Quantum Mechanics Flashcards
(143 cards)
1
Q
Pauli Exclusion Principle
A
No 2 electrons in an atom can have the same set of 4 quantum numbers
2
Q
Principal Quantum Number, n
A
Gives shell/orbital- tells average distance of electron from the nucleus
N=1 closest to nucleus
N=3- further from nucleus
Goes to infinity n#s
3
Q
Angular Momentum Quantum #
A
L - tells the shape of the orbital,
& gives subshell
Whatever N is, L includes all #s, including zero, b4 the # N
L= 0 to n-1
Ex. N=1, L can only be 0
N=2, L= 0 or 1
N=3, L= 0, 1, or 2
4
Q
L orbital numbers correspond to what letters?
A
0 = S
1 = P
2 = D
3 = F
4 = G
5 = H
5
Q
Angular Momentum Orbital, L shape of orbital
N=1
A
N=1, L=0 (S) (1 value)
N=2, L= 0 or 1 (S or P) (2 values)
N=3, L= 0, 1, or 2 (S, P, or D) (3 values)
6
Q
3rd Quantum #
A
Magnetic Quantum #, m
Tells spacial orientation & gives available energy levels within a subshell
MsubL ranges from -L to +L
For N=3, L can equal:
L=0, & ML= 0
L=1, & ML=-1,0,1
L=2, & ML=-2,-1,0,1,2
7
Q
If n=1,
L=?
mL=?
A
L=0
ML= 0
8
Q
If n=2
L= ?
ML=?
A
L= 0 ML = 0
L= 1 ML= -1, 0, 1
N=3
L= 0, 1, 2
ML= 0
ML=-1, 0, 1
ML= -2, -1, 0, 1, 2
9
Q
- Last Quantum #
Invented by and year?
What is it?
A
Electron Spin Quantum #
1924 Stern & Gerlach
Electrons have two different spins = clockwise and counterclockwise
which equal +1/2 and -1/2
10
Q
What is electromagnetic theory?
A
The spinning charges of electrons generate magnetic fields
+1/2 clockwise
-1/2 counter clockwise
11
Q
Stern & Gerlach
A
+1/2 & -1/2 electron spin determination
1924 Shot gaseous electrons, which passed through a slit (so the beam would be focused), and then they pass it through a magnet
They noticed 1/2 electrons went up and 1/2 electrons went downwards
So they determined electrons have two different spins, one is clockwise & one is counterclockwise
12
Q
Electromagnetic theory
A
Stern & Gerlach
Spinning charges generate magnetic field
+1/2 & -1/2
13
Q
Each box on orbitals can hold how many electrons? And what does each electron represent?
A
Each box holds 2 electrons that represent either +1/2 or -1/2 magnetic spin of the electron to represent one spinning clockwise and the other counter clockwise
14
Q
Summary of quantum #s
Principal = ?
Angular Momentum = ?
Magnetic = ?
Spin = ?
A
Principal = N (1,2,3,4,…)
Angular Momentum = L (n-1)
Magnetic = ML= (-L, …, -1,0,1, ….L)
Spin = +1/2, -1/2
15
Q
Quantum number chart
A
Picture at bottom
16
Q
-Predict the number of subshells in the fourth shell, which would be n=4
-Give the label for each of these subshells
-How many orbitals are in each of these subshells?
A
4 different subshells- 0, 1, 2, 3
Labels = 4s, 4p, 4d, 4f
4s=1orbital
4p=3 orbitals
4d=5 orbitals
4f=7orbitals
16 total orbitals = 32 electrons
17
Q
Give all possible subshells, and designations when n=3
A
L= 0, 1, or 2 which = 3s, 3p, 3d
18
Q
How many orbitals and electrons in an atom can have the following sets of quantum numbers?
A) n=2
B) n=2, L=0
C) n=2, L=2
D) n=10
A
A) 2s, 2p, 2p, 2p = 4 orbitals, 8 electrons
B) 2s orbital 2 electrons in there
C) L can’t be 2 so 0 orbitals, 0 electrons
D) square it = 100 orbitals 200 electrons
19
Q
How many electrons can have the quantum numbers:
N=2
L=0
ML=0
Ms= -1/2
A
1 electron, only 1 spin
20
Q
List all the quantum numbers for the designation 3D
A
N=3
L= 0, 1, or 2 (s, p, d)
L has to be 2
ML= -2,-1,0,1,2
Ms= +1/2 or -1/2
21
Q
List all of the sets of quantum numbers for 1s
A
1 orbital , 2 electrons , 2 sets of quantum #’s = (1,0,0,+1/2) and (1,0,0,-1/2)
22
Q
Give the possible subshell designations, mL values (is the # of orbitals in shell), & total # of orbitals for:
1) n=2
2) n=3
3) n=4 square this # to get the #of orbitals
A
1) N=2, mL= -1, 0, 1 / Subshells 2s=0 has 1 orbital, 2p=-1,0,1 has 3 orbitals
2) N=3, mL=-2,-1,0,1,2 / Subshells 3s=0, 3p=-1,0,1, 3d=-2,-1,0,1,2 has 5 orbitals
3) N=4, mL=-3,-2,-1,0,1,2,3 / 4s=0, 4p=-1,0,1, 4d=-2,-1,0,1,2, 4f=-3,-2,-1,0,1,2,3
23
Q
The most stable arrangement of electrons is the one with the most parallel spins is a description of what?
A
Hund’s Rule
24
Q
When you have a value for n, how can you easily find the # of orbitals ?
A
Square the n value to get the number of orbitals
25
What does L =
0 to n-1
26
What subshell belongs to L=0?
S subshell
27
What subshell belongs to L= 3?
F subshell
28
What subshell belongs to L=2?
D Subshell
29
What subshell belongs to L=1?
P Subshell
30
How many orbitals are in the 4f subshell?
7 orbitals listed below
-3,-2,-1,0,1,2,3
31
How many orbitals are in the P subshell?
3 orbitals
With #s -1,0,1
32
What’s the total number of orbitals in n=3?
9
Square it
33
What tells us that electrons fill lowest energy orbitals first?
Aufbau Principal
34
Orbital filling diagram
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
35
How do you write the Noble Gas Shortcut Electron Configuration?
Our previous Noble Gas in Brackets and continue with the configuration of the row the element resides in
36
Electron configuration exceptions happen because why?
Because half filled and completely filled D orbitals are very stable
A filled D orbital would have 10 electrons, 1/2 of that is 5.
Example Cr should be [Ar]4s2 3d4 but in order to half fill the D orbital, it takes 1 from s and moves it to D for [Ar]4s1 3d5
Same will happen with Mo, W, Cu, Ag, Au
37
What is an Orbital Diagram
Picture representation of electron configuration with
BOXES representing ORBITALS, and ARROWS representing ELECTRONS
38
Hydrogen Orbital Diagram
39
Orbital Diagrams for
He &
Li
40
Orbital Diagram for Phosphorus
Electrons fill all boxes with 1 before doubling up to 2 electrons and the arrows point up first and go back to fill the down arrows after all orbitals are spun up first
41
Magnetism- what are the 2 types?
Paramagnetic= attracted by a magnet bc electrons are UNPAIRED - (para =1)
Diamagnetic = repelled by magnet bc electrons are all PAIRED (dia=2)
NEED ORBITAL DIAGRAM TO DETERMINE
42
Examples of diamagnetic and paramagnetic diagrams
Beryllium is Diamagnetic
All odd number electron elements are diamagnetic
but even numbers, we must draw the diagram to know bc how they fill orbitals before pairing.
43
19th Century Scientists arranged elements by?
Atomic Mass
44
19th century scientists believe mass was related to
Chemical Behavior
45
Heavier element has a larger? And higher?
A heavier element has a larger DENSITY and higher BOILING POINT
46
Who claimed that every 8th element had a similar property?
What year? And what was it called?
1864 John Newlands- “Law of Octaves”
—-In quotes bc it’s not a law, doesn’t work for everything
Works up to right b4 aluminum
47
Who redid the periodic table and left spaces for future elements, and predicted unknown future elements properties?
Dmitri Mendeleev
48
Who was Henry Moseley?
1913 found relationship between atomic # and frequency of x-ray generated when elements are bombarded with electrons
Hydrogen had pattern #1
Helium had pattern #2
And that’s how the atomic number came about
49
Elements have similar properties when
They have the same number of electrons in the “valence” (outermost) shell
- which puts them in the same group #
50
If we form a CATION we _____ electrons, so we have 1________electron?
If we form a CATION we LOOSE electrons, so we have 1 LESS electron!
51
If we form an ANION we _____ electrons, so we have 1________electron?
If we form an ANION we GAIN electrons, so we have 1 MORE electron!
52
FLOURINE electron configuration is
1s2 2s2 2p5
If we form a an ANION we get
F minus
F— 1s2 2s2 2p6 which becomes Neon
That’s why flourine has a charge of -1 as an ion that wants to gain or loose charges
53
Na 1s2 2s2 2p6 3s1
Wants to do what??
Loose an electron to become a positively charged cation
It becomes Na+1 with 1 less electron
-Then it becomes 1s2 2s2 2p6 - behaving like Neon
Which makes Sodium and Neon isoelectronic with each other
54
When elements have the same configuration b/c they lost or gained electrons to become the same electron configuration as another element, we say they are ________ with each other?
Isoelectronic
F-1, Ne, Na+1 are isoelectronic with each other
*Fluorine and Neon are not isoelectronic, F- and Neon are.
55
____ electrons become anions
_____ electrons become Cations
Gain electrons to become anions
Loose electrons to become cations
56
How are electrons removed from transition metals?
First the s orbital, then the d orbital
Ex Mn 4s2 3d5 - the outer shell is 4s so we removed 1 from there first so Mn +1 would be [Ar] 4s1 3d5
If we take a second electron, is still comes from the 4s shell before the 3d shell so Mn+2 is [Ar] 3d5
57
What is the atomic radius of an atom?
1/2 the distance between 2 nuclei in 2 adjacent atoms.
This is determined by attraction of nucleus & valence electrons to the protons in the nucleus
58
A stronger attraction makes a _____ atomic radius ?
Smaller
59
A weaker attraction makes a ____ atomic radius?
Larger
60
The radius of elements on the periodic table gets larger as we go
Left and down
61
As you go down in groups, the atomic mass becomes significantly _____ and has more _____?
Larger and has more cells
62
The atomic radius decreases as we go ____ on the periodic table?
And why?
Right bc we get more protons and electrons attracting each other closely with a stronger pull to the nucleus making it pulled together smaller
63
What affects the physical and chemical properties of ionic compounds?
Ionic radius - radius of an ion
(anion or cation)
64
If an atom forms an anion, the radius gets?
If an atoms forms a negatively charged ANION, The radius gets LARGER
that means putting more electrons in the outer shell which repel electrons already there making the radius larger
65
If an atom forms a CATION, we ____electrons from the outer shell
This goes what to the atomic radius?
If an atom forms a CATION, we REMOVE electrons from the outer shell
Less repulsion, smaller radius
66
Smallest elements are located where on the periodic table?
Top right
Largest on the bottom left
67
The largest atomic radius on the table would be the 1 furthest
Bottom left
68
What is the minimum amount of energy needed to remove an electron from a gas atom in its ground state?
Ionization energy
69
To take an element and remove an electron we would display it how and what is it called?
X(g) —> X(g)+1 + e- 1st Ionization Energy
Charges must add to 0 on the first side
70
How would we display the removal of a 2nd electron and what would it be called?
X+1(g) -> X+2(g) + e- 2nd ionization energy
X plus 1 YIELDS x plus 2 plus 1 electron
71
How would you display the 100 ionization energy?
X+99(g) —-> X+100(g) + e-
72
How would you display the fifth ionization energy of phosphorus?
P+4(g)—-> P+5(g) + e-
73
Ionization energy increases how as we move along the periodic table?
It increase from bottom to top, left to right
As we go right, we get closer to a full shell and as we get closer to becoming a noble gas, it gets harder to remove an electron
74
The higher we are in a group, we have _____ shells which makes the electrons ____ to the protons in the nucleus which has what kind of hold on those electrons?
The higher we are in a group, we have LESS shells which makes the electrons CLOSER to the protons in the nucleus and has a STRONGER hold on those electrons
75
Ionization energy is harder to break what to loose electrons?
Higher ionization energy means it wants to loose electrons more
It’s harder to break full shells so if an electron configuration includes a full outer shell, that configuration wants to loose electrons less than a configuration who’s shell is partially full so the one that wants to loose electrons more is the higher ionization energy
76
The energy change associated by the acceptance of an electronic in the gas state is called what?
Electron Affinity
77
As ionization energy _________ electrons, electron affinity ________ electrons.
As ionization energy LOOSES, electrons electron affinity, GAINS electrons
78
Examples of Electron Affinity of X
Show 100th electron Affinity
X(g) + e- —> X-1(g)
X-99(g) + e- ——> X-100(g)
Further right, closer to a full shell, they want to gain an electron more.
HIGHER electron affinity means WANTS TO GAIN ELECTRONS
79
Highest 1st election affinity means
It wants 1 electron the most!
80
Lowest electron affinity means?
Which one wants An electron the least?
81
Lowest 4th electron affinity means
Lowest 6th?
Which does not want 4 electrons
It would be nitrogen
Lowest 6th would be boron
82
When n= a number, how do you find how many orbitals are in that n shell?
N² = total # of ORBITALS in N Shell
N² *(2)= # of ELECTRONS in Shell
You square the number that n= and the number n=squared (n^2) is the total number of orbitals in that shell
Multiply the # of ORBITALS BY 2 to get the # of ELECTRONS in shell
83
How many total electrons in N=1? And who do they belong to?
N=1 has 2 electrons,
1 belongs to
Hydrogen & 1 to Helium
84
Total # of electrons in an orbital =
The # of elements in that orbital
85
What shape is a 1S orbital?
Spherical
86
What shape is the 3P orbital and where do they lie?
3p orbitals are dumbbell shaped and lie -horizontally on the X-Axis
-vertically on the Z-Axis
-& the Y-Axis at pi/4 going in and out of
The sheet of paper
87
The most stable arrangement of electrons is the one with the most _____________?
What makes this statement?
HUND’S RULE
The most stable arrangement of electrons is the one with the most parallel spins
88
What does the Aufbau Principle State?
Electrons fill lowest energy orbitals first
89
Aufbau comes from?
German Word - “BUILDING UP”
90
Which orbital is closest to the nucleus?
N=1
91
What does is take for an electron to move from N=1 to N=3?
More energy to move an electron to further out, higher up n= orbitals
Values for N increases as the distance from the nucleus increases and the energy required to move them further out increases.
92
What increases as N= orbital values increase?
— distance from the nucleus increases
— energy required to move electrons further out increases
93
How do electrons fill orbitals?
Lowest Energy Orbitals FIRST!
3D is more energy than 4s (more
electrons)
4P is more energy than 4s and 3D
94
Electron Configurations
Hydrogen 1s1
Helium 1s2
Lithium 1s2 2s1
Beryllium 1s2 2s2
95
Atomic Number is equal to
The # of PROTONS &
The # of ELECTRONS
96
What number should tell you how many electrons should be in your electron configuration?
Atomic #
97
Boron electron configuration
Carbon ?
Neon?
Argon?
Barium?
Scanium?
Barium?
B - 1s2 2s2 2p1
C - 1s2 2s2 2p2
N - 1s2 2s2 2p6
Ar [Ne] 3s2 3p6
Ba [Xe] 6s2
Sc [Ar] 4s2 3d1
Ba [Xe] 6s2
98
When do electron configurations change from the normal structure
When they are close to being
- 1/2 filled d orbitals (5 electrons)
- completely filled d orbitals (10 electrons)
Then electrons will shift from s orbital to d orbital leaving s shells semi filled
Electron config=more stable-full or 1/2 full
99
Give an example of the electron configuration exception with d orbitals
Cu expected as [Ar] 4s2 3d9
Observed as Cu [Ar] 4s1 3d10
100
Sample electron configuration drawings
101
The most stable arrangement of electrons is the one with the most parallel spins is an example of what?
Hund’s Rule
102
Valence Electrons =
Electrons in the Outer Shell (also group #)
Reminder: if an ion has a negative charge, that adds the extra number of negative electrons to the ion to make it negative
103
What makes a substance paramagnetic ?
Paramagnetic- attracted to a magnet-
happens when they have electrons that are unpaired - spin is only upwards!
Basically electrons filled every orbital, but didn’t go back and fill every orbital a second time for the counter clockwise spin, so some electrons are left unpaired and that makes this substance paramagnetic cause it wants to pair !
104
What does it mean when a substance is repelled by a magnet?
It is Diamagnetic
Died magnetism - none- all paired up
(Because all electrons are paired so its ability to pair up has died)
105
Para/diamagnetic
Para - 1 electron (for a magnet)
Dia- 2 (dos magnetic - 2 electrons each)
106
What elements are paramagnetic in general and applies to all in this category?
ODD Atomic # = paramagnetic
All elements who’s atomic # is an odd # is a diamagnetic element bc there is no way an odd number can all be paired up
EVEN Atomic # must draw orbital diagram
Because we fill orbitals completely before giving electrons a partner
107
Carbon has an even atomic number, is it diamagnetic or paramagnetic?
Paramagnetic
2p orbital has 2 electrons and it has to have 3 by themselves before before it goes back around and gives each 1 a second electron so therefore, its paramagnetic unless the orbital is filled
108
What says that spinning charges generate magnetic fields
Electromagnetic theory
Stern & Gerlach (1924)
Electron spin, quantum # Ms
109
Horizontal rows on the periodic table are called
Periods
110
Vertical columns on the periodic table are called?
Groups
111
What are substances with unpaired electrons that are attracted by a magnet?
Paramagnetic
112
What are substances who’s electrons are all paired and they are repelled by a magnet
Diamagnetic
113
19th century scientist arranged elements on the periodic table by?
Atomic mass
114
Atomic mass is related to what?
Chemical behavior
115
Every eighth element on the periodic table has similar properties according to what?
“ law of octaves” John Newlands
116
Who improved John Newland‘s “law of octaves” by designing the periodic table
1869 Dimitri Mendeleev
- Group 66 elements by property
- left spaces for future elements
- predicted the properties of unknown future elements
117
Eka-aluminum (Ea) is an element that was left a space for on the periodic table when it was designed by Mendeleev.
What was it discovered as and named four years later?
Gallium
118
Henry Moseley 1913
Discovered a relationship between atomic number and frequency of x-rays generated when elements are bombarded with electrons
How the atomic number came about - discovered a unique pattern that was generated for every element, that was unique to that element
119
When elements have the same number of electrons in their valence shell, which is their outermost shell, these elements all have what?
Similar properties
The valance shell is the outermost which consists of let’s say it’s the three shell. It consists of everything in the 3 shell, the S and the P orbital= 3s and 3p are the outer most 3 shell
120
If we form a cation, we _______ electrons?
We lose electrons
1 less electron gives us 1 more proton, which means we have a + charge.
121
If we form a anion, we _______ electrons?
We become negative by GAINING electrons
When we form an anion, we have 1 MORE ELECTRON
122
The Flourine atom can form an anion F- by gaining an electron. What does this configuration look like?
F 1s2 2s2 2p5 becomes
F- 1s2 2s2 2p6 with a full p shell
Neon has this same electron configuration and elements want to behave like noble gases so they are looking to do this exact thing
Flourine wants to become Neon (noble) and when this happens, they become isoelectronic with each other
F- and Neon are isoelectronic and have the same electron configuration- F minus is isoelectronic with Neon not Flourine!
123
For transition metals, how are electrons removed?
From the S orbital first and then the D orbital
124
Electrons are removed from the outermost shelf first so if an electron ends in 4S2 3D5, where do we take the electron from first?
4S2 bc 4 is outer more than 3
if we take a 2nd electron it comes from the 2nd electron in the 4S orbital because that still comes before the 3D orbital
125
What does the electron configuration of MN+1 look like?
The removed electron moved to the left side as a a +1 and it took away 1 from the right side
Mn+1 = [Ar] 4s1 3d5
Mn would be [Ar] 4s2 3d5
The +1 means it lost an electron that has a negative charge so now it has a +1 charge and the electron was lost from the 4S orbital because 3d5 is 1/2 full & we don’t break 1/2 full shells
126
How does atomic radius increase?
To the LEFT and DOWN
127
What element on the periodic table would have the smallest atomic radius?
Top right
128
If an atom forms an anion, what happens to the atomic radius?
It increases because you have more e- which = more repulsion/bigger atomic radius
129
If an atom forms a cation, what happens to the radius?
It decreases because less e- means less repulsion, which = a tighter circle
130
The minimum amount of energy needed to remove an electron from a gas atom and its ground state is called what?
Ionization energy
You start with a gas atom and it goes to the same gas atom with a plus charge plus an electron and whatever charge that new atom has, is the level of ionization energy you are at , so if it’s X+2(g) you’re on the second ionization energy
131
Energy change associated by the acceptance of an electron in the gas state is called what?
Electron affinity
X(g) + e- —> X-1(g)
Adding electrons to the left side of the arrow which leads to a negative charge on the right side
132
Ionization energy leads to a positive charge on the right side plus a negative electron
Electron affinity - adds a negative electron to the left side of the arrow and leads to an atom with a negative charge on the right side
Ionization energy happens on the right side of the arrow and deals with positive charges
Electronic affinity happens on the left side of the arrow and deals with negative charges
133
Ionization energy ______ electrons
Electron affinity _________ electrons
Ionization energy loses, electrons
Electron affinity gains electrons
134
Electronic affinity and
ionization energy
both increased as
The element gets closer to a full shell
Aka they increase up and right
135
High electronic affinity means?
Wants to gain electrons
The closer you get to the right side of the table, the more the element wants to gain electrons to get to period 8 Noble
136
How do you calculate the formal charge of an atom in a Lewis Structure?
Total # of valence electrons in atom
- total # of non-bonding electrons
- 1/2(total # of bonding electrons)
Neutral molecule = 0 formal charge
Cations = + formal charge (positive)
Anions = - formal charge (negative)
137
1 of 2 or more Lewis Structures for a single molecule that cannot be accurately represented by only 1 Lewis Structure is known as?
Resonance - the use of 2 or more Lewis Structures to represent a molecule
Ex. O3- Ozone
138
Octet Rule Exceptions
What is an incomplete octet and an example of?
In some compounds, the # of electrons surrounding the central atom is < 8
Example BeH2
H—Be—H
139
Odd Electron Molecule Example:
NO & NO2
NO = 5+6=11 NO2 = 5+6+6=17
.:N=O:: ::O=N.—O:::
Reasonance: :::O—N.=O::
N has an unpaired electron on all structures for both of these
140
What is an expanded Octet?
In some compounds, the # of electrons surrounding the central atom is more than 8
Ex. SF6
6+7(6)=48
141
What is the enthalpy change required to break a particular bond in 1 mole of gaseous molecules called?
Bond Dissociation Energy- which is a measure of stability
🔼H= Esum(bond energy of reactants)-
- Esum(bond energy of products)
Total energy input - total energy released
Bonds broken - bonds formed
142
Calculate the enthalpy change for the reaction: H2(g) + F2(g) —> 2HF(g)
H—H + :::F—F::: —> H—F:::
Bonds Broken Bonds Formed
143
2H2 + O2 —> 2H2O
Show bond dissociation energy
2 H—H + ::O=O:: —> 2 H—:O:—H
