p Block Elements class 12 Notes

p block elements are present on the right side of the Periodic table. As a result, the periodic table has six groups of p–block elements numbered 13 to 18.

The families falling into this block belong to boron, carbon, nitrogen, oxygen, fluorine, and helium. Their electronic valence shell configuration is ns2np1-6 (except for He). However, the inner core of the electrical arrangement may differ.

The final electron in a P block element can enter any of the three p-orbitals of its corresponding shell.
Because a p-subshell has three degenerate p-orbitals, each of which may take two electrons, there are six groups of p-block elements in total.
Because of their susceptibility to losing an electron, P block elements are bright and typically excellent conductors of electricity and heat.
Elements with extraordinary qualities can be found in a P-block element like gallium.
It’s a metal that you can melt in your hand.
Silicon is also an important metalloid of the p-block group since it is a component of glass.

General Characteristics of p block elements

ns2np1-6 is the typical electrical configuration of p-block components (except He). The inner core electrical configuration, on the other hand, may vary. Because of this variation in the inner core, the physical and chemical characteristics of the elements vary.
When the entire number of valence electrons, i.e. the sum of S and P electrons, is equal to a total number of valence electrons, the oxidation state of elements in the p – block is maximal.
One of the most intriguing aspects of the p-block elements is the presence of both non-metals and metalloids.

The first member of the p block components varies from the others in two ways:

i)The first is the size, followed by every attribute that is dependent on the size.
ii)The second distinction only pertains to the p-block element that occurs as a result of changes in the d orbitals present in the outer shelf heavy elements.

Nitrogen Family Group 15 elements

Group 15 includes nitrogen, phosphorus, arsenic, antimony, and bismuth. As we advance along with the group, there occurs a transition from nonmetallic to metallic via the metalloid characteristic.
Nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, and bismuth is a typical metal.
The electrical configuration of the valence shells of these elements is ns2np3. In these elements, the s orbital is filled, but the p orbital is only half-filled, resulting in an extraordinarily stable electronic structure.
Because of the additional stable half-filled p-orbital electronic configuration and smaller size, the ionization enthalpy of group 15 elements is significantly larger than that of group 14 elements in similar periods. The sequence of sequential ionization enthalpies is, as Δ1H1<Δ1H2<Δ1H3.
Physical Properties: The elements in this category are all polyatomic. Except for dinitrogen, which is a diatomic gas, all other elements are solids. As the group evolves, its metallic quality becomes more prominent. Bismuth is classified as a metalloid, whereas nitrogen and phosphorus are classified as nonmetals. Bismuth is a metalloid, while arsenic and antimony are metals. This is due to a decrease in ionization enthalpy and an increase in atomic size. In general, boiling points in the group climb from top to bottom, but melting points rise until arsenic and then decline until bismuth. All elements, except nitrogen, display allotropy.
Chemical Properties: The most frequent oxidation states for these elements are – 3, + 3, and + 5. The possibility of exhibiting a -3 oxidation state decreases as one moves down the group, and bismuth seldom creates any –3 oxidation state compounds. As you advance through the group, the stability of the + 5 oxidation state decreases. BiF5 Bi F 5 is the only Bi Bi (V) molecule that has received extensive research.
Almost all intermediate phosphorus oxidation levels disproportionate into +5 and –3 in alkali and acid. In the case of arsenic, antimony, and bismuth, however, the +3-oxidation state becomes progressively stable in terms of disproportionation.

p Block elements class 12 notes stretch a bit more on Nitrogen Family in detail. Let us know about the behavior of this element.

Anomalous behavior of Nitrogen Nitrogen has the unique capacity to form p–p multiple bonds with itself and other small, electronegativity-rich elements (C, O). Heavy elements in this group do not form p – p bonds because their atomic orbitals are too wide and dispersed to overlap efficiently. As a result, nitrogen is a diatomic molecule with a triple link between its two atoms (one s and two p). As a result, the bond enthalpy (941.1 KJ mol-1 ) is condition.
A single N – N link is weaker than a single P – P connection. Due to this, the catenation property tends to tube weaker in nitrogen. Another factor contributes to the absence of d orbitals in its exterior shell. Nitrogen, unlike heavier elements, cannot form d–p bonds, e.g., R3P = 0 R 3 P = 0, because its covalency is restricted to four.
Reaction with Hydrogen The Nitrogen family members make halides in the form of EH₃ where E is the nitrogen family. From NH3 to BiH3, the bond dissociation enthalpy of hydrides falls, suggesting that their stability diminishes. As a result, the reducing character of the hydrides improves. Although ammonia is a poor reducing agent, BiH3 is the most effective hydride reducing agent. The order of basicity is also decreasing. NH3 > PH3 > AsH3 > BiH3
Reaction with Oxygen These elements produce two types of oxides: E2O3 and E2O5. The oxide in the higher oxidation state of the element is more acidic than the oxide in the lower oxidation state. As they pass through the group, their acidity decreases. Nitrogen and phosphorus oxides are completely acidic, whereas arsenic and antimony oxides are amphoteric and bismuth oxides are mostly basic.
Reactivity to Halogens When these elements combine, they form two halide series: EH3and EX5. It does not form pentahalide because the d orbitals in nitrogen’s valence shell are unavailable. Pentahalides have stronger covalent bonds than trihalides. Except for the trihalides of nitrogen, all of these elements’ trihalides are stable.
Metallic Reactivity These elements react with metals to form binary compounds with an oxidation state of –3, such as Ca3N2N (calcium nitride), Ca3P2 (calcium phosphide), Na3As2 sodium arsenide), Zn3Sb2 (zinc antimonide), and Mg3Bi2 (magnesium bismuthide).

Oxygen Family Group 16 Elements

Group 16 of the Periodic Table contains the elements oxygen (8O), sulfur (16S), selenium (34Se), tellurium (52Te), and polonium (84Po). 10. The elements’ valence shells have the electrical structure ns2np4. Because of its tiny size and strong electronegativity, the initial element of group 16 has a different chemical behavior than the other members of the group.

As the atomic number grows, so does the metallic nature. The first four elements have a non-metallic nature. Nonmetallic character is strongest in O and S, weaker in Se and Te, and metallic in Po.
In the relevant periods, elements in group 16 have lower ionization enthalpy values than elements in group 15. This is because group 15 elements have extremely stable half-filled p-orbital electronic structures.
Because of the compact structure of the oxygen atom, oxygen has a lower negative electron gain enthalpy than sulfur.
After F, O has the greatest electronegativity value among the elements. Within the group, electronegativity falls as atomic number increases.
As we move down the group, the likelihood of catenation reduces.
As the number of shells rises, the atomic and ionic radii increase from top to bottom.
As the group size grows, the ionization enthalpy reduces.
As we move down the group, the likelihood of catenation reduces.
The group’s components all form volatile hydrides.
From water to hydrogen sulfide, the volatility rises and subsequently falls. Their boiling point demonstrates this. H2S <H2Se<H2Te <H20 is the hydride with the highest rank of boiling points. The group boiling point decreases as molecular weight increases, increasing the van der Waal’s forces of 28. interaction. Because of hydrogen bonding, H20 has an extremely high b.p.
Hydride thermal stability declines in the following order: H20>H2S>H2Se>H2Te>H2Po.
The acidity of the hydrides increases in the following order: H20 forms a variety of oxo-acids. Sulphuric acid is the most significant of S’s oxo-acids.
The thermal stability of hydrides diminishes in the following order: H20>H2S>H2Se>H2Te>H2Po.
The acidity of the hydrides increases in the following order: H20<H₂S<H₂Se<H₂Te.

Halogen Family Group 17 elements

The Periodic Table Group 17 contains fluorine (9F), chlorine (17Cl), bromine (35Br), iodine (53I), and astatine (55At). For valence shells, the electronic arrangement is ns2 np5.

Because of maximal effective nuclear change, halogens have the shortest atomic radii in their respective periods.
The early ionization energies are quite high but diminish as one progresses through the group. Iodine can shed one electron and produce the I+ ion.
Electron affinity increases as Cl > F > Br > I.
Reactivity increases with F2 > Cl2 > Br2 > I2
The ionic character order in the M – X bond is M-F>M-Cl>M-Br>M-I.
Hydrohalic strength Acids are classified as follows: HF HOBr > HOI
Acid strength increases as the number of O-atoms for a given halogen atom increases. HOCl < HClO₂ < HClO₃ < HClO₄
The most electronegative element known is F. Electronegativity declines as one moves along the group. Halogens are excellent oxidising agfents. The oxidizing power reduces as one moves along the group.

Noble Gas Family Group 18 Elements

Noble Elements contain Helium (2He), neon (10Ne), argon (18Ar), krypton (36Kr), xenon (54Xe), and radon (86Rn) are elements found in Periodic Group 18. They are referred to collectively as noble gases.

Noble gases are found after each cycle. Their valence shell orbitals are filled.
They are monoatomic and just slightly soluble in water.
Xe produces the fluorides XeF2, XeF4, and XeF6.
Xe03 has a trigonal pyramidal structure, whereas XeOF4 has a square pyramidal shape.
Applications:
Because He is a non-flammable gas that is lighter than air, He is utilized to fill balloons for meteorological measurements.
Ar is utilized to create an inert environment.
Their positive electron gain enthalpy is greater than their ionization enthalpy.
Noble gases’ reactivity has been explored on a few occasions since their discovery, but all attempts to induce them to react to form compounds have failed for many years. In March 1962, Neil Bartlett of the University of British Columbia discovered a noble gas reaction.
Helium is compressed and liquefied to form He() liquid at 4.2 K. This liquid, like any other, is a standard liquid. However, if it is cooled further, He() at 2.2 K is formed, which is known as a superfluid since it is a liquid with gas characteristics. It bursts through the walls of the container and emerges. It is also low in viscosity.

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p Block Elements class 12 Notes – Chemistry Chapter 7 ncert