element has atoms
  2. Atoms of an element have the same mass
  3. Atoms combine in ratios to form compounds
  4. Atoms cannot become other elements
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36
Q

J.J. Thomson

A
  • Cathode ray experiment
  • Plum pudding model
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37
Q

Thomson’s Conclusions

A
  • All atoms have electrons
  • Atoms contain a positive charge
  • Charge to mass ratio of electrons
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38
Q

Rutherford

A
  • Alpha particles sent at gold foil
  • Nuclear theory atomic model
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39
Q

Rutherford’s Conclusions

A
  • Dense positive core (nucleus)
  • Open space where electrons are dispersed
  • Atoms is electrically neutral (# electrons = # protons)
  • Atom has additional mass
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40
Q

Millikan

A

Oil drop experiment

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41
Q

Millikan’s Conclusions

A
  • Electron charge: -1.6*10^-19 C
  • Electron mass: 9.10*10^-28 g
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42
Q

Plum Pudding Model

A

Electrons in a sphere of positive charge

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43
Q

Proton

A
  • Large
  • Positive charge
  • Nucleus
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44
Q

Neutron

A
  • Large
  • Neutral charge
  • Nucleus
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45
Q

Electron

A
  • Small
  • Negative charge
  • Electron cloud
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46
Q

Isotope

A

Same number of protons, different number of neutrons

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47
Q

Atomic Mass Number

A

Total number of protons and neutrons present

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48
Q

Atomic Number

A

Number of protons

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49
Q

Atomic Weight

A

Mass of a single atom

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50
Q

Cation

A
  • Positively charged
  • Loss of electrons
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51
Q

Anion

A
  • Negatively charged
  • Gain of electrons
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52
Q

Mole

A

1 mol = 6.02214 * 10^23 units (Avogadro’s Number)

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53
Q

Energy

A

The ability to do work

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54
Q

Frequency

A

Number of cycles per second

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55
Q

Wavelength

A

Distance from crest to crest

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56
Q

Electromagnetic radiation from lowest to highest?

A
  • Radio
  • Microwave
  • Infrared
  • Visible
  • Ultraviolet
  • X ray
  • Gamma ray
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57
Q

Quantum

A

Energy “packet”

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58
Q

Photon

A

Quantum of electromagnetic radiation

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59
Q

Photoelectric Effect

A

Emission of electrons from a material

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60
Q

Atomic Theory

A

Matter is composed of particles (atoms)

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61
Q

Dual Nature of Light

A

Light is both a wave and a particle (photon)

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62
Q

Bohr’s Model

A

Derived energy levels available to the electrons in the hydrogen atom

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63
Q

Principal Quantum Number

A
  • N
  • Size and energy
  • Shell, principal level
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64
Q

Ground State

A
  • Lowest level
  • Formed by releasing energy (emission spectra)
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65
Q

Excited State

A
  • Any level above ground state
  • Formed by absorbing energy
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66
Q

Emission Spectra of Hydrogen

A
  • Distinct, colored lines when viewed through a prism, representing specific wavelengths of light
  • Emitted by excited hydrogen atoms
  • As their electrons transition from higher to lower energy levels
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67
Q

Absorption Spectra of Hydrogen

A

Results of emission spectra interaction

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68
Q

DeBroglie Wavelength

A
  • γ = h / mv
  • All matter has duality
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69
Q

Heisenberg Uncertainty Principle

A

Δx = Δ(mv) ≥ h / 4π

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70
Q

Complementary Properties

A

Exclude one another

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71
Q

Determinacy

A

Definite predictable future

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72
Q

Indeterminacy

A

Can only predict future using probability

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73
Q

Angular Momentum Quantum Number

A
  • L
  • Shape
  • Subshell, Sublevel
  • Probability distribution (nodes = 0)
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74
Q

Magnetic Quantum Number

A
  • M(l)
  • Orientation
  • -1 … 0 … +1
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75
Q

Electron Spin Quantum Number

A
  • Ms
  • Spin direction
  • +1/2 (up) and -1/2 (down)
  • Must be opposite
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76
Q

S orbital

A
  • Spherical
  • Lowest
  • 1 per level
  • Gets larger with energy
  • No phase
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77
Q

P Orbital

A
  • Peanut/Dumbbell
  • Nodes separate lobes
  • Has phase
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78
Q

D Orbital

A

Four-leaf clover

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79
Q

F Orbital

A

Tetrahedral

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80
Q

Dimitri Mendeleev

A
  • Ordered elements by atomic weight
  • Columns: similar properties
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81
Q

Metals

A

Lose electrons to form cations

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82
Q

Nonmetals

A

Gain electrons to form anions

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83
Q

Metalloids

A
  • Metallic / nonmetallic behavior
  • Forms cations / anions
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84
Q

Pauli Exclusion Principle

A

No two electrons can have the same four quantum numbers

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85
Q

Shielding

A
  • Outer electrons repelled by core electrons
  • Nuclear charge decreases
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86
Q

Penetration

A

Ability of an electron to get closer to the nucleus

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87
Q

Effective nuclear charge

A

Zeff = Z - S

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88
Q

Noble Gas

A
  • Full shell
  • Stable
  • Nonreactive
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89
Q

Alkali Metals

A
  • Lose one electron
  • Cation (+1)
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90
Q

Alkaline Earth Metals

A
  • Lose two electrons
  • Cation (+2)
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91
Q

Transition Metals

A

Lose electrons to form cations

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92
Q

P Block Metals

A

Lose P shell and S shell to form cations

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93
Q

Halogens

A
  • Metals: gain one electron to form anions
  • Nonmetals: share electrons
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94
Q

Hund’s Rule

A

Electrons fill singly before being paired (degenerate)

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95
Q

Degenerate Orbitals

A

Same energy level but differ in spatial orientation

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96
Q

Valence Electrons

A

Electrons in the outermost shell of an atom

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97
Q

Monoatomic Ion

A

A charged single atom with a different number of electrons and protons

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98
Q

Atomic Radius

A

Average bonding radii

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99
Q

Atomic Radius Trends

A
  • Decreases across rows
  • Increases down groups
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100
Q

Ionic Radius

A
  • Cations are small
  • Anions are large
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101
Q

Ionization Energy

A

Energy required to remove an electron

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102
Q

Ionization Energy Trend

A
  • Increases across rows
  • Decreases down groups
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103
Q

Electron Affinity

A

The amount of energy released when an electron attaches to a gas

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104
Q

Electron Affinity Trend

A
  • Increases across rows
  • Decreases down groups
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105
Q

Paramagnetic

A
  • Unpaired electrons
  • Weakly attracted to magnetic fields
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106
Q

Diamagnetic

A
  • Paired electrons
  • Slightly repelled to magnetic fields
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107
Q

Ionic Bond

A
  • Cations and anions
  • Metal and nonmetal
  • Complete transfer of electrons
  • Lewis model
  • Energy released to form crystal lattices
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108
Q

Covalent Bond

A
  • Sharing 2 electrons
  • Molecular compounds
  • Lewis structure
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109
Q

Structural Formulas

A
  • Lewis structures
  • Shows structure
  • Most informative
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110
Q

Molecular Formulas

A

Actual numbers and types of atoms

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111
Q

Empirical Formulas

A

Relative numbers and types of atoms

112
Q

Octet Rule

A
  • Filled shells
  • 8 valence electrons
  • 3rd row and below, can expand
113
Q

Duet Rule

A
  • 2 valence electrons
  • H and He
114
Q

Monoatomic Ions

A
  • Li +, Na +, K +, Rb +, Cs +, Ag +
  • Mg 2+, Ca 2+, Sr 2+, Ba 2+
  • Al 3+, Ga 3+, Zn 2+, Cd 2+
  • Cr 2+, Cr 3+, Mn 2+, Mn 3+, Fe 2+, Fe 3+, Co 2+, Co 3+
  • Cu +, Cu 2+, Hg 2+, Hg2 2+, Sn 2+, Sn 4+, Pb 2+, Pb 4+
  • F -, Cl -, Br -, I -, O2 -, S2 -, N3 -
115
Q

Acetate

A

C2H3O2 -

116
Q

Carbonate

A

CO3, 2-

117
Q

Hydrogen Carbonate (Bicarbonate)

A

HCO3 -

118
Q

Hydroxide

A

OH -

119
Q

Nitrite

A

NO2 -

120
Q

Nitrate

A

NO3 -

121
Q

Chromate

A

CrO4, 2-

122
Q

Dichromate

A

Cr2O7, 2-

123
Q

Phosphate

A

PO4, 3-

124
Q

Hydrogen Phosphate

A

HPO4, 2-

125
Q

Dihydrogen Phosphate

A

H2PO4 -

126
Q

Ammonium

A

NH4 +

127
Q

Hypochlorite

A

ClO -

128
Q

Chlorite

A

ClO2 -

129
Q

Chlorate

A

ClO3 -

130
Q

Perchlorate

A

ClO4 -

131
Q

Permanganate

A

MnO4 -

132
Q

Sulfite

A

SO3 2-

133
Q

Hydrogen Suflite (Bisulfite)

A

HSO3 -

134
Q

Sulfate

A

SO4 2-

135
Q

Hydrogen Sulfate (Bisulfate)

A

HSO4 -

136
Q

Cyanide

A

CN

137
Q

Peroxide

A

O2 2-

138
Q

Diatomic Atoms

A
  • H2
  • N2
  • O2
  • F2
  • Cl2
  • Br2
  • I2
139
Q

Naming one type of ion

A

Cation + Anion-ide

140
Q

Naming many ions

A

Cation (charge) Anion-ide

141
Q

Naming nonmetals

A

Prefix (1st element) + Prefix (2nd element)

142
Q

Naming (-ide) acids

A

Hydro____ic Acid

143
Q

Naming (-ate) acids

A

____ic Acid

144
Q

Naming (-ite) acids

A

____ous Acid

145
Q

Molar Mass

A

Molecular Weight

146
Q

Mass Percent

A

(# atoms + atomic weight) / molecular weight

147
Q

Electronegativity

A

Atoms that attract electrons

148
Q

Electronegativity Trends

A
  • Group: decreases down
  • Period: increases across
149
Q

Polar Bonds

A
  • Unequal sharing of electrons
  • 0.4 - 2.0
150
Q

Nonpolar Bonds

A
  • Equal sharing of electrons
  • 0.0 - 0.4
151
Q

Lewis Structure

A
  1. Total number of valence electrons
  2. Draw skeletal structure with single bonds
  3. Complete octets
  4. Are all the valence electrons filled?
  5. Complete central octet
152
Q

Exceptions to Octet Rule

A
  • Be, B, Al (fewer than 8 valence electrons)
  • 3rd row and below exceeds octet with empty d orbitals
153
Q

Resonance

A

More than one valid lewis structure

154
Q

Bond Length

A
  • Distance between nuclei at minimum energy
  • Triple bond = shortest
155
Q

Bond energy

A
  • Energy required to break a covalent bond
  • Bond length decreases = Bond energy increases
156
Q

Formal Charges

A
  • Overall charge
  • # valence e - [ 1/2 (# shared e) + (# nonbonding e)]
157
Q

2 Electrons

A

Linear, 180 degrees

158
Q

3 Electrons

A

Trigonal Planar, 120 degrees
* 1 lone pair = bent, < 120 degrees

159
Q

4 Electrons

A

Tetrahedral, 109.5 degrees
* 1 lone pair = trigonal bipyramidal, < 109.5 degrees
* 2 lone pairs = bent, < 109.5 degrees

160
Q

5 Electrons

A

Trigonal Bipyramidal, 120 or 90 degrees
* 1 lone pair = seesaw, < 120 degrees (equatorial) and 90 degrees (axis)
* 2 lone pairs = t-shaped, < 90 degrees
* 3 lone pairs = linear, 180 degrees

161
Q

6 Electrons

A

Octahedral, 90 degrees
* 1 lone pair = square pyramidal, < 90 degrees
* 2 lone pairs = square planar, 90 degrees

162
Q

Polarity

A
  1. Predict geometry of molecule
  2. Predict dipole moment of each bond
  3. Add bond dipole moments based on geometry
  4. Calculate overall molecular dipole moment
163
Q

Chemical bond

A

Overlap of two half-filled orbitals and spin pairing of two valence electrons

164
Q

Hybridization

A

Results in specific electron geometry

165
Q

SP2

A

Trigonal planar, 3 electrons
* 3 SP2 hybrid orbitals
* 1 unhybridized P orbital

166
Q

SP

A

Linear, 2 electrons
* 2 SP hybrid orbitals
2 unhybridized P orbitals

167
Q

SP3

A

Tetrahedral, 4 electrons

168
Q

SP3D

A

Trigonal bipyramidal, 5 electrons
* 5 SP3D hybrid orbitals
* 4 unhybridized D orbitals

169
Q

SP3D2

A

Octahedral, 6 electrons
* 6 SP3D2 hybrid orbitals
* 3 unhybridized D orbitals

170
Q

Sigma Bond

A

Head-on-Overlap
* Single bonds
* Forms along interatomic axis
* Stronger than pi bonds

171
Q

Pi Bond

A

Sideways Overlap
* Unhybridized P orbitals
* Restricted rotation around pi bonds

172
Q

Valence Bond Theory

A

Where are valence electrons located?
* Aligned along the axis between atoms
* Parallel to each other
* Perpendicular to interatomic axis

173
Q

Molecular Orbital Theory

A

Combine constructively -> MO energy < AO’s energy
* Bonding molecular orbital
* Sigma, pi (more stable)

Combine destructively -> MO energy > AO’s energy
* Antibonding molecular orbital
* Sigma, pi (less stable)

174
Q

Highest Occupied Molecular Orbital (HOMO)

A

Highest energy electron orbital within a molecule that is fully occupied by electrons

175
Q

Lowest Unoccupied Molecular Orbital

A

Lowest energy orbital that is not occupied by electrons

176
Q

Chemical Change

A

Rearranging atoms to create new substances

177
Q

Physical Change

A

Change in state or appearance

178
Q

Chemical Equation

A
  • Reactants -> products
  • Bonds broken/formed
  • Atoms are neither created nor destroyed
179
Q

Combustion Reaction

A

React with O2 to form oxygen-containing compounds, often including water, and emitting heat

180
Q

Obtaining Amount Steps

A
  1. Balance chemical reaction
  2. Mass -> moles
  3. Moles -> moles (mole ratio)
  4. Moles -> mass
181
Q

Stoichiometry Steps

A
  1. Balanced chemical reaction
  2. Known masses -> moles
  3. Determine limiting reagent
  4. Find moles of desired substance
  5. Moles -> grams
182
Q

Limiting Reagent

A
  • Completely consumed
  • Limits product
183
Q

Reactant in Excess

A

Greater reactant than required

184
Q

Theoretical Yield

A

Amount of product made based on limiting reagent

185
Q

Actual Yield

A

Amount of product actually produced

186
Q

Percent Yield

A
  • Actual yield is usually less than theoretical yield
  • ( Actual / theoretical ) * 100%
187
Q

Molarity

A

moles of solute / liters of solute

188
Q

Dilution

A
  • Adding more solvent to reduce molarity
  • M dec, mole = same, L inc
  • M1V1 = M2V2
189
Q

Hydration

A
  • Salt dissolved in water
  • Ionic compounds “break up” into cations and anions
  • Negative ends of water appeal to cations / positive ends of water appeal to anions
190
Q

Weak Electrolyte

A

Only a small degree of ionization in water

191
Q

Strong Electrolytes

A
  • Completely ionize
  • Soluble salts and strong acids/bases
192
Q

Non-Electrolytes

A
  • Dissolves in water
  • No ions produced
193
Q

Li +, Na +, K +, NH4 +

A

Soluble

194
Q

NO3 - & C2H3O2 -

A

Soluble

195
Q

Cl -, Br-, I-

A

Soluble

196
Q

What are Cl -, Br-, I- insoluble with?

A

Ag +, Hg2 2+, Pb 2+

197
Q

SO4 2-

A

Soluble

198
Q

What are SO4 2- insoluble with?

A

Sr 2+, Ba 2+, Pb 2+, Ag +, Ca 2+

199
Q

OH - & S 2-

A

Insoluble

200
Q

What are OH - & S 2- soluble with?

A

Li +, Na +, K+, NH4 +, Ca 2+, Sr 2+, Ba 2+

201
Q

CO3 2- & PO4 3-

A

Insoluble

202
Q

What are CO3 2- & PO4 3- soluble with?

A

Li +, Na +, K+, NH4 +

203
Q

Precipitation Reaction

A

Exchange reactions involve exchanging ions in solution

204
Q

Complete Ionic Equation

A

Strong electrolytes written as ions

205
Q

Net Ionic Equation

A

Only reacting components included (no spectator ions)

206
Q

Acid

A

Proton (H+) donor

207
Q

Bases

A

Proton acceptor (OH-)

208
Q

Binary Acids

A

An acid with only two elements

209
Q

Oxyacids

A

An acid that contains oxygen

210
Q

Strong Acids

A

HCl, HBr, HI, HNO3, H2SO4, HClO4

211
Q

Strong Bases

A
  • Alkali metal hydroxides
  • NaOH, LiOH, KOH, Ca(OH)2, Ba(OH)2
212
Q

Neutralization

A

Just enough base to react with an acid

213
Q

Net Ionic Acid/Base Reaction Equation

A

H+ (aq) + OH- (aq) -> H2O (l)

214
Q

Acid/Base Reaction Equation

A

Acid + Base -> Salt + Water

215
Q

Oxidation

A

Loss of electron

216
Q

Reduction

A

Gain of electrons

217
Q

Oxidizing Agent

A
  • Reduction
  • Gains electrons
218
Q

Reducing Agent

A
  • Oxidation
  • Loss of electrons
219
Q

Redox Reactions

A

Transfer of electrons

220
Q

Oxidation Numbers

A
  1. Atoms in an element = 0
  2. Monatomic = charge of ion
  3. Sum = 0 in compounds
  4. Sum = charge of ion in ions
  5. Oxygen = -2 (peroxide = -1)
  6. Hydrogen = +1
  7. Fluorine = -1
221
Q

Units of Energy

A
  • Calorie
  • Joule
222
Q

Specific Heat Capacity

A

Heat required to raise 1 gram of a substance by 1 degree (s)

223
Q

Internal Energy Equation

A

ΔE = q + w

224
Q

Heat Equation

A

q = mcΔT

225
Q

Work Equation

A

w = -pΔv

226
Q

Thermal Energy Transfer

A
  • Heat- temperature difference
  • Work- force over a distance
227
Q

Endothermic

A
  • Heat absorbed
  • Surroundings -> system
  • ΔH is positive
  • Feels cold
228
Q

Exothermic

A
  • Heat released
  • System -> surroundings
  • ΔH is negative
  • Feels hot
229
Q

Enthalpy Change

A

Heat that passes into or out of the system during a reaction

230
Q

Calorimetry

A

Science of measuring heat
* Calorimeter- measures heat flow
* Isolated system
* Insulated
q(rxn) = ΔH (rxn)

231
Q

Hess’ Law

A

ΔH is the sum of individual ΔH steps
* Enthalpy is extensive (magnitude of ΔH is directly proportional to amount)
* Reverse reaction has an opposite sign for ΔH
* Change in enthalpy depends on physical state

232
Q

Standard States

A
  • Gas- 1 atm
  • Solution- 1 M
  • ΔHf° = 0 at atm and 25° C
233
Q

Standard Enthalpies of Formation

A

When 1 mole is formed from standard states (ΔHf°)

234
Q

Bond Energies

A

Breaking bonds
* Requires energy
* Endothermic (+)

Forming bonds
* Releases energy
* Exothermic (-)

235
Q

Solid

A
  • High density
  • Definite shape
  • Definite volume
  • Strong IMFs
236
Q

Liquid

A
  • High density
  • Indefinite shape
  • Definite volume
  • Moderate IMFs
237
Q

Gas

A
  • Low density
  • Indefinite shape
  • Indefinite volume
  • Weak IMFs
238
Q

Dipole-Dipole

A

Electrostatic Attraction forming a permanent dipole (polar)

239
Q

London Dispersion Forces

A
  • Forces between an instantaneous and an induced dipole
  • Present in all molecules and atoms
240
Q

Ion-Dipole

A
  • Attraction between an ion and a permanent dipole
  • Dissolves ions in a solution
241
Q

Hydrogen Bond

A
  • H bonded to F, O, and N
  • Unusually strong
  • Unusually high boiling and melting points
242
Q

Sublimation

A
  • Endothermic
  • Solid to gas
243
Q

Deposition

A
  • Exothermic
  • Gas to solid
244
Q

Melting

A
  • Endothermic
  • Solid to liquid
245
Q

Freezing

A
  • Exothermic
  • Liquid to solid
246
Q

Vaporization

A
  • Endothermic
  • Liquid to gas
247
Q

Condensation

A
  • Exothermic
  • Gas to liquid
248
Q

Enthalpy of Fusion

A

ΔH (fus)

249
Q

Enthalpy of Vaporization

A

ΔH (vap) = -ΔH (condensation)

250
Q

Surface Tension

A

Resistance to increase its surface area
* IMFs inc = ST inc

251
Q

Solubility

A

“Like dissolves like”

252
Q

Miscible

A

Substances mix

253
Q

Immiscible

A

Substances do not mix

254
Q

Hydrophobic

A

Repels water

255
Q

Hydrophilic

A

Attracts water

256
Q

Viscosity

A

Resistance to flow
* IMFs inc = Viscosity inc
* Temp inc = Viscosity dec

257
Q

Capillary Action

A

Rising of a liquid in a narrow tube

258
Q

Cohesive

A

IMFs between like substances

259
Q

Adhesive

A

IMFs between unlike substances

260
Q

Boiling Point

A

Vaporization point is equal to 1 atm

261
Q

Vapor Pressure

A

Pressure of a vapor in dynamic equilibrium
* Temp inc = VP inc
* IMFs inc = VP dec
* VP dec = BP dec

262
Q

Triple Point

A

All 3 states exist simultaneously

263
Q

Critical Point

A

Liquid and gas phases are indistinguishable
* Critical temperature
* Critical pressure

264
Q

Pressure Conversion Factors

A

1 atm = 760 mmHg = 760 torr

265
Q

Graham’s Law

A
  • Occurs only at same temperature
  • rate (A) / rate (B) = √ M (B) / M (A)
266
Q

Boyle’s Law

A

P ∝ 1/V

267
Q

Amonton’s Law

A

P ∝ T

268
Q

Charles’ Law

A

V ∝ T

269
Q

Avogadro’s Law

A

V ∝ n

270
Q

Ideal Gas Law

A

PV = nRT

271
Q

STP Conditions

A
  • P = 1 atm
  • T = 0° C
  • Molar volume = 22.4 L
272
Q

Dalton’s Law of Partial Pressure

A

P (total) = Pn

273
Q

Kinetic Molecular Theory

A
  • Volume of a gas particle is small
  • Gas particles have no IMFs
  • KE (avg) ∝ Kelvin
  • Constant motion causes elastic collisions
274
Q

Ideal Gas Behavior

A
  • Gases behave ideally
  • Real gases behave ideally at low pressures and high temperatures
275
Q

Real Gas Behavior

A

Corrected at a high pressure (small volume) and low temperature (IMFs)

Fundamentals of Chemistry I (1 decks)

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