P-Block Elements: Groups 13-18

• The properties of p-block elements vary due to the influence of d and f electrons in the inner core of heavier elements
• In p-block elements, the last electron occupies the outermost p orbital
• Each p orbital can hold a maximum of two electrons
• There are six groups of p-block elements, numbered 13 to 18

Group 13-18 Elements

• Boron
• Carbon
• Nitrogen
• Oxygen
• Fluorine
• Helium (heads the groups)
• The valence shell electronic configuration is ns²np¹⁻⁶, with the exception of helium
• Inner core electronic configuration differences have influence on physical and chemical properties

Oxidation

• The maximum oxidation state matches the total valence electrons, increasing from left to right on the periodic table
• P-block elements exhibit varying oxidation states, typically differing by two units from the total valence electrons
• Elements within Boron, Carbon and Nitrogen families have group oxidation state that is most stable but only for lighter elements in the groups
• Progression down the group sees a group oxidation state that is two units less becoming progressively more stable for the heavier elements due to the 'inert pair effect'

Stability

• The stability of group oxidation states and those two units less vary across groups

Metals vs Non-Metals

• Non-metals and metalloids are exclusive to the p-block
• Descending the group sees a decrease in non-metallic properties, the heaviest element in each p-block group leans metallic
• This transition from non-metallic to metallic affects the chemistry of elements in said groups
• As a rule of thumb, non-metals tends to have higher electronegativities and ionization enthalpies than metals

Cations vs Anions

• Unlike metals which readily form cations, non-metals form anions
• High-reactivity non-metals form ionic compounds with high-reactivity metals because of the differences in electronegativity
• Compounds form between non-metals, largely covalent, due to small electronegativity differences
• The change from non-metallic to metallic is illustrated by oxide types, non-metal oxides are acidic/neutral while metallic ones are basic

P-Block Elements

• The first p-block element differs from its group members
• This difference is explained by properties and size
• Lightest p-block elements share differences seen in light s-block elements like lithium and beryllium
• The second difference found applies only to p-block elements which stems from the effect of d-orbitals in heavier elements and their absence in the second period elements

Covalence

• Second-period p-group elements, starting with boron, is restricted to a maximum covalence of four through using 2s and three 2p orbitals
• Third period elements in the p-groups may have a 3s²3pⁿ electronic configuration where vacant 3d orbitals lie between the 3p and 4s energy levels, the use of d-orbitals can expand their covalence above four
• While Boron forms only [BF₄]⁻, aluminium gives [AlF₆]³⁻, the presence of d-orbitals influences the chemistry of heavier elements
• The ability of elements to form π bonds is considerably influenced by size and availability of d orbitals
• The first member of a group differs from heavier members in its ability to from pπ - pπ multiple bonds to itself

Types of Pπ - Pπ bonds

• C=C
• C≡C
• N≡N
• With more, second-row elements include;
  • C=O
  • C=N
  • C≡N
  • N=O
• These types of π - bonding are not strong for the heavier p-block elements

Heavier Elements

• Heavier elements can form π bonds using d orbitals
• Consisting of;
  • dπ - pπ
  • dπ –dπ
• Dorbitals are higher in energy than p orbitals, and contribute less to the overall stability of molecules than pr - pr bonding of second-row elements
• Coordination number in species of heavier elements may be higher than for the first element in the same oxidation state.
• N and P in a +5 oxidation exhibits oxoanions, NO₃⁻ (three-coordination with π – bond involving one nitrogen p-orbital) and PO₄³⁻(four-coordination involving s, p and d orbitals contributing to the π – bond).

Group 13

• Properties vary widely
  • Boron is a typical non-metal
  • Aluminium is a metal with chemical similarities to Boron
  • Gallium, Indium, Thallium, and Nihonium are almost exclusively metallic

Boron

• Boron is a rare element, mainly present as:
  • Orthoboric acid (H₃BO₃)
  • Borax (Na₂B₄O₇·10H₂O)
  • Kernite (Na₂B₄O₇·4H₂O)
• Borax is located in areas of;
  • Puga Valley (Ladakh)
  • Sambhar Lake (Rajasthan)
• Boron has a mass abundance of less than 0.0001% in earths crust
• Boron comes in two isotopic structures; ¹⁰B (19%) and ¹¹B (81%)

Aluminium

• Aluminium is the most abundant metal, comprising 8.3% of Earth's crust.
• Most of the Earth crust is comprised of Oxygen (45.5%) and silicone (27.7%)
• Most important minerals of aluminum;
  • Bauxite (Al₂O₃·2H₂O)
  • Cryolite (Na₃AlF₆)
• India contains aluminum located in
  • Madhya Pradesh
  • Karnataka
  • Orissa
  • Jammu
• Gallium, indium, and thallium are less abundant

Nihonium

• Nihonium(Nh) has:
  • Symbol Nh
  • Atomic number 113
  • Atomic mass 286 g mol⁻¹
  • Electronic configuration [Rn] 5f¹⁴ 6d¹⁰ 7s² 7p²
• Has been prepared in a small amount
• The most stable isotope has a half-life of 20 seconds
• Chemistry is not well-established due to these reasons
• Nihonium is a synthetic radioactive element
• Atomic, physical, and chemical properties of group elements other than nihonium are discussed below

Electronic Configuration

• The outer electronic configuration is ns²np¹
• Boron and aluminum have noble gas cores
• Gallium/Indium have noble gas plus 10 d-electrons
• Thallium has noble gas + 14 f-electrons + 10 d-electrons cores, thus the electronic structures of these elements are more complex than for the first two groups of elements
• Electronic structure difference affects other properties and chemistry

Atomic Radius

• Atomic radius is expected to increase down the group as each success member adds an extra shell of electrons
• Gallium has a lower atomic radius than aluminium due to variance of inner core configuration
• Additional 10 d-electrons have a poor screening effect, that does not compensate for the increased nuclear charge in gallium
• Gallium's atomic radius is 135 pm, less than aluminium's 143 pm.

Ionization Enthalpy

• Due to general trends, ionisation enthalpy values do not decrease smoothly down
• The decrease from B to Al is associated with increase in size.
• Discontinuity in values observed between Al and Ga, and In and Tl
• Inability of d- and f-electrons, with their low screening effect, is to blame
• Increase in nuclear charge cannot be compensated for in this case
• The order of ionisation enthalpies is (as expected) ∆iH₁