Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

CHEM 119 > Unit 2: Atomic Structure > Flashcards

Unit 2: Atomic Structure Flashcards

1
Q

define wavelength

A
  • the distance between identical points on successive waves
  • distance from peak to peak
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

what are the units for wavelength

A
  • nanometers (nm)
  • 1nm = 10^-9m
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

define frequency

A
  • number of wavelengths that pass through a particular point in one second
  • wavelengths per second
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

what are the units for frequency

A
  • hertz (Hz)
  • also denoted as s^-1
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

which electromagnetic rays are more dangerous to humans

A
  • ultraviolet
  • x-ray
  • gamma rays
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

what is the order of visible colors from longer to shorter wavelenghts

A

ROYGBIV

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

what is the wavelength in nm of the visible color red

A

700 nm

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

what is the wavelength in nm of the visible color violet

A

400 nm

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

what is the equation including speed of light (c) wavelength (λ) and frequency (ν)

A

λν=c

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

what is the speed of light in m/s

A

2.9979 x 10^8 m/s

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

how are wavelength and frequency related

A
  • inversely proportional
  • as one increases, the other decreases
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

is the color of light caused by wavelength or amplitude

A

wavelength

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

is the brightness of light caused by wavelength or amplitude

A

amplitude

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

what is the value of Planck’s constant (h)

A

6.63E-34 J*s

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

what is the equation relating to energy, frequency, and wavelength

A

E=hv or E=hc/λ

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

what did Max Planck discover

A

energy from radiation emitted from hot objects is quantized

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

what are the two ways light can behave

A
  • particle/photon
  • wave
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

what happens to energy as frequency increases

A

energy increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

what happens to energy as frequency decreases

A

energy decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

what happens to energy as wavelength increases

A

energy decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

what happens to energy as wavelength decreases

A

energy increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

how are wavelength and energy related

A

inversely proportional

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

how are frequency and energy related

A

directly proportional

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

what do we see when we look at a continuum source of light through a prism

A
  • continuous spectrum
  • all colors of the rainbow
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

what do we see when we look at a continuum source of light through a cloud of gas and then through a prism

A
  • absorption spectrum
  • see all colors except those absorbed by the gas
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

describe the absorption spectrum

A
  • 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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

describe the emission spectrum

A
  • opposite of absorption spectrum
  • the wavelengths that a gas will emit
  • produced when atoms release energy
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

why is atomic emission/absorption spectroscopy useful

A

can help determine what elements are in something

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

explain why orbital levels are consider quantized

A

electrons can only occupy the specific orbitals and can’t be in between them

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

where are the lowest and highest energy levels of electron orbitals

A
  • lowest energy: closest to nucleus
  • highest energy: furthest from nucleus
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

what happens to the space between electron orbitals as you move further from the nucleus

A

space between orbitals gets smaller as energy levels increase

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

which direction do electrons move when the atom absorbs energy (absorption)

A

electrons move from low energy levels to high energy levels

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

which direction do electrons move when the atom releases energy (emission)

A

electrons move from high energy levels to low energy levels

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

define ground state

A

lowest possible energy configuration

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

define excited state

A

electron occupies a higher energy level

42
Q

do electrons always want to fall to ground level or to excited levels

A

ground level

43
Q

where are electrons moving during absorption in terms of ground and excited states

A

ground state to excited state

44
Q

where are electrons moving during emission in terms of ground and excited states

A

excited state to ground state

45
Q

what energy levels of electrons moving can we see visible light

A

2 through 6

46
Q

what happens to the energy of an electron as you move up energy levels

A

becomes more positive

47
Q

are higher energy movement of electrons bigger or smaller gaps

A

bigger gaps between energy levels have more energy

48
Q

what is a node in a wave

A

a spot where there is no amplitude

49
Q

what is the de broglie wavelength equation with units

A

λ (wavelength in m) = h / m (mass in kg) * v (velocity in m*s)

50
Q

are wavelength and mass directly or inversely proportional in the de broglie wavelength equation

A

inversely

51
Q

describe the heisenberg uncertainty principle

A
  • 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
Q

what did Erwin Schrodinger develop

A

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
Q

how many quantum numbers are there

A

4

54
Q

what are the 4 quantum numbers

A
  • principal quantum number (n)
  • azimuthal quantum number (l)
  • magnetic quantum number (ml)
  • spin magnetic quantum number (ms)
55
Q

which quantum numbers describe electron orbitals

A
  • principal (n)
  • azimuthal (l)
  • magnetic (m)
56
Q

describe the principal quantum number

A
  • n
  • defines the energy level of the orbital
  • n=1, 2, 3, …
57
Q

describe the azimuthal quantum number

A
  • l
  • defines the shape of the orbital
  • l=0, 1, 2, …, n-1
58
Q

what is the highest value for the azimuthal quantum number (l)

A

n-1

59
Q

describe the magnetic quantum number

A
  • ml
  • defines the orientation of the orbital
  • ml= -l, …, 0, …, l
60
Q

describe the s orbital (shape, azimuthal quantum number, and magnetic quantum number)

A
  • circular shape
  • no directionality
  • l=0
  • ml=0
61
Q

describe the p orbital (shape, azimuthal quantum number, and magnetic quantum number)

A
  • infinity sign shape
  • 3 orientations
  • l=1
  • ml=-1, 0, 1
62
Q

describe the d orbital (shape, azimuthal quantum number, and magnetic quantum number)

A
  • 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
Q

how many electrons can be in one orbital

A

2

64
Q

describe the spin magnetic quantum number

A
  • ms
  • describes the electron spin
  • ms= +1/2 OR -1/2
65
Q

describe the pauli exclusion principle

A
  • 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
Q

what happens to the energy of orbitals as they move away from the nucleus

A

increase in energy

67
Q

list the s, p, d, and f orbitals in order of lowest energy to highest energy

A

s, p, d, f

68
Q

describe degenerate orbitals

A
  • orbitals with the same energy
  • 4s and 3d
69
Q

explain the notation of electron configuration

A
  • (number denoting energy level)(letter denoting type of orbital)^superscript denoting the number of electrons in these orbitals
  • 4p^5
70
Q

define hund’s rule

A
  • 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
Q

what is the order of energy for electron orbitals

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p

72
Q

describe how you would write a condensed electron configuration

A
  • find the previous noble gas and add orbitals that are missing
  • Li = [He]2s^1
73
Q

what is the electron configuration of Cr (chromium)

A

[Ar]4s^13d^5

74
Q

what is the electron configuration of Cu (copper)

A

[Ar]4s^13d^10

75
Q

why do chromium and copper have different electron configurations than expected

A
  • 4s and 3d orbitals are degenerate
  • electrons spread between the two instead of filling up 4s first
76
Q

are ion electron configurations more or less stable than the paretn atom

A

more stable

77
Q

what elements are ion electron configurations similar to

A

nearest noble gas configuration

78
Q

define isoelectronic

A

same number of electrons

79
Q

explain how to add and take away electrons for ions in the transition metals

A
  • start adding at the 4s orbital
  • start taking away at the 4s orbital
80
Q

define paramagnetic materials

A
  • substances attracted to a magnetic field
  • have at least one unpaired electrons
81
Q

what is an example of a paramagnetic material

A

iron

82
Q

define diamagnetic materials

A
  • substances that are repelled by a magnetic field
  • have no unpaired electrons
83
Q

how do you determine whether an element or ion has any unpaired electrons

A
  • draw the picture with the valence orbitals
  • any slots with only one electron are considered unpaired
84
Q

describe the nonbonding atomic radius

A

distance between the nucleus of an atom and the outside of the electron cloud

85
Q

describe the bonding atomic radius

A

half the distance between the two nuclei of bonded atoms

86
Q

why do elements in larger periods/rows have increasing atomic radii

A

the energy levels are increasing which makes the atom larger

87
Q

describe the box drawing that indicates which elements have larger atomic radii

A
  • increasing as you move down periods/rows
  • increasing as you move to the left of groups/columns
88
Q

why does atomic radii increase as you move to the left of the periodic table groups/columns

A
  • 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
Q

are cations smaller or larger than their parent atom and why

A
  • smaller than parent atom
  • more protons and less electrons
  • more of a positive pull towards the nucleus that shrinks the atom
90
Q

are anions smaller or larger than their parent atom and why

A
  • 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
Q

define ionization energy

A

amount of energy required to remove an electron from the ground state

92
Q

how does ionization energy change as you remove more electrons

A
  • it takes more energy as you go
  • electrons closer to the nucleus are harder to pull away
93
Q

are you more or less likely to lose electrons as you move down the periodic table

A
  • more likely
  • lower ionization energy
  • bigger atom = electrons farther away from nucleus = easier to pull away
94
Q

are you more or less likely to lose electrons as you move across (from left to right) the periodic table

A
  • less likely
  • higher ionization energy
  • harder to pull electrons away as you increase groups
95
Q

describe the box drawing that indicates which elements have higher ionization energy

A
  • 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
Q

why is there a drop in ionization energy between beryllium and boron and between nitrogen and oxygen

A
  • 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
Q

define electron affinity

A

energy change accompanying addition of an electron

98
Q

is energy required or release for electron affinity

A

released

99
Q

describe the box drawing that indicates which elements have higher electron affinity

A
  • 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
Q

what does a more negative number indicate for electron affinity

A

more negative = higher affinity = more favorable to add electron

CHEM 119 (8 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions