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Group 13 Flashcards

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1
Q

Where does the last electron enter in p-block elements?

A

The last electron enters the outermost p orbital.

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2
Q

How many p orbitals are there in an atom?

A

There are three p orbitals.

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3
Q

What is the maximum number of electrons that can be accommodated in a set of p orbitals?

A

Six electrons.

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4
Q

How many groups of p-block elements are there in the periodic table?

A

There are six groups, numbered from 13 to 18.

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5
Q

Which elements head the groups of p-block elements?

A

Boron, carbon, nitrogen, oxygen, fluorine, and helium.

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6
Q

What is the valence shell electronic configuration of p-block elements?

A

ns² np¹⁻⁶ (except for helium).

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7
Q

What part of an element’s electronic configuration influences its physical and chemical properties?

A

The inner core of the electronic configuration.

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8
Q

What properties are influenced by the inner core of an element’s electronic configuration?

A

Atomic and ionic radii, ionization enthalpy, and chemical properties.

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9
Q

How does the variation in the inner core affect p-block elements in a group?

A

It causes a lot of variation in their properties.

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10
Q

What is the maximum oxidation state shown by a p-block element?

A

It is equal to the total number of valence electrons (sum of s and p electrons).

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11
Q

How does the number of possible oxidation states change across the periodic table?

A

It increases towards the right of the periodic table.

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12
Q

What is the ‘group oxidation state’ of p-block elements?

A

It is the oxidation state equal to the total number of valence electrons.

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13
Q

Can p-block elements show oxidation states other than the group oxidation state?

A

Yes, they may show oxidation states differing by a unit of two.

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14
Q

In which p-block element families is the group oxidation state most stable for lighter elements?

A

Boron, carbon, and nitrogen families.

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15
Q

What trend is observed in oxidation states for heavier elements in p-block groups?

A

The oxidation state two units less than the group oxidation state becomes more stable.

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16
Q

What is the ‘inert pair effect’?

A

It is the tendency of heavier p-block elements to favor oxidation states two units less than the group oxidation state.

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17
Q

How do the relative stabilities of oxidation states vary in p-block elements?

A

The relative stabilities of the group oxidation state and the oxidation state two units less may vary from group to group.

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18
Q

Where do non-metals and metalloids exist in the periodic table?

A

They exist only in the p-block.

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19
Q

How does non-metallic character change down a group in p-block elements?

A

It decreases down the group.

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20
Q

Which element in each p-block group is the most metallic?

A

The heaviest element in each p-block group is the most metallic.

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21
Q

What causes diversity in the chemistry of p-block elements?

A

The change from non-metallic to metallic character down the group.

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22
Q

How do ionization enthalpies and electronegativities compare between non-metals and metals?

A

Non-metals have higher ionization enthalpies and higher electronegativities than metals.

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23
Q

How do metals and non-metals differ in ion formation?

A

Metals readily form cations, whereas non-metals readily form anions.

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24
Q

Why are compounds formed between highly reactive metals and non-metals usually ionic?

A

Due to the large differences in their electronegativities.

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25
Why are compounds formed between non-metals usually covalent?
Because of the small differences in their electronegativities.
26
How does the change from non-metallic to metallic character reflect in oxides?
Non-metal oxides are acidic or neutral, while metal oxides are basic in nature.
27
Why does the first member of the p-block differ from the rest of its group?
Due to its smaller size and the absence of d-orbitals in its valence shell.
28
Which s-block elements show similar differences as the first member of the p-block?
Lithium and Beryllium.
29
Why do second-period p-block elements have a maximum covalence of four?
They use only 2s and three 2p orbitals, lacking d-orbitals.
30
How can third-period p-block elements expand their covalence beyond four?
By utilizing vacant 3d orbitals.
31
Give an example of different covalence capabilities between second- and third-period p-block elements.
Boron forms [BF₄]⁻, while aluminum can form [AlF₆]³⁻.
32
How does the presence of d-orbitals influence the chemistry of heavier p-block elements?
It affects their bonding, coordination number, and ability to form π bonds.
33
How does π bonding differ between lighter and heavier p-block elements?
Lighter elements form strong pπ-pπ bonds, while heavier ones rely on weaker dπ-pπ or dπ-dπ bonds.
34
Why are dπ-pπ and dπ-dπ bonds weaker than pπ-pπ bonds?
d-orbitals are higher in energy, contributing less to bond stability.
35
How does the coordination number vary in heavier p-block elements?
It is often higher due to the involvement of d-orbitals.
36
Give an example of different coordination numbers in the same oxidation state.
NO₃⁻ has a three-coordination with p-orbital π bonding, while PO₄³⁻ has a four-coordination using s, p, and d orbitals.
37
What type of multiple bonds can second-period p-block elements form?
They can form strong pπ-pπ multiple bonds like C=C, C≡C, N≡N, C=O, C=N, and N=O.
38
Why do heavier p-block elements struggle to form strong π bonds?
Their d-orbitals are involved in π bonding (dπ-pπ or dπ-dπ), which is weaker than pπ-pπ bonding.
39
What is a key factor influencing bond strength in p-block elements?
The relative energy and effectiveness of p-orbitals versus d-orbitals in π bonding.
40
Why does phosphorus form PO₄³⁻ while nitrogen forms NO₃⁻?
Phosphorus can use d-orbitals for expanded coordination, while nitrogen is limited to p-orbitals.
41
How does size influence the bonding behavior of p-block elements?
Smaller elements form stronger π bonds, while larger elements rely on d-orbitals, leading to weaker π bonding.
42
How do the elements of this group vary in properties?
Boron is a non-metal, aluminum is a metal with some similarities to boron, while gallium, indium, thallium, and nihonium are metallic.
43
In what forms does boron mainly occur in nature?
Orthoboric acid (H₃BO₃), borax (Na₂B₄O₇·10H₂O), kernite (Na₂B₄O₇·4H₂O).
44
What is the abundance of boron in the Earth's crust?
Less than 0.0001% by mass.
45
What are the two isotopes of boron?
10B (19%) and 11B (81%).
46
Which is the most abundant metal in the Earth's crust?
Aluminium (8.3% by mass).
47
What are the most abundant elements in the Earth's crust?
Oxygen (45.5%), Silicon (27.7%), and Aluminium (8.3%).
48
Name two important minerals of aluminum.
Bauxite (Al₂O₃·2H₂O) and cryolite (Na₃AlF₆).
49
How abundant are gallium, indium, and thallium in nature?
They are relatively rare.
50
What is the atomic number and symbol of Nihonium?
Atomic number 113, symbol Nh.
51
What is the general outer electronic configuration of group 13 elements?
ns² np¹
52
How do the electronic cores of boron and aluminum differ from those of gallium, indium, and thallium?
Boron and aluminum have a noble gas core. Gallium and indium have a noble gas core plus 10 d-electrons. Thallium has a noble gas core plus 14 f-electrons and 10 d-electrons.
53
What trend is generally expected for atomic radius down a group?
It should increase due to the addition of extra electron shells.
54
Why is the atomic radius of gallium (135 pm) smaller than that of aluminum (143 pm)?
Due to poor shielding by the 10 d-electrons, which cannot effectively counteract the increased nuclear charge.
55
How does poor shielding by d-electrons affect atomic size?
It causes a stronger attraction between the nucleus and outer electrons, reducing atomic size.
56
Ionisation Enthalpy order
B>Tl>Ga>Al>In
57
Do ionisation enthalpy values decrease smoothly down the group?
No, the ionisation enthalpy values do not decrease smoothly down the group due to general trends.
58
What causes discontinuities in ionisation enthalpy values between Al-Ga and In-Tl?
The inability of d- and f-electrons, which have low screening effects, to compensate for the increase in nuclear charge.
59
What is the general order of ionisation enthalpies?
The order is ΔiH1 < ΔiH2 < ΔiH3.
60
Why is the sum of the first three ionisation enthalpies very high?
Because removing three electrons requires significant energy, affecting their chemical properties.
61
How does electronegativity change down the group?
Electronegativity first decreases from B to Al, then increases marginally due to discrepancies in atomic size.
62
What is the nature of boron?
Boron is non-metallic, extremely hard, and a black-colored solid.
63
In how many allotropic forms does boron exist?
Boron exists in many allotropic forms.
64
Why does boron have an unusually high melting point?
Due to its very strong crystalline lattice.
65
How do the properties of the rest of the group members compare to boron?
The rest are soft metals with low melting points and high electrical conductivity.
66
Why can gallium exist in liquid form during summer?
Because it has an unusually low melting point of 303 K.
67
Why is gallium useful for measuring high temperatures?
Because it has a high boiling point of 2676 K.
68
How does the density of group elements change from boron to thallium?
Density increases down the group from boron to thallium.
69
Why does boron not form +3 ions?
Due to its small size, the sum of its first three ionization enthalpies is very high, preventing the formation of B³⁺ ions.
70
What type of compounds does boron form?
Boron forms only covalent compounds.
71
Why can aluminum form Al³⁺ ions?
The sum of its first three ionization enthalpies is considerably lower than that of boron, making Al³⁺ formation possible.
72
Why does the inert pair effect occur down the group?
Due to the poor shielding effect of intervening d- and f-orbitals, increased nuclear charge holds ns electrons tightly, restricting their participation in bonding.
73
Which orbitals participate in bonding in Ga, In, and Tl?
Only p-orbital electrons may be involved in bonding due to the inert pair effect.
74
What oxidation states are observed in Ga, In, and Tl?
Both +1 and +3 oxidation states are observed.
75
How does the stability of the +1 oxidation state change down the group?
The stability increases as Al < Ga < In < Tl.
76
Which oxidation state is predominant in thallium?
The +1 oxidation state is predominant, while the +3 oxidation state is highly oxidizing.
77
How do the ionic characteristics of +1 and +3 oxidation states compare?
Compounds in the +1 oxidation state are more ionic than those in the +3 oxidation state.
78
What are the common oxidation states in Group 13 elements?
The common oxidation states are +3 and +1.
79
Why is the +3 oxidation state more stable in lighter elements like B and Al?
In lighter elements, the effective nuclear charge is lower, allowing all valence electrons to participate in bonding, favoring the +3 state.
80
Why does the +1 oxidation state become more stable in heavier elements?
Due to the inert pair effect, ns² electrons are held more tightly and do not easily participate in bonding, making the +1 state more stable.
81
What is the trend in the stability of the +1 oxidation state down the group?
The stability of the +1 oxidation state increases down the group: Al < Ga < In < Tl.
82
Why is thallium predominantly found in the +1 oxidation state?
The inert pair effect is strongest in thallium, making the +1 state more stable than +3.
83
Why is the +3 oxidation state highly oxidizing in thallium?
The +3 state readily gains electrons to revert to the more stable +1 state, making it highly oxidizing.
84
How does the ionic nature of compounds change with oxidation state?
Compounds in the +1 oxidation state are more ionic, while those in the +3 state are more covalent.
85
What is the role of d- and f-electrons in oxidation states?
Poor shielding by d- and f-electrons increases effective nuclear charge, enhancing the inert pair effect and stabilizing the +1 oxidation state.
86
How does the tendency to behave as a Lewis acid change down the group?
It decreases with the increase in atomic size.
87
In the trivalent state, how many electrons surround the central atom in electron-deficient molecules like BF₃?
Six electrons.
88
What complex ion does aluminum chloride form in acidified aqueous solution?
[Al(H₂O)₆]³⁺
89
What does the negative E° value of Al³⁺/Al indicate about aluminum?
Aluminum has a high tendency to form Al³⁺ ions in solution.
90
What does the positive E° value of Tl³⁺/Tl suggest about thallium?
Tl³⁺ is unstable in solution and acts as a strong oxidizing agent.
91
What happens to the trichlorides of group 13 elements when hydrolyzed in water?
They form tetrahedral [M(OH)₄]⁻ species.
92
What happens when BCl₃ reacts with ammonia (NH₃)?
It forms a complex BCl₃⋅NH₃ by accepting a lone pair of electrons.
93
What is the hybridization of aluminum in the complex ion [Al(H₂O)₆]³⁺?
sp³d² hybridization.
94
What is the hybridization state of element M in [M(OH)₄]⁻?
sp³ hybridization.
95
Which metal is more electropositive, aluminum or thallium?
Aluminum, because it readily forms Al³⁺ ions.
96
Which oxidation state of thallium is more stable in solution?
Tl⁺ is more stable than Tl³⁺.
97
Why do electron-deficient molecules like BF₃ behave as Lewis acids?
They have a tendency to accept a pair of electrons to achieve a stable electronic configuration.
98
Electronegativity order ( Reverse order)
B>Tl>In>Ga>Al
99
Electronegativity order ( Reverse order)
B>Tl>In>Ga>Al

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