Gases & particle behaviour Flashcards
(202 cards)
1
Q
Which elements are gases at room temperature?
A
Hydrogen (H₂), Nitrogen (N₂), Oxygen (O₂), Fluorine (F₂), Chlorine (Cl₂), and all Group 8 noble gases.
2
Q
Which elements are liquids at room temperature?
A
Bromine (Br₂) and Mercury (Hg).
3
Q
What are diatomic gases? List them.
A
Molecules composed of two atoms: H₂, N₂, O₂, F₂, Cl₂, I₂.
3
Q
What are monoatomic gases? List them.
A
Noble gases: He, Ne, Ar, Kr, Xe, Rn.
4
Q
Name common compounds that exist as gases at room temperature.
A
Methane (CH₄), ethane (C₂H₆), carbon dioxide (CO₂).
5
Q
What is the composition of dry air at sea level?
A
78.09% N₂, 20.95% O₂, 0.93% Ar, 0.039% CO₂.
6
Q
Why are gases important in biology and medicine?
A
They are essential for respiration, lung function, and solubility in blood.
7
Q
What is the formula for volume in terms of cross-sectional area and height?
A
V = A × h
8
Q
What is the formula for pressure?
A
p = F / A
9
Q
What is the SI unit of pressure?
A
Pascal (Pa), where 1 Pa = 1 N/m²
10
Q
Convert 1 bar to other pressure units.
A
1 bar = 100 kPa = 10⁵ Pa = 14.504 psi
11
Q
Convert 1 atm to other units.
A
1 atm = 1.0133 bar = 101.325 kPa = 760 Torr
12
Q
What is the conversion factor between Torr and Pa?
A
1 Torr = 133.32 Pa
13
Q
What is the conversion factor between PSI and Pa?
A
1 PSI = 6895 Pa
14
Q
What are SATP conditions?
A
25°C (298.15 K), 1 bar (100 kPa)
15
Q
What are STP conditions?
A
0°C (273.15 K), 1 atm (101.325 kPa)
16
Q
State Avogadro’s Law.
A
Volume is directly proportional to the number of moles (n) at constant T and P.
17
Q
What is the implication of Avogadro’s Law?
A
Equal volumes of gases at the same T and P contain equal numbers of particles.
18
Q
What is Avogadro’s constant?
A
Nₐ = 6.022 × 10²³ mol⁻¹
19
Q
What is the volume of 1 mol of gas at SATP?
A
24.8 L
20
Q
State Boyle’s Law.
A
Volume is inversely proportional to pressure at constant T and n.
21
Q
Write the mathematical form of Boyle’s Law.
A
p ∝ 1 / V
22
Q
State Charles’ Law.
A
Volume is directly proportional to absolute temperature at constant P and n.
23
Q
Write the mathematical form of Charles’ Law.
A
V ∝ T
24
At what temperature do all gases extrapolate to zero volume?
-273.15°C or 0 K
25
How do you convert Celsius to Kelvin?
T(K) = T(°C) + 273.15
26
What is the ideal gas equation?
pV=nRT
27
Define each variable in the ideal gas law.
p = pressure (Pa)
V = volume (m³)
n = moles of gas
R = gas constant = 8.314 J·K⁻¹·mol⁻¹
T = temperature (K)
28
What are the different values of the gas constant R?
R = 8.314 J·K⁻¹·mol⁻¹
R = 0.08206 L·atm·K⁻¹·mol⁻¹
R = 0.083144626 L·bar·K⁻¹·mol⁻¹
29
What is the formula for density?
n = m / M
30
What is Dalton’s Law of Partial Pressures?
The total pressure of a gas mixture is equal to the sum of the partial pressures of the individual gases.
31
What is the formula for the total pressure of a gas mixture?
pₜ = (nₐ + nᵦ + n𝚌 + ...) × R × T / V
or
pₜ = pₐ + pᵦ + p𝚌 + ...
32
What is the mole fraction of a gas A?
xₐ = nₐ / nₜ
33
How are partial pressure and mole fraction related for an ideal gas?
pₐ = xₐ × pₜ
34
What does the mole fraction represent?
The fraction of total moles that is made up by gas A.
35
What are the key assumptions of the kinetic molecular theory?
Gas molecules are point masses with negligible volume.
Collisions are elastic (no energy lost).
There are no intermolecular interactions.
Gas molecules are in constant, random motion.
36
How is pressure created according to the kinetic molecular theory?
Pressure is created by collisions of gas molecules with the container walls.
37
What is the equation for kinetic energy of a single particle?
Eₖ = (1/2) × m × u²
38
How is temperature related to kinetic energy?
Average kinetic energy is proportional to temperature.
39
What does the Maxwell-Boltzmann distribution describe?
The distribution of speeds among gas particles — most are at intermediate speed, some slow, some fast.
40
How does molar mass affect molecular speed?
Higher molar mass leads to lower average speed and a narrower distribution.
41
How does temperature affect molecular speed?
Higher temperature increases average speed and broadens the distribution.
42
What is the equation for average molar kinetic energy?
43
What is the equation for molecular kinetic energy in terms of RMS speed?
44
What is effusion?
45
What is diffusion?
46
What causes effusion and diffusion to occur?
47
What is the equation for concentration of a gas A?
48
What are the units for gas concentration?
49
What direction does diffusion move in terms of concentration?
50
How does temperature affect RMS speed and diffusion?
51
Why do real gases deviate from ideal behavior?
52
What equation describes the behavior of real gases?
53
What does the 'a' term in the Van der Waals equation correct for?
54
What does the 'b' term in the Van der Waals equation correct for?
55
Where is most of an atom's mass concentrated?
56
What determines the size of an atom?
57
What is the role of the electron cloud in atomic structure?
58
How do electrons move according to the Rutherford model?
59
What generates the magnetic field in a wave?
60
In what way do the electric and magnetic fields in a wave relate to each other?
61
What is the definition of wavelength (λ)?
62
How is frequency (ν) defined and measured?
63
What is the relationship between the speed of light (c), wavelength (λ), and frequency (ν)?
64
How is light intensity related to wave properties?
65
What factor determines the color of light?
66
What happens when two waves are out of phase with each other?
67
What is constructive interference in wave behavior?
68
What is destructive interference in wave behavior?
69
List the categories of the electromagnetic spectrum in order from smallest to largest wavelength.
70
What does it mean for energy to be quantised?
71
According to Planck’s theory, why is electromagnetic radiation quantised?
72
What is the equation for photon energy?
73
How does classical physics describe the behavior of objects?
74
What does quantum mechanics describe about the behavior of tiny particles?
75
How can tiny particles behave in both wave-like and particle-like ways?
76
Why is it impossible to determine the exact location of a tiny particle?
77
Write the wave equation that relates speed of light, wavelength, and frequency.
78
Write the equation that relates photon energy to frequency and wavelength.
79
What is the value and unit of Planck's constant?
80
What is the electron cloud?
81
How does constructive interference differ from destructive interference?
82
What does the quantisation of energy imply about the way energy changes?
83
What does Planck's constant define in relation to photons?
84
What does wave-particle duality describe?
85
What is the main difference between absorption and emission spectroscopy?
86
How is energy absorbed in absorption spectroscopy?
87
What is indicated by the black lines in an absorption spectrum?
88
In absorption spectroscopy, what determines the specific energy states the sample transitions between?
89
How does the emission spectrum differ from the absorption spectrum in terms of the input and output?
90
What process does a sample undergo during emission spectroscopy?
91
What do the colored lines in an emission spectrum represent?
92
How do the emission lines in an emission spectrum compare to those in an absorption spectrum?
93
How can the atomic emission spectra be used to identify chemical elements?
94
Why are atomic emission spectra important in astrophysics and astrochemistry?
95
How can the emission spectrum of hydrogen help us understand the structure of the hydrogen atom?
96
What does the emission spectrum of hydrogen consist of, and what do these lines represent?
97
How does the energy level of an electron in hydrogen relate to the emitted light in the emission spectrum?
98
How can the Rydberg equation be used to predict the position of emission lines in the hydrogen spectrum?
99
What does the Rydberg frequency constant represent, and what is its value?
100
How are the emission lines in the hydrogen spectrum spaced at different frequencies?
101
How do the quantum numbers
𝑛initial and 𝑛final relate to absorption and emission?
102
How does the energy change (ΔE) for absorption compare to emission, according to the Rydberg equation?
103
What is the significance of the atomic number
𝑍
Z in the Rydberg equation?
104
How does Bohr’s model explain the discrete energy levels of electrons in hydrogen?
105
How does the Bohr model describe the process of energy absorption and emission for an electron?
106
What are the limitations of Bohr’s model?
107
Why does Bohr’s model work well for single-electron atoms but not for multi-electron atoms?
108
How is the Rydberg equation used to calculate the frequency of light emitted during an electronic transition?
109
What is the significance of Planck’s equation E=hν in relation to photon energy?
110
Define the electron ground state and its importance in atomic transitions.
111
What is the role of quantum numbers in determining electron energy levels?
112
What is the Rydberg constant, and how is it used in spectroscopy?
113
What is the Bohr model, and how does it describe electron orbits and energy levels?
114
How is momentum defined and calculated?
115
What is the law of conservation of momentum in collisions?
116
How does momentum behave in a collision?
117
What happens when EM radiation hits a metal surface?
118
Under what conditions are electrons ejected in the photoelectric effect?
119
How is the kinetic energy of ejected electrons related to the frequency of incoming light?
120
What happens to the number of electrons ejected when the intensity of light increases beyond the threshold frequency?
121
How did Einstein explain the photoelectric effect using Planck's concept?
122
What is the role of critical energy in the photoelectric effect?
123
List the characteristics of waves.
124
What is the energy of one photon and how is it calculated?
125
How is the number of photons in a light beam related to its intensity?
126
What is the result of the double slit experiment when light is used?
127
What was de Broglie’s hypothesis regarding matter and wavelengths?
128
What is the de Broglie wavelength formula and how is it derived?
129
Why are the de Broglie wavelengths of heavy particles difficult to observe?
130
What does Heisenberg’s Uncertainty Principle state?
131
How does the uncertainty in position (Δx) affect the uncertainty in momentum (Δp)?
132
What is the formula for the Uncertainty Principle?
133
How does the Uncertainty Principle challenge Bohr’s assumption of electrons moving in well-defined orbits?
134
How does the Uncertainty Principle apply to both macroscopic and microscopic particles?
135
What is Schrödinger’s equation used to model?
136
What is the wavefunction (Ψ) and what does it represent?
137
How does the wavefunction (Ψ) relate to the probability of finding an electron at a specific location?
138
What is electron density (ED) and how is it calculated from Ψ²?
139
How are electron densities measured experimentally?
140
What is the Schrödinger equation and how is it written?
141
What does solving Schrödinger’s equation provide for an electron or particle in a system?
142
What is the equation for kinetic energy (Ek)?
143
What is the equation for momentum (p)?
144
What is the energy equation for ejected electrons in the photoelectric effect (Ek(e⁻))?
145
What is the equation for Heisenberg’s Uncertainty Principle (Δx Δp ≥ h / 4π)?
146
What is the Schrödinger equation ($\hat{H}\Psi = E\Psi$)?
147
What does Ek stand for and how is it defined?
148
What is wavelength (λ) and how does it relate to the energy of light?
149
What is the wavefunction (Ψ) and why is it important in quantum mechanics?
150
What is electron density (ED) and how is it used in quantum mechanics?
151
What do Δx, Δp, and Δv represent in the context of the Uncertainty Principle?
152
What are the two terms in the Hamiltonian of the Schrödinger equation for a hydrogen atom?
153
How does the principal quantum number (n) relate to the energy and size of an electron's orbital in the hydrogen atom?
154
Define the term energetically degenerate orbitals. Why do orbitals with the same energy level have this property?
155
What is the significance of the angular momentum quantum number (ℓ) in determining the shape of orbitals?
156
Why is the Bohr model unable to predict the existence of orbitals for higher quantum numbers (n > 1)?
157
How do the radial and angular nodes differ, and how are they related to the quantum numbers n and ℓ?
158
Write down the Schrödinger equation solution for a hydrogen atom in terms of radial (R) and angular (Y) wavefunctions.
159
What is the radial distribution function? How does it relate to the probability of finding an electron at a given distance?
160
Why do orbitals with higher values of n become larger and have more nodes?
161
How does the magnetic quantum number (mₗ) affect the orientation of orbitals within a subshell?
162
What is the shape of an s orbital?
163
How does the size of an s orbital change with increasing n?
164
What is the shape of a p orbital?
165
What feature distinguishes p orbitals from s orbitals?
166
How many orientations do d orbitals have?
167
How many orientations do f orbitals have?
168
What does the principal quantum number (n) define in terms of an orbital?
169
What does the angular momentum quantum number (ℓ) determine?
170
What does the magnetic quantum number (mₗ) define for an orbital?
171
Why can't the Schrödinger Equation be solved exactly for multi-electron atoms?
172
What type of atomic orbitals are used to approximate solutions for multi-electron atoms?
173
Does orbital degeneracy apply to multi-electron atoms?
174
Are the three p orbitals within a subshell degenerate or non-degenerate?
175
Which orbital (s or p) has higher energy in the same principal shell?
176
What are the two possible electron spin states?
177
What are the two possible values for the spin magnetic quantum number?
178
According to the Pauli Exclusion Principle, how many electrons can occupy an orbital?
179
What does the Aufbau Principle state about how electrons fill orbitals?
180
What does Hund’s Rule state about how electrons are distributed in degenerate orbitals?
181
What do electrons in singly occupied orbitals have in common, according to Hund's Rule?
182
What are valence electrons, and why are they important?
183
What are core electrons, and are they involved in bonding?
184
What is shielding in relation to electron repulsion?
185
How does shielding affect the nuclear attraction on valence electrons?
186
How does effective nuclear charge (Z_eff) trend across a period and down a group?
187
How does shielding explain the energy difference between p and s orbitals in the same shell?
188
Why do 3d orbitals have higher energy than 4s orbitals?
189
How many electrons can 3d orbitals accommodate?
190
What element has a configuration of [Ar]4s¹3d⁵, and why is it stable?
191
What element has a configuration of [Ar]4s¹3d¹⁰, and why is it stable?
192
In which block are 3d and 4d transition metals found?
193
Where do lanthanide electrons fill, and which element marks the beginning of this block?
194
Where do actinide electrons fill, and which element marks the beginning of this block?
195
How do atomic radii trend across a period and down a group?
196
How do ion radii change when an atom forms a cation or an anion?
197
What is the trend for ionization energy across a period and down a group?
198
How is ionization energy related to atomic size and shielding?
199
What is electron affinity, and how does it change across a period and down a group?
200
How does electronegativity trend across a period and down a group in the periodic table?
201
What is the range of electronegativity values on the Pauling scale, and which elements are at the extremes?
