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By Shailendra Singh
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Updated on 11 Nov 2025, 18:21 IST
The periodic classification of elements is one of the most fundamental organizing principles in chemistry. Before the 18th century, when only about 30 elements were known, studying and remembering their properties was relatively simple. However, as more elements were discovered, the need arose for a systematic method to organize them based on their properties.
The periodic table arranges all known elements according to their properties in such a way that elements with similar properties are grouped together in a tabular form. This arrangement has evolved through several stages, from early attempts like Dobereiner's triads and Newlands' law of octaves to Mendeleev's periodic table, and finally to the modern periodic table we use today.
German chemist Johann Wolfgang Dobereiner arranged elements in groups of three (triads) where the atomic mass of the middle element was roughly the average of the other two elements.
Example:
Limitations:
J.A.R. Newlands arranged elements in order of increasing atomic mass and observed that every eighth element had properties similar to the first, like musical octaves.
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Limitations:
Dmitri Ivanovich Mendeleev arranged 63 known elements in increasing order of atomic mass and grouped them based on chemical properties.
Achievements:
Drawbacks:
In 1913, Henry Moseley discovered that properties of elements depend on their atomic number rather than atomic mass. This led to the Modern Periodic Law:
"The physical and chemical properties of elements are periodic functions of their atomic numbers."
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The modern periodic law recognizes that chemical properties depend on the number of valence electrons, which in turn depends on atomic number. When elements are arranged in increasing order of atomic numbers, elements with the same number of valence electrons appear at regular intervals, creating periodicity.
The modern periodic table, also called the long form or Bohr, Bury & Rang, Werner periodic table, consists of:
| Period | Principal Quantum Number (n) | Subshells Filled | Number of Elements | Range | Name |
| 1 | 1 | 1s | 2 | H to He | Shortest |
| 2 | 2 | 2s, 2p | 8 | Li to Ne | Short |
| 3 | 3 | 3s, 3p | 8 | Na to Ar | Short |
| 4 | 4 | 4s, 3d, 4p | 18 | K to Kr | Long |
| 5 | 5 | 5s, 4d, 5p | 18 | Rb to Xe | Long |
| 6 | 6 | 6s, 4f, 5d, 6p | 32 | Cs to Rn | Longest |
| 7 | 7 | 7s, 5f, 6d | 26 (incomplete) | Fr to Uub | Incomplete |
Period = Number of shells in the atom
Groups are numbered from 1 to 18 according to IUPAC recommendations.
Key Groups:
Group Number = Number of valence electrons (for main group elements)
Elements are classified into four blocks based on which subshell receives the last electron:
Electronic Configuration: ns¹⁻²
Characteristics:
General Configuration:
Electronic Configuration: ns² np¹⁻⁶
Characteristics:
Valence Electrons: Number of electrons in outermost s and p subshells
Group Number = 10 + number of valence electrons
Electronic Configuration: (n-1)d¹⁻¹⁰ ns⁰⁻²
Characteristics:
Group Number = Number of electrons in (n-1)d + number in ns
Note: Zn, Cd, Hg with d¹⁰ configuration are sometimes not considered true transition elements because d subshell is completely filled.
Electronic Configuration: (n-2)f¹⁻¹⁴ (n-1)d⁰⁻¹ ns²
Characteristics:
Properties:
| Aspect | Mendeleev's Periodic Table | Modern Periodic Table |
| Basis of Classification | Increasing atomic mass | Increasing atomic number |
| Periodic Law | Properties are periodic function of atomic mass | Properties are periodic function of atomic number |
| Number of Groups | 8 groups (divided into A and B subgroups) | 18 groups |
| Number of Periods | 6 periods | 7 periods |
| Position of Elements | Based on atomic mass and properties | Based on electronic configuration |
| Isotopes | No proper position (should be placed separately) | Same position (same atomic number) |
| Hydrogen | Uncertain position | Placed above alkali metals |
| Noble Gases | Not present initially (added later) | Group 18 (0 group) |
| Anomalous Pairs | Present (Ar-K, Co-Ni, Te-I) | Resolved by atomic number arrangement |
| Transition Elements | Placed in group VIII outside main table | Separate d-block (Groups 3-12) |
| Metals and Non-metals | No clear distinction | Clearly separated |
| Explanation of Periodicity | Not explained | Explained by electronic configuration |
| Predictions | Left gaps for undiscovered elements | No gaps (except in period 7) |
Periodic properties are characteristics that show regular variations when moving across a period or down a group.
Definition: Distance from the center of nucleus to the outermost shell containing electrons.
Types:
Trends:
Across a Period (Left to Right):
Down a Group (Top to Bottom):
Cationic Radius:
Anionic Radius:
Isoelectronic Species:
Definition: Minimum energy required to remove the most loosely bound electron from an isolated gaseous neutral atom.
Equation: M(g) + Energy → M⁺(g) + e⁻
Units: kJ/mol or eV
Successive Ionization Energies:
Trends:
Across a Period:
Down a Group:
Factors Affecting IE:
Definition: Energy released when an isolated gaseous atom accepts an electron to form a negative ion.
Equation: X(g) + e⁻ → X⁻(g) + Energy
Units: kJ/mol (usually reported as negative values)
Trends:
Across a Period:
Down a Group:
Special Cases:
Definition: Tendency of an atom to attract shared pair of electrons towards itself in a covalent bond.
Scales:
Trends:
Across a Period:
Down a Group:
Most Electronegative: Fluorine (4.0) Least Electronegative: Francium/Cesium (~0.7)
Metallic Character: Tendency to lose electrons and form positive ions.
Trends:
Non-metallic Character: Tendency to gain electrons and form negative ions.
Trends:
Definition: Combining capacity of an element.
Trends:
Across a Period:
Down a Group:
| Property | Across Period (→) | Down Group (↓) |
| Atomic Size | Decreases | Increases |
| Ionization Energy | Increases | Decreases |
| Electron Affinity | Increases (more negative) | Decreases (less negative) |
| Electronegativity | Increases | Decreases |
| Metallic Character | Decreases | Increases |
| Non-metallic Character | Increases | Decreases |
| Valency | 1→4→1 (with H), 1→7 (with O) | Constant |
| Chemical Reactivity (Metals) | Decreases | Increases |
| Chemical Reactivity (Non-metals) | Increases | Decreases |
Expected: [Ar] 3d⁴ 4s² Actual: [Ar] 3d⁵ 4s¹
Reason: Half-filled d subshell (d⁵) is more stable than d⁴
Expected: [Ar] 3d⁹ 4s² Actual: [Ar] 3d¹⁰ 4s¹
Reason: Fully-filled d subshell (d¹⁰) is more stable than d⁹
Expected: [Kr] 4d⁸ 5s² Actual: [Kr] 4d¹⁰ 5s⁰
Reason: Completely filled d¹⁰ configuration is exceptionally stable
| Property | Formula/Relationship | Explanation |
| Mendeleev's Triad | Atomic mass of middle element = (First + Third)/2 | Average of first and third element equals middle element |
| Modern Periodic Law | Properties = f(Atomic Number) | Properties are periodic function of atomic number |
| Group Number (p-block) | Group Number = 10 + valence electrons | For elements in p-block |
| Group Number (d-block) | Group Number = (n-1)d electrons + ns electrons | For transition elements |
| Period Number | Period = Number of shells (n) | Highest principal quantum number |
| Van der Waals Radius | r(VdW) = 2 × r(covalent) | Relationship between radii types |
| Atomic Size Trend | Size ∝ 1/Zeff | Inversely proportional to effective nuclear charge |
| Ionization Energy | M(g) → M⁺(g) + e⁻ ; ΔH = IE | Energy required to remove electron |
| Successive IE | IE₁ < IE₂ < IE₃ ... | Each successive ionization requires more energy |
| Electron Affinity | X(g) + e⁻ → X⁻(g) + Energy | Energy released when electron is added |
| Mulliken Electronegativity | χ = (IE + EA)/2 | Average of ionization energy and electron affinity |
| Ionic vs Atomic Radius | r(cation) < r(atom) < r(anion) | Relative size relationship |
| Isoelectronic Series | Size ∝ 1/Z | Inversely proportional to nuclear charge |
| Effective Nuclear Charge | Zeff = Z - S | Nuclear charge minus shielding |
The periodic classification of elements represents one of chemistry's greatest organizational achievements. The journey from Dobereiner's simple triads to the sophisticated modern periodic table demonstrates the evolution of scientific understanding.
The modern periodic law, based on atomic number rather than atomic mass, provides a rational basis for understanding the systematic variation of properties. The division into s, p, d, and f blocks based on electronic configuration offers deep insights into chemical behavior.
Understanding periodic trends how atomic size, ionization energy, electron affinity, and electronegativity vary enables prediction of element properties and reactivity. These trends are not arbitrary but arise logically from electronic structure and effective nuclear charge.
Mastery of the periodic table is essential for all further study in chemistry, as it serves as the fundamental framework for understanding chemical behavior, bonding, and reactions.
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The periodic table is a systematic arrangement of all known elements according to their properties, where elements with similar characteristics are grouped together in a tabular format. It was developed because as the number of discovered elements increased beyond 30 in the eighteenth century, scientists needed a simpler method to study and remember their properties. This classification system groups similar elements together while separating different ones, making it easier to predict element behavior and understand chemical relationships.
Dmitri Ivanovich Mendeleev created the most influential periodic table in 1869, arranging 63 known elements by increasing atomic mass and chemical property similarities. However, the modern periodic table is based on the work of Henry Moseley (1913), who demonstrated that atomic number not atomic mass is the fundamental property determining element characteristics. This led to the modern periodic law: "The physical and chemical properties of elements are periodic functions of their atomic numbers."
Dobereiner's Triads (1817) grouped elements in sets of three where the atomic mass of the middle element was roughly the average of the other two. For example, lithium (7), sodium (23), and potassium (39) form a triad where (7+39)/2 = 23. The limitation was that this classification only worked for a small number of elements only a few triads could be identified at the time, making it insufficient for organizing all known elements.
Newlands' Law of Octaves (1866) stated that when elements are arranged by increasing atomic mass, every eighth element has properties similar to the first, like musical octaves. This pattern worked well only up to calcium and with lighter elements. It failed because it couldn't accommodate elements discovered later, placed dissimilar elements together (like cobalt and nickel with halogens), and separated similar elements (like iron from cobalt and nickel).
The modern periodic table contains 18 vertical columns called groups and 7 horizontal rows called periods. Elements are arranged by increasing atomic number, with those having the same number of valence electrons placed in the same group. The table includes 118 confirmed elements and distinguishes between s-block, p-block, d-block (transition elements), and f-block (lanthanides and actinides) elements based on which subshell receives the last electron.
Periods are horizontal rows in the periodic table. Period 1 is the shortest with 2 elements (H, He). Periods 2 and 3 are short periods with 8 elements each. Periods 4 and 5 are long periods with 18 elements each. Period 6 is the longest with 32 elements (including lanthanides). Period 7 is incomplete with 26 elements discovered so far (including actinides). The number of elements in each period corresponds to the number of electrons that can occupy the subshells being filled.
Groups are vertical columns containing elements with similar properties because they have the same number of valence electrons. According to IUPAC recommendations, there are 18 groups numbered 1-18. Group 1 contains alkali metals, Group 2 has alkaline earth metals, Groups 3-12 are transition elements, Groups 13-17 contain representative elements, and Group 18 consists of noble gases. Elements in the same group typically show similar chemical behavior.
These classifications indicate which atomic subshell receives the last electron. S-block elements (Groups 1-2) have their outermost electron in an s orbital with general configuration ns¹⁻². P-block elements (Groups 13-18) fill p orbitals with configuration ns²np¹⁻⁶. D-block elements (Groups 3-12, transition metals) have configuration (n-1)d¹⁻¹⁰ns⁰⁻². F-block elements (lanthanides and actinides) have configuration (n-2)f¹⁻¹⁴(n-1)d⁰⁻¹ns². This classification helps predict chemical properties and reactivity patterns.
Effective nuclear charge (Z_eff) is the net positive charge experienced by valence electrons after accounting for shielding by inner electrons. As you move across a period, Z_eff increases because protons are added but electrons enter the same shell with similar shielding. Down a group, Z_eff remains relatively constant as both nuclear charge and shielding increase proportionally. Higher Z_eff causes smaller atomic size, higher ionization energy, and greater electronegativity explaining most periodic trends.
Lanthanide contraction refers to the greater-than-expected decrease in atomic and ionic radii of lanthanide elements (atomic numbers 58-71) and subsequent elements. It occurs because 4f electrons shield outer electrons imperfectly from the increasing nuclear charge. As 4f orbitals fill across the series, each additional proton pulls the entire electron cloud closer, causing size to decrease more than expected. This affects properties of elements following the lanthanides, making 5d transition elements similar in size to 4d elements.
Oxide character changes from basic to acidic as you move from left to right across a period. Metals on the left form basic oxides (like Na₂O, MgO) that react with acids. Metalloids in the middle form amphoteric oxides (like Al₂O₃) that react with both acids and bases. Non-metals on the right form acidic oxides (like SO₂, P₂O₅, Cl₂O₇) that react with bases. This trend reflects the gradual change from metallic to non-metallic character across the period.
Alkali metals (Group 1: Li, Na, K, Rb, Cs, Fr) have one electron in their outermost shell (ns¹ configuration), making them highly reactive metals with low ionization energies. They readily lose this electron to form +1 ions. Metallic character and reactivity increase down the group, with cesium being the most metallic and reactive (excluding radioactive francium). They all react with water to produce hydrogen gas and form strongly basic hydroxides.
Transition elements (d-block, Groups 3-12) have their last electron entering (n-1)d orbitals with general configuration (n-1)d¹⁻¹⁰ns⁰⁻². They show variable valency because both ns and (n-1)d electrons can participate in bonding. These elements typically form colored compounds, exhibit catalytic properties, and can form complex ions. They show less dramatic property changes across a period compared to representative elements due to similar outer shell configurations.
Atomic size (atomic radius) is the distance from the nucleus to the outermost electrons. In a group, atomic size increases from top to bottom as new electron shells are added, increasing the distance from the nucleus. Across a period from left to right, atomic size decreases because the nuclear charge increases while electrons are added to the same shell, pulling electrons closer. For example, in Period 2, lithium is the largest atom and fluorine is the smallest.