Ch 5 Sectjon 3 Flashcards
1
Q
Size of an atom can’t be defined by edge of orbital because this boundary is
A
Fuzzy and caries under different conditions
2
Q
To estimate he rice of an atom the conditions under with the atom exists must be
A
Specified
3
Q
One way to express an atoms radius is to measure the distance between the
A
Nuclei of two identical atoms that are chemically bonded together, then divide this distance by two
4
Q
Atomic radius may be defined as one half the durance between the
A
Nuclei of identical atoms that are bonded together
5
Q
There is a gradual decrease in atomic radii from
A
Across the second period to neon
6
Q
The trend to smaller atoms across a period is caused by the
A
Increasing positive charge of the nucleus
7
Q
As electrons add to s and p sublevels in the same main energy level they are gradually
A
Pulled closer to the more highly charged nucleus
8
Q
increased pull results in a
A
Decrease in atomic radii
9
Q
the attraction of the nucleus is somewhat offset by
A
repulsion among the increased number of electrons in the same outer energy level
10
Q
the difference in radii between neighboring atoms in each period grows
A
Smaller
11
Q
The radio of the elements
A
Increase as you read down the group
12
Q
As electrons occupy sublevels in successively higher main energy levels located farther from the nucleus the sixes of the atoms
A
Increase
13
Q
In general the atomic radii of the main group elements
A
Increase down a group
14
Q
The expected increase in fallouts readies caused by the filling of the fourth main remedy level is outweighed by a
A
Shrinking of the electron cloud caused by s nuclear charge that is considerably higher than that of aluminum
15
Q
An electron can be removed from an atom if enough
A
Energy is supplied
16
Q
A + energy–>
A
A^+ + e^-
17
Q
The A^+ represents an ion of element a with a
A
Single positive charged referred to as a 1+ ion
18
Q
An ion is an atom or group of bonded atoms that has a
A
Positive or negative charge
19
Q
Ionization: any oroceds that
A
Results in the formation of an ion
20
Q
To compare the ease with which atoms to different elements give up electrons chemists compare
A
Ionization energies
21
Q
Ionization energy (first ionization energy)
A
The energy required to remove one electron from a neutral atom
22
Q
To avoid the influence of nearby atoms measurements of ionization energies are made on
A
Isolated atoms in the gas phase
23
Q
Group 1 metals have the lowest
A
First ionization energies in their respective periods
24
Q
Group 1 metals lose electrons
A
Most easily
25
Ease of electron loss is the major reason for
Hugh reactivity of the alkali metals
26
Group 18 elements have the highest
Ionization energies
27
Group 18 elements do not
Lose electrons easily
28
The low reactivity of the noble gases is partly based on thisb
Difficulty of electron removal
29
In general, ionization energies of main group elements
Increase across each period
30
Increase in ionization energies across periods is caused by
Increasing nuclear charge
31
A higher charge more strongly attracts
Electrons in the same energy level
32
Increasing nuclear charge is responsible for
Both increasing ionization energy and decreasing radii across the periods
33
In general nobmetals have higher ionization energies than
Metals do
34
In each period the element of group 1 has the lowest
Ionization energy and the element of group 18 has the highest ionization energy
35
Among the main group elements ionization energies generally
Decrease down the groups
36
Electrons removed from atoms of each succeeding element in s group are in
Higher energies and are therefore removed more easily
37
As atomic number increase going down a group more electrons lie between the
Nucleus and the electrons in the highest occupied energy levels
38
(More electrons lie between the...) this partially shields the
Outer electrons from the effect of the nuclear charge
39
These influences overcome the attest toon of the electrons to
Increasing nuclear charge
40
With sufficient energy electrons can be removed from
Positive ions as well as from neutral atoms
41
The energies for removal of additional electrons from an atom are referred to as the
Second ionization energy, third ionization energy and so on
42
The second ionization energy is always
Higher than the first
43
The third ionization energy is always
Higher than the second
44
The reason second ie is higher than first, etc is because as electrons are removed in successive ionizations fewer
Electrons remain within the atom to shield the attractive fierce of the nucleus
45
Each successive electron removed from an ion feels an increasingly
Stronger effective nuclear charge (the nuclear charge minus the electron shielding)
46
Removing a single electron from an atom in group 18 elements is more difficult than removing an electron from atoms of
Other elements in the same period
47
The special stability of the noble gas configuration also applies to ions that have
Noble gas configurations
48
The jump in ionization energy occurs when an ion assumes a
Noble gas configuration
49
Electrons affinity is the beefy change that occurs when
An electron is acquired by a neutral atom
50
Most atoms... Energy when they acquire an electron
Release
51
Release of energy when atoms acquire an electron) A + e^- --->
A^- + energy
52
Some atoms must be... To gain an..:
Forced to gain an electron by the addition of energy
53
(Forced gain of electron) A + e^_ + energy --->
A^-
54
The quantity of energy absorbed would be represented by s positive number but ions produced in this way are
Very unstable and hence the electron affinity for them is very difficult to determine
55
An ion produced in this way will be
Unstable and will lose the added electron spontaneously
56
The halogens gain electrons lost
Readily
57
The ease with which halogen atoms gain electrons is a major reason for the
High deactivated of the group 17 elements
58
In general, as electrons add to the same o sublevel of atoms with increasing nuclear charge electron affinities become
More negative across each period within the p block (exception between groups 14 and 15)
59
Trends for electron affinities within groups are not as regular as
Trends for ionization energies
60
As a general rule electrons add with greater... Down a grouo
Difficulty
61
This pattern(greater difficulty in adding electrons down group) is a result of two
Compete ring factors
62
The first competing factor is a slight increase in
Effective nuclear charge down a group, which increases electron affinities
63
There is a rough correlation between the arrangement of
Elements and their electron configurations
64
The second competing factor is an increase in atomic
radius down a group which decreases electron affinities
65
In general the size effect
Predominates (exceptions. Esp among heavy transition metals which tend to be the same size or even decrease in radius down a group)
66
For an isolated Jon in the gas phase it is always more difficult to add a
Second electron to an already negatively charged ion
67
Second electron affinities are all
Positive
68
Certain p block Nonmetals tend to form
Ions that have noble gas configurations
69
The halogens form ions that have noble gas configurations by
Adding one electron
70
Adding another electron to cl is so difficult that cl^2-
Never occurs
71
Atoms of group 16 elements are present in many
Compounds as 2- ions
72
Cation
Positive ion
73
Anion
Negative ion
74
Formation of a cation by the loss of one or more electrons always leads to a decrease in
Atomic radius because removal of highest energy level electrons results in smaller electron cloud
75
Remaining electrons are drawn
Closer to the nucleus by its unbalanced positive charge
76
Formation of an anion by the addition of one or more electrons always leads to
An increase in atomic radius because total positive charge of the nucleus remains unchanged when an electron is added to an atom or an ion
77
Electrons are not drawn to the nucleus as strongly as they were before
The addition of the extra electron
78
Electron cloud also spreads out because of greater
Repulsion between the increased number of electrons
79
Within each period of the periodic table the metals at the left tend to form
Cations
80
Nonmetals at the upper right of periodic table tend to form
Anions
81
Carbonic radii decrease across a
Period
82
Carbonic radii decrease across period because the electron cloud
Shrinks due to increasing nuclear charge acting on the electrons in the same main energy level
83
Starting with group 15 anions are more common than
Cations
84
An ionic radii decrease
Across each period for elements in groups 15-18
85
Reasons for the an ionic radii trend are the same as the reasons that car ionic radii decrease from
Left to right across a period
86
The outer electrons in both cations and anions are in
Higher energy levels as one reads down a group
87
Just as there is a gradual increase of atomic radii down a group there is also a gradual
Increase of ionic radii
88
Chemical compounds form because electrons are
Lost, gained, or shared between atoms
89
The valence electrons are the electrons most subject to the influence of
Nearby atoms or ions
90
Valence electrons: the electrons available to be
Lost, gained, or shared in the formation of chemical compounds
91
Valence electrons are often located in
Incompletely filled main energy levels
92
The inner electrons are in filled energy levels and are held too tightly by the nucleus to be involved in
Compound formation
93
The group 1 and group 2 elements have
One and 2 valence electrons
94
Elements of groups 13-18 have a number of valence electrons equal to the
Group number minus 10
95
In some cases, both the s and p sublevel valence electrons of the o block elements are involved in
Compound formation
96
Valence electrons hold atoms together in
Chemical compounds
97
In many compounds the negative charge of the valence electrons is concentrated closer to
One atom than to another
98
This uneven concentration of charge has a significant effect on the
Chemical properties of s compound
99
It is useful to have a measure of how strongly one atom attracts the
Electrons of another atom within a compound
100
Linus Pauling devised a scale of
Numerical values reflecting the tendency of an atom to attract electrons
101
Electro negativity is a measure of the ability of an atom in a
Chemical compound to attract electrons from another atom in the compound
102
The most electronegative element, fluorine, is arbitrarily assigned an electronegative ti value of
Four
103
Electro negativity values for the other elements are then calculated in
Relation to the value of fluorine
104
Electronegative stems to ... Across each period (excretions apply)
Increase
105
Alkali and alkaline earth metals are the least
Electronegative elements
106
Nitrogen oxygen and halogens are the most
Electronegative elements
107
Atoms of nitrogen oxygen and halogens attract electrons
Strongly in compounds
108
Electronegativities tend to either decrease
Down a group or remain about the same
109
Noble gases are unusual in that some do not form
Compounds and thus can't be assigned electronegativities
110
When s noble gas does form s compound its electronegativities is rather
High
| Similar to values for the halogens
111
The combination of the period and group trends in electronegativities results in the
Highest values belonging to the elements in the upper right of the periodic table
112
The properties of the d block elements baru less and with less regularity than those of the
Main group elements
113
Atoms of the d block elements contain from zero to two
Electrons in the s orbital of their highest occupied energy level
114
Atoms of the d block elements contain from one to ten electrons in the
D sublevel of the next lower energy level
115
Electrons in both the ns sublevel and the (n-1)d sublevel are available to
Interact with their surroundings
116
Electrons in the incompletely filled d sublevels are responsible for many
Characteristic properties of the d block elements
117
The atomic radii of the d block elements generally
Decrease across the periods
118
Decrease in atomic radii of the d block elements is
Less than that for the main group elements
119
Decrease in atomic radii of d block elements is less than that for main group elements because the electrons added to the
(n-1)d sublevel shield the outer electrons from the nucleus
120
The radii dip to a low and then increase slightly across each of the
Four period that contain d block elements
121
As the number of electrons in the d sublevel increase, the radii ... Because of...
Increase because of repulsion among the electrons
122
Because of increase in atomic number that occurs from lanthanum to hafnium, the atomic radius of hafnium is actually
Slightly less than that of zirconium (element immediately above it)
123
Radii of elements following hafnium in the sixth period Gary with
Increasing atomic number in the usual manner
124
Ionization defied of he d block and f block elements generally
Increase across the periods
125
The first ionization energies of the d block elements generally
Increase down each group
126
The reason IE 1 of d block elements increases down each group is because the electrons available for ionization in the outer s sublevels are
Less shielded from the increasing nuclear charge by the electrons in the incomplete (n-1)d sublevels
127
Amin all atoms of the d block and f block elements electrons in the highest occupied sublevel are always
Removed first
128
For the d block elemts this means that although newly added electrons occupy the d sublevels the
First electrons to be removed are those in the outermost sublevels
129
Most d block elements commonly form 2+ ions in
Compounds
130
Some d block elements (e. g. Iron and chromium) commonly form
3+ ions
131
Group 3 elements form only ions with a
3+ charge
132
Cations have smaller
Radii than the atoms do
133
Comparing 2+ ions across the periods shows a ... In size that parallels the... In atomic radius
Decrease; decrease
134
D block elements all have electronegativities between
1.1 and 2.54
135
Only the active metals of groups 1 and 2 have lower electronegativities than
D block elements
136
D block elements also follow the general trend for
Electronegativity values to increase as radii decrease and vice versa
137
F block elements all have similar electronegativities which range from
1.1 to 1.5
