Unit 2: Atomic Structure Flashcards
1
Q
define wavelength
A
- the distance between identical points on successive waves
- distance from peak to peak
2
Q
what are the units for wavelength
A
- nanometers (nm)
- 1nm = 10^-9m
3
Q
define frequency
A
- number of wavelengths that pass through a particular point in one second
- wavelengths per second
4
Q
what are the units for frequency
A
- hertz (Hz)
- also denoted as s^-1
5
Q
what is the order of the electromagnetic spectrum from low frequency/long wavelength to high frequency/short wavelength
A
- radio
- microwave
- infrared
- visible
- ultraviolet
- x-ray
- gamma rays
6
Q
which electromagnetic rays are more dangerous to humans
A
- ultraviolet
- x-ray
- gamma rays
7
Q
what is the order of visible colors from longer to shorter wavelenghts
A
ROYGBIV
8
Q
what is the wavelength in nm of the visible color red
A
700 nm
9
Q
what is the wavelength in nm of the visible color violet
A
400 nm
10
Q
what is the equation including speed of light (c) wavelength (λ) and frequency (ν)
A
λν=c
11
Q
what is the speed of light in m/s
A
2.9979 x 10^8 m/s
12
Q
how are wavelength and frequency related
A
- inversely proportional
- as one increases, the other decreases
13
Q
is the color of light caused by wavelength or amplitude
A
wavelength
14
Q
is the brightness of light caused by wavelength or amplitude
A
amplitude
15
Q
what is the value of Planck’s constant (h)
A
6.63E-34 J*s
16
Q
what is the equation relating to energy, frequency, and wavelength
A
E=hv or E=hc/λ
17
Q
what did Max Planck discover
A
energy from radiation emitted from hot objects is quantized
18
Q
explain the word quantized
A
- very specific values that relate to something
- ex: potential energy on stairs is a very specific amount for each stair you’re on, you can’t be in between steps
- ex: energy of electrons on each electron ring, electrons can’t be in between rings
19
Q
explain how Einstein discovered how to measure the amount of energy needed to eject an electron from a metal (photoelectric experiment)
A
- he would shine a light on the metal surface in an evacuated chamber
- he would increase the frequency of the light until the current indicator would spike which indicates that electrons have been ejected
- he discovered that different metals required different amounts of energy to eject electrons
20
Q
explain binding energy
A
- amount of energy used to keep electrons orbiting the nucleus
- an equal amount of energy acting on the atom can eject the electron
21
Q
write the equation for energy in terms of frequency and it terms binding and kinetic energy
A
E=hv=E(binding) + E(kinetic)
22
Q
explain what happens if the energy acting on an electron is greater than the binding energy
A
- will create kinetic energy
- enough of the energy will be used to eject the electron and the rest will act as kinetic energy
23
Q
what are the two ways light can behave
A
- particle/photon
- wave
24
Q
what happens to energy as frequency increases
A
energy increases
25
what happens to energy as frequency decreases
energy decreases
26
what happens to energy as wavelength increases
energy decreases
27
what happens to energy as wavelength decreases
energy increases
28
how are wavelength and energy related
inversely proportional
29
how are frequency and energy related
directly proportional
30
what do we see when we look at a continuum source of light through a prism
- continuous spectrum
- all colors of the rainbow
31
what do we see when we look at a continuum source of light through a cloud of gas and then through a prism
- absorption spectrum
- see all colors except those absorbed by the gas
32
describe the absorption spectrum
- spectrum of visible light that is seen through a cloud of gas and a prism
- specific wavelengths (colors) are missing and have been absorbed by the gas
- produced when atoms absorb energy
33
describe the emission spectrum
- opposite of absorption spectrum
- the wavelengths that a gas will emit
- produced when atoms release energy
34
why is atomic emission/absorption spectroscopy useful
can help determine what elements are in something
35
explain why orbital levels are consider quantized
electrons can only occupy the specific orbitals and can't be in between them
36
where are the lowest and highest energy levels of electron orbitals
- lowest energy: closest to nucleus
- highest energy: furthest from nucleus
37
what happens to the space between electron orbitals as you move further from the nucleus
space between orbitals gets smaller as energy levels increase
38
which direction do electrons move when the atom absorbs energy (absorption)
electrons move from low energy levels to high energy levels
39
which direction do electrons move when the atom releases energy (emission)
electrons move from high energy levels to low energy levels
40
define ground state
lowest possible energy configuration
41
define excited state
electron occupies a higher energy level
42
do electrons always want to fall to ground level or to excited levels
ground level
43
where are electrons moving during absorption in terms of ground and excited states
ground state to excited state
44
where are electrons moving during emission in terms of ground and excited states
excited state to ground state
45
what energy levels of electrons moving can we see visible light
2 through 6
46
what happens to the energy of an electron as you move up energy levels
becomes more positive
47
are higher energy movement of electrons bigger or smaller gaps
bigger gaps between energy levels have more energy
48
what is a node in a wave
a spot where there is no amplitude
49
what is the de broglie wavelength equation with units
λ (wavelength in m) = h / m (mass in kg) * v (velocity in m*s)
50
are wavelength and mass directly or inversely proportional in the de broglie wavelength equation
inversely
51
describe the heisenberg uncertainty principle
- the momentum and the position of an electron cannot be simultaneously determined
- if you know one, then you can't know the other
- uncertainty for light objects in large while the uncertainty for heavy objects is small
52
what did Erwin Schrodinger develop
a mathematical treatment that explains both the wave and particle nature of matter and how to find the probability of where an electron is at any given instant in time
53
how many quantum numbers are there
4
54
what are the 4 quantum numbers
- principal quantum number (n)
- azimuthal quantum number (l)
- magnetic quantum number (ml)
- spin magnetic quantum number (ms)
55
which quantum numbers describe electron orbitals
- principal (n)
- azimuthal (l)
- magnetic (m)
56
describe the principal quantum number
- n
- defines the energy level of the orbital
- n=1, 2, 3, ...
57
describe the azimuthal quantum number
- l
- defines the shape of the orbital
- l=0, 1, 2, ..., n-1
58
what is the highest value for the azimuthal quantum number (l)
n-1
59
describe the magnetic quantum number
- ml
- defines the orientation of the orbital
- ml= -l, ..., 0, ..., l
60
describe the s orbital (shape, azimuthal quantum number, and magnetic quantum number)
- circular shape
- no directionality
- l=0
- ml=0
61
describe the p orbital (shape, azimuthal quantum number, and magnetic quantum number)
- infinity sign shape
- 3 orientations
- l=1
- ml=-1, 0, 1
62
describe the d orbital (shape, azimuthal quantum number, and magnetic quantum number)
- butterfly shape; except dz^2 which is infinity shaped with a circle around the middle
- 5 orientations
- l=2
- ml=-2, -1, 0, 1, 2
63
how many electrons can be in one orbital
2
64
describe the spin magnetic quantum number
- ms
- describes the electron spin
- ms= +1/2 OR -1/2
65
describe the pauli exclusion principle
- no two electrons in the same atom can have the same 4 quantum numbers
- each electron has to have a different "address"
- the two electrons in the same orbital must have opposite spin magnetic quantum numbers
66
what happens to the energy of orbitals as they move away from the nucleus
increase in energy
67
list the s, p, d, and f orbitals in order of lowest energy to highest energy
s, p, d, f
68
describe degenerate orbitals
- orbitals with the same energy
- 4s and 3d
69
explain the notation of electron configuration
- (number denoting energy level)(letter denoting type of orbital)^superscript denoting the number of electrons in these orbitals
- 4p^5
70
define hund's rule
- for degenerate orbitals (ex: all orbitals in p level) the lowest energy is attained when the number of electrons with the same spin is maximized
- electrons want to be in their own orbital if they are in orbitals with the same energy
71
what is the order of energy for electron orbitals
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
72
describe how you would write a condensed electron configuration
- find the previous noble gas and add orbitals that are missing
- Li = [He]2s^1
73
what is the electron configuration of Cr (chromium)
[Ar]4s^13d^5
74
what is the electron configuration of Cu (copper)
[Ar]4s^13d^10
75
why do chromium and copper have different electron configurations than expected
- 4s and 3d orbitals are degenerate
- electrons spread between the two instead of filling up 4s first
76
are ion electron configurations more or less stable than the paretn atom
more stable
77
what elements are ion electron configurations similar to
nearest noble gas configuration
78
define isoelectronic
same number of electrons
79
explain how to add and take away electrons for ions in the transition metals
- start adding at the 4s orbital
- start taking away at the 4s orbital
80
define paramagnetic materials
- substances attracted to a magnetic field
- have at least one unpaired electrons
81
what is an example of a paramagnetic material
iron
82
define diamagnetic materials
- substances that are repelled by a magnetic field
- have no unpaired electrons
83
how do you determine whether an element or ion has any unpaired electrons
- draw the picture with the valence orbitals
- any slots with only one electron are considered unpaired
84
describe the nonbonding atomic radius
distance between the nucleus of an atom and the outside of the electron cloud
85
describe the bonding atomic radius
half the distance between the two nuclei of bonded atoms
86
why do elements in larger periods/rows have increasing atomic radii
the energy levels are increasing which makes the atom larger
87
describe the box drawing that indicates which elements have larger atomic radii
- increasing as you move down periods/rows
- increasing as you move to the left of groups/columns
88
why does atomic radii increase as you move to the left of the periodic table groups/columns
- less electrons in the outer shell as you move to the left and more protons
- more pull from positive protons in the nucleus to shrink the atom
89
are cations smaller or larger than their parent atom and why
- smaller than parent atom
- more protons and less electrons
- more of a positive pull towards the nucleus that shrinks the atom
90
are anions smaller or larger than their parent atom and why
- larger than parent atom
- less protons and more electrons
- less of a positive charge pull towards the nucleus allows the electrons to spread out more
91
define ionization energy
amount of energy required to remove an electron from the ground state
92
how does ionization energy change as you remove more electrons
- it takes more energy as you go
- electrons closer to the nucleus are harder to pull away
93
are you more or less likely to lose electrons as you move down the periodic table
- more likely
- lower ionization energy
- bigger atom = electrons farther away from nucleus = easier to pull away
94
are you more or less likely to lose electrons as you move across (from left to right) the periodic table
- less likely
- higher ionization energy
- harder to pull electrons away as you increase groups
95
describe the box drawing that indicates which elements have higher ionization energy
- increasing as you move up periods/rows (lower periods have higher energy)
- increasing as you move right across the groups (group 7 has higher energy than group 2)
96
why is there a drop in ionization energy between beryllium and boron and between nitrogen and oxygen
- boron and oxygen both have a single unpaired electron in their p orbital that is easier to remove
- beryllium and nitrogen have very stable electron configurations that makes it harder to remove electrons
97
define electron affinity
energy change accompanying addition of an electron
98
is energy required or release for electron affinity
released
99
describe the box drawing that indicates which elements have higher electron affinity
- same as ionization energy (arrow going up and arrow going right)
- x out the noble gases column; they do not want to add an electron
100
what does a more negative number indicate for electron affinity
more negative = higher affinity = more favorable to add electron
