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AP Chemistry > Chapter 7: Electronic Structure > Flashcards

Chapter 7: Electronic Structure Flashcards

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1
Q

Rutherford’s model did not address

A

how the electrons occupied the space around the nucleus

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2
Q

initially, it was assumed that electrons held … about the nucleus, leading to the … model of the atom

A

fixed orbits; planetary

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3
Q

under certain conditions, atoms and molecules emit and absorb energy in the form of

A

light

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4
Q

in the late 19th century, physicists knew that light could be described as

A

waves

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5
Q

waves are … in nature: they…

A

periodic; repeat at regular intervals of both time and distance

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6
Q

any wave is described by its

A

wavelength, frequency, and amplitude

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7
Q

wavelength (λ, …) is the …

A

lambda; distance between one peak to the next

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8
Q

in SI system, wavelength is measured in

A

meters

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9
Q

frequency (v, ..) is the number of …

A

nu; waves that pass a fixed point in 1 second

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10
Q

the SI unit for frequency si … and is called ..

A

s^-1; hertz (Hz)

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11
Q

amplitude is the … of a wave

A

maximum height

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12
Q

height of a wave varies between

A

+Amax and -Amax

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13
Q

light waves are called… because they consist of …, which are perpendicular to … and to the direction of …

A

electromagnetic radiation; oscillating electric and magnetic fields; each other; propagation;

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14
Q

the speed at which a wave travels is the product of its

A

wavelength and frequency

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15
Q

the periodic nature of wave motion is not always

A

easily seen

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16
Q

the experimentally measured speed of light shows that all electromagnetic radiation travels at the

A

same speed in a vacuum, no matter what its wavelength

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17
Q

the speed of light in a vacuum is one of the fundamental constants of nature:

A

3.00 x 10^8 m/s

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18
Q

c=

A

λv

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19
Q

the amplitude is the vertical …

A

displacement from the undisturbed medium

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20
Q

the length of time that it takes for one complete wave to pass a point is

A

deltat

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21
Q

wavelength and frequency are

A

inversely proportional

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22
Q

the human eye can detect only a very small part of the electromagnetic range, called

A

visible light

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23
Q

visible light includes wavelengths from … to … nm

A

400; 700

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24
Q

order of colors for visible spectrum:

A

VIBGYOR

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25
Q

shorter wavelengths are

A

higher in energy

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26
Q

amplitude indicates the .. of light source

A

brightness

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27
Q

gamma and x rays are the

A

shortest wavelengths (highest in energy)

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28
Q

radio and microwaves are the

A

longest wavelengths (smallest in energy)

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29
Q

at temperatures greater than absolute zero, matter emits …., and the emission is referred to as a …

A

electromagnetic radiation of all wavelengths; continuum

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30
Q

not all wavelengths of light are emitted with

A

equal intensity

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31
Q

the distribution of the intensity of different wavelengths changes with

A

temperature

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32
Q

in 1900, max planck proposed an explanation of the wavelengths emitted by … that was based on an assumption that violated the classical models of physics

A

heated objects

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33
Q

planck assumed that the particles of matter in the heated objects were …., and that the amount of energy the particles had was proportional to the …

A

vibrating back and forth; frequency at which the particles vibrated

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34
Q

planck’s equation

A

E= hv

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35
Q

h, …., is =

A

planck’s constant; 6.626 x 10^-34 J * s

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36
Q

because the energy of the vibrating particle has a specific energy, the energy is considered to be

A

quantized

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37
Q

Einstein applied planck’s equation to … and proposed that light behaves as a …. whose value is directly proportional to the …

A

light; particle of energy; frequency of the light

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38
Q

Einstein essentially proposed that the energy of light was …, meaning that it could only be a …

A

quantized; certain amount

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39
Q

photoelectric effect is the process in which electrons are ….

A

ejected from solid metal when it is exposed to light

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40
Q

each metal has a characteristic …., that is necessary before any …

A

minimum frequency; electrons are emitted

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41
Q

as the frequency of light increases from the minimum, the kinetic energy of the ejected electrons

A

alos increases

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42
Q

more intense light does not increase the ….., but it does increase the …

A

kinetic energy of the electrons; number of electrons emitted

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43
Q

Einstein suggested that light not only had wave properties, but could be viewed as a stream of tiny particles, referred to as

A

photons

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44
Q

a … with an energy of hv must provide enough energy to …

A

SINGLE photon; eject an electron

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45
Q

some of the energy, hv0, must be used to overcome the ….

A

attraction the solid has for the electrons

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46
Q

equation with hvs and ke

A

hv = hv0 + KE

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47
Q

one photon of light can eject

A

one electron

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48
Q

increasing the intensity of the light source produces more …, because the number of photons is …

A

ejected electrons of the same KE; proportional to the intensity

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49
Q

if the energy of the absorbed photon is less than hv0, even with an increased intensity it will only

A

heat the metal

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50
Q

Einstein’s explanation in conjunction with planck’s theory suggested that each …. was carried by a …

A

quantum of energy; particle of light or a photon

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51
Q

in the interpretation of the photoelectric effect, electromagnetic radiation is treated as … instead of ..

A

particles of light (photons); waves

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52
Q

when energy in the form of heat or an electric discharge is added to a sample of gaseous atoms in a process called …, the atoms can emit some of the …

A

excitation; added energy as light

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53
Q

spectrum: the … of the light as a function of …

A

intensity; wavelength

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54
Q

a heated solid produces a …., one in which all … are present

A

continuous spectrum; wavelengths

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55
Q

the light emitted by excited atoms is called a … because it contains light only at …

A

line spectrum; specific wavelengths

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56
Q

each element produces a line spectrum that is … and different from the spectrum of …

A

characteristic of that element; any other element

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57
Q

wavelengths of all lines in the spectrums given by the

A

Rydberg equation

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58
Q

Rydberg equation:

A

1/λ=R_H (1/(n_1^2 )-1/(n_2^2 ))

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59
Q

n1 and n2 are positive integers with n1 …

A

< n2

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60
Q

Rh, called the .., has the value of …

A

Rydberg constant; 1.097 x10^7 m^-1

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61
Q

they hydrogen atom spectrum consists of series of lines that are named after those who discovered them:

A 
Lyman (n1 = 1) 
Balmer (n1 = 2) 
Paschen (n1 = 3)
Brackett (n1 = 4)
Pfund (n1 = 5)
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62
Q

The Rydberg equation accurately predicts the wavelengths ofall observed lines in the

A

spectrum of hydrogen atoms

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63
Q

the discrete line spectra of atoms suggested that electrons exist in

A

only certain allowed energy levels

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64
Q

niels bohr proposed a model for the … that accounted for the …

A

hydrogen atom; observed spectrum of hydrogen

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65
Q

bohr assumed that the electron moved in … around the nucleus

A

circular orbits

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66
Q

bohr assumed that the electron could have only certain values of

A

angular momentum

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67
Q

angular momentum:

A

momentum of a mass moving in a circle

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68
Q

bohr also found that the allowed … and …. are also quantized

A

radii; energies

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69
Q

En=

A

(-(2π^2me^4)/h^2)*1/n^2 = -B/ n^2

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70
Q

for En, m is the …., e is the …, h is …, and n is a positive integer that indicates ..

A

mass; charge of electrons; planck’s constant; electron’s energy level

71
Q

B=

A

2.18 x 10^-18 J

72
Q

bohr concluded that the energy levels of the electron in a hydrogen atom are

A

quantized

73
Q

bohr realized that the light emitted by the atom must have energy (hv) that is exactly equal to the difference between the

A

energies of two of its allowed levels

74
Q

(bohr) Elight= E2 - E1 =

A

B/n1^2 - B/n2^2

75
Q

as the electron and the H+ nucleus move closer together, the atom becomes…, so the energies of all the allowed states have a …

A

more stable (lower in energy); negative sign

76
Q

the energy of an allowed state is proportional to …, so the energies of the allowed states get … as n increases

A

1/n^2; closer together

77
Q

when an electron goes from one quantized energy state to a lower one, the

A

difference in energy is released as a single photon

78
Q

all of the energy released when an atom goes from one allowed energy state to a lower one is contained in a

A

single photon of light

79
Q

the electron in the hydrogen atom can also be excited to higher levels by the …

A

absorption of a photon

80
Q

the only photons absorbed are those with energy identical to the

A

energy difference between two allowed states of the atom

81
Q

the grounds state of an atom is its lowest

A

quantized energy state

82
Q

at normal temperatures, nearly all hydrogen atoms are present in the

A

ground state

83
Q

thus, the only lines observed in the absorption spectrum of hydrogen atoms are those in the

A

lyman series

84
Q

momentum=p=

A

h/λ

85
Q

p =

A

mv

86
Q

equation–> known as … λ (in terms of momentum)=

A

de Broglie equation; h/p = h/mv (m is mass, v is velocity)

87
Q

de Broglie predicted that a particle of matter would have a wavelength that is

A

inversely proportional to the mass

88
Q

the davisson-germer experiment, which demonstrated that matter exhibited its …, was a significant step forward in understanding the …

A

wave properties and particle properties; properties of the electron

89
Q

diffraction is a property of

A

waves

90
Q

de Broglie’s equation explained the assumption of quantized angular momentum of the electron in the hydrogen atom by suggesting that the electron “wave” in an atom must be a …., which is a wave that …

A

standing wave; stays in a constant position

91
Q

de Broglie’s restriction for standing wave expressed in this equation:

A

2πr= nλ

92
Q

bohr’s treatment of the hydrogen atom cannot be extended to

A

larger atoms

93
Q

Erwin schrodinger devised a wave model to describe the

A

behavior of the electrons in atoms

94
Q

the electron wave can be described by a mathematical function that gives the …. this function is called a …, represented by the greek letter psi

A

amplitude of the wave at any point in space; wave function

95
Q

the square of the wave function gives the … of finding the electron at any point in space

A

probability

96
Q

the wave model doesn’t conflict with the Heisenberg uncertainty principle because it does not precisely define the

A

location of the electron

97
Q

he values of quantum numbers are related to the … and … of the electron wave and the location of the …

A

shape; size; electron in 3d space

98
Q

it is possible to calculate the energy of an electron having each possible

A

wave function

99
Q

the angular momentum of the electron is also quantized but this is a …

A

natural consequence of the wave function

100
Q

schrodinger’s wave model is a fundamental idea in the theory called

A

quantum mechanics

101
Q

the best current description of the electronic structure of the atom treats the electron as a

A

wave

102
Q

quantum numbers represented by

A

n, l, and msubl

103
Q

the 3d wave function of an electron is called an

A

atomic orbital

104
Q

each of the quantum numbers is restricted to certain

A

whole-number values

105
Q

n refers to the

A

principal quantum number

106
Q

the principal quantum number gives info about the

A

distance of the electron from the nucleus

107
Q

all orbitals that have the same value of n are in the same

A

principal shell

108
Q

the term principal shell (or simply, shell) refers to all atomic orbitals that hae the same value of n, because they all have approx. the same

A

average distance from the nucleus

109
Q

l refers to the

A

angular momentum quantum number

110
Q

the possible values of l for a given n are all positive integers from zero up to

A

n-1

111
Q

angular momentum quantum number describes the

A

shape of the orbital

112
Q

a subshell contains all orbitals that have the

A

same values for n and l

113
Q

the notation for a subshell consists of a …, which is the value of the …, followed by a … that identifies the value of the l quantum number

A

number; n quantum number; lower case letter (s, p, d, or f

114
Q

s, p, d, and f stand for

A

sharp, principal, diffuse, and fundamental

115
Q

ml refers to the

A

magnetic quantum number

116
Q

allowed values for ml

A

-l to l

117
Q

ml quantum number provides info about the

A

orientation in space of the atomic orbital

118
Q

each subshell consists of one or more

A

atomic orbitals

119
Q

the fourth quantum number does not directly come from the

A

wave model

120
Q

ms refers to the

A

electron spin quantum number

121
Q

the allowed values of m are

A

+1/2 and -1/2

122
Q

the electron spin does not depend on the values of any of the

A

other quantum numbers

123
Q

two electrons that have the same spins are …, whereas electrons with different spins are …

A

parallel; paired

124
Q

nodes: regions in space where the probability of finding an electron is exactly

A

zero

125
Q

as the principal quantum number increases in value, the average distance of the electron from the nucleus …, and thus the size of the … increases

A

increases; contour surface

126
Q

if the principal quantum number is the same, no matter which subshell or orbital the electron occupies, the hydrogen atom has exactly the same

A

energy

127
Q

the energy of each wave function for any atomic species containing only one electron is given by

A

En = -(Z^2B)/n^2 (Z is nuclear charge)

128
Q

the different subshells within the same shell of a multielectron atom

A

do not have the same energy

129
Q

because charges of opposite sign attract each other, the energy of the atom decreases (the atom becomes more …) as the electron

A

stable; gets closer to the nucleus

130
Q

energies in multielectron atoms depend on the values of …

A

both the n and l quantum numbers

131
Q

in one-electron atoms and ions, the energy depends only on the

A

value of the n quantum number

132
Q

the single electron in any one-electron species, regardless of its location or the orbital it occupies, is attracted by the

A

nuclear charge

133
Q

the electron-electron repulsions, known as .., reduce the effect of the positive charge of the nucleus on each electron, thus …

A

interelectronic repulsions; influencing its energy

134
Q

the net attraction of the nucleus for n electron at any distance r is reduced, or shielded, by the ….

A

repulsive forces from the electrons between it and the nucleus

135
Q

effective nuclear charge, Zeff, is the weighted average of the …., after correction for the shielding of nuclear charge by … and the …

A

nuclear charge that affects an electron in the atom; inner electrons; interelectronic repulsions

136
Q

electron shielding is the result of the influence of …. on the …

A

inner electrons; effective nuclear charge

137
Q

to determine the effective nuclear charge for each electron, we need to know whether the other electrons in the atom are between

A

it and the nucleus

138
Q

although the average distances of the 2s and 2p electrons are about the same, the probability that the electron is close to the nucleus is greater for the … than for the …

A

2s electron; 2p electron

139
Q

the 2s electron penetrates the electron density of the filled 1s shell more than does the 2p electron, so it is influenced by a

A

greater effective nuclear charge

140
Q

the energy of an s electron is … than the energy of a p electron in the same shell

A

lower

141
Q

within any shell, the penetration of the s orbital is always greater than that of the … orbitals, which in turn, is greater than that of the … orbitals

A

p; d

142
Q

within any shell, the subshells increase in energy in the order of

A

increasing value of the quantum number l

143
Q

different interelectronic repulsive forces affect electrons in different subshells, so the energy of an atom depends on

A

which subshells are occupied

144
Q

the overlap in the energy of different shells becomes more common as

A

n increases

145
Q

energy separation between subshells gets quite small in the higher shells, so small changes int eh shielding effects may cause the energy order to

A

change from one element to the next

146
Q

because the energies of different orbitals depend only on the values of the n and l quantum numbers and not on the value of ml, all of the orbitals in a subshell have

A

exactly the same energy

147
Q

when orbitals are of exactly the same energy, they are referred to as

A

degenerate orbitals

148
Q

pauli exclusion principle: no two electrons in the same atom can have teh

A

same set of all four quantum numbers

149
Q

the pauli exclusion principle is the quantum-mechanical equivalent of saying that two objects cannot

A

occupy the same space at the same time

150
Q

using the pauli exclusion principle, we find that the maximum number of electrons that can share a single orbital in an atom is

A

two

151
Q

two electrons in the same orbital are referred to as an … or …, because they must have different spins

A

electron pair; paired electrons

152
Q

when a single electron is in an orbital, it is called an

A

unpaired electron

153
Q

the restrictions on the quantum numbers and the pauli exclusion principle determines the capacities of …, …, and …

A

orbitals; subshells; principal shells

154
Q

aufbau principle states that electrons are added to the atom one at a time until the

A

proper number is present

155
Q

(aufbau principle) as each electron is added, it is assigned the quantum numbers of the

A

lowest energy orbital available

156
Q

practically all of the atoms in a sample are in the …at normal temperatures

A

ground state

157
Q

if one or more of the electrons is in any other allowed orbital of the diagram, the atom is in an

A

excited state

158
Q

the excited state is of …, and the atom tends to return to its ground state by …., often by emitting a …

A

higher energy; losing energy; photon of light

159
Q

an excited state is not equivalent to an impossible state, in which ….. of quantum numbers are present

A

forbidden combinations

160
Q

two electrons in the same orbital must always have …. , represented by “…” and “…” …

A

opposite spins; up; down arrows

161
Q

in an orbital diagram, each orbital is represented by a …, with orbitals in the same subshell shown as …

A

box; grouped boxes

162
Q

both … and … are used to represent the electrons in atoms

A

energy-level diagrams; orbital diagrams

163
Q

the electron configuration of an atom is …; it does not contain the detailed information about … that an orbital diagram provides

A

compact; electron spins

164
Q

whenever electrons are added to a subshell that contains more than one orbital, the electrons enter →…

A

separate orbitals until there is one electron in each; hund’s rule

165
Q

two electrons in the same orbital are closer together than they would be if they were in separate orbitals, and they therefore

A

repel each other more strongly

166
Q

hund’s rule states that degenerate orbitals are filled with one electron in each before

A

any electrons are paired

167
Q

we use hund’s rule to write the ground-state electron configurations and orbital diagrams for the elements iwth atomic numbers

A

7 through 10

168
Q

abbreviated electron configuration: uses … to represent the partial electron configuration up to the number of electrons for that gas

A

noble ass

169
Q

abbreviated electron configurations are more convenient for expressing the electron configurations of

A

heavier atoms

170
Q

anomalous electron configuration are configurations that are the

A

exceptions

171
Q

the anomalous electron configuration leads to a combination of either … or a … and a …

A

two half-filled subshells; half-filled; completely filled subshell

172
Q

we cannot predit which atoms will ahve

A

anamolous electron configurations in advance

173
Q

the total electronic energy of the atom is … in an anomalous configuration in comparison with a …

A

lower; “normal” electron configuration

AP Chemistry (24 decks)

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