AP Chem - Chapter 7 - Electrons and Periodicity Flashcards
1
Q
what is wavelength
A
symbol = lambda
distance between 2 consecutive peaks or troughs in a wave
2
Q
what is frequency
A
symbol = nu
number of waves (cycles per second) that pass a given point in space
3
Q
do all types of electromagnetic radiation travel the same speed
A
yes
4
Q
how fast do all types of electromagnetic radiation travel
A
at the speed of light
5
Q
what is the relationship between wavelength and frequency
A
inversely proportional
6
Q
what is the relationship between wavelength and energy
A
inversely proportional
7
Q
what is the relationship between frequency and energy
A
directly proportional
8
Q
what does it mean that energy is quantized
A
it can only occur in discrete units called quanta
9
Q
what do we call atoms that have absorbed more energy
A
excited
10
Q
what happens when atoms are excited
A
they contain excess energy which they release by emitting light of various wavelengths to produce an emission spectrum
11
Q
what is the ground state
A
lowest possible energy state
12
Q
as the electron is brought closer to the nucleus, energy is
A
removed from the system
13
Q
what is the Heisenberg uncertainty principle
A
we cannot know both the position and the momentum of an electron simultaneously
14
Q
what is an orbital
A
a specific wave function
15
Q
what is the principal quantum number (n)
A
has positive integral values (1, 2, 3…)
relates to the size and energy of the orbital
often called energy level or shell
16
Q
what does a bigger n (principal quantum number) mean
A
farther from nucleus and electrons are less tightly held and increase in energy
17
Q
what is the angular momentum quantum number (l)
A
has integral values from 0 (n-1 for each n value) related to the shape of atomic orbitals l = 0 S shape l = 1 P shape l = 2 D shape l = 3 F shape
18
Q
what is the magnetic quantum number (m sub l)
A
has integral values from l to -l
relates the the kind of orbital and how many there are of it
19
Q
what is a node
A
an area of an orbital having zero electron probability
20
Q
what shape are s orbitals
A
sphere
21
Q
how many orbitals does s have
A
1
22
Q
what shape are p orbitals
A
infinity shape
23
Q
how many orbitals does p have
A
3
24
Q
how many orbitals does d have
A
5
25
how many orbitals does f have
7
26
how is the hydrogen electron viewed in the quantum mechanical model
as a standing wave, which leads us to a series of wave functions (orbitals) that describe the possible energies and spatial distributions available to the electron
27
what is the electron spin quantum number
electrons can either spin clockwise or counter clockwise
| represents one of the 2 possible values for the electron spin; either 1/2 or -1/2
28
can any 2 electrons in an atom have the same set of 4 quantum numbers?
no
29
what does degenerate mean
having the same energy
30
if all 2p orbitals are empty, why does it not matter which orbital the electron goes to
because they all have the same energy (are degenerate)
31
what is Hund's rule
The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals
32
what are the core electrons
the inner electrons
33
what are valence electrons
the electrons in the outermost principal quantum level of an atom
34
what is the electron configuration for chromium
4s1 3d5
35
what is the electron configuration for copper
4s1 3d10
36
use effective nuclear charge (Zeff) to justify trends
across a period
37
use increased distance (greater value of n) to justify trends -
down a group
38
what are the 4 arguments for justifying periodic trends
1. effective nuclear charge
2. Distance
3. Shielding
4. Electron/electron repulsions
39
what is effective nuclear charge
essentially equal to the group number
the higher the Zeff, the more positive the nucleus, the more attractive force emanating from the nucleus drawing electrons in or holding them in place
40
what is distance
attractive forces dissipate w increased distance
| distant electrons are held loosely and thus are more easily removed
41
what is shielding
electrons in the "core" effectively shield the nucleus' attractive force the the valence electrons
use this ONLY when going UP AND DOWN the table NOT ACROSS
42
what is minimize electron/electron repulsions
puts the atom at a lower energy state, which is more stable
| typically good for explaining weird exceptions
43
what is the trend for atomic radius
increases going down the table, decreases going across the table
44
what is the trend for ionization energy
decreases going down the table, increases going across
45
what is atomic radius
refers to the distance between the nucleus and the outer edge of the electron cloud
46
what is atomic radius influenced by
nuclear pull and number of energy levels
47
why does atomic radius decrease moving across the table
```
Zeff increases (more protons for the same number of energy levels) as we move across the table
nucleus has a greater positive charge so the entire electron cloud is more strongly and "shrinks" (until the point at which electron/electron repulsions overcome the nuclear attractions and stop the contraction of the electron cloud
```
48
why does atomic radius increase going down the table
the principal level (n) determines the size of an atom, so you add another principal level and the atoms get a much larger radii
increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons
full energy levels provide some shielding between the nucleus and the valence electrons
49
What is ionization energy
the energy needed to remove an electron from a gaseous atom or ion, i.e. an isolated one, not part of a solid, liquid, or a molecule. Always endothermic
removing each subsequent electron requires more energy
50
why does ionization energy increase across the table
Increasing Zeff increases the attraction of the nucleus and therefore holds the electrons more tightly
51
What exceptions exist in regards to ionization energy
1) a drop in IE occurs between groups II and III because the p electrons do not penetrate the nuclear region as greatly as s electrons do and are therefore not tightly held
2) a drop in IE occurs between groups 5 and 6 because the increased repulsion created by the first pairing of electrons outweighs the increase in Zeff and thus less energy is required to remove the electron
52
why does ionization energy decrease down the table
increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons
full energy levels provide some shielding between the nucleus and valence electrons
53
What is electron affinity
involves the addition of an electron to a gaseous atom or ion (can be exo or endothermic)
54
what is the trend for electron affinity
decreases as you go down the table (becomes less negative, giving off less energy)
increases across the table (becomes more negative, giving off more energy)
55
why does electron affinity decrease as you go down the table
increased distance from the nucleus w each increasing principal E level. The nucleus is farther from the valence level and more shielded
56
why does electron affinity increase as you go across the table
increasing Zeff more strongly attracts the electrons
57
what are the exceptions with electron affinity
F and Cl bc they only need 1 more electron to achieve noble gas configuration so they will readily accept it
K and Na but for the opposite reason
58
what is electronegativity
the ability of an atom IN A MOLECULE to attract shared electrons to itself
59
what is the trend for electronegativity
increases across table and decreases down table
60
why does electronegativity increase across the table
as Zeff increases, the nucleus becomes more strongly attracted to the electrons
61
why does electronegativity decrease down the table
increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons
Full energy levels provide some shielding between the nucleus and valence electrons
62
what is ionic radius
the distance from the nucleus to the outer edge of the electron cloud in a charged ion
63
what is the ionic radii trend
increase going down the table
cations shrink
anions expand
64
why do cations shrink
cations result from the loss of valence electrons so in many cases this means the farthest electrons are now in a smaller principal energy level than the original neutral atom
as electrons are lost, the ratio of protons to electrons increases and thus the electrons are held closest and with more strength
65
why do anions expand
the nucleus is not attracting more electrons than there are protons
enchanted electron/electron repulsions
66
whats the deal w isoelectric ions
consider the number of protons to determine size (i.e F has one more proton than O which further attracts the electron cloud so its smaller)
67
what is the trend with reactivity
metals are more reactive as you move down a column
| non metals are more reactive as you move up a column
68
why are metals more reactive as you move down a column
b/c metals react by losing electrons, a loosely held electron will result in a more reactive metal. With an increase number of energy levels (n) comes increased distance from the nuclear attraction and thus a more loosely held electron available for reacting
69
why are nonmetals more reactive as you move up a column
b/c nonmetals tend to gain electrons, a strong nuclear attraction will result in a more reactive non-metal. This means an atom with the highest Zeff and the least number of energy levels should be the most reactive nonmetal b/c its nucleus exerts the strongest pull
