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Chemistry 1b > Chapter 8 > Flashcards

Chapter 8 Flashcards

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1
Q

Atoms in the same group __________ in size down the column.

A

Increase

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2
Q

Atomic radii of __________ __________ roughly the same size across the d block.

A

Transition metals

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3
Q

__________ form when the atom loses electrons from the valence shell.

A

Cations

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4
Q

When transition metals form cations, the first electrons removed are __________ __________, even though other electrons were added after.

A

Valence electrons

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5
Q

Electrons may also be removed from the sublevel __________ to the valence shell after the valence electrons.

A

Closest

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6
Q

The iron atom has _____ valence electrons.

A

2

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7
Q

When ion forms a cation, it first __________ it’s valence electrons.

A

Loses

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8
Q

Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field. This is called __________.

A

Paramagnetism

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9
Q

Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field. This is called __________.

A

Diamagnetism

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10
Q

Atom or ion will be attracted to a magnetic field in:

A

Paramagnetism

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11
Q

Atom or ion will be slightly repelled by a magnetic field in:

A

Diamagnetism

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12
Q

Ions in the same group have the same __________.

A

Charge

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13
Q

Ion size _________ down the column.

A

Increases

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14
Q

Cations are __________ than neutral atoms.

A

Smaller

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15
Q

Anions are __________ than neutral atoms.

A

Larger

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16
Q

Cations are __________ than anions.

A

Smaller

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17
Q

Larger positive charge = __________ cation.

A

Smaller

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18
Q

Larger __________ charge = smaller cation.

A

Positive

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19
Q

Larger negative charge = __________ anion.

A

Larger

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20
Q

Larger __________ charge = larger anion.

A

Negative

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21
Q

Same electron configuration:

A

Isoelectronic

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22
Q

When anions form cations, the valence electrons are __________.

A

Removed

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23
Q

When anions form cations, the __________ __________ are removed.

A

Valence electrons

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24
Q

When atoms form anions, electrons are __________ to the valence shell.

A

Added

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25
Q

Minimum energy needed to remove an electron from an atom or ion in the gaseous state:

A

Ionization energy

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26
Q

The larger the effective nuclear charge on the electron, the __________ energy it takes to remove it. (In first IE)

A

More

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27
Q

The __________ the effective nuclear charge on the electron, the more energy it takes to remove it. (In first IE)

A

Larger

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28
Q

The farther the most probable distance the electron is from the nucleus, the __________ energy it takes to remove it. (In first IE)

A

Less

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29
Q

The _________ the most probable distance the electron is from the nucleus, the less energy it takes to remove it. (In first IE)

A

Farther

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30
Q

First ionization __________ down the group.

A

Decreases

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31
Q

First ionization generally __________ across the period.

A

Increases

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32
Q

Ionization energy generally __________ down a column.

A

Decreases

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33
Q

Ionization energy generally _________ as we move to the right across a period (or row).

A

Increases

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34
Q

O and Cl __________ __________ __________ __________ in terms of ionization energy.

A

Cancel each other out

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35
Q

The _________ _________ of an atom or ion is the energy change associated with the gaining of an electron by the atom in the gaseous state.

A

Electron affinity (EA)

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36
Q

The electron affinity is usually __________.

A

Negative

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37
Q

We except electron affinity to become __________ __________ as we move down a column.

A

More positive

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38
Q

As we move to right across a period, metallic character __________.

A

Decreases

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39
Q

As we move down a column, metallic character __________.

A

Increases

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40
Q

With the exception of potassium, density __________ as we move down the column.

A

Increases

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41
Q

Alkali metals generally have ______ ionization energies.

A

Low

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42
Q

The relative reactivities of the alkali metals tend to __________ as we move down the column.

A

Increases

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43
Q

Transition elements all have _________ valence electrons.

A

2

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44
Q

Chemical bonds form because they lower the __________ __________ between the charged particles that compose atoms.

A

Potential energy

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45
Q

Metal + nonmetal =

A

Ionic bond

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46
Q

Nonmetal + nonmetal =

A

Covalent bond

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47
Q

Metal + metal =

A

Metallic bond

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48
Q

Characteristic of ionic bond: electrons are __________.

A

Transferred

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49
Q

Characteristic of covalent bond: electrons are __________.

A

Shared

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50
Q

Characteristic of of metallic bond: electrons are __________.

A

Pooled

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51
Q

The __________ __________ is a hypothetical series of steps that represents the formation of an ionic compound from its constituent elements.

A

Born-Haber cycle

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52
Q

The ability of an atom to attract electrons to itself in a chemical bond is called:

A

Electronegativity

53
Q

Sodium ions are pumped _____ _____ cells.

A

Out of

54
Q

Potassium ions are pumped _____ _____ cells.

A

In to

55
Q

Predictable based on an element’s position within the periodic table.

A

Periodic property

56
Q

The first attempt to organize elements according to similarities in their properties was made by:

A

Johann Dobereiner

57
Q

Johann Dobereiner grouped elements into __________.

A

Triads

58
Q

__________ __________ organized elements into octaves.

A

John Newlands

59
Q

The modern periodic table is credited primarily to:

A

Dmitri Mendeleev

60
Q

An organization similar to the periodic table had been suggested by:

A

Julius Lothar Meyer

61
Q

States that when elements are arranged in order of increasing mass, certain properties recur periodically.

A

Periodic law

62
Q

Laws __________ behavior.

A

Summarize

63
Q

Theories __________ behavior.

A

Explain

64
Q

The theory that explains the reasons behind the periodic law is the:

A

Quantum-mechanic theory

65
Q

Describes the behavior of electrons in atoms:

A

Quantum-mechanic theory

66
Q

An _________ __________ for atom shows the particular orbitals that are occupied for that atom.

A

Electron configuration

67
Q

Lowest energy state

A

Ground state

68
Q

Electrons generally occupy the __________ energy orbitals available.

A

Lowest

69
Q

The _________ __________ cannot be solved exactly.

A

Schrodinger equation

70
Q

A fundamental property of all electrons that affects the number of electrons allowed in one orbital.

A

Electron spin

71
Q

Determines the order of orbital filling within a level.

A

Sublevel energy splitting

72
Q

Symbolizes an electron as an arrow and the orbital as a box.

A

Orbital diagram

73
Q

In orbital diagrams, the direction of the arrow represents __________ __________.

A

Electron spin

74
Q

Electron spin was demonstrated experimentally in 1922 by the ____________ experiment.

A

Stern-Gerlach

75
Q

All electrons have _____ _____ amount of spin.

A

The same

76
Q

The spin of an electron is specified by a fourth quantum number called the:

A

Spin quantum number

77
Q

The possible values of the spin quantum number are:

A

+1/2 and -1/2

78
Q

+1/2 =

A

Spin up

79
Q

-1/2 =

A

Spin down

80
Q

No two electrons in an atom can have the same four quantum numbers, according to:

A

The Pauli exclusion principle

81
Q

Each orbital can have a maximum of only _____ electrons, with __________ spins.

A

2, opposing

82
Q

Two electrons occupying the same orbital have 3 __________ quantum numbers (m, l, and m1), and they must have __________ spin quantum numbers.

A

Identical, different

83
Q

A major difference in the approximate solutions for the Schrodinger equation for multi-electron atoms compared to the solutions for the hydrogen atom is the:

A

Energy ordering of the orbitals

84
Q

In the hydrogen atom, the energy of an orbital depends only on:

A

The principal quantum number (n)

85
Q

A term describing two or more electron orbitals with the same value of n that have the same energy (Ex: 3s, 3d, 3p).

A

Degenerate

86
Q

In general, the __________ the value of l within a principal level, the lower the energy of the corresponding orbital.

A

Lower

87
Q

When orbitals are not degenerate, the energies of their sublevels are __________.

A

Split

88
Q

The orbitals within a principal level of a __________ __________ are not degenerate.

A

Multielectron atom

89
Q

Describes the interactions between charged particles.

A

Coulomb’s law

90
Q

Describes how one electron can shield another electron from the full charge of the nucleus.

A

Shielding

91
Q

Describes how one atomic orbital can overlap spatially with another, thus penetrating into a region that is close to the nucleus (and therefore less shielded from nuclear charge).

A

Penetration

92
Q

The attractions and repulsions between charged particles are described by:

A

Coulomb’s law

93
Q

States that the potential energy (E) of two charged particles depends on their charges (q1 and q2) and on their separation (r).

A

Coulomb’s law

94
Q

E0=

A

8.85 x 10^-12

95
Q

Equation for Coulomb’s law:

A

E = [(1)/(4(pi)(E0))][(q1q2)/(r)]

96
Q

The __________ of the PE depends inversely on the separation between the charged particles.

A

Magnitude

97
Q

For like charges, the potential energy (E) is __________ and __________ as the particles get farther apart (as r increases).

A

Positive, decreases

98
Q

Like charges __________ each other.

A

Repel

99
Q

For opposite charges, the potential energy (E) is _________ and becomes more __________ as the particles get closer together (as r decreases).

A

Negative,negative

100
Q

Opposite charges __________ each other.

A

Attract

101
Q

The magnitude of the interaction between charged particles __________ as the charges of the particles __________.

A

Increases, increases

102
Q

An electron with a charge of 1- is __________ __________ attracted to a nucleus with a charge of 2+ than it would be to a nucleus with a charge to 1+.

A

More strongly

103
Q

For __________ atoms, any one electron experiences both the positive charge of the nucleus (which is attractive) and the negative charges of other electrons (which are repulsive).

A

Multielectron

104
Q

Another word for shielding:

A

Screening

105
Q

We can think of the repulsion of one electron by other electrons as __________ that electron from the full effects of the nuclear charge.

A

Shielding

106
Q

The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons:

A

Effective nuclear charge (Z,eff)

107
Q

_________ __________ _____ _________ is a result of the spatial distributions of electrons within a sublevel.

A

Energy splitting of sublevels

108
Q

Because of __________, the sublevels of each principal level are not degenerate for multielectron atoms.

A

Penetration

109
Q

An orbital that penetrates into the region occupied by core electrons is _____ shielded from nuclear charge than an orbital that does not penetrate and will therefore have a _____ energy.

A

Less, lower

110
Q

Indicates the pattern of orbital filling in an atom.

A

The aufbau principle

111
Q

“Build up”.

A

Aufbau

112
Q

States that when electrons fill degenerate orbitals they first fill them singly with parallel spins.

A

Hund’s rule

113
Q

_________ is a result of an atom’s tendency to find the lowest energy state possible.

A

Hund’s rule

114
Q

Orbitals fill in the following order:

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s

115
Q

Electron configuration shortcut is called:

A

Inner electron configuration

116
Q

The s sublevel has _____ orbital(s) and can hold _____ electrons.

A

One, 2

117
Q

The p sublevel has _____ orbital(s) and can hold _____ electrons.

A

Three, 6

118
Q

The d sublevel has _____ orbital(s) and can hold _____ electrons.

A

Five, 10

119
Q

The f sublevel has _____ orbital(s) and can hold _____ electrons.

A

Seven, 14

120
Q

As we move down a column, the # of electrons in the outermost principal energy level:

A

Remains the same

121
Q

Electrons important in chemical bonding.

A

Valence electrons

122
Q

Elements in the same column of the periodic table have the:

A

Same # of valence electrons

123
Q

Electrons in a complete principal energy level and those in complete d and f sublevels.

A

Core electrons

124
Q

The group number of a main-group element is equal to the # of _________ _________ for that element.

A

Valence electrons

125
Q

The row number of a main-group element is equal to the __________ __________ __________ number of that element.

A

Highest principal quantum

126
Q

The chemical properties of elements are largely determined by:

A

The # of valence electrons they contain

127
Q

The noble gases are the most __________ __________ and relatively _________ family in the periodic table.

A

Chemically stable, unreactive

128
Q

Elements with electron configurations close to those of the noble gases are the __________ __________ because they can attain noble gas electron configurations by losing or gaining a small # of electrons.

A

Most reactive

Chemistry 1b (3 decks)

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