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Complete Notes & Study Material
Master the Periodic Table — from Döbereiner's triads to modern periodic law. Covering all trends, properties, and exam-specific insights for JEE, NEET & Board exams.
This chapter is the backbone of inorganic chemistry — virtually every concept in Class 12 and competitive exams ties back to periodic trends.
Based on last 10 years of papers across major entrance exams.
| Property | Across Period → | Down Group ↓ | Notable Exception |
|---|---|---|---|
| Atomic Radius | Decreases | Increases | Noble gases (Van der Waals) |
| Ionisation Enthalpy | Increases | Decreases | B < Be; O < N |
| Electron Gain Enthalpy | More Negative | Less Negative | O < S; F < Cl (EGE) |
| Electronegativity | Increases | Decreases | Noble gases: no value |
| Metallic Character | Decreases | Increases | — |
| Non-metallic Character | Increases | Decreases | — |
| Oxidising Power | Increases | Decreases | F is strongest oxidiser |
| Basic nature of oxides | Decreases | Increases | Amphoteric oxides (Al₂O₃) |
As chemists began discovering more and more elements during the nineteenth century, a serious problem emerged: how could this growing collection be organized meaningfully? Each element possessed its own mass, reactivity, and combining power, yet certain properties appeared to repeat in a regular manner. The search for order within this apparent complexity led to the birth of periodic classification.
The earliest systematic attempt was made by Johann Wolfgang Döbereiner, who grouped elements into triads. In each triad, the atomic mass of the middle element was approximately the average of the other two, and their chemical behavior showed close resemblance. Although limited in scope, this idea hinted that elemental properties might depend on atomic mass.
A broader pattern was proposed by John Newlands. He arranged elements in increasing order of atomic mass and observed that every eighth element showed similar properties—an analogy he called the Law of Octaves. While this worked reasonably well for lighter elements, it failed for heavier ones, and several mismatches reduced its acceptance.
A major breakthrough came with Dmitri Mendeleev, who constructed a table based on atomic masses and chemical similarities. His genius lay in leaving gaps for undiscovered elements and boldly predicting their properties. When these elements were later found with characteristics close to his predictions, confidence in periodic classification grew immensely.
The final refinement arrived through Henry Moseley, who demonstrated that atomic number—not atomic mass—is the fundamental basis of periodicity. This resolved earlier anomalies and firmly established the modern periodic law: the properties of elements are periodic functions of their atomic numbers.
Thus, periodic classification evolved step by step, from simple groupings to a powerful scientific framework—revealing that the diversity of elements is governed by an underlying order. This journey reflects how careful observation, logical arrangement, and experimental evidence together shaped one of chemistry’s most important organizing principles.
The arrangement of elements into a systematic framework reached its scientific maturity with the formulation of the Modern Periodic Law. Careful experimental investigations revealed that the fundamental property distinguishing one element from another is not its atomic mass but its atomic number. The atomic number represents the number of protons in the nucleus and, consequently, determines the number of electrons in a neutral atom. Since chemical behavior depends primarily on electronic configuration, periodicity in properties must arise from periodic repetition of similar electronic arrangements.
The Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers. When elements are arranged in increasing order of atomic number, elements with similar valence shell configurations recur at regular intervals. This repetition explains similarities in valency, reactivity, atomic size, and other properties.
The present form of the periodic table, often called the long form, is based directly on atomic number and electronic configuration. It consists of:
Each period corresponds to the filling of a principal energy level. For example, the first period contains only two elements because only the 1\(s\) orbital is available for electron occupancy. As we move to higher periods, additional subshells become available, increasing the number of elements in each row.
A distinctive feature of the present table is its division into blocks according to the type of subshell being filled:
This block-wise arrangement reflects the quantum mechanical understanding of atomic structure and directly links periodic classification with electronic configuration.
The present periodic table is not merely a list of elements; it is a structured representation of electronic architecture. It provides predictive power: once the position of an element is known, its likely properties can be inferred. Thus, the Modern Periodic Law establishes a direct connection between atomic structure and observable chemical behavior, making the periodic table a central organizing principle of chemistry.
The modern arrangement of elements is deeply rooted in their electronic configurations. Every atom contains electrons distributed in different energy levels and sublevels (orbitals). This pattern of distribution is not random—it follows definite rules—and it is this orderly filling of electrons that gives rise to the structure of the periodic table and the regular repetition of chemical properties.
Electrons occupy orbitals in increasing order of energy. Each principal energy level \((n=1,2,3,\ldots)\) contains subshells such as \(s,\ p,\ d\) and \(f\), which differ in shape and energy. As atomic number increases from one element to the next, electrons are added successively to these orbitals. Because similar outer electronic arrangements recur at regular intervals, elements with comparable properties appear periodically.
Each period in the periodic table corresponds to the filling of a principal energy level.
Thus, the length of a period is directly related to the number of orbitals being filled at that stage.
Elements placed in the same group possess similar valence shell electronic configurations. This similarity explains why members of a group show closely related chemical behavior.
Hence, group properties arise from identical patterns in the outermost electrons.
A distinctive feature of the present periodic table is its division into four blocks, determined by the subshell receiving the last electron:
This block arrangement provides a clear link between quantum mechanical principles and the periodic classification of elements.
Because electronic configuration governs atomic size, valency, ionization energy, and electronegativity, these properties also vary periodically across the table. As electrons are added to the same shell across a period, nuclear attraction increases, leading to systematic trends. Down a group, the addition of new shells causes predictable changes in atomic dimensions and reactivity.
The elements placed in Groups 1 and 2 of the periodic table constitute the s-block elements. They are so named because the differentiating electron in these elements enters the outermost s-orbital. The general outer electronic configuration of these elements is \(ns^1\) for Group 1 and \(ns^2\) for Group 2, where nrepresents the principal quantum number of the valence shell. This simple electronic arrangement is responsible for their characteristic properties and high chemical reactivity.
The s-block occupies the extreme left portion of the periodic table. Group 1 elements are known as the alkali metals, while Group 2 elements are referred to as the alkaline earth metals. Hydrogen, though placed above Group 1 due to its electronic configuration \(1s^1\), exhibits properties that are distinct from both alkali metals and halogens.
Because the valence shell contains only one or two electrons, these elements readily lose electrons to attain a stable noble gas configuration. As a result, they are strong reducing agents and form predominantly ionic compounds.
The reactivity of s-block elements can be directly correlated with their electronic configuration.
Down a group, the atomic size increases due to the addition of new electron shells. Consequently, the outermost electrons experience less effective nuclear attraction and are more easily removed. This explains why reactivity increases from top to bottom in both groups.
Across a group, metallic character becomes more pronounced. Melting and boiling points are relatively low compared to many other metals. Alkali metals are soft and can often be cut with a knife. Alkaline earth metals are comparatively harder and denser.
The s-block elements play a vital role in both industrial and biological systems. Sodium and potassium ions are essential for nerve impulse transmission and maintenance of osmotic balance in living organisms. Calcium and magnesium are fundamental components of bones, teeth, and chlorophyll, respectively.
The elements belonging to Groups 13 to 18 occupy the right-hand portion of the periodic table and are collectively known as the p-block elements. In these elements, the differentiating electron enters the outermost p-orbital. The general outer electronic configuration of these elements is \(ns^2np^{1-6}\), where nrepresents the principal quantum number of the valence shell. This progressive filling of the p-subshell gives rise to a remarkable diversity in physical and chemical properties.
The p-block begins from Group 13 and extends up to Group 18. As we move across a period in this region, electrons are successively added to the three \(p\)-orbitals of the same energy level. Since each \(p\)-subshell can accommodate a maximum of six electrons, there are six groups in this block.
The number of valence electrons increases from three in Group 13 to eight in Group 18. This steady increase in valence electrons explains the gradual transition from metallic to non-metallic character across a period.
One of the distinctive features of the p-block is the presence of metals, metalloids, and non-metals within the same block.
Thus, the \(p\)-block illustrates the gradual change in properties across a period.
The chemical properties of p-block elements are largely governed by the number of valence electrons.
Down a group, the tendency to exhibit lower oxidation states becomes more pronounced in heavier elements. This behavior arises due to the increasing stability of the \(ns^2\) electron pair in heavier atoms.
Across a period, atomic size decreases because of increasing nuclear charge, while ionization enthalpy and electronegativity generally increase. This results in a shift from electropositive metallic behavior to electronegative non-metallic behavior.
Down a group, atomic size increases and metallic character becomes more dominant. Bonding patterns also change: lighter elements tend to form covalent compounds, whereas heavier members may exhibit more ionic character.
The \(p\)-block elements include many substances essential to life and industry. Carbon forms the basis of organic chemistry. Nitrogen and phosphorus are vital components of biological molecules. Oxygen supports respiration and combustion. The halogens are widely used in disinfectants and industrial processes. Noble gases find application in lighting and protective atmospheres.
The elements occupying Groups 3 to 12 in the periodic table form the d-block. In these elements, the differentiating electron enters the \((n-1)\) dsubshell, while the outermost shell generally contains one or two \(s\)-electrons. Their general electronic configuration may be represented as \((n-1)d^{1–10}ns^{1–2}\).
A large number of these elements are known as transition elements. They are so called because they represent a gradual transition in properties between the highly reactive \(s\)-block metals on the left and the \(p\)-block elements on the right. Strictly speaking, transition elements are those which possess partially filled \(d\)-orbitals either in their atoms or in at least one of their common oxidation states.
The \(d\)-block extends across the centre of the periodic table and is spread over four long periods. As atomic number increases within a series, electrons are progressively added to the \((n-1)d\) orbitals. Because these d-orbitals are close in energy to the outer nsorbitals, both sets of electrons often participate in bonding. This unique electronic arrangement is responsible for many characteristic features of transition elements.
Although individual members show differences, the \(d\)-block elements display several common properties:
Across a transition series, atomic size decreases slightly due to increasing nuclear charge, though the change is much less pronounced than in \(s-\) or \(p-\)blocks. Ionization enthalpy shows only a gradual increase. Metallic character remains strong throughout, but subtle changes in reactivity and bonding behavior occur as the \(d\)-subshell becomes progressively filled.
Transition elements are of immense practical value. Iron forms the backbone of construction materials, copper is essential for electrical wiring, and nickel and chromium are used in corrosion-resistant alloys. Many biological systems also depend on transition metals—for example, iron in haemoglobin and cobalt in certain vitamins.
The elements placed separately at the bottom of the periodic table form the f-block, also known as the inner-transition elements. In these elements, the differentiating electron enters the \((n-2)f\) subshell. They are arranged in two horizontal rows: the lanthanoids (from atomic number 57 to 71) and the actinoids (from atomic number 89 to 103). Although shown apart from the main body of the table, they actually belong to periods 6 and 7.
Their general electronic configuration may be represented as \((n-2)f^{1–14}(n-1)d^{0–1}ns^2\). This gradual filling of the \(f\)-orbitals gives rise to many characteristic properties that distinguish these elements from the \(s-,\ p-\), and \(d\)-blocks.
The \(f\)-block lies between Groups 2 and 3 of the periodic table. For convenience and compactness, these elements are displayed below the main table. In the lanthanoids, electrons are added to the \(4f\) subshell, while in the actinoids, filling occurs in the 5fsubshell.
\(s,\ p,\) and \(d\) orbitals. As a result, the chemical behavior of many \(f\)-block elements appears quite similar within each series.
The lanthanoids are a group of fourteen elements following lanthanum. They are often referred to as rare earth metals, although many of them are not truly rare in nature.
General characteristics of lanthanoids include:
A notable feature of this series is the lanthanoid contraction, a gradual decrease in atomic and ionic radii from the first to the last member. This occurs because the added \(4f\) electrons do not shield nuclear charge effectively, leading to a steady increase in attraction between the nucleus and the outer electrons.
The actinoids follow actinium and involve filling of the \(5f\) orbitals. In contrast to lanthanoids, most actinoids are radioactive, and several are synthetic.
Their chemical behavior is more complex because both \(5f\) and \(6d\) electrons can participate in bonding. As a result, actinoids show a wider range of oxidation states, especially in the earlier members of the series.
Key features of actinoids include:
Across both series, atomic size decreases gradually due to poor shielding by \(f\)-electrons. Many \(f\)-block elements form coloured compounds and exhibit paramagnetism, arising from unpaired electrons in \(f\)-orbitals. Their compounds are largely ionic, and most metals are chemically reactive.
Despite being less familiar than other blocks, \(f\)-block elements are technologically significant. Lanthanoids are widely used in magnets, phosphors, lasers, and electronic devices. Certain actinoids play crucial roles in nuclear energy production and scientific research.
One of the most useful ways of understanding the periodic table is by grouping elements into metals, non-metals, and metalloids. This broad classification is based on differences in physical appearance, electrical behavior, and chemical reactivity. Although each element has its own identity, clear patterns emerge when they are arranged according to atomic number, revealing a gradual change in character across periods and down groups.
Metals occupy the left-hand side and the central region of the periodic table. Most elements belong to this category.
General properties of metals include:
Because metals readily give up electrons, they are described as electropositive. Their oxides are generally basic in nature, and many metals react with acids to liberate hydrogen gas. Down a group, metallic character increases as atomic size becomes larger and outer electrons are held less strongly by the nucleus.
Metals play an essential role in daily life and industry, forming structural materials, electrical conductors, machinery components, and countless alloys.
Non-metals are found mainly on the right-hand side of the periodic table, along with a few at the top.
Typical characteristics of non-metals are:
Non-metals are described as electronegative because of their tendency to attract electrons. Their oxides are usually acidic or neutral. Across a period, as nuclear charge increases and atomic size decreases, elements gradually change from metallic to non-metallic behavior.
Non-metals are fundamental to life and chemistry: they form the backbone of organic compounds, atmospheric gases, and many biologically important molecules.
Between metals and non-metals lies a narrow zigzag region occupied by a small group of elements known as metalloids. These elements exhibit properties intermediate between the two extremes.
Metalloids typically:
Because of this intermediate nature, metalloids are especially important in electronics. Their controlled conductivity makes them suitable for use in semiconductors and modern technological devices.
The distribution of metals, non-metals, and metalloids is not random. It reflects underlying trends in electronic configuration:
This gradual shift illustrates how changes in atomic size and electron-holding ability govern chemical behavior.
The systematic arrangement of elements in the periodic table reveals that many physical properties do not change randomly. Instead, they vary in a regular and predictable manner from left to right across a period and from top to bottom within a group. These recurring variations are known as periodic trends. They arise mainly due to changes in atomic size, nuclear charge, and electronic configuration.
The atomic radius refers to the size of an atom, usually expressed as half the distance between the nuclei of two identical bonded atoms. Although atoms do not have sharply defined boundaries, this measure provides a useful comparison of relative sizes.
Across a period
atomic radius generally decreases. As we move from
left to right, electrons are added to
the same principal energy level, while nuclear charge steadily increases. The stronger attraction
between the nucleus and electrons pulls the electron cloud closer, resulting in smaller atomic size.
Down a group
atomic radius increases. Each successive element
possesses an additional electron shell. The
outer electrons are therefore farther from the nucleus and experience greater shielding by inner
electrons, leading to a larger atomic radius.
When atoms lose or gain electrons, they form ions, and their sizes change accordingly. The size of an ion is called the ionic radius.
Cations (positive ions)
formed by loss of electrons, are smaller than
their parent atoms. This reduction
occurs because removal of electrons decreases electron–electron repulsion and allows the nucleus to pull
the
remaining electrons closer.
Anions (negative ions)
formed by gain of electrons, are larger than
their parent atoms. The added electrons
increase repulsion in the outer shell, causing expansion of the electron cloud.
Ionization enthalpy is the energy required to remove the most loosely held electron from an isolated gaseous atom. It reflects how strongly an atom holds its outer electrons.
Across a period
ionization enthalpy generally increases. Atomic size
decreases and nuclear attraction
increases, making it more difficult to remove an electron.
Down a group
ionization enthalpy decreases. The outer electrons are
farther from the nucleus and more
effectively shielded by inner shells, so they are removed more easily.
Small irregularities occur due to differences in subshell stability and electron pairing, but the overall trend remains consistent.
Electron gain enthalpy represents the energy change when an electron is added to a neutral gaseous atom. It indicates the tendency of an atom to accept an electron.
Across a period
electron gain enthalpy becomes more negative (more
energy is released), as increasing
nuclear charge enhances the attraction for incoming electrons. Non-metals on the right side of the table
generally show high negative values.
Down a group
the trend is less regular. Although atomic size
increases, reduced attraction for added
electrons and increased shielding can offset this effect, leading to deviations in expected values.
Noble gases show little tendency to gain electrons because their valence shells are already complete.
Electronegativity is the tendency of an atom to attract shared electrons towards itself in a chemical bond. It is not measured directly but expressed on relative scales.
Across a period
electronegativity increases due to decreasing atomic
size and increasing nuclear charge,
which strengthen the pull on bonding electrons.
Down a group
electronegativity decreases as atomic size increases and
valence electrons are held less
firmly.
Elements with high electronegativity readily attract electrons, while those with low electronegativity tend to lose them. This difference governs bond polarity and the nature of chemical bonding.
Chemical properties of elements show a definite periodic pattern when the elements are arranged in increasing order of atomic number. This regularity arises from the repetition of similar valence shell electronic configurations. Among the most important chemical trends are the variation in valence or oxidation states and the distinctive behavior of second period elements.
The valence of an element refers to its combining capacity, while the oxidation state represents the apparent charge on an atom in a compound. Both are closely linked to the number of valence electrons.
As we move from left to right across a period, the number of valence electrons increases from one to eight. Consequently:
For example, elements near the middle of a period commonly show higher oxidation states, while those near the ends exhibit lower or fixed values. This systematic change reflects the gradual filling of the same outer shell.
Elements belonging to the same group possess similar valence shell configurations and therefore exhibit similar valence and oxidation states. However, as atomic size increases down a group, additional trends appear:
Transition elements (d-block) display a wide range of oxidation states because both nsand (n-1)delectrons can participate in bonding. In contrast, s-block elements generally show fixed oxidation states corresponding to the loss of one or two electrons.
Thus, periodicity in oxidation states arises directly from electronic configuration and the relative ease with which electrons are lost or shared.
The elements of the second period—lithium to neon—often differ markedly from their heavier congeners in the same groups. This unusual behavior is known as their anomalous character.
The main reasons for this anomaly are:
Due to these factors, second period elements show several distinctive features:
For example, carbon readily forms long chains and multiple bonds, while silicon, though in the same group, shows far less tendency for such behavior.
Chemical reactivity refers to the tendency of an element to undergo chemical change. When elements are arranged in the periodic table according to increasing atomic number, their reactivity does not vary randomly. Instead, it follows definite periodic trends, arising from systematic changes in electronic configuration, atomic size, ionization enthalpy, electron gain enthalpy, and electronegativity. These factors together determine how readily an atom loses, gains, or shares electrons during chemical reactions.
At the atomic level, chemical reactions involve only the valence electrons. Elements react in order to attain a more stable electronic arrangement, usually resembling that of a noble gas. Whether an element tends to lose electrons, gain electrons, or share electrons depends primarily on:
As these quantities change periodically across periods and down groups, chemical reactivity also exhibits regular variation.
Metals generally react by losing electrons to form positive ions. Their reactivity is therefore linked to how easily they can release valence electrons.
From left to right across a period, metallic character decreases. Atomic size becomes smaller and ionization enthalpy increases, making electron removal more difficult. Consequently, metallic reactivity decreases across a period.
Down a group, atomic size increases and outer electrons are held less firmly. Ionization enthalpy decreases, so electrons are lost more easily. Hence, reactivity of metals increases down a group.
This trend is clearly observed in alkali metals, where each successive element reacts more vigorously than the one above it.
Non-metals typically react by gaining electrons or by sharing electrons. Their reactivity depends on electron gain enthalpy and electronegativity.
Across a period, electronegativity and tendency to gain electrons increase. As a result, non-metallic reactivity increases from left to right, reaching a maximum near the halogens.
Down a group of non-metals, atomic size increases and the attraction for incoming electrons weakens. Therefore, reactivity of non-metals generally decreases down the group.
For example, among halogens, the element at the top is the most reactive, while heavier members show progressively lower reactivity.
Elements with completely filled valence shells exhibit extremely low chemical reactivity. Their electronic configuration is already stable, so they show little tendency to gain, lose, or share electrons. This explains the chemical inertness associated with this group.
Chemical reactivity is closely connected with other periodic trends:
Thus, trends in atomic size, ionization enthalpy, electron gain enthalpy, and electronegativity together shape the overall pattern of reactivity.
A broad view of the periodic table reveals that:
This diagonal distribution highlights the balance between electron-losing and electron-gaining tendencies across the table.
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