Lectures 1-8 Flashcards
(169 cards)
1
Q
gases at room temperature
A
HEFONC & noble gases
2
Q
1 L
A
1 dm^3
3
Q
1 Pa
A
1 N/m^2
4
Q
1 bar
A
1 x 10^5 Pa
5
Q
1 kPa
A
1000 Pa
6
Q
equal volumes of gas at constant T and p contain
A
equal number of particles
7
Q
volumes of all gases extrapolate to zero at
A
0 K
8
Q
R
A
8.314 J K^-1 mol^-1
9
Q
V is proportional to
A
1/p
10
Q
V is proportional to
A
T
11
Q
V is proportional to
A
n
12
Q
density
A
pM/RT
13
Q
to find partial pressure
A
n(a) x (RT/V)
14
Q
to find total pressure
A
(n(a) + n(b) + n(c)) x (RT/V)
15
Q
to find mole fraction
A
n(a)/n(total)
16
Q
to find mole fraction
A
p(a)/p(total)
17
Q
to find partial pressure
A
x(a)/p(total)
18
Q
kinetic molecular theory assumptions
A
negligible particle size, elastic collisions, no interaction
19
Q
average kinetic energy of particles
A
3/2kT where k is R/N(A)
20
Q
RMS speeds related to
A
temperature and molar mass
21
Q
all gases at same temperature have same
A
kinetic energy, but not same average speed
22
Q
effusion
A
gas escapes through a hole
23
Q
diffusion
A
different gases mix
24
Q
diffusion occurs from regions of
A
high to low concentration
25
rate of effusion is inversely proportional to
square root of molar mass
26
rate of effusion
number of molecules passing through a hole per second
27
particles attract each other
p decreases
28
particles repel each other
p increases
29
wavelength
distance from maximum to maximum
30
frequency
how many cycles a second
31
wavelength and frequency are
anti-proportional
32
wave number
inverse of wavelength; number of wavelengths per unit distance
33
light intensity is related to
amplitude
34
light colour is related to
wavelength
35
constructive interference
when two waves in phase interfere
36
destructive interference
when two waves out of phase interfere (amplitude = 0)
37
wavelength and energy content are
anti-proportional
38
same temperature
same kinetic energy
39
to calculate RMS, molar mass must be in
kg/mol
40
RMS
(3RT/M)^1/2
41
rate of effusion of gas 1/rate of effusion of gas 2
SQUARE ROOT OF: molar mass gas 2/molar mass gas 1
42
p1V1/T1 =
p2V2/T2
43
frequency =
c/wavelength
44
photons
energy packets of light
45
energy of a single photon (E)
hv
46
absorption spectroscopy
input light, analyse what was absorbed
47
in absorption spectroscopy, the sample transitions from a
lower to higher energy state
48
emission spectroscopy
input heat, analyse what was emitted
49
in emission spectroscopy, sample transitions from
higher to lower energy state
50
deltaE
E2-E1
51
predicting positions of lines in emission spectra
RH(1/n^2final - 1/n^2initial)
52
energy of emitted light
-hv or -hRydberg equation
53
quantum numbers
n(initial) and n(final)
54
emission quantum numbers
n(final) < n(initial) and deltaE is negative
55
absorption quantum numbers
n(final) > n(initial) and deltaE is positive
56
the Bohr model does not work because
emission spectra of many-electron atoms cannot be described, does not explain intensity of spectral lines, does not take into account particles as waves
57
photoelectric effect
electromagnetic radiation hits a metal surface and electrons are emitted
58
electrons are only ejected if light reaches
critical/threshold frequency
59
intensity increase causes
higher number of electrons to be ejected
60
photons must overcome both
binding energy and threshold frequency
61
kinetic energy of electron
hv(photon) - Ebind
62
kinetic energy of electron
hv(photon) - h x threshold frequency
63
Hz
1/s
64
de Broglie wavelength of matter
wavelength = h/p = h/mv
65
wavelength and mass are
anti-proportional
66
fast-moving particles have
shorter wavelengths
67
slow-moving particles have
longer wavelengths
68
Heisenberg's Uncertainty Principle
it is not possible to accurately detect the exact position of a particle and its momentum at the same time
69
Heisenberg's Uncertainty Principle
deltaxdeltap >/ h/4pi
70
wavefunction
used to describe the standing wave for an electron
71
for each energy state of an electron/particle, we assign a different
wavefunction
72
regions with positive and negative signs for each wavefunction
phases (there is a negative phase and a positive phase)
73
regions where wavefunction is zero
nodes
74
wavefunction^2
related to probability of finding a particle at a particular point in space
75
when wavefunctions are squared, nodes are
unchanged
76
point with maximum probability of finding a particle
maximum of wavefunction^2
77
point where probability of finding particle is zero
node
78
electron density
calculated through wavefunction^2, gives us probability of finding an electron in area around nucleus
79
orbital
a one-electron wavefunction in three-dimensional space
80
energy levels are always
negative
81
orbitals with the same energy are called
energetically degenerate
82
angular momentum quantum number; l
determines type & shape of orbital; tells us subshell/letter
83
limit for angular momentum quantum number (l)
n-1
84
magnetic momentum quantum number; ml
determines orientation of orbital; allows us to distinguish between each individual p, d orbital
85
limit to magnetic momentum quantum number (ml)
-l and l
86
principle quantum number; n
determines energy for that shell
87
s orbitals are
spherical
88
p orbitals are
dumbbell-shaped
89
two lobes of p orbitals have different
signs
90
two lobes of p orbitals separated by
angular node/nodal plane (zero)
91
keep radius fixed and
plot angles to explore shape
92
keep angles fixed and
play around with radius (can be graphed)
93
higher n,
higher number of nodes (for example, 1s orbital has no nodes, 2s orbital has one node, 3s orbital has two nodes)
94
the first time an orbital appears there are
no nodes
95
shell volume increases with
radius
96
wavefunction and probability decrease with
r
97
for larger shells, you're more likely to find an electron
further away from the nucleus
98
to find wavelength when given mass
wavelength = h/mv
99
in many electron atoms, orbital degeneracy
disappears
100
electron spin is
quantised; two different angular momenta
101
spin magnetic quantum number ms for spin up/clockwise
+1/2
102
spin magnetic quantum number ms for spin down/anti-clockwise
-1/2
103
in an atom, no two electrons can have the same
four quantum numbers
104
the same orbital can only hold two electrons if they have
opposite spins
105
orbitals cannot hold more than
two electrons
106
fill up lower-lying orbitals
first
107
only fill next orbital once orbital below is
occupied
108
principal quantum number determines
row
109
superscripted number
number of electrons in that orbital
110
two degenerate orbitals want to be
singly occupied first
111
number of p orbitals in a shell
3 (6 electrons)
112
number of d orbitals in a shell
5 (10 electrons)
113
same group, same
electronic configuration in outer shell
114
core electrons
electrons in shells that are energetically lower than outer shell
115
nucleus has
positive charge Z
116
shielding is a
consequence of electron repulsion, which reduces net interaction between the positive nucleus and valence electrons
117
valence electrons do not experience full Z, only
effective charge Zeff
118
why do p orbitals have higher energies than s orbitals for many-electron atoms?
shielding
119
because 3d orbitals are shielded by 3s and 3p orbitals, they are
less stable and have higher energies than 4s orbitals
120
3d is filled after
4s
121
4d is filled after
5s
122
4f is filled after
6s
123
5d is filled after
4f
124
number of f orbitals in shell
7; 14 electrons
125
first transition metals
3d metals (fourth period)
126
4d transition metals appear after
5s orbital has been filled (fifth period)
127
chromium electronic configuration
[Ar]4s^13d^5
128
additional stabilisation in chromium is due to
half-filled 3d subshell with all 5 electrons having parallel spin
129
copper electronic configuration
[Ar]4s^13d^10
130
stabilisation in copper due to
fully-filled 3d subshell
131
after lanthanum, we fill
4f orbitals before remaining 5d orbitals
132
after actinium, we fill
5f orbitals before remaining 5d orbitals
133
after actinium, we fill
5f orbitals before remaining 5d orbitals
134
ionisation energy for a single H atom
13.6 eV
135
covalent radius
atomic radius estimated by halving distance between two chemically bound nuclei of the same type
136
size increases
down a group
137
as we go down a group,
principal quantum number (n) increases
138
size decreases
across a period
139
across a period, electrons from outer shell are drawn closer to nucleus due to
increasing positive charge
140
across a period, electron-electron repulsion
increases as we increase the number of electrons
141
cations are always
smaller than the parent atom
142
anions are always
larger than the parent atom
143
isoelectric
two species that contain the same number of electrons
144
for an isoelectric series, size decreases as
number of protons increases/nuclear charge increases
145
ionisation energy (IE)
energy required to remove the highest energy (most loosely held) electron from an atom/ion
146
removing an electron is
endothermic
147
first ionisation energy < second < third because
consecutively removed electrons cause a reduction in electron-electron repulsion and experience a greater attraction to Z
148
IE decreases
down a group
149
IE increases
across a period
150
across a period, nuclear charge
increases but the number of core electrons remains the same
151
high ionisation energy
small size
152
low ionisation energy
large size
153
electron affinity (EA)
the energy change associated with the addition of an electron to a gaseous atom
154
electron affinity is
exothermic, generally
155
the closer the vacant orbital is to the nucleus, the more favourable the addition of an electron and the
more negative EA; more energy released
156
EA decreases/is less negative
down a group
157
EA increases/is more negative
across a period
158
EA is not always exothermic due to
electron-electron repulsion
159
electronegativity
the ability of an atom to attract shared electrons to itself
160
electronegativity numerically
IE + EA
161
electronegativity increases
across a period
162
electronegativity decreases
down a group
163
metallic character increases
down a group
164
metallic character decreases
across a period
165
metals have relatively low
IEs
166
total number of orbitals in a shell
n^2
167
substance that has one or more unpaired electrons
paramagnetic
168
substance that has no unpaired electrons
diamagnetic
169
substance that has unpaired electrons that are aligned in a particular direction
ferromagnetic
