CH 2.6-2.12 Flashcards
(127 cards)
1
Q
Principal quantum number
A
The quantum number relating to the SIZE & ENERGY of an ORBITAL
- n
- shell
- has integer values: 1, 2, 3, …
2
Q
As n increases, the orbital becomes ____ & the electron spends more time _____ from the nucleus
A
larger; farther
3
Q
An increase in n = _____ energy, because the electron is ______ tightly bound to the nucleus, & the energy is _____ negative
A
higher; less; less
4
Q
orbital angular momentum quantum number
A
Quantum # relating to the SHAPE of an atomic ORBITAL
- Integral value from 0 -> n-1
- L
- subshell
- l = 0 is s
- l = 1 is p
- l = 2 is d
- l = 3 is f
5
Q
Magnetic quantum number
A
Quantum # relating to the ORIENTATION (direction) of an atomic ORBITAL in space relative to the other orbitals in the atom
- m sub l
- orbitals of subshell
- Integral values between +L and -L, including 0
6
Q
Nodes
A
- aka nodal surfaces
- an area of an orbital having 0 electron probability
7
Q
The number of nodes increases as _____ increases
A
n
8
Q
For s orbitals, the number of nodes is given by
A
n-1
9
Q
Function for s orbital
A
positive everywhere in 3D space
- when the see orbital function is evaluated at any point in space, its results are a positive #
10
Q
Function for p sub z orbital
A
has a positive sign in all regions of space in which z is positive & negative sign for when z is negative
- similar to sine wave with alternating positive and negative phases
11
Q
The d orbitals first occur in level
A
n = 3
12
Q
5 d orbitals
A
dxz, dyz, dxy, d(x^-y^2), d(z^2)
- x,y,z are subscripts
13
Q
dxy orbital
A
centered in the xy plane
- lie between the axes
14
Q
d(x^2-y^2)
A
- centered in the xy plane
- lies along the x and y axes
15
Q
d(z^2)
A
two lobes along z axis & a belt centered in the xy plane
16
Q
The f orbitals occur in level
A
n = 4
17
Q
All orbitals with the same value of n have the ______
A
same energy
18
Q
Degenerate
A
A group of orbitals with the same energy
19
Q
Summary of the Hydrogen Atom
A
- In the quantum (wave) mechanical model, the electron is viewed as a standing wave. This representation leads to a series of wave functions (orbitals) that describe the possible energies and spatial distributions available to the electron
- In agreement with the Heisenberg uncertainty principle, the model cannot specify the detailed electron motions. Instead, the square of the wave function represents the probability distribution of the electron in that orbital. This allows us to picture orbitals in terms of probability distributions, or electron density maps.
- The size of an orbital is arbitrarily defined as the surface that contains 90% of the total electron probability.
- The hydrogen atom has many types of orbitals. In the ground state, the single electron resides in the orbital. The electron can be excited to higher-energy orbitals if energy is put into the atom.
20
Q
Electron spin
A
a quantum # representing 1 of 2 possible values for the electron spin; either +1/2 (spin up) or -1/2 (spin down)
-ms
- developed by Samuel Goudsmit & George Uhlenbeck
- connected with Pauli’s postulate
- Magnetic field induced by the moving electric charge of the electron as it spins
- spins are opposite regardless of orientation
21
Q
Pauli Exclusion Principle
A
In a given atom no two electrons can have the same set of 4 quantum numbers (n, l, ml, ms)
- since ms has only 2 values, an orbital can only hold 2 electrons & they have opposite spins
- If there are 2 electrons in an orbital ONE MUST MAVE +1/2 spin and the OTHER MUST HAVE -1/2 SPIN
22
Q
Fraunhofer
A
- widened wavelength rainbow spectrum
- Saw blank & black spaces btwn colors (they weren’t continuous)
23
Q
Visible light characteristics
A
- White light contains all wavelengths (when passed through a prism, the different wavelengths are separated
24
Q
Line spectrum
A
- aka hydrogen emission spectrum
- Light from an electrical discharge through gaseous element doesn’t contain all wavelengths
- Spectrum is discontinuous (big gaps)
25
Continuous spectrum
occurs when white light is passed through a prism
26
Balmer-Rydberg Equation
1/lambda = R (1/(n1^2) - 1/(n2^2))
- (n1
27
Neils Bohr (1913)
- proposed model of an atom w/ discrete (quantum) states
- Explained how atoms emit light quanta & their stability
- Combined postulates of Planck & Einstein
28
Postulates of Bohr's theory
- Each atom has a # of discrete energy levels (orbits) (stationary states) that exist w/out emitting/absorbing EM radiations
- As the orbital radius decreases, so does the energy of the electron
- An electron may move from one energy level (orbit) to another, but in so doing, monochromatic radiation is EMITTED (TO LOWER ENERGY) or ABSORBED (TO HIGHER ENERGY)
- an electron "revolves" in a circular orbit about the nucleus & its motion is STILL governed by the ordinary laws of mechanics & electrostatics, with the restriction that its ANGULAR MOMENTUM IS QUANTIZED & CONSERVED
29
Angular momentum equation
angular momentum = m x v x r = nh/(2pi)
- m = mass of electron
- v = velocity of electron
- r = radius of orbit
- n = energy levels
- h = Planck's constant
30
Energy of states of electrons equation
- A/ n^2
- A = constant
- n = quantum #
31
Bohr equation
E = -2.178 x 10^-18 J (Z^2/n^2)
- n = integer
- z = nuclear charge
32
What trend occurs as you go up in energy levels
Distance between the orbitals decrease
33
Radius of Bohr Hydrogen orbit equation
r = n^2 (5.29 x 10^-11 m)
34
Any one-electron atom radius equation
r = (n^2 x a0)/Z
35
Ionization energy
The energy required to remove the electron from an atom
- each ionization energy is removing 1 highest-energy electron
- express in kJ per mole of atoms (or ions)
36
Change in energy of ionization of an atom equation
delta E = -AZ^2 (1/(nf^2) - 1/(ni^2))
- delta E > 0 means energy is absorbed
- delta E < 0 means energy is emitted
37
The change in energy for one atom of hydrogen
- 2.178 x 10^-18 J
- Ionization for 1 mol = 2.178x10^-18 J x 6.022x10^23= 13.12 x 10^5 J mol^-1
38
As energy is brought closer to the nucleus, energy is _______ the system --> ______ atomic stability --> energy becomes more ______ relative to the free electron
released from; increased; negative
38
As energy is brought closer to the nucleus, energy is _______ the system --> ______ atomic stability --> energy becomes more ______ relative to the free electron
released from; increased; negative
39
Limitations of Bohr model
- only explains hydrogen
- neither accounts for intensities nor the fine structure of the spectral lines for hydrogen
- couldn't explain binding of atoms
- unexplained quantum jumps (not good physics)
40
evidence for wave-nature of light
- diffraction & interference
- optical microscopes
- EM wave
41
evidence for particle-nature of light
- photoelectric effect
- compton effect
- black body effect
- photons
- convert light to electric current
42
Compton Effect
X-ray scattering by e- in atoms- photons have momentum
43
Number of photons is = to
energy density (square of EM field strength)
44
What is the wave aspect of matter
radiation (light)
45
as momentum increases, wavelength gets _____
shorter
46
as momentum decreases, wavelength gets ______
longer
47
heavy particles of matter are mainly _____-like
particle
48
extremely light particles of matter are mainly _____-like
wave
49
Particles in order from most particle-like to wave-like
proton, electron, photon
50
does the wave nature of a particle have a visibility threshold
yes,
the wave nature of a particle will only show up when the scale p is comparable (or smaller) with the size of h
51
wave nature of electrons
- matter waves
- electron microscope
- discrete (quantum) states of confined systems, such as atoms
52
particle nature of electrons
- electric current
- photon-electron collisions
53
parts of a wavefunction
ψ = R(r) Y(φ, θ)
- R(r) = radial part, gives you range of position
- Y(φ, θ) angular part, gives you shape of
54
What does the radial part of a wave function give you
- the principal quantum number (circumference), n
- the orbital angular momentum quantum number, l (max # determined by n-1)
55
What does the angular part of a wave function give you
- the orbital angular momentum quantum number, l
- the magnetic quantum number, m sub l (determined by +-l)
56
An orbital (wavefunction) is completely characterized by..
quantum numbers n, l, & m sub l
- defines orbit of electron
- solves Bohr model
- consistent w/spectrum lines & Bohr model
57
What is special about the orbitals of the H-atom
all orbitals of the same value of n have the same energy
58
equation for # of orbitals
n^2
59
ψ
value of the amplitude of the electron wave at any position in space
60
ψ^2
corresponds to the probability of finding the electron at any point in space
- probability density
61
Electron density map
the more time the electron visits a particular point, the darker the negative becomes
62
Radial probability distribution grap
plots the total probability of finding an electron in each spherical shell vs the distance from the nucleus`
63
size of 1s orbital
radius of the sphere that encloses 90% of the total electron probability
64
The larger the # of radial nodes in an orbital, the _____ the energy of the orbital
higher
65
# of radial nodes =
n - l - 1
66
# of planar (angular) nodes increases with ______
l
67
total # of nodes =
(radial + angular) = n - 1
68
orbital
- a wavefunction defined by quantum number n, l, & m sub l
- a region of space 'occupied' by an electron
- have energies, shapes, & specific orientations in space
69
d orbitals
- clover shaped
- important for metals
-
70
f orbitals
- not involved in bonding in any compounds
71
____ are said to be isoelectronic with the H-atom
Ions
72
The attractive Coulomb force is much bigger due to the higher _____ charge of the nuclei
Z+, (more protons)
73
Increasing Z charge has what trend
atomic orbitals contract & come closer to the nucleus
74
What is the effect of the Z^2 term
to increase the energy level spacing
75
What is a flaw of the schrodinger equation
fails if atom has > 1 atom
76
What fields does an electron have
an intrinsic magnetic field that interacts with its orbital magnetic field
77
Dirac
- tried to solve Schrodinger's equation
- extended wave mechanics to include special relativity & an extra coordinate-- time
- yielded a new QN- an intrinsic angular momentum of the electron (ELECTRON SPIN)
78
greater nuclear charge _____ sublevel energy
lowers
79
Orbital approximation
- make a polyelectronic atom like a 1e- atom
- assumes every electron to be on its own experiencing an eff nuc charge
- then the orbitals for the electrons take the form of those in H but their energy & sizes are modified by simply using eff nuc charge
80
electrons in the same level lead to a _____ lower Zeff
slightly
81
Electrons in an inner level lead to a _____ lower Zeff
much
82
Electrons in _____ energy levels shield the outer electrons very effectively
inner
83
Penetration
- to get close to the nucleus
- energy does down when it gets closer to the nucleus; also ^ negative
- arises from angular nodes in wave function
- occurs when electron gets excited
84
screening/shielding
to block the view of other electrons of the nucleus
85
Penetration _____ nuclear attraction & ____ shielding
increases; decreases
86
higher electron energy =
- easy to remove electron
- farther from nucleus
87
splitting
- caused by penetration & its effect on shielding
88
order of sublevel energies
s < p < d < f
89
order of effective nuclear charge in a polyelectronic atom
zeff (s) > zeff (p) > zeff (d)
90
Energies of atomic orbitals are affected by
- nuclear charge
- shielding by other electrons
91
A higher nuclear charge
- increases nucleus-electron interactions
- lowers sublevel energy
92
Shielding by other electrons _____ the full nuclear charge to an eff nuclear charge
reduces
93
orbital shape affects what energy
sublevel energy
94
Effective nuclear charge
lower than actual nuclear charge
- increases toward nucleus (ns > np > nd > nf)
- Zeff = Z - S
- Z = protons
- S = electrons shielding, # core electrons
95
General energies of subshells of the same principal QN
s < p < d < f
96
electron-electron repulsion
- removes degeneracy
- electrons less tightly bound to the nucleus
97
Quantum Wave Mechanics of the Atom explains...
electron configuration of the atoms
98
Aufbau rule
- aka building up principle
- determines electron config
- electrons are always placed in the lowest energy sublevel available
1. within a shell (n) the filling order is s>p>d>f
2. within a subshell (l), lowest energy arrangement has the highest # of unpaired spin (Hund's)
3. The (n+1)s orbitals always fill before the nd orbitals
4. after lanthanum, the 4f orbitals are filled
5. after actinium, the 5f orbitals are filled
99
core electrons
electrons that are in filled, inner shells & not involved in chem rxn
100
Maximizing the # of parallel spins- the exchange interaction
- each electron pair w/ parallel spins leads to a lowering of the electronic energy of the atom
101
p subshell orbitals
3 orbitals w/ same energy
102
d subshell orbitals
5 orbitals w/same energy
103
f subshell orbitals
7 orbitals w/same energy
104
Hund's rule
when orbitals of equal energy are available, the lowest energy electron config has the max # of unpaired electrons w/parallel spins
- parallel spin config of lower energy for degenerate orbitals
- every orbital has to occupy one electron first before doubling up
105
Elements in the same group of the PT....
- have the same outer electron config
- exhibit similar chemical behavior
- therefore... similar outer electron configs correlate w/similar chem behavior
- are isoelectronic w/respect to # of valence electrons
106
Ground state of Chromium
[Ar] 4s1 3d5
- because of 2 half-filled subshells
107
Ground state of Copper
[Ar] 4s1 3d10
- because of half-filled subshell & a filled subshell
108
filled subshells and # electrons
s: 2 e-
p: 6 e-
d: 10 e-
f: 14 e-
109
Isoelectronic
atoms/ions that have identical #'s & electorn configs
110
cation
positive ion formed by removing electrons from atoms
111
anion
negative ion formed by adding electrons to atoms
112
removal of electrons from transition metals
remove ns electrons & then (n-1)d electrons
113
paramagnetic
- attracted to a magnet
- atoms (or ions) w/unpaired electorns
114
diamagnetic
- repelled by a magnet
- atoms (or ions) w/out unpaired electrons
115
Elements in the same row...
- show regular trends in their properties due to the continuing increase in # of valence electrons until a shell is filled
116
atomic radii
- Half the distance btwn nearest atoms in element
- for nonmetallic atoms; estimated form their covalent solids (elements/compounds) since they don't form diatomic molecules
- for metallic atoms; calculating half of distance btwn metal atoms in solid metal crystals\
- for ions;interatomic distance in ionic crystals
117
Valence electrons
- major role in periodic properties of atoms
- simultaneously attracted to + nucleus & - electrons
118
as shielding increases, effective nuclear charge...
decreases
119
Trends across a period
- # protons increase
- # core electrons stays same
- valences drawn closer to nucleus
- Zeff increases
- atomic & ionic radius decreases
- Ionization energy increase
120
Trends down a group
- atomic & ionic radius increases
- ionization energy decrease
121
Ions and their relative sizes
- cation smaller than neutral atom (cores exposed, zeff incr, remaining e more tightly bounded)
- anion bigger than neutral atom (increase e-e repulsion, no change of zeff)
122
a larger atom has ___ potential energy
lower
123
ionization energy levels in order
l1 < l2 < l3
- removing an electron from a positive ion is more difficult than removing it form a neutral atom
124
Metals
- form cations
- have smaller # of electrons in valence shells
- low ionization energies
- s, d, f , and some p block
125
metallic character
- related to atomic radius & ionization energy
- increases to left of period & down
126
nonmetal
- larger numbers of electrons in valence shells
- form anions
- all in p-block
