Chapter 8 Flashcards
1
Q
Atoms in the same group __________ in size down the column.
A
Increase
2
Q
Atomic radii of __________ __________ roughly the same size across the d block.
A
Transition metals
3
Q
__________ form when the atom loses electrons from the valence shell.
A
Cations
4
Q
When transition metals form cations, the first electrons removed are __________ __________, even though other electrons were added after.
A
Valence electrons
5
Q
Electrons may also be removed from the sublevel __________ to the valence shell after the valence electrons.
A
Closest
6
Q
The iron atom has _____ valence electrons.
A
2
7
Q
When ion forms a cation, it first __________ it’s valence electrons.
A
Loses
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
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
10
Q
Atom or ion will be attracted to a magnetic field in:
A
Paramagnetism
11
Q
Atom or ion will be slightly repelled by a magnetic field in:
A
Diamagnetism
12
Q
Ions in the same group have the same __________.
A
Charge
13
Q
Ion size _________ down the column.
A
Increases
14
Q
Cations are __________ than neutral atoms.
A
Smaller
15
Q
Anions are __________ than neutral atoms.
A
Larger
16
Q
Cations are __________ than anions.
A
Smaller
17
Q
Larger positive charge = __________ cation.
A
Smaller
18
Q
Larger __________ charge = smaller cation.
A
Positive
19
Q
Larger negative charge = __________ anion.
A
Larger
20
Q
Larger __________ charge = larger anion.
A
Negative
21
Q
Same electron configuration:
A
Isoelectronic
22
Q
When anions form cations, the valence electrons are __________.
A
Removed
23
Q
When anions form cations, the __________ __________ are removed.
A
Valence electrons
24
Q
When atoms form anions, electrons are __________ to the valence shell.
A
Added
25
Minimum energy needed to remove an electron from an atom or ion in the gaseous state:
Ionization energy
26
The larger the effective nuclear charge on the electron, the __________ energy it takes to remove it. (In first IE)
More
27
The __________ the effective nuclear charge on the electron, the more energy it takes to remove it. (In first IE)
Larger
28
The farther the most probable distance the electron is from the nucleus, the __________ energy it takes to remove it. (In first IE)
Less
29
The _________ the most probable distance the electron is from the nucleus, the less energy it takes to remove it. (In first IE)
Farther
30
First ionization __________ down the group.
Decreases
31
First ionization generally __________ across the period.
Increases
32
Ionization energy generally __________ down a column.
Decreases
33
Ionization energy generally _________ as we move to the right across a period (or row).
Increases
34
O and Cl __________ __________ __________ __________ in terms of ionization energy.
Cancel each other out
35
The _________ _________ of an atom or ion is the energy change associated with the gaining of an electron by the atom in the gaseous state.
Electron affinity (EA)
36
The electron affinity is usually __________.
Negative
37
We except electron affinity to become __________ __________ as we move down a column.
More positive
38
As we move to right across a period, metallic character __________.
Decreases
39
As we move down a column, metallic character __________.
Increases
40
With the exception of potassium, density __________ as we move down the column.
Increases
41
Alkali metals generally have ______ ionization energies.
Low
42
The relative reactivities of the alkali metals tend to __________ as we move down the column.
Increases
43
Transition elements all have _________ valence electrons.
2
44
Chemical bonds form because they lower the __________ __________ between the charged particles that compose atoms.
Potential energy
45
Metal + nonmetal =
Ionic bond
46
Nonmetal + nonmetal =
Covalent bond
47
Metal + metal =
Metallic bond
48
Characteristic of ionic bond: electrons are __________.
Transferred
49
Characteristic of covalent bond: electrons are __________.
Shared
50
Characteristic of of metallic bond: electrons are __________.
Pooled
51
The __________ __________ is a hypothetical series of steps that represents the formation of an ionic compound from its constituent elements.
Born-Haber cycle
52
The ability of an atom to attract electrons to itself in a chemical bond is called:
Electronegativity
53
Sodium ions are pumped _____ _____ cells.
Out of
54
Potassium ions are pumped _____ _____ cells.
In to
55
Predictable based on an element's position within the periodic table.
Periodic property
56
The first attempt to organize elements according to similarities in their properties was made by:
Johann Dobereiner
57
Johann Dobereiner grouped elements into __________.
Triads
58
__________ __________ organized elements into octaves.
John Newlands
59
The modern periodic table is credited primarily to:
Dmitri Mendeleev
60
An organization similar to the periodic table had been suggested by:
Julius Lothar Meyer
61
States that when elements are arranged in order of increasing mass, certain properties recur periodically.
Periodic law
62
Laws __________ behavior.
Summarize
63
Theories __________ behavior.
Explain
64
The theory that explains the reasons behind the periodic law is the:
Quantum-mechanic theory
65
Describes the behavior of electrons in atoms:
Quantum-mechanic theory
66
An _________ __________ for atom shows the particular orbitals that are occupied for that atom.
Electron configuration
67
Lowest energy state
Ground state
68
Electrons generally occupy the __________ energy orbitals available.
Lowest
69
The _________ __________ cannot be solved exactly.
Schrodinger equation
70
A fundamental property of all electrons that affects the number of electrons allowed in one orbital.
Electron spin
71
Determines the order of orbital filling within a level.
Sublevel energy splitting
72
Symbolizes an electron as an arrow and the orbital as a box.
Orbital diagram
73
In orbital diagrams, the direction of the arrow represents __________ __________.
Electron spin
74
Electron spin was demonstrated experimentally in 1922 by the ____________ experiment.
Stern-Gerlach
75
All electrons have _____ _____ amount of spin.
The same
76
The spin of an electron is specified by a fourth quantum number called the:
Spin quantum number
77
The possible values of the spin quantum number are:
+1/2 and -1/2
78
+1/2 =
Spin up
79
-1/2 =
Spin down
80
No two electrons in an atom can have the same four quantum numbers, according to:
The Pauli exclusion principle
81
Each orbital can have a maximum of only _____ electrons, with __________ spins.
2, opposing
82
Two electrons occupying the same orbital have 3 __________ quantum numbers (m, l, and m1), and they must have __________ spin quantum numbers.
Identical, different
83
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:
Energy ordering of the orbitals
84
In the hydrogen atom, the energy of an orbital depends only on:
The principal quantum number (n)
85
A term describing two or more electron orbitals with the same value of n that have the same energy (Ex: 3s, 3d, 3p).
Degenerate
86
In general, the __________ the value of l within a principal level, the lower the energy of the corresponding orbital.
Lower
87
When orbitals are not degenerate, the energies of their sublevels are __________.
Split
88
The orbitals within a principal level of a __________ __________ are not degenerate.
Multielectron atom
89
Describes the interactions between charged particles.
Coulomb's law
90
Describes how one electron can shield another electron from the full charge of the nucleus.
Shielding
91
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).
Penetration
92
The attractions and repulsions between charged particles are described by:
Coulomb's law
93
States that the potential energy (E) of two charged particles depends on their charges (q1 and q2) and on their separation (r).
Coulomb's law
94
E0=
8.85 x 10^-12
95
Equation for Coulomb's law:
E = [(1)/(4(pi)(E0))][(q1q2)/(r)]
96
The __________ of the PE depends inversely on the separation between the charged particles.
Magnitude
97
For like charges, the potential energy (E) is __________ and __________ as the particles get farther apart (as r increases).
Positive, decreases
98
Like charges __________ each other.
Repel
99
For opposite charges, the potential energy (E) is _________ and becomes more __________ as the particles get closer together (as r decreases).
Negative,negative
100
Opposite charges __________ each other.
Attract
101
The magnitude of the interaction between charged particles __________ as the charges of the particles __________.
Increases, increases
102
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+.
More strongly
103
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).
Multielectron
104
Another word for shielding:
Screening
105
We can think of the repulsion of one electron by other electrons as __________ that electron from the full effects of the nuclear charge.
Shielding
106
The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons:
Effective nuclear charge (Z,eff)
107
_________ __________ _____ _________ is a result of the spatial distributions of electrons within a sublevel.
Energy splitting of sublevels
108
Because of __________, the sublevels of each principal level are not degenerate for multielectron atoms.
Penetration
109
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.
Less, lower
110
Indicates the pattern of orbital filling in an atom.
The aufbau principle
111
"Build up".
Aufbau
112
States that when electrons fill degenerate orbitals they first fill them singly with parallel spins.
Hund's rule
113
_________ is a result of an atom's tendency to find the lowest energy state possible.
Hund's rule
114
Orbitals fill in the following order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s
115
Electron configuration shortcut is called:
Inner electron configuration
116
The s sublevel has _____ orbital(s) and can hold _____ electrons.
One, 2
117
The p sublevel has _____ orbital(s) and can hold _____ electrons.
Three, 6
118
The d sublevel has _____ orbital(s) and can hold _____ electrons.
Five, 10
119
The f sublevel has _____ orbital(s) and can hold _____ electrons.
Seven, 14
120
As we move down a column, the # of electrons in the outermost principal energy level:
Remains the same
121
Electrons important in chemical bonding.
Valence electrons
122
Elements in the same column of the periodic table have the:
Same # of valence electrons
123
Electrons in a complete principal energy level and those in complete d and f sublevels.
Core electrons
124
The group number of a main-group element is equal to the # of _________ _________ for that element.
Valence electrons
125
The row number of a main-group element is equal to the __________ __________ __________ number of that element.
Highest principal quantum
126
The chemical properties of elements are largely determined by:
The # of valence electrons they contain
127
The noble gases are the most __________ __________ and relatively _________ family in the periodic table.
Chemically stable, unreactive
128
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.
Most reactive
