Periodic Properties And P-Block Elements Notes for NEET by PW

Periodic Properties and P-Block Elements can be challenging when several trends have exceptions or depend on more than one factor. For example, atomic size generally decreases across a period, but ionisation energy does not increase uniformly. Similarly, the stability of oxidation states changes down the p-block groups because of the inert-pair effect.

For NEET, focus on connecting each periodic trend with electronic configuration, nuclear charge and the shielding effect. PW supports Class 11, Class 12 and dropper aspirants through live and recorded classes, handwritten notes, Daily Practice Problems (DPPs), mock tests, previous years' questions (PYQs) and regular assessments. These resources help you revise NCERT concepts, practise trend-based questions and identify exceptions that need more attention.

Modern Periodic Table

The modern periodic table arranges elements in increasing order of atomic number, which determines the number of protons in an atom's nucleus. This arrangement reflects the electronic configurations of elements and helps explain similarities in their chemical properties.

Key points to remember:

• Atomic number equals the number of protons in an atom.

• Elements are arranged in increasing atomic-number order.

• The periodic table has seven periods and eighteen groups.

• The Aufbau principle and the (n+l) rule help explain the general order of orbital filling.

Periods And Groups

The table has seven periods and eighteen groups.

Elements in the same group generally show similar properties because they have similar outer electronic configurations.

Blocks

The block is decided by the subshell receiving the last electron:

• s-block: Groups 1 and 2

• p-block: Groups 13 to 18

• d-block: Groups 3 to 12

• f-block: Inner transition elements

Hydrogen is treated separately because of its unique properties. Helium is placed in the p-block by position, although its configuration is 1s2.

Periodic Trends

Periodic trends arise from changes in nuclear charge, the number of occupied shells, shielding and electronic configuration. Understanding these factors helps explain why atomic size, ionisation energy and electronegativity change across periods and down groups.

Screening Effect And Effective Nuclear Charge

Inner-shell electrons repel outer electrons and reduce the attraction of the nucleus. This is called the screening effect.

The shielding ability of subshells follows: s > p > d > f

The effective nuclear charge is the net nuclear attraction experienced by an electron:

Effective nuclear charge = nuclear charge − screening effect

When effective nuclear charge increases, attraction increases and atomic size decreases.

Poor shielding by d and f electrons produces contraction:

• Transition or scandide contraction: poor shielding by d-electrons

• Lanthanoid contraction: poor shielding by 4f-electrons

• Actinoid contraction: poor shielding by 5f-electrons

Atomic And Ionic Radius

Across a period, atomic radius decreases because effective nuclear charge increases. Down a group, atomic radius increases because new shells are added.

For the same element:

Cation < neutral atom < anion

For isoelectronic species, all species have the same number of electrons. The species with more protons is smaller because it has greater nuclear attraction.

Ionisation Energy

Ionisation energy is the minimum energy required to remove an electron from an isolated gaseous atom in its ground state. It generally increases across a period and decreases down a group, although electronic configuration creates important exceptions.

Key points:

• Across a period, increasing effective nuclear charge generally makes electron removal more difficult.

• Down a group, additional shells and shielding generally make the outermost electron easier to remove.

• Helium has the highest first ionisation energy.

• Caesium has a very low first ionisation energy among the naturally occurring elements commonly discussed in introductory periodic trends.

For second-period elements, two important exceptions are:

• Be and B: Boron has a lower first ionisation energy than beryllium because the electron removed from boron occupies a higher-energy 2p subshell.

• N and O: Oxygen has a lower first ionisation energy than nitrogen because electron pairing in an oxygen 2p orbital introduces additional repulsion, whereas nitrogen has a relatively stable half-filled 2p subshell.

Important exceptions in the second period are: B < Be < C < O < N < F < Ne

The Be-B exception is caused by greater penetration of the 2s electron in beryllium. The N-O exception is caused by the extra stability of nitrogen’s half-filled p3 configuration.

Successive ionization energies always increase. A sudden large jump shows that all valence electrons have been removed. The position of the jump gives the number of valence electrons.

Electron-Gain Enthalpy

Electron-gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom. A more negative value indicates a more exothermic process.

Important trends and exceptions include:

• Halogens generally have highly negative first electron-gain enthalpies because gaining one electron completes their valence shell.

• Chlorine has a more negative first electron-gain enthalpy than fluorine. Fluorine's compact 2p orbitals create greater repulsion for the incoming electron.

• Noble gases generally have positive first electron-gain enthalpies because the added electron must enter a higher-energy shell.

• The second electron-gain enthalpy of an element is positive because an electron is added to an already negatively charged ion, requiring energy to overcome electrostatic repulsion.

Electronegativity

Electronegativity describes the tendency of an atom in a chemical bond to attract the shared electron pair towards itself. Unlike ionisation energy, it is a relative property of bonded atoms rather than the energy required to remove an electron from an isolated atom.

The general trends are:

• Across a period, electronegativity generally increases.

• Down a group, electronegativity generally decreases.

• Fluorine is the most electronegative element.

• Higher electronegativity generally corresponds to stronger non-metallic character.

For carbon hybrid orbitals: sp > sp2 > sp3

Higher oxidation state also generally increases electronegativity.

Acidic And Basic Character Of Oxides

The acid-base character of oxides is linked to the element's position in the periodic table and its oxidation state. Across a period, oxides generally become more acidic as metallic character decreases. Down a group, basic character often becomes more prominent as metallic character increases.

Examples include:

• N₂O₅: Acidic oxide.

• CaO: Basic oxide.

• B₂O₃: Acidic oxide.

• Al₂O₃ and Ga₂O₃: Amphoteric oxides.

• In₂O₃ and Tl₂O₃: Predominantly basic oxides in standard introductory classifications.

Amphoteric oxides react with both acids and bases. Common examples include aluminium oxide and zinc oxide.

Neutral oxides commonly listed at this level include NO, N₂O and CO. Water is also often classified as a neutral oxide in introductory acid-base discussions.

The oxidation state matters: oxides of the same element can show different acid-base behaviour at different oxidation states.

Important P-Block Concepts

The p-block contains Groups 13 to 18. Its elements display a wide range of bonding patterns, oxidation states and chemical behaviour because their valence electrons occupy p-subshells, with the ns² pair also influencing many heavier elements.

Inert-Pair Effect

The inert-pair effect is the tendency of the valence-shell (ns^2) electron pair in heavier p-block elements to participate less readily in bonding.

As a result, oxidation states two units lower than the group oxidation state often become more stable down a group.

Examples include:

• Tl⁺ is more stable than Tl³⁺ in many common compounds.

• Pb²⁺ is more stable than Pb⁴⁺ in many common compounds.

• Bi³⁺ is more stable than Bi⁵⁺ in many common compounds.

Lead(IV) compounds can act as oxidising agents because Pb⁴⁺ may be reduced to the more stable Pb²⁺ state. Lead tetraiodide, PbI₄, is unstable because iodide can reduce Pb⁴⁺ to Pb²⁺ while itself being oxidised to iodine.

Anomalous Behaviour of Second-Period Elements

Second-period elements differ from their heavier group members because they have small atomic sizes, relatively high ionisation energies, high electronegativities and no available valence-shell d-orbitals.

These features explain several important differences:

• Carbon forms strong C–C bonds and shows extensive catenation.

• Nitrogen has a maximum covalency of four in standard introductory examples and does not form ordinary stable pentahalides analogous to those of phosphorus.

• The N–N single bond is relatively weak because of strong repulsion between lone pairs on adjacent nitrogen atoms.

• Nitrogen, oxygen and fluorine can form hydrogen bonds when bonded to hydrogen, unlike their heavier group members under comparable conditions.

• Second-period elements can form effective pπ–pπ bonds, as seen in molecules containing multiple bonds.

These differences are frequently tested through comparisons between the first and later members of a group.

Group-Wise P-Block Revision

The p-block groups have characteristic valence-shell configurations and recurring trends. Use the following points for revision, but pair each trend with the relevant NCERT examples and exceptions.

Group 13: Boron Family

General valence-shell configuration: (ns2 np1)

• The common oxidation states are +3 and +1; the +1 state becomes more stable down the group due to the inert-pair effect.

• Boron compounds are predominantly covalent because removing three electrons to form B³⁺ requires substantial energy.

• Boron does not form conventional six-coordinate covalent compounds by expanding its octet using vacant 2d orbitals; such orbitals are not available in the second shell.

• B₂O₃ is acidic, while Al₂O₃ and Ga₂O₃ are amphoteric.

• BF₃, BCl₃, BBr₃ and BI₃ are electron-deficient Lewis acids, but their relative Lewis acidity is influenced by back bonding.

Back-bonding order: BF3 > BCl3 > BBr3 > BI3

Lewis acidity order: BF3 < BCl3 < BBr3 < BI3

Group 14: Carbon Family

General valence-shell configuration: (ns2 np2)

• The common oxidation states are +4 and +2.

• The +2 oxidation state becomes more stable down the group due to the inert-pair effect.

• Carbon and silicon form acidic oxides, while tin and lead form oxides that can show amphoteric behaviour.

• Catenation generally decreases down the group because the strength of element-element bonds decreases.

• Pb²⁺ is more stable than Pb⁴⁺ in many common compounds.

• Carbon-14 is radioactive and is used in radiocarbon dating.

Group 15: Nitrogen Family

General valence-shell configuration: (ns2 np3)

Common oxidation states include −3, +3 and +5.

• The stability of the +5 oxidation state generally decreases down the group.

• The +3 oxidation state becomes more stable down the group due to the inert-pair effect.

• Nitrogen has a maximum covalency of four in standard introductory chemistry examples.

• Nitrogen forms several oxides with oxidation states ranging from +1 to +5.

• Some compounds containing nitrogen in intermediate oxidation states can undergo disproportionation.

• NF₃ is more stable than several other nitrogen trihalides because the N–F bond is strong.

Group 16: Oxygen Family

General valence-shell configuration: (ns2 np4)

• Common positive oxidation states for heavier members include +4 and +6.

• The stability of the +6 oxidation state generally decreases down the group, while the +4 state becomes more stable.

• Oxygen displays anomalous behaviour because of its small size and high electronegativity.

• Sulfur shows significant catenation, although carbon forms especially strong chains.

• The bond strength of hydrogen chalcogenides generally decreases down the group, so their thermal stability also decreases.

Group 17: Halogens

General valence-shell configuration: (ns2 np5)

Halogens readily gain electrons and commonly act as oxidising agents. Their oxidising power in aqueous solution follows: F2 > Cl2 > Br2 > I2

Fluorine is the strongest oxidising agent in this series. Its behaviour is influenced by its high reduction potential, strong hydration of fluoride ions and the relatively low F–F bond dissociation enthalpy compared with what its small size might otherwise suggest.

Hydrogen halides show the following general trends down the group:

• H–X bond strength decreases.

• Thermal stability decreases.

• Acidic strength in water increases.

• Reducing character generally increases.

• HI has the longest H–X bond among HF, HCl, HBr and HI.

Chlorine and bromine can disproportionate in water. Fluorine oxidises water, while iodine reacts only to a very limited extent under ordinary conditions.

Group 18: Noble Gases

General valence-shell configuration: (ns2 np6), except helium, which has (1s^2).

Noble gases have filled valence shells and are generally much less reactive than other elements. They are monoatomic gases under ordinary conditions and interact mainly through weak London dispersion forces.

Down the group:

• Atomic size increases.

• Ionisation energy generally decreases.

• Density generally increases.

• Melting and boiling points generally increase.

• Solubility in water generally increases.

• First electron-gain enthalpy is generally positive.

Xenon forms several compounds with highly electronegative elements such as fluorine and oxygen. The discovery of xenon compounds followed Neil Bartlett's preparation of a xenon-containing compound in 1962, challenging the earlier assumption that noble gases could not form compounds.

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