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ENGCHEM 111 LEC > THE BOHR ATOM > Flashcards

THE BOHR ATOM Flashcards

1
Q

Although the Bohr model is not completely accurate, it can be used to explain

A

absorption and emission

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2
Q

Electrons move from low energy to higher energy orbits by

A

absorbing energy

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3
Q

Electrons move from high energy to lower energy orbits by

A

emitting energy

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4
Q

Electron energy is

A

quantized

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5
Q

the higher the wavelength, the lower the energy

A

the higher the energy, the lower the wavelength

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6
Q

n

A

principal quantum numbers

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7
Q

describes the energy level on which the orbital residue
- numbers greater than 0

A

Quantum Numbers

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8
Q

this quantum number defines the shape of the orbital
- allowed values of ranging 0 to n − 1

A

Azimuthal Quantum Number (l)

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9
Q

describes the three-dimensional orientation of the orbital
- values of integers ranging from -l to l

A

Magnetic Quantum Number (ml)

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10
Q

electrons in the same orbital do not have the same energy.
- upward or downward

A

Spin Quantum Number (ms)

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11
Q

sphere around the nucleus

A

1s orbital

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12
Q

the one that tells you that the electron is in the orbital closest to the nucleus

A

1s orbital

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13
Q

similar to 1s except the electron is most likely in the region farther from the nucleus

A

2s orbital

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14
Q

at the first energy level, there is only the 1s orbital, after the second energy level there are

A

2p orbitals

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15
Q

look like dumbells

A

p orbitals

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16
Q

distribution of all electrons in an atom

A

electron configurations

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17
Q

4p^5

what is 5?

A

superscript denoting the number of electrons in those orbitals

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18
Q

4p^5

what is p?

A

letter denoting the type of orbital

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19
Q

4p^5

what is 4?

A

number denoting the energy level

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20
Q

“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

A

Hund’s Rule

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21
Q

Filling Rules of Electron Orbitals

Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.

A

Aufbau Principle

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22
Q

Filling Rules for Electron Orbitals

An orbital can hold a maximum of two electrons, to occupy the same orbital, two electrons must spin in opposite directions

A

Pauli Exclusion Principle

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23
Q

Filling Rules for Electron Orbitals

Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.

A

Hund’s Rule

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24
Q

when n = 1?
what are the possible values of l, subshell designation, ml, numbers of orbitals in subshell, total numbers of orbitals?

A

l = 0
subshell designation = 1s
ml = 0
numbers of orbitals in shell = 1
total number of orbs = 1

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25
Q

when n = 2?
what are the possible values of l, subshell designation, ml, numbers of orbitals in subshell, total numbers of orbitals?

A

l = 0,1
subshell designation = 2s, 2p
ml = 0 1, 0 , -1
numbers of orbitals in shell = 1, 3
total number of orbs = 4

26
Q

when n = 3?
what are the possible values of l, subshell designation, ml, numbers of orbitals in subshell, total numbers of orbitals?

A

l = 0,1,2
subshell designation = 3s 3p 3d
ml = 0 1, 0, -1 2, 1, 0, -1, -2
numbers of orbitals in shell = 1, 3, 5
total number of orbs = 9

27
Q

when n = 4?
what are the possible values of l, subshell designation, ml, numbers of orbitals in subshell, total numbers of orbitals?

A

l = 0, 1, 2, 3
subshell designation = 4s 4p 4d 4f
ml = 0 1, 0, -1 2, 1, 0, -1, -2 3, 2, 1, 0, -1, -2, -3
numbers of orbitals in shell = 1, 3, 5, 7
total number of orbs = 16

28
Q

It is the grouping of electrons that is not at the lowest possible energy state

A

Excited state

29
Q

It is the grouping of electrons that is at the lowest possible energy state

A

Ground state

30
Q

Atoms return to ground state by

A

emitting radiation

31
Q

To understand the electronic structure of atoms, one must understand the nature of

A

electromagnetic radiation

32
Q

The number of waves passing a given point per unit of time is the

A

frequency (v)

33
Q

For waves traveling at the same velocity, the _____the wavelength,
the ______the frequency

A

longer , smaller

34
Q

in, E = hv, what is h?

A

Planck’s constant (6.626 x 10 J*s)

35
Q

One does not observe a continuous
spectrum, as one gets from a white light source.

A

Only a line spectrum of discrete wavelengths is observed.

36
Q

developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.

A

Erwin Schrödinger

quantum mechanics

37
Q

is the amount of energy needed to move from one energy level to another

A

quantum

38
Q

The wave equation is designated with a

A

lower case Greek psi ()

39
Q

gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time

A

square of the wave equation, 2

40
Q
  • Describes the three-dimensional orientation of the orbital
A

Magnetic Quantum Number, ml

41
Q

Distribution of Electrons

No two electrons in the same atom can have exactly the same energy

A

Pauli Exclusion Principle

42
Q

no two electrons in the same atom can have identical sets of .

A

quantum numbers

43
Q

Distribution of Electrons

Electrons occupy the positions of the lowest energy

A

Aufbau Principle

44
Q

Distribution of Electrons

Electrons in the same sublevel occupy empty orbitals rather than pair up

A

Hund’s Rule

45
Q

no two electrons in an atom have the same four quantum number’s

A

Pauli exclusion principle

46
Q

Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.

A

Aufbau Principle

47
Q

An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.

Type of principle

A

Pauli Exclusion Principle

48
Q

Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.

A

Hund’s Rule

49
Q

Electrons behave as

A

particle and wave

50
Q

describes the probability of the location of the electrons

A

Shapes and Orientations of Orbitals

51
Q

are useful way of the describing the elements.

A

Electronic Configuration

52
Q

adopted Planck’s assumption and explained these phenomena in this way

A

Niels Bohr

53
Q

The Nature of Energy

A
  1. Electrons in an atom can only occupy certain orbits
  2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
  3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = hv
54
Q

RH is the

A

Rydberg constant, 2.18 x 10^−18 J

55
Q

posited that if light can have material properties, the matter should exhibit wave properties.

A

Louis de Broglie

56
Q

Louis de Broglie demonstrated that the relationship between mass and wavelength was

A

λ =h/mv

57
Q

Chromium can be detected in atomic absorption spectroscopy by
monitoring the absorbance of UV light at a wavelength of 357.8 nm
➢ What is the energy of a photon of this light?

A
58
Q

replaced the Bohr model of the atom

A

Quantum mechanical model

59
Q

Quantum mechanical model depicts electrons as waves spread out or delocalized through a region of space called

A

orbital

60
Q

Wave behavior was described using a wave function, called as the

A

Schrödinger equation

ENGCHEM 111 LEC (13 decks)

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