Periodic Classification of Elements

Table of Contents
1. Need for the Periodic Classification of Elements
2. Periodic Table
3. Classification of Elements as Metals, Non‐metals & Metalloids
4. Classification of Elements on the Basis of Electronic Configuration
5. IUPAC Nomenclature for Elements with (Z > 100)
6. Effective Nuclear Charge (Z*)
7. Empirical Rules for Estimating Effective Nuclear Charge: Slater's Rules
8. Periodic Trends and Their Explanations
9. Applications and Importance of the Periodic Classification
10. References for Further Reading
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Need for the Periodic Classification of Elements

All existing matter is made of basic units called elements. Early in the 19th century, only a few elements were known; by 1800 about 31 elements had been discovered. As the number of known elements increased (now 118), it became impractical to study each element individually. Scientists therefore sought an organised scheme to group elements so that their properties could be studied systematically and predicted for undiscovered or less-known elements. This led to the development of the periodic classification of elements, commonly presented as a periodic table.

• When elements are arranged according to certain regularities in their properties, similar elements appear in the same column and periodic trends become apparent.
• A well-constructed periodic table simplifies the study of physical and chemical behaviour and enables prediction of properties for new elements or compounds.

Periodic Table

The periodic table is the systematic arrangement of elements in a tabular form so that elements with similar properties fall in the same vertical column (called a group) and elements with similar electronic energy levels lie in the same horizontal row (called a period).

Modern Periodic Law

The modern periodic law states: The physical and chemical properties of the elements are periodic functions of their atomic numbers.

• If elements are arranged in increasing order of atomic numbers, there is a repetition of properties at regular intervals (periodicity).
• The pattern of repetition corresponds to completion of electron shells; commonly noted periodicities are 2, 8, 18 and 32 electrons for shell filling - these are sometimes referred to as magic numbers.

Long Form of the Periodic Table

The long form of the periodic table organises elements primarily according to their electronic configuration so that the table aligns electronic structure with physical and chemical properties.

• Elements are arranged in increasing order of atomic number (Z).
• Horizontal rows are called periods and vertical columns are called groups or families.
• The periodic table is divided into four main blocks: s, p, d and f. This block division is based on the differentiating (last) electron's subshell.

Electronic Configuration According to Different Blocks

• Representative elements: Most s and p block elements (groups 1, 2 and 13-18) are called representative or main‐group elements (except the zero group of noble gases when context applies).
• Typical elements: Elements of the second and third periods are often called typical elements because their properties closely follow periodic trends.

Classification of Elements as Metals, Non‐metals & Metalloids

1. Metals

• About 80% of the known elements are metals.
• Metals are generally lustrous, malleable, ductile, good conductors of heat and electricity, and have relatively high densities.
• Most metals have high melting and boiling points and are solid at room temperature.
• Mercury (Hg) is the only common metal that is liquid at room temperature; gallium (Ga) has a low melting point close to 303 K and can melt in the hand.

Location of Metals, Non-metals and Metalloids in Periodic Table

2. Non‐Metals

• Non‐metals are fewer in number compared to metals.
• They usually have low melting and boiling points and may be solids, liquids or gases at room temperature.
• Non‐metals are neither malleable nor ductile and are generally poor conductors of heat and electricity.
• The non‐metallic character increases across a period from left to right; the metallic character increases down a group.

3. Metalloids (Semi‐metals)

• There is no abrupt dividing line between metals and non‐metals. A zig‐zag borderline on the periodic table separates metallic and non‐metallic regions.
• Elements on the borderline (for example silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te)) show intermediate properties and are called metalloids or semi‐metals.

Classification of Elements on the Basis of Electronic Configuration

Elements can be grouped according to which subshell receives the last electron (the differentiating electron). This leads to the s, p, d and f block classification.

1. s‐block Elements

• The last electron enters an s subshell.
• They occupy the leftmost portion of the periodic table and include group 1 (alkali metals) and group 2 (alkaline earth metals).
• All s‐block elements are metals.
• General valence electronic configuration: ns1-2 (n = 1 to 7).

2. p‐block Elements

• The last electron enters a p subshell.
• They are located on the right side of the periodic table and comprise groups 13-18.
• Most p‐block elements include non‐metals and metalloids; some are metals.
• General valence electronic configuration: ns2 np1-6 (n = 2 to 7).
• A filled ns2 np6 configuration corresponds to a noble gas (stable) configuration. For example, He has configuration 1s2.

3. d‐block Elements (Transition Elements)

• The last electron enters a d subshell (the differentiating electron enters an (n-1)d orbital).
• They lie in the central block of the periodic table (groups 3-12).
• All d‐block elements are metals and often show variable oxidation states and formation of coloured compounds.
• Outermost electronic configuration: (n-1)d1-10 ns1-2 (n = 4 to 7).
• Major series of d‐block elements:
  3d series - Sc (21) to Zn (30)
  4d series - Y (39) to Cd (48)
  5d series - La (57), Hf (72) to Hg (80)

4. f‐block Elements (Inner Transition Elements)

• These elements are also called inner transition elements and the differentiating electron enters an (n - 2)f orbital.
• All f‐block elements are metals and are placed separately below the main table as two rows.
• Lanthanides: Differentiating electron enters the 4f orbital; these 14 elements run from cerium (Ce) to lutetium (Lu) and are so named because they follow lanthanum (La).
• Actinides: Differentiating electron enters the 5f orbital; these 14 elements run from thorium (Th) to lawrencium (Lr) and follow actinium (Ac).
• General electronic configuration:
  Lanthanides: [Xe] 4f1-14 5d0-1 6s2
  Actinides: [Rn] 5f1-14 6d0-1 7s2

IUPAC Nomenclature for Elements with (Z > 100)

• Elements beyond uranium (Z = 92) are synthetic and are called transuranium elements. Elements beyond fermium (Z = 100) are often called transfermium elements (Z ≥ 101).
• Some early transfermium elements such as fermium (Z = 100), mendelevium (Z = 101), nobelium (Z = 102) and lawrencium (Z = 103) were named after eminent scientists.
• Nomenclature (IUPAC temporary systematic names): For an element with atomic number represented by three digits abc, the name is built by using Latin/Greek numerical roots for each digit followed by the suffix "‐ium". The roots for digits 0-9 are standardised (shown in the table represented in the image below).

• Example: Atomic number 109 → digits 1, 0, 9 correspond to roots une‐nil‐ennium; temporary name given in simplified form as ununennium (systematic temporary name conventions are applied until a permanent name is approved).

Table: Names of elements with atomic number above 100 are summarised in the following reference image.

Effective Nuclear Charge (Z*)

Effective nuclear charge (Z*) is the net positive charge experienced by an electron in a multi‐electron atom after accounting for the screening (shielding) effect of other electrons. In simple terms, it measures the attractive force exerted on a given electron by the nucleus, reduced by repulsion from inner electrons.

• Electrons in inner shells shield outer electrons from the full nuclear charge; as shielding increases, the effective nuclear charge experienced by valence electrons decreases.
• Within the same principal quantum number n, subshell penetration influences Z*: electrons in subshells that penetrate closer to the nucleus (for example an s electron) experience a higher Z* than those in subshells that penetrate less (for example a p or d electron).
• Typical orbital energy ordering based on penetration is: 2s < 2p, 3s < 3p < 3d, and 4s < 4p < 4d < 4f. As atomic number increases, energy separations between orbitals of the same n value decrease.

Screening (Shielding) Effect

The screening effect (or shielding) arises because inner shell electrons repel outer electrons, reducing the full attractive force of the nucleus felt by the outer electrons.

• Outer electrons are therefore screened from the full nuclear charge and experience a reduced effective nuclear charge Z*.
• An increase in the number of inner electrons increases shielding and thus reduces Z*, which typically lowers ionisation energy for outer electrons.

Penetrating Power of Electrons

The ability of an electron in a given subshell to approach the nucleus is called penetration. Penetration decreases in the order: s > p > d > f.

• Greater penetration means the electron is held more tightly by the nucleus and is harder to remove; therefore ionisation energy correlates with penetrating power of the subshell.
• Penetration helps explain subtle anomalies in orbital energy order and chemical behaviour (for example why 4s is filled before 3d in neutral atoms).

Empirical Rules for Estimating Effective Nuclear Charge: Slater's Rules

Slater's rules provide an approximate method to estimate the shielding constant S and hence the effective nuclear charge Z* experienced by an electron. The relation is:

Z* = Z - S

• Write the electronic configuration grouped as: (1s), (2s 2p), (3s 3p) (3d), (4s 4p), (4d) (4f), (5s 5p), and so on.
• Electrons in groups to the right of the electron under consideration contribute nothing to S.
• Rules for an electron in an ns or np orbital:
  (a) Each other electron in the same (ns, np) group contributes 0.35 to S.
  (b) Each electron in the (n-1) shell contributes 0.85 to S.
  (c) Each electron in the (n-2) or lower shells contributes 1.00 to S.
• For an electron in an nd or nf orbital:
  Each other electron in the same (nd) or (nf) group contributes 0.35 to S.
  All electrons in groups lower than (nd) or (nf) contribute 1.00 each.

Example (Slater's Rules): Consider the potassium atom. Electronic configuration in the grouped form is (1s2)(2s2 2p6)(3s2 3p6)(4s1).

The goal is to estimate the effective nuclear charge Z* experienced by the 4s electron.

Z = 19 (nuclear charge)

Determine the shielding constant S using Slater's rules:

Electrons in the same group (4s) other than the electron under consideration: none → contribution = 0 × 0.35 = 0

Electrons in the (n - 1) shell (n = 4 so (n - 1) = 3): there are 8 electrons (3s2 3p6) → contribution = 8 × 0.85 = 6.80

Electrons in (n - 2) or lower shells (1s2 and 2s22p6): total 10 electrons → contribution = 10 × 1.00 = 10.00

Therefore, S = 0 + 6.80 + 10.00 = 16.80

Z* = Z - S = 19 - 16.80 = 2.20

Thus the 4s electron experiences an effective nuclear charge of approximately 2.20, which helps explain why 4s is occupied in the ground state of potassium.

Periodic Trends and Their Explanations

Several important properties show regular trends across periods and down groups. These trends follow from changes in atomic size, effective nuclear charge and electron shielding.

Atomic Radius

• Across a period (left → right): Atomic radius generally decreases due to increasing Z* that pulls electrons closer to the nucleus.
• Down a group: Atomic radius increases because electrons occupy higher principal shells (larger n) even though Z increases; increased shielding dominates.

Ionic Radius

• Cations (< neutral atom) are smaller due to loss of outer electron(s) and increased effective nuclear attraction on remaining electrons.
• Anions (> neutral atom) are larger because added electron(s) increase electron-electron repulsion and reduce Z* per electron.

Ionisation Energy (I.E.)

• First ionisation energy is the energy required to remove the most loosely bound electron from a neutral atom in the gas phase.
• Across a period: I.E. generally increases due to increasing Z* and decreasing atomic radius.
• Down a group: I.E. decreases because outer electrons are farther from the nucleus and experience greater shielding.

Electron Affinity and Electronegativity

• Electron affinity is the energy change when an atom in the gas phase accepts an electron; trends are more complex but generally become more negative (more favourable) across a period.
• Electronegativity is a measure of an atom's tendency to attract electrons in a chemical bond; it generally increases across a period and decreases down a group.

Metallic Character and Valency

• Metallic character increases down a group and decreases across a period.
• Valency (combining capacity) usually follows the electronic configuration: valency tends to increase across a period up to the half‐filled subshell and then decrease; representative elements show valencies predictable from their group number.

Applications and Importance of the Periodic Classification

• Predicting chemical reactivity and bonding preferences of elements and their compounds.
• Understanding trends in physical properties such as atomic/ionic sizes, melting and boiling points, densities.
• Designing materials and alloys by selecting elements with desirable electronic and structural properties (important in engineering disciplines such as civil, electrical and materials engineering).
• Guiding the synthesis of new elements and isotopes and predicting their likely chemical behaviour.

References for Further Reading

• Standard school chemistry textbooks that follow the modern periodic law and long‐form periodic table (NCERT/CBSE style content).
• Advanced inorganic chemistry texts for detailed treatments of electronic structure, Slater rules and inner transition elements.