chapter 7 Flashcards
(109 cards)
1
Q
how was the behavior of matter seen before the 20th century?
A
deterministic (present set of conditions used to predict future behavior)
2
Q
are subatomic particles really deterministic?
A
no, the act of measuring/observing very small particles affects the system in unpredictable ways
3
Q
what are quantum mechanics?
A
a set of guidelines subatomic particles follow
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.
5
Q
what is electromagnetic radiation?
A
the energy created from oscillating electric and magnetic fields that are perpendicular to each other
6
Q
what is amplitude? determines?
A
height of trough, determines the intensity of the waves
7
Q
what is wave length? what does it determine?
A
distance between two troughs/crests in the visible spectrum. determines the color
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)
9
Q
formula for frequency
A
c (velocity of light) / wavelength
10
Q
what does EM span from?
A
10^-6 m (gamma rays) to 10^5m (radio waves)
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
12
Q
what did einstein predict about light?
A
it is quantized, meaning light energy is delivered in packets called quanta or photons
13
Q
what is the energy of a photon proportional to?
A
its frequency
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
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
16
Q
what two things are the energy of orbits?
A
quantized or fixed, and proportional to the distance from the nucleus
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
18
Q
(hydrogen atom) what is the lowest energy orbit called?
A
n=1, the ground state
19
Q
how can an electron go into excited state?
A
electron absorbs energy equal to the energy difference between orbits
20
Q
what happens if enough energy is absorbed?
A
electron can be ejected from hydrogen
21
Q
(atomic spectra) how long does the excited state last?
A
very short lifetime
22
Q
what happens when the electron relaxes to ground state
A
light is emitted
23
Q
emitted light has the same energy as ______?
A
absorbed energy
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
25
equation for transitions in HYDROGEN
ΔEelectron = Efinal state − Einitial state
26
where does an electron move when energy is emitted? is delta E positive or negative?
closer to the nucleus, s (nfinal>ninitial), so ΔE is positive
27
when an atom absorbs energy where does it move? is delta E positive or negative?
electron moves further from the nucleus (nfinal>ninitial), so ΔE is positive
28
small energy to big energy means?
absorption (usually)
29
what are the three problems with bohrs model?
○ 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
30
what did Louis de Broglie suggest?
Wave particle duality ;suggest that if energy is particle like, perhaps matter is wavelike
• Electrons interfering with themselves
31
what is the Heisenberg Uncertainty Principle?
you cannot observe the particle and wave nature of an electron at the same time
32
what is position?
particle nature
33
what is velocity?
wave nature
34
what is probability distribution?
electrons wont always hit the same place (five hole)
35
what is the Schrodinger equation? what does the answer give us?
• 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
36
how can standing waves exist?
A standing wave can also exist in a circular form.
| Like linear standing waves, only certain wavelengths are allowed
37
what did Broglie and Schrodinger about Bohr's model?
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
38
what is a set of quantum numbers used for?
A set of three quantum numbers is used to describe each orbital
39
what is the principal quantum number? what does it describe?
n: describes the size and energy of the orbital
40
what is the angular momentum quantum number? what does it describe?
l: describes the shape of the orbital
41
what is the magnetic quantum number? what does it describe?
ml: describes the 3D orientation of the orbital
42
what is the magnetic spin number?
ms: describes the direction of electron spin; clockwise or counterclockwise
43
what numbers can n be? what happens as n gets larger?
N can be any integer greater than or equal to 1, the larger the more energy. the space between orbitals becomes smaller.
44
what values can l have?
values from 0 to n - 1, for each n value
45
what is used to represent l values? from L=0 to L=3 list the designated letters.
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designated letters.
l=0 is s
l=1 is p
l=2 is d
l=3 is f
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46
what can values for ml be? what is each l (subshell) made up of?
```
any integer (including zero) from -l to +l
Each l (subshell) is made up of (2l+1) orbitals
```
47
what is the number of integral values equal to?
the number of orbitals in the subshell (l)
48
how many orientations are possible for a p orbital? what about an s orbital?
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
which ways can electrons spin?
Electrons can spin either clockwise or counterclockwise
50
define level. what is it given by?
an atom’s energy levels (or shells) are given by n.
51
define sublevel. what is it given by?
an atoms level is divided into sublevels, or subshells,
| that are given by the value of l. generally denoted based on letters
52
define orbital
Each combination of n, l, ml species the size (energy),
| shape, and spatial orientation of one of the atom’s orbitals.
53
which ways can electrons spin?
Electrons can spin either clockwise (+1/2) or counterclockwise (-1/2)
54
how many electrons can each orbital hold?
Each orbital can hold two electrons
No two electrons will have the same 4 quantum
numbers
55
how are orbitals named?
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
what is a group of orbitals with the same value of n called?
a shell. all orbitals in the same quantum level share the same shell
57
what is a group of orbitals with the same value of l called?
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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
can atoms have the same four quantum numbers?
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
what do the first three quantum numbers provide?
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
why is the shape of orbitals important?
* Described by the angular quantum momentum number, l
| * The shape of these orbitals is important in discussions of bonding and molecular shapes
61
what shape are s orbitals? how many orbitals does each principal energy level have?
* Spherical shape
| * Each principal level (n) has one orbital
62
what orbitals have the lowest energy? where does it first appear at?
s. appear at n= 1, growing with increasing n
63
hoe many p orbitals are there? shape? l=? ml=?
(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
how do the 3 p orbitals differ?
| where do they lie?
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
how many d orbitals does each principal energy state have? l=? ml=?
five. l=2. ml=-2,-1,0,1,2
66
what are d orbitals named according to?
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
d orbital shapes?
* Mainly four-lobbed
| * One is 2-lobed with a toroid
68
when do f orbitals occur? l=? n=? ml=?
(l=3)
• • Each principal energy state above n = 4 has seven f orbitals
• - ml = −3, −2, −1, 0, +1, +2, +3
69
• l is always less than n, ml ≥ -l and ≤+l
kk
70
list the amount of orbitals each designated letter has
s=1 orbital
p = 3
d = 5
f = 7
71
how is the periodic table organized? what does this highlight?
by similar chemical properties, highlighting similarity in electron configurations
72
what is determined by an elements electron configuration?
its chemical and physical properties. explains bonding properties, reactivity, and periodic trends
73
what is Aufbau's principle?
fill the orbitals with electrons in increasing order of the principal quantum number, n
74
what is the Pauli exclusion principle?
no two electrons in an atom can have the same four quantum numbers
75
what is Hund's rule?
within a subshell, degenerate (same energy) orbitals are each filled with one electron before being paired
76
how do we fill hydrogen?
* 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
how do we make the orbital diagram for helium?
* Helium has two electrons
* We add the electron into the lowest energy orbital
* Pauli exclusion principle
* Half arrows up and down= opposite spin
78
what are degenerate orbitals?
orbitals with the same principal quantum number, n. have the same energy level
79
how is the energy of an orbital determined?
determined by the electron proton attraction
80
what three energy contributions must be considered when electrons occupy orbitals?
• KE of electrons
• PE of attraction between electrons and the nucleus
PE of repulsion between two electrons
81
what happens with subshell splitting?
* 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
three reasons for subshell splitting?
S (l=0) < p (l-1) < d (l=2) < f (l=3)
| coulomb's law, shielding, penetration
83
what does coulombs law do? does a system maximize or minimize PE?
* Calculates the PE between two charged particles
| * A system always wants to minimize its PE
84
do particles with like charges have positive or negative PE?
opposite charges?
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
what happens to the magnitude of interaction?
○ the magnitude of the interaction increases as the charge of the particles increase
86
what forces do electrons feel? electrons higher in energy?
* 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
which electrons are more shielded? 2D radial probability?
* 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
what does penetration do to orbitals with the same principle level?
penetration causes orbitals with the same principal level to have different energy levels (no longer degenerate)
• Only different subshells (ex not within 3p)
89
what effect does penetration have on multi electron atoms?
Penetration has even more drastic effects at the higher principle levels
90
describe spacing and order in multi electron atoms.
• 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
are orbitals within a subshell degenerate? what does smaller energy differences mean?
• Orbitals within a subshell are degenerate
| Smaller energy differences at the high levels means you get swapping depending on configuration
92
how does period # relate to subhsells?
Period # is the same as subshell
93
describe Hund's rule
* 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
what would happen to row 2 with addition or subtraction of electrons?
• • Neon is a stable, non-reactive gas: filled n=2 subshell • Addition or subtraction of e- would increase energy
95
how are inner electron configurations done?
* 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
how are transition metals filled?
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
what are the two exceptions to hunds rule?
chromium and copper
98
actual ground state of chromium. why?
[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
actual ground state of copper. why?
s [Ar]3d104s1
| • A full 3d orbital is more energetically favorable that a full 4s orbital and partially filled 3d orbital
100
how do we fill 4p elements? which are they?
• From Ga to Kr, to obtain correct electron configuration in each case, we simply fill up the 4 p orbitals following hunds rule.
101
what do quantum mechanics predict about 8 valence electrons?
○ 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
what are the most reactive metals and non metals?
○ Halogens are the most reactive non metals
| ○ Alkali metals are the most reactive metals
103
what metals don't form/have noble gas configurations what do they form?
• 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
when is pseudo noble gas config formed?
* 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
when is inert pair config formed?
○ 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
do main group ions with a noble gas config have high or low reactivity?
* Low reactivity due to filled energy level (ns^2np^6)
| * Metals and non metal main group elements ex: mg, P
107
describe electron configs of transition metals
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
define paramagnetic. attracted or repelled?
unpaired electrons, metals are attracted to a magnetic field
109
define diamagnetic. attracted or repelled?
paired electrons, repelled
