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Chemistry 121 > Exam 4 > Flashcards

Exam 4 Flashcards

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1
Q

Excited atom makes what

A

Light emitted

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2
Q

Atom plus light absorption makes what

A

Excited atom

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3
Q

Distance between adjacent peaks on the wave (m)

A

Wavelength (λ)

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4
Q

Number of waves that pass a point per second (Hz or 1/s)

A

Frequency (v)

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5
Q

Vertical distance from the mid-line to the peak

A

Amplitude (intensity)

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6
Q

Speed of light (c)

A

3.00 * 10^8 m/s (in a vacuum)

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7
Q

Wavelength & frequency have what relationship

A

Inverse

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8
Q

Who developed a theory for light, proposing that light waves consists of 2 components; a vibrating electric field & a vibrating magnetic field

A

Maxwell

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9
Q

Who studied radiation emitted by matter & postulated that energy in matter is quantized

A

Max Planck

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10
Q

Smallest increment of energy that can be emitted or absorbed in the form of electromagnetic radiation

A

Quantum of energy

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11
Q

Energy equation of a single quantum

A

E=hv (E-energy J, v-frequency 1/2, h-Planck’s constant)

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12
Q

Plancks constant

A

6.6310^-34 Js

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13
Q

Light can cause electrons to be ejected from a metal surface

A

Photoelectric effect

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14
Q

If light has a frequency less than the threshold frequency, what happens

A

No electrons ejected

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15
Q

If light has a frequency greater than the threshold frequency, what happens

A

Number of electrons ejected is proportional to the light intensity

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16
Q

Who considered that light is composed of a stream of particles (photons- quantum packets of energy), extended Planck’s ideas

A

Albert Einstein

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17
Q

Electromagnetic radiation has what properties

A

Wave-like & particle-like properties

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18
Q

E=hv can also be written as

A

E=hc/λ

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19
Q

Short wavelengths have high energy and are

A

Damaging to life, like X rays

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20
Q

Long wavelengths have low energy and are

A

Safe, like radio waves

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21
Q

Atom that has the simplest atomic spectrum

A

Hydrogen atom

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22
Q

Each atom has its own unique

A

Atomic spectrum

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23
Q

Who studied the lines in the visible portion of the hydrogen spectrum (400-700nm)

A

JJ Balmer

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24
Q

Rydberg equation for calculating the emission lines in the hydrogen spectrum

A

1/λ=Rh[(1/n^2i) - (1/n^2f)]

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25
Q

Rydberg’s constant

A

109,678cm^-1

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26
Q

nf=1; ni=2,3,4… (ultraviolet)

A

Lyman series

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27
Q

nf=2; ni=3,4,5…(visible)

A

Balmer series

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28
Q

nf=3; ni=4,5,6…(infrared)

A

Paschen series

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29
Q

nf=4; ni=5,6,7… (infrared)

A

Brackett series

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30
Q

Danish physicist who developed a theoretical model for the hydrogen atom that restriceted on the sizes of the orbits & the energy that the electron can have in a given orbit

A

Niels Bohr

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31
Q

In Bohr’s model, lowest energy atoms are what

A

n=1, E<0

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32
Q

In Bohr’s model, highest energy atoms are what

A

n=∞, E=0

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33
Q

Electron in ground state absorbs energy & does what

A

Moves to a higher energy orbit, excited state

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34
Q

Photon emitted, meaning what

A

ni>nf, ΔE is -

35
Q

Photon absorbed, meaning what

A

ni

36
Q

Hypothesis that suggested that if electromagnetic radiation can have particle-like properties, then electrons may also possess wave-like properties

A

de Broglie’s hypothesis

37
Q

Combining Einstein’s theory with de Broglie’s hypothesis, you get what equation

A

λ=h/mc (mc is momentum- mass * velocity)

38
Q

Two waves add together when waves have the same frequency & are in phase with each other

A

Constructive interference

39
Q

Two waves with same frequency but are out of phase

A

Destructive interference

40
Q

Who applies wave properties of matter to explain atomic properties, created wave mechanics

A

Erwin Schrodinger

41
Q

Electrons exist only where what

A

Waves reinforce each otherq

42
Q

An atomic orbital is characterized by what

A

Energy, size, shape, orientation in space

43
Q

s orbitals have what shape

A

Spherical

44
Q

p orbitals have what shape

A

Two lobes about the nucleus

45
Q

p orbitals shape

A

Balloons on each side of nucleus, vary in orientation

46
Q

d orbitals shape

A

Complex, four lobes of electron density, differ in orientation

47
Q

f orbitals shape

A

Very complex, 8 lobes of electrons about nucleus

48
Q

Quantum number that determines the energy & size of an orbital

A

Principal quantum # (n=1,2,3… ∞)

49
Q

Quantum number that determines the shape of the orbital

A

Angular momentum quantum # (l=0,1,2,3…n-1)

50
Q

l=0

A

s

51
Q

l=1

A

p

52
Q

l=2

A

d

53
Q

l=3

A

f

54
Q

l=4

A

g

55
Q

l=5

A

h

56
Q

Quantum number that describes the orientation in space relative to other orbitals in that same subshell

A

Magnetic quantum # (ml= -l, -l+1,….+l)

57
Q

s orbitals have a spherical shape

A

n=1 to ∞, l=0, ml=0

58
Q

p orbitals have the same shape, size, & energy- differ only in orientation in space, consist of two lobes of electron density about the nucleus

A

n≥2, l=1, ml=-1,0,1.

59
Q

d orbitals have 5 individual orbitals, complex shape, differ only in orientation in space

A

n≥3, l=2, ml=-2, -1,0,1,2

60
Q

f orbitals have 7 individual orbitals, complex shape consisting of 8 lobes of electron density, differ only in orientation in space

A

n≥4, l=3, ml=-3,-2,-1,0,1,2,3

61
Q

Quantum number describing an electrons spin

A

Electron spin quantum # (ms=+1/2 or -1/2

62
Q

Principle stating that no two electrons in an atom can have the same set of 4 quantum #’s

A

Pauli exclusion principle

63
Q

Ground state electron configuration

A

Lowest energy arrangement of electrons

64
Q

Electron configuration that lists subshells that contain electrons & designates their populations by what

A

Appropriate superscripts

65
Q

Rule that states that for a partially filled subshell, the lowest energy arrangement will have the max # of unpaired electrons with the same spin state

A

Hund’s rule

66
Q

If all electrons in an atom are paired, the atom is not attracted by a magnet & is considered

A

Diamagnetic

67
Q

If an atom has unpaired electrons, then the atom is attracted to a magnet & is considered

A

Paramagnetic

68
Q

Principle that says you need to put electrons in the lowest energy orbitals before moving up

A

Aufbau principle

69
Q

Periods arranged by

A

Increasing mass

70
Q

Groups arranged by

A

Similar properties

71
Q

Periodic table created by

A

Mendeleev & Meyer

72
Q

Atomic number or mass more important?

A

Atomic number

73
Q

Group 1A-7A, incompletely filled outer s or p shells

A

Representative elements

74
Q

Group 8A, filled outer subshells (outer s and p shells filled)

A

Noble gases

75
Q

Groups 1B and 3B-8B, incompletely filled outer d subshells

A

Transition metals

76
Q

Neither representative nor transition metal elements, no special name for this group

A

Group 2B

77
Q

Elements set off below the table, have incompletely filled f subshell

A

Lanthanides & Actinides

78
Q

Ions & atoms with the same number of e- and the same ground state are called what

A

Isoelectronic

79
Q

1/2 the distance between atomic nuclei

A

Atomic radius

80
Q

Negatively charged e- are attracted by the + charged nucleus

A

Effective nuclear charge

81
Q

Core electrons shield the outer shell electrons from what

A

+ nucleus

82
Q

How to find Zeff

A

Z (#of protons) - S (#of core e-)

83
Q

Minimum energy required to remove an e- from the ground state of a gaseous atom

A

Ionization energy

84
Q

Greater e- affinity means a greater what

A

Tendency to gain an e-

Chemistry 121 (5 decks)

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