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Chem > Ch 5 Sectjon 3 > Flashcards

Ch 5 Sectjon 3 Flashcards

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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

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2
Q

To estimate he rice of an atom the conditions under with the atom exists must be

A

Specified

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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

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4
Q

Atomic radius may be defined as one half the durance between the

A

Nuclei of identical atoms that are bonded together

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5
Q

There is a gradual decrease in atomic radii from

A

Across the second period to neon

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6
Q

The trend to smaller atoms across a period is caused by the

A

Increasing positive charge of the nucleus

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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

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8
Q

increased pull results in a

A

Decrease in atomic radii

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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

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10
Q

the difference in radii between neighboring atoms in each period grows

A

Smaller

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11
Q

The radio of the elements

A

Increase as you read down the group

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12
Q

As electrons occupy sublevels in successively higher main energy levels located farther from the nucleus the sixes of the atoms

A

Increase

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13
Q

In general the atomic radii of the main group elements

A

Increase down a group

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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

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15
Q

An electron can be removed from an atom if enough

A

Energy is supplied

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16
Q

A + energy–>

A

A^+ + e^-

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17
Q

The A^+ represents an ion of element a with a

A

Single positive charged referred to as a 1+ ion

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18
Q

An ion is an atom or group of bonded atoms that has a

A

Positive or negative charge

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19
Q

Ionization: any oroceds that

A

Results in the formation of an ion

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20
Q

To compare the ease with which atoms to different elements give up electrons chemists compare

A

Ionization energies

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21
Q

Ionization energy (first ionization energy)

A

The energy required to remove one electron from a neutral atom

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22
Q

To avoid the influence of nearby atoms measurements of ionization energies are made on

A

Isolated atoms in the gas phase

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23
Q

Group 1 metals have the lowest

A

First ionization energies in their respective periods

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24
Q

Group 1 metals lose electrons

A

Most easily

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25
Q

Ease of electron loss is the major reason for

A

Hugh reactivity of the alkali metals

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26
Q

Group 18 elements have the highest

A

Ionization energies

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27
Q

Group 18 elements do not

A

Lose electrons easily

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28
Q

The low reactivity of the noble gases is partly based on thisb

A

Difficulty of electron removal

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29
Q

In general, ionization energies of main group elements

A

Increase across each period

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30
Q

Increase in ionization energies across periods is caused by

A

Increasing nuclear charge

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31
Q

A higher charge more strongly attracts

A

Electrons in the same energy level

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32
Q

Increasing nuclear charge is responsible for

A

Both increasing ionization energy and decreasing radii across the periods

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33
Q

In general nobmetals have higher ionization energies than

A

Metals do

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34
Q

In each period the element of group 1 has the lowest

A

Ionization energy and the element of group 18 has the highest ionization energy

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35
Q

Among the main group elements ionization energies generally

A

Decrease down the groups

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36
Q

Electrons removed from atoms of each succeeding element in s group are in

A

Higher energies and are therefore removed more easily

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37
Q

As atomic number increase going down a group more electrons lie between the

A

Nucleus and the electrons in the highest occupied energy levels

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38
Q

(More electrons lie between the…) this partially shields the

A

Outer electrons from the effect of the nuclear charge

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39
Q

These influences overcome the attest toon of the electrons to

A

Increasing nuclear charge

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40
Q

With sufficient energy electrons can be removed from

A

Positive ions as well as from neutral atoms

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41
Q

The energies for removal of additional electrons from an atom are referred to as the

A

Second ionization energy, third ionization energy and so on

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42
Q

The second ionization energy is always

A

Higher than the first

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43
Q

The third ionization energy is always

A

Higher than the second

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44
Q

The reason second ie is higher than first, etc is because as electrons are removed in successive ionizations fewer

A

Electrons remain within the atom to shield the attractive fierce of the nucleus

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45
Q

Each successive electron removed from an ion feels an increasingly

A

Stronger effective nuclear charge (the nuclear charge minus the electron shielding)

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46
Q

Removing a single electron from an atom in group 18 elements is more difficult than removing an electron from atoms of

A

Other elements in the same period

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47
Q

The special stability of the noble gas configuration also applies to ions that have

A

Noble gas configurations

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48
Q

The jump in ionization energy occurs when an ion assumes a

A

Noble gas configuration

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49
Q

Electrons affinity is the beefy change that occurs when

A

An electron is acquired by a neutral atom

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50
Q

Most atoms… Energy when they acquire an electron

A

Release

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51
Q

Release of energy when atoms acquire an electron) A + e^- —>

A

A^- + energy

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52
Q

Some atoms must be… To gain an..:

A

Forced to gain an electron by the addition of energy

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53
Q

(Forced gain of electron) A + e^_ + energy —>

A

A^-

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54
Q

The quantity of energy absorbed would be represented by s positive number but ions produced in this way are

A

Very unstable and hence the electron affinity for them is very difficult to determine

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55
Q

An ion produced in this way will be

A

Unstable and will lose the added electron spontaneously

56
Q

The halogens gain electrons lost

A

Readily

57
Q

The ease with which halogen atoms gain electrons is a major reason for the

A

High deactivated of the group 17 elements

58
Q

In general, as electrons add to the same o sublevel of atoms with increasing nuclear charge electron affinities become

A

More negative across each period within the p block (exception between groups 14 and 15)

59
Q

Trends for electron affinities within groups are not as regular as

A

Trends for ionization energies

60
Q

As a general rule electrons add with greater… Down a grouo

A

Difficulty

61
Q

This pattern(greater difficulty in adding electrons down group) is a result of two

A

Compete ring factors

62
Q

The first competing factor is a slight increase in

A

Effective nuclear charge down a group, which increases electron affinities

63
Q

There is a rough correlation between the arrangement of

A

Elements and their electron configurations

64
Q

The second competing factor is an increase in atomic

A

radius down a group which decreases electron affinities

65
Q

In general the size effect

A

Predominates (exceptions. Esp among heavy transition metals which tend to be the same size or even decrease in radius down a group)

66
Q

For an isolated Jon in the gas phase it is always more difficult to add a

A

Second electron to an already negatively charged ion

67
Q

Second electron affinities are all

A

Positive

68
Q

Certain p block Nonmetals tend to form

A

Ions that have noble gas configurations

69
Q

The halogens form ions that have noble gas configurations by

A

Adding one electron

70
Q

Adding another electron to cl is so difficult that cl^2-

A

Never occurs

71
Q

Atoms of group 16 elements are present in many

A

Compounds as 2- ions

72
Q

Cation

A

Positive ion

73
Q

Anion

A

Negative ion

74
Q

Formation of a cation by the loss of one or more electrons always leads to a decrease in

A

Atomic radius because removal of highest energy level electrons results in smaller electron cloud

75
Q

Remaining electrons are drawn

A

Closer to the nucleus by its unbalanced positive charge

76
Q

Formation of an anion by the addition of one or more electrons always leads to

A

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
Q

Electrons are not drawn to the nucleus as strongly as they were before

A

The addition of the extra electron

78
Q

Electron cloud also spreads out because of greater

A

Repulsion between the increased number of electrons

79
Q

Within each period of the periodic table the metals at the left tend to form

A

Cations

80
Q

Nonmetals at the upper right of periodic table tend to form

A

Anions

81
Q

Carbonic radii decrease across a

A

Period

82
Q

Carbonic radii decrease across period because the electron cloud

A

Shrinks due to increasing nuclear charge acting on the electrons in the same main energy level

83
Q

Starting with group 15 anions are more common than

A

Cations

84
Q

An ionic radii decrease

A

Across each period for elements in groups 15-18

85
Q

Reasons for the an ionic radii trend are the same as the reasons that car ionic radii decrease from

A

Left to right across a period

86
Q

The outer electrons in both cations and anions are in

A

Higher energy levels as one reads down a group

87
Q

Just as there is a gradual increase of atomic radii down a group there is also a gradual

A

Increase of ionic radii

88
Q

Chemical compounds form because electrons are

A

Lost, gained, or shared between atoms

89
Q

The valence electrons are the electrons most subject to the influence of

A

Nearby atoms or ions

90
Q

Valence electrons: the electrons available to be

A

Lost, gained, or shared in the formation of chemical compounds

91
Q

Valence electrons are often located in

A

Incompletely filled main energy levels

92
Q

The inner electrons are in filled energy levels and are held too tightly by the nucleus to be involved in

A

Compound formation

93
Q

The group 1 and group 2 elements have

A

One and 2 valence electrons

94
Q

Elements of groups 13-18 have a number of valence electrons equal to the

A

Group number minus 10

95
Q

In some cases, both the s and p sublevel valence electrons of the o block elements are involved in

A

Compound formation

96
Q

Valence electrons hold atoms together in

A

Chemical compounds

97
Q

In many compounds the negative charge of the valence electrons is concentrated closer to

A

One atom than to another

98
Q

This uneven concentration of charge has a significant effect on the

A

Chemical properties of s compound

99
Q

It is useful to have a measure of how strongly one atom attracts the

A

Electrons of another atom within a compound

100
Q

Linus Pauling devised a scale of

A

Numerical values reflecting the tendency of an atom to attract electrons

101
Q

Electro negativity is a measure of the ability of an atom in a

A

Chemical compound to attract electrons from another atom in the compound

102
Q

The most electronegative element, fluorine, is arbitrarily assigned an electronegative ti value of

A

Four

103
Q

Electro negativity values for the other elements are then calculated in

A

Relation to the value of fluorine

104
Q

Electronegative stems to … Across each period (excretions apply)

A

Increase

105
Q

Alkali and alkaline earth metals are the least

A

Electronegative elements

106
Q

Nitrogen oxygen and halogens are the most

A

Electronegative elements

107
Q

Atoms of nitrogen oxygen and halogens attract electrons

A

Strongly in compounds

108
Q

Electronegativities tend to either decrease

A

Down a group or remain about the same

109
Q

Noble gases are unusual in that some do not form

A

Compounds and thus can’t be assigned electronegativities

110
Q

When s noble gas does form s compound its electronegativities is rather

A

High

Similar to values for the halogens

111
Q

The combination of the period and group trends in electronegativities results in the

A

Highest values belonging to the elements in the upper right of the periodic table

112
Q

The properties of the d block elements baru less and with less regularity than those of the

A

Main group elements

113
Q

Atoms of the d block elements contain from zero to two

A

Electrons in the s orbital of their highest occupied energy level

114
Q

Atoms of the d block elements contain from one to ten electrons in the

A

D sublevel of the next lower energy level

115
Q

Electrons in both the ns sublevel and the (n-1)d sublevel are available to

A

Interact with their surroundings

116
Q

Electrons in the incompletely filled d sublevels are responsible for many

A

Characteristic properties of the d block elements

117
Q

The atomic radii of the d block elements generally

A

Decrease across the periods

118
Q

Decrease in atomic radii of the d block elements is

A

Less than that for the main group elements

119
Q

Decrease in atomic radii of d block elements is less than that for main group elements because the electrons added to the

A

(n-1)d sublevel shield the outer electrons from the nucleus

120
Q

The radii dip to a low and then increase slightly across each of the

A

Four period that contain d block elements

121
Q

As the number of electrons in the d sublevel increase, the radii … Because of…

A

Increase because of repulsion among the electrons

122
Q

Because of increase in atomic number that occurs from lanthanum to hafnium, the atomic radius of hafnium is actually

A

Slightly less than that of zirconium (element immediately above it)

123
Q

Radii of elements following hafnium in the sixth period Gary with

A

Increasing atomic number in the usual manner

124
Q

Ionization defied of he d block and f block elements generally

A

Increase across the periods

125
Q

The first ionization energies of the d block elements generally

A

Increase down each group

126
Q

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

A

Less shielded from the increasing nuclear charge by the electrons in the incomplete (n-1)d sublevels

127
Q

Amin all atoms of the d block and f block elements electrons in the highest occupied sublevel are always

A

Removed first

128
Q

For the d block elemts this means that although newly added electrons occupy the d sublevels the

A

First electrons to be removed are those in the outermost sublevels

129
Q

Most d block elements commonly form 2+ ions in

A

Compounds

130
Q

Some d block elements (e. g. Iron and chromium) commonly form

A

3+ ions

131
Q

Group 3 elements form only ions with a

A

3+ charge

132
Q

Cations have smaller

A

Radii than the atoms do

133
Q

Comparing 2+ ions across the periods shows a … In size that parallels the… In atomic radius

A

Decrease; decrease

134
Q

D block elements all have electronegativities between

A

1.1 and 2.54

135
Q

Only the active metals of groups 1 and 2 have lower electronegativities than

A

D block elements

136
Q

D block elements also follow the general trend for

A

Electronegativity values to increase as radii decrease and vice versa

137
Q

F block elements all have similar electronegativities which range from

A

1.1 to 1.5

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