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CHEM 1311 > CH 2.6-2.12 > Flashcards

CH 2.6-2.12 Flashcards

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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, …

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2
Q

As n increases, the orbital becomes ____ & the electron spends more time _____ from the nucleus

A

larger; farther

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3
Q

An increase in n = _____ energy, because the electron is ______ tightly bound to the nucleus, & the energy is _____ negative

A

higher; less; less

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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

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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

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6
Q

Nodes

A
  • aka nodal surfaces
  • an area of an orbital having 0 electron probability
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7
Q

The number of nodes increases as _____ increases

A

n

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8
Q

For s orbitals, the number of nodes is given by

A

n-1

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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 #

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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

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11
Q

The d orbitals first occur in level

A

n = 3

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12
Q

5 d orbitals

A

dxz, dyz, dxy, d(x^-y^2), d(z^2)
- x,y,z are subscripts

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13
Q

dxy orbital

A

centered in the xy plane
- lie between the axes

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14
Q

d(x^2-y^2)

A
  • centered in the xy plane
  • lies along the x and y axes
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15
Q

d(z^2)

A

two lobes along z axis & a belt centered in the xy plane

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16
Q

The f orbitals occur in level

A

n = 4

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17
Q

All orbitals with the same value of n have the ______

A

same energy

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18
Q

Degenerate

A

A group of orbitals with the same energy

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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.
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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

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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

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22
Q

Fraunhofer

A
  • widened wavelength rainbow spectrum
  • Saw blank & black spaces btwn colors (they weren’t continuous)
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23
Q

Visible light characteristics

A
  • White light contains all wavelengths (when passed through a prism, the different wavelengths are separated
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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)
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25
Q

Continuous spectrum

A

occurs when white light is passed through a prism

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26
Q

Balmer-Rydberg Equation

A

1/lambda = R (1/(n1^2) - 1/(n2^2))
- (n1

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27
Q

Neils Bohr (1913)

A
  • proposed model of an atom w/ discrete (quantum) states
  • Explained how atoms emit light quanta & their stability
  • Combined postulates of Planck & Einstein
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28
Q

Postulates of Bohr’s theory

A
  • 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
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29
Q

Angular momentum equation

A

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

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30
Q

Energy of states of electrons equation

A
  • A/ n^2
  • A = constant
  • n = quantum #
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31
Q

Bohr equation

A

E = -2.178 x 10^-18 J (Z^2/n^2)
- n = integer
- z = nuclear charge

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32
Q

What trend occurs as you go up in energy levels

A

Distance between the orbitals decrease

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33
Q

Radius of Bohr Hydrogen orbit equation

A

r = n^2 (5.29 x 10^-11 m)

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34
Q

Any one-electron atom radius equation

A

r = (n^2 x a0)/Z

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35
Q

Ionization energy

A

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)

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36
Q

Change in energy of ionization of an atom equation

A

delta E = -AZ^2 (1/(nf^2) - 1/(ni^2))
- delta E > 0 means energy is absorbed
- delta E < 0 means energy is emitted

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37
Q

The change in energy for one atom of hydrogen

A
  • 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
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38
Q

As energy is brought closer to the nucleus, energy is _______ the system –> ______ atomic stability –> energy becomes more ______ relative to the free electron

A

released from; increased; negative

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38
Q

As energy is brought closer to the nucleus, energy is _______ the system –> ______ atomic stability –> energy becomes more ______ relative to the free electron

A

released from; increased; negative

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39
Q

Limitations of Bohr model

A
  • 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)
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40
Q

evidence for wave-nature of light

A
  • diffraction & interference
  • optical microscopes
  • EM wave
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41
Q

evidence for particle-nature of light

A
  • photoelectric effect
  • compton effect
  • black body effect
  • photons
  • convert light to electric current
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42
Q

Compton Effect

A

X-ray scattering by e- in atoms- photons have momentum

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43
Q

Number of photons is = to

A

energy density (square of EM field strength)

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44
Q

What is the wave aspect of matter

A

radiation (light)

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45
Q

as momentum increases, wavelength gets _____

A

shorter

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46
Q

as momentum decreases, wavelength gets ______

A

longer

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47
Q

heavy particles of matter are mainly _____-like

A

particle

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48
Q

extremely light particles of matter are mainly _____-like

A

wave

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49
Q

Particles in order from most particle-like to wave-like

A

proton, electron, photon

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50
Q

does the wave nature of a particle have a visibility threshold

A

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
Q

wave nature of electrons

A
  • matter waves
  • electron microscope
  • discrete (quantum) states of confined systems, such as atoms
52
Q

particle nature of electrons

A
  • electric current
  • photon-electron collisions
53
Q

parts of a wavefunction

A

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
Q

What does the radial part of a wave function give you

A
  • the principal quantum number (circumference), n
  • the orbital angular momentum quantum number, l (max # determined by n-1)
55
Q

What does the angular part of a wave function give you

A
  • the orbital angular momentum quantum number, l
  • the magnetic quantum number, m sub l (determined by +-l)
56
Q

An orbital (wavefunction) is completely characterized by..

A

quantum numbers n, l, & m sub l
- defines orbit of electron
- solves Bohr model
- consistent w/spectrum lines & Bohr model

57
Q

What is special about the orbitals of the H-atom

A

all orbitals of the same value of n have the same energy

58
Q

equation for # of orbitals

A

n^2

59
Q

ψ

A

value of the amplitude of the electron wave at any position in space

60
Q

ψ^2

A

corresponds to the probability of finding the electron at any point in space
- probability density

61
Q

Electron density map

A

the more time the electron visits a particular point, the darker the negative becomes

62
Q

Radial probability distribution grap

A

plots the total probability of finding an electron in each spherical shell vs the distance from the nucleus`

63
Q

size of 1s orbital

A

radius of the sphere that encloses 90% of the total electron probability

64
Q

The larger the # of radial nodes in an orbital, the _____ the energy of the orbital

A

higher

65
Q

of radial nodes =

A

n - l - 1

66
Q

of planar (angular) nodes increases with ______

A

l

67
Q

total # of nodes =

A

(radial + angular) = n - 1

68
Q

orbital

A
  • 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
Q

d orbitals

A
  • clover shaped
  • ## important for metals
70
Q

f orbitals

A
  • not involved in bonding in any compounds
71
Q

____ are said to be isoelectronic with the H-atom

A

Ions

72
Q

The attractive Coulomb force is much bigger due to the higher _____ charge of the nuclei

A

Z+, (more protons)

73
Q

Increasing Z charge has what trend

A

atomic orbitals contract & come closer to the nucleus

74
Q

What is the effect of the Z^2 term

A

to increase the energy level spacing

75
Q

What is a flaw of the schrodinger equation

A

fails if atom has > 1 atom

76
Q

What fields does an electron have

A

an intrinsic magnetic field that interacts with its orbital magnetic field

77
Q

Dirac

A
  • 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
Q

greater nuclear charge _____ sublevel energy

A

lowers

79
Q

Orbital approximation

A
  • 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
Q

electrons in the same level lead to a _____ lower Zeff

A

slightly

81
Q

Electrons in an inner level lead to a _____ lower Zeff

A

much

82
Q

Electrons in _____ energy levels shield the outer electrons very effectively

A

inner

83
Q

Penetration

A
  • 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
Q

screening/shielding

A

to block the view of other electrons of the nucleus

85
Q

Penetration _____ nuclear attraction & ____ shielding

A

increases; decreases

86
Q

higher electron energy =

A
  • easy to remove electron
  • farther from nucleus
87
Q

splitting

A
  • caused by penetration & its effect on shielding
88
Q

order of sublevel energies

A

s < p < d < f

89
Q

order of effective nuclear charge in a polyelectronic atom

A

zeff (s) > zeff (p) > zeff (d)

90
Q

Energies of atomic orbitals are affected by

A
  • nuclear charge
  • shielding by other electrons
91
Q

A higher nuclear charge

A
  • increases nucleus-electron interactions
  • lowers sublevel energy
92
Q

Shielding by other electrons _____ the full nuclear charge to an eff nuclear charge

A

reduces

93
Q

orbital shape affects what energy

A

sublevel energy

94
Q

Effective nuclear charge

A

lower than actual nuclear charge
- increases toward nucleus (ns > np > nd > nf)
- Zeff = Z - S
- Z = protons
- S = electrons shielding, # core electrons

95
Q

General energies of subshells of the same principal QN

A

s < p < d < f

96
Q

electron-electron repulsion

A
  • removes degeneracy
  • electrons less tightly bound to the nucleus
97
Q

Quantum Wave Mechanics of the Atom explains…

A

electron configuration of the atoms

98
Q

Aufbau rule

A
  • 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
Q

core electrons

A

electrons that are in filled, inner shells & not involved in chem rxn

100
Q

Maximizing the # of parallel spins- the exchange interaction

A
  • each electron pair w/ parallel spins leads to a lowering of the electronic energy of the atom
101
Q

p subshell orbitals

A

3 orbitals w/ same energy

102
Q

d subshell orbitals

A

5 orbitals w/same energy

103
Q

f subshell orbitals

A

7 orbitals w/same energy

104
Q

Hund’s rule

A

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
Q

Elements in the same group of the PT….

A
  • 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
Q

Ground state of Chromium

A

[Ar] 4s1 3d5
- because of 2 half-filled subshells

107
Q

Ground state of Copper

A

[Ar] 4s1 3d10
- because of half-filled subshell & a filled subshell

108
Q

filled subshells and # electrons

A

s: 2 e-
p: 6 e-
d: 10 e-
f: 14 e-

109
Q

Isoelectronic

A

atoms/ions that have identical #’s & electorn configs

110
Q

cation

A

positive ion formed by removing electrons from atoms

111
Q

anion

A

negative ion formed by adding electrons to atoms

112
Q

removal of electrons from transition metals

A

remove ns electrons & then (n-1)d electrons

113
Q

paramagnetic

A
  • attracted to a magnet
  • atoms (or ions) w/unpaired electorns
114
Q

diamagnetic

A
  • repelled by a magnet
  • atoms (or ions) w/out unpaired electrons
115
Q

Elements in the same row…

A
  • show regular trends in their properties due to the continuing increase in # of valence electrons until a shell is filled
116
Q

atomic radii

A
  • 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
Q

Valence electrons

A
  • major role in periodic properties of atoms
  • simultaneously attracted to + nucleus & - electrons
118
Q

as shielding increases, effective nuclear charge…

A

decreases

119
Q

Trends across a period

A
  • # protons increase
  • # core electrons stays same
  • valences drawn closer to nucleus
  • Zeff increases
  • atomic & ionic radius decreases
  • Ionization energy increase
120
Q

Trends down a group

A
  • atomic & ionic radius increases
  • ionization energy decrease
121
Q

Ions and their relative sizes

A
  • 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
Q

a larger atom has ___ potential energy

A

lower

123
Q

ionization energy levels in order

A

l1 < l2 < l3
- removing an electron from a positive ion is more difficult than removing it form a neutral atom

124
Q

Metals

A
  • form cations
  • have smaller # of electrons in valence shells
  • low ionization energies
  • s, d, f , and some p block
125
Q

metallic character

A
  • related to atomic radius & ionization energy
  • increases to left of period & down
126
Q

nonmetal

A
  • larger numbers of electrons in valence shells
  • form anions
  • all in p-block

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