Ch 5 Sectjon 3 Flashcards Preview

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Flashcards in Ch 5 Sectjon 3 Deck (137):
1

Size of an atom can't be defined by edge of orbital because this boundary is

Fuzzy and caries under different conditions

2

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

Specified

3

One way to express an atoms radius is to measure the distance between the

Nuclei of two identical atoms that are chemically bonded together, then divide this distance by two

4

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

Nuclei of identical atoms that are bonded together

5

There is a gradual decrease in atomic radii from

Across the second period to neon

6

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

Increasing positive charge of the nucleus

7

As electrons add to s and p sublevels in the same main energy level they are gradually

Pulled closer to the more highly charged nucleus

8

increased pull results in a

Decrease in atomic radii

9

the attraction of the nucleus is somewhat offset by

repulsion among the increased number of electrons in the same outer energy level

10

the difference in radii between neighboring atoms in each period grows

Smaller

11

The radio of the elements

Increase as you read down the group

12

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

Increase

13

In general the atomic radii of the main group elements

Increase down a group

14

The expected increase in fallouts readies caused by the filling of the fourth main remedy level is outweighed by a

Shrinking of the electron cloud caused by s nuclear charge that is considerably higher than that of aluminum

15

An electron can be removed from an atom if enough

Energy is supplied

16

A + energy-->

A^+ + e^-

17

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

Single positive charged referred to as a 1+ ion

18

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

Positive or negative charge

19

Ionization: any oroceds that

Results in the formation of an ion

20

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

Ionization energies

21

Ionization energy (first ionization energy)

The energy required to remove one electron from a neutral atom

22

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

Isolated atoms in the gas phase

23

Group 1 metals have the lowest

First ionization energies in their respective periods

24

Group 1 metals lose electrons

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