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chapter 7 Flashcards

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1
Q

how was the behavior of matter seen before the 20th century?

A

deterministic (present set of conditions used to predict future behavior)

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2
Q

are subatomic particles really deterministic?

A

no, the act of measuring/observing very small particles affects the system in unpredictable ways

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3
Q

what are quantum mechanics?

A

a set of guidelines subatomic particles follow

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4
Q

what is wave particle duality? does light have it?

A

a concept that every quantum entity can be described as either a particle or a wave. yes.

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5
Q

what is electromagnetic radiation?

A

the energy created from oscillating electric and magnetic fields that are perpendicular to each other

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6
Q

what is amplitude? determines?

A

height of trough, determines the intensity of the waves

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7
Q

what is wave length? what does it determine?

A

distance between two troughs/crests in the visible spectrum. determines the color

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8
Q

what is frequency?

A

the number of cycles that pass through a stationary point in a given period of time (Hertz, cycle per second)

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9
Q

formula for frequency

A

c (velocity of light) / wavelength

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10
Q

what does EM span from?

A

10^-6 m (gamma rays) to 10^5m (radio waves)

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11
Q

what is photoelectric effect?

A

the observation that metals can emit electrons when light shines upon them. classic EM theory could not account for it

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12
Q

what did einstein predict about light?

A

it is quantized, meaning light energy is delivered in packets called quanta or photons

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13
Q

what is the energy of a photon proportional to?

A

its frequency

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14
Q

what is the proportional constant?

A

proportionality constant is called Planck’s Constant, (h) and has the value 6.626 x 10−34 J∙s

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15
Q

how can orbits exist within bohrs model?

A

• Each of these orbits could only exist at a specific, fixed distance from the nucleus

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16
Q

what two things are the energy of orbits?

A

quantized or fixed, and proportional to the distance from the nucleus

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17
Q

how is the energy of orbits ranked in relation to the nucleus?

A
  • orbits closer to the nucleus are lower in energy than those further away,
    • En=1 < En=2
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18
Q

(hydrogen atom) what is the lowest energy orbit called?

A

n=1, the ground state

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19
Q

how can an electron go into excited state?

A

electron absorbs energy equal to the energy difference between orbits

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20
Q

what happens if enough energy is absorbed?

A

electron can be ejected from hydrogen

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21
Q

(atomic spectra) how long does the excited state last?

A

very short lifetime

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22
Q

what happens when the electron relaxes to ground state

A

light is emitted

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23
Q

emitted light has the same energy as ______?

A

absorbed energy

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24
Q

what did analogy with the photoelectric effect suggest?

A

suggested the light emission originated with the electrons in atoms. This suggests that an electron in an atom can only have certain energies - the electron energy levels are quantized

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25
Q

equation for transitions in HYDROGEN

A

ΔEelectron = Efinal state − Einitial state

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26
Q

where does an electron move when energy is emitted? is delta E positive or negative?

A

closer to the nucleus, s (nfinal>ninitial), so ΔE is positive

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27
Q

when an atom absorbs energy where does it move? is delta E positive or negative?

A

electron moves further from the nucleus (nfinal>ninitial), so ΔE is positive

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28
Q

small energy to big energy means?

A

absorption (usually)

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29
Q

what are the three problems with bohrs model?

A

○ Could not explain the emission spectra of multi electron elements
○ Electrons do not move in fixed orbitals
○According to classic physics electrons should emit ER and collapse into the nucleus

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30
Q

what did Louis de Broglie suggest?

A

Wave particle duality ;suggest that if energy is particle like, perhaps matter is wavelike
• Electrons interfering with themselves

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31
Q

what is the Heisenberg Uncertainty Principle?

A

you cannot observe the particle and wave nature of an electron at the same time

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32
Q

what is position?

A

particle nature

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33
Q

what is velocity?

A

wave nature

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34
Q

what is probability distribution?

A

electrons wont always hit the same place (five hole)

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35
Q

what is the Schrodinger equation? what does the answer give us?

A

• Solutions to the equation for a hydrogen atom yields many different solutions
The resulting functions are used to generate 3D plots to describe where an electron is found

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36
Q

how can standing waves exist?

A

A standing wave can also exist in a circular form.

Like linear standing waves, only certain wavelengths are allowed

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37
Q

what did Broglie and Schrodinger about Bohr’s model?

A

the Bohrs allowed electron orbits were the ones that could contain an integral number of wavelengths. • Other orbits produce destructive interference, and do not exist

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38
Q

what is a set of quantum numbers used for?

A

A set of three quantum numbers is used to describe each orbital

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39
Q

what is the principal quantum number? what does it describe?

A

n: describes the size and energy of the orbital

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40
Q

what is the angular momentum quantum number? what does it describe?

A

l: describes the shape of the orbital

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41
Q

what is the magnetic quantum number? what does it describe?

A

ml: describes the 3D orientation of the orbital

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42
Q

what is the magnetic spin number?

A

ms: describes the direction of electron spin; clockwise or counterclockwise

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43
Q

what numbers can n be? what happens as n gets larger?

A

N can be any integer greater than or equal to 1, the larger the more energy. the space between orbitals becomes smaller.

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44
Q

what values can l have?

A

values from 0 to n - 1, for each n value

45
Q

what is used to represent l values? from L=0 to L=3 list the designated letters.

A 
designated letters.  
l=0   is   s 
l=1    is   p
l=2   is   d
l=3   is    f
46
Q

what can values for ml be? what is each l (subshell) made up of?

A 
any integer (including zero) from -l to +l 
Each l (subshell) is made up of (2l+1) orbitals
47
Q

what is the number of integral values equal to?

A

the number of orbitals in the subshell (l)

48
Q

how many orientations are possible for a p orbital? what about an s orbital?

A

For a p-orbital (l=1), three orientations are possible (ml=-1, 0, +1)
For an s-orbital (l=0), only one orientation is possible (ml=0)

49
Q

which ways can electrons spin?

A

Electrons can spin either clockwise or counterclockwise

50
Q

define level. what is it given by?

A

an atom’s energy levels (or shells) are given by n.

51
Q

define sublevel. what is it given by?

A

an atoms level is divided into sublevels, or subshells,

that are given by the value of l. generally denoted based on letters

52
Q

define orbital

A

Each combination of n, l, ml species the size (energy),

shape, and spatial orientation of one of the atom’s orbitals.

53
Q

which ways can electrons spin?

A

Electrons can spin either clockwise (+1/2) or counterclockwise (-1/2)

54
Q

how many electrons can each orbital hold?

A

Each orbital can hold two electrons
No two electrons will have the same 4 quantum
numbers

55
Q

how are orbitals named?

A

We name orbitals based on their principle quantum number and subshell letter designation. ex: • n=2, l=1 would be called 2p

  • n=1, l=0 1s
  • n=2, l=0 2s
  • n=3, l=0, 1 or 2 3s, 3p, 3d
56
Q

what is a group of orbitals with the same value of n called?

A

a shell. all orbitals in the same quantum level share the same shell

57
Q

what is a group of orbitals with the same value of l called?

A 
a subshell ( e.g. the five d-orbitals together constitute a subshell).
   All orbitals in the same subshell have the same energy (they are termed degenerate).
58
Q

can atoms have the same four quantum numbers?

A

no two electrons can have the same set of four quantum numbers (i.e. two electrons in a filled orbital must have opposite spins and the net spin for the filled orbital must be zero (one must have spin +1/2 and the other spin -1/2).

59
Q

what do the first three quantum numbers provide?

A

The first three quantum numbers—n, l and ml—provide an approximate address for each electron of specified energy.
(l=1 for all p-orbitals)
(l=2 for all d-orbitals)
(l=0 for all s-orbitals)

60
Q

why is the shape of orbitals important?

A
  • Described by the angular quantum momentum number, l

* The shape of these orbitals is important in discussions of bonding and molecular shapes

61
Q

what shape are s orbitals? how many orbitals does each principal energy level have?

A
  • Spherical shape

* Each principal level (n) has one orbital

62
Q

what orbitals have the lowest energy? where does it first appear at?

A

s. appear at n= 1, growing with increasing n

63
Q

hoe many p orbitals are there? shape? l=? ml=?

A

(l=1)There are 3 P-orbitals dumbbell shaped, with two identical lobes separated by a node at the nucleus.
ml= -1, 0, and 1.

64
Q

how do the 3 p orbitals differ?

where do they lie?

A

differ only in their orientation, and are perpendicular to one another. labelled as Px, Py, and Pz, according to the axis.
grow larger as the value of n increases

65
Q

how many d orbitals does each principal energy state have? l=? ml=?

A

five. l=2. ml=-2,-1,0,1,2

66
Q

what are d orbitals named according to?

A

named according to the planes that the orbitals lie along.
• Four of the five orbitals are aligned in a different plane.
The fifth is aligned with the z axis

67
Q

d orbital shapes?

A
  • Mainly four-lobbed

* One is 2-lobed with a toroid

68
Q

when do f orbitals occur? l=? n=? ml=?

A

(l=3)
• • Each principal energy state above n = 4 has seven f orbitals
• - ml = −3, −2, −1, 0, +1, +2, +3

69
Q

• l is always less than n, ml ≥ -l and ≤+l

A

kk

70
Q

list the amount of orbitals each designated letter has

A

s=1 orbital
p = 3
d = 5
f = 7

71
Q

how is the periodic table organized? what does this highlight?

A

by similar chemical properties, highlighting similarity in electron configurations

72
Q

what is determined by an elements electron configuration?

A

its chemical and physical properties. explains bonding properties, reactivity, and periodic trends

73
Q

what is Aufbau’s principle?

A

fill the orbitals with electrons in increasing order of the principal quantum number, n

74
Q

what is the Pauli exclusion principle?

A

no two electrons in an atom can have the same four quantum numbers

75
Q

what is Hund’s rule?

A

within a subshell, degenerate (same energy) orbitals are each filled with one electron before being paired

76
Q

how do we fill hydrogen?

A
  • According to Aufbau principle, electrons fill orbitals in the order of lowest energy to greatest energy
    • Hydrogen has 1 electron
    • The lowest possible energy orbital for this electron to be in the s orbital in the n=1 shell
77
Q

how do we make the orbital diagram for helium?

A
  • Helium has two electrons
    • We add the electron into the lowest energy orbital
    • Pauli exclusion principle
    • Half arrows up and down= opposite spin
78
Q

what are degenerate orbitals?

A

orbitals with the same principal quantum number, n. have the same energy level

79
Q

how is the energy of an orbital determined?

A

determined by the electron proton attraction

80
Q

what three energy contributions must be considered when electrons occupy orbitals?

A

• KE of electrons
• PE of attraction between electrons and the nucleus
PE of repulsion between two electrons

81
Q

what happens with subshell splitting?

A
  • Principle levels with different l values are no longer degenerate
    • The lower the value of the l quantum number, the lower the energy of the subshell
82
Q

three reasons for subshell splitting?

A

S (l=0) < p (l-1) < d (l=2) < f (l=3)

coulomb’s law, shielding, penetration

83
Q

what does coulombs law do? does a system maximize or minimize PE?

A
  • Calculates the PE between two charged particles

* A system always wants to minimize its PE

84
Q

do particles with like charges have positive or negative PE?
opposite charges?

A

positive PE which decreases as the r increases

with opposite charges have a negative PE, which increases (becomes more negative) as the r shrinks

85
Q

what happens to the magnitude of interaction?

A

○ the magnitude of the interaction increases as the charge of the particles increase

86
Q

what forces do electrons feel? electrons higher in energy?

A
  • Any one electron feels attractive force from the positively charged nucleus and repulsive forces from other electrons
    • Electrons higher in energy shields the electron from feeling the full force of the nucleus
87
Q

which electrons are more shielded? 2D radial probability?

A
  • Electrons higher in energy are more shielded
    • In the 2D radial probability plots some orbitals overlap
    • Higher energy orbitals overlap into the same space as lower energy orbitals
    • Orbitals that penetrate feel less shielding from electrons in lower energy orbitals
88
Q

what does penetration do to orbitals with the same principle level?

A

penetration causes orbitals with the same principal level to have different energy levels (no longer degenerate)
• Only different subshells (ex not within 3p)

89
Q

what effect does penetration have on multi electron atoms?

A

Penetration has even more drastic effects at the higher principle levels

90
Q

describe spacing and order in multi electron atoms.

A

• Very small spacing at higher energy levels
• The ordering can therefore vary among electrons
• The ordering can therefore vary among elements
among variations present for transition metals and their ions

91
Q

are orbitals within a subshell degenerate? what does smaller energy differences mean?

A

• Orbitals within a subshell are degenerate

Smaller energy differences at the high levels means you get swapping depending on configuration

92
Q

how does period # relate to subhsells?

A

Period # is the same as subshell

93
Q

describe Hund’s rule

A
  • For adding electrons to degenerate orbitals of a subshell: electrons are added with parallel spins until each of the orbitals has one electron before a second electron is placed in one of these orbitals
    • As a result, when filling up the max number of unpaired electrons (and hence maximum spin state) is assured
94
Q

what would happen to row 2 with addition or subtraction of electrons?

A

• • Neon is a stable, non-reactive gas: filled n=2 subshell • Addition or subtraction of e- would increase energy

95
Q

how are inner electron configurations done?

A
  • Once a principal energy level has been filled, the electron configuration of elements is often abbreviated using the nearest noble gas in brackets
    • Superscripts is the # of electrons in each orbital
    • Big number is the energy level (number on table going across)
96
Q

how are transition metals filled?

A

Transition metals are elements that have a ground state configuration with partially filled d-orbitals, or oxidation states that give rise to partially filled d-orbitals
• Scandium, Sc, is the first transition metal
Aside from two notable exceptions, the rest of the period 4 elements fill their shells as we would expect from Hund’s Rule

97
Q

what are the two exceptions to hunds rule?

A

chromium and copper

98
Q

actual ground state of chromium. why?

A

[Ar]4s13d5
• More energetically favorable to have 5 unpaired d-electrons (half-filled) than to have paired 4s electrons and partially filled d-orbitals

99
Q

actual ground state of copper. why?

A

s [Ar]3d104s1

• A full 3d orbital is more energetically favorable that a full 4s orbital and partially filled 3d orbital

100
Q

how do we fill 4p elements? which are they?

A

• From Ga to Kr, to obtain correct electron configuration in each case, we simply fill up the 4 p orbitals following hunds rule.

101
Q

what do quantum mechanics predict about 8 valence electrons?

A

○ Quantum mechanics calculations predict that 8 valence electrons should produce a very unreactive atom
Conversely, atoms with 1 more or 1 fewer electrons should be very reactive

102
Q

what are the most reactive metals and non metals?

A

○ Halogens are the most reactive non metals

○ Alkali metals are the most reactive metals

103
Q

what metals don’t form/have noble gas configurations what do they form?

A

• Except aluminium, metals of group 13-15 do not form ions with noble gas configurations
Form ions with two different stable configuration; pseudo noble gas or inert pair configuration

104
Q

when is pseudo noble gas config formed?

A
  • If all electrons are emptied from its highest energy level
    • Loses all ns and np electrons, still has filled inner (n-10)d^10 electrons (tin, Sn)
105
Q

when is inert pair config formed?

A

○ If the metal loses just its np electrons, it can attain a stable configuration with filled ns and (n-1)d subshells
Retained ns^2 electrons are called an inert pair

106
Q

do main group ions with a noble gas config have high or low reactivity?

A
  • Low reactivity due to filled energy level (ns^2np^6)

* Metals and non metal main group elements ex: mg, P

107
Q

describe electron configs of transition metals

A

Elements in the d-block are normally unable to form noble gas electron configurations
• The next best configuration is a closed d-subshell
l • s-electrons are lost before the d-electrons

108
Q

define paramagnetic. attracted or repelled?

A

unpaired electrons, metals are attracted to a magnetic field

109
Q

define diamagnetic. attracted or repelled?

A

paired electrons, repelled

chem 101 (13 decks)

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