Studying s-block, hydrogen, and p-block elements for JEE Mains requires a systematic approach that combines understanding, memorization, and consistent practice. Begin with a strong grasp of the periodic table layout, focusing on how groups and periods define properties. Start with s-block (Group 1 and 2 elements), followed by hydrogen and then p-block (Groups 13 to 18).For s-block and p-block, first understand their general electronic configurations and periodic trends (like atomic size, ionization energy, reactivity, and oxidation states). Don’t just memorize; try to relate trends to electronic structure and nuclear charge. For each group, make concise notes on key properties, important compounds, and anomalous behaviors. For hydrogen, focus on its unique position, isotopes, and reactions with other elements.Use the NCERT textboo k as your primary reference and complement it with well-organized notes and concept maps. After studying theory, practice as many questions as possible including NCERT in-text and exercise problems, previous years’ JEE questions, and online quizzes. Make a list of frequently asked reactions (like preparation of NaOH, properties of hydrogen peroxide, trends in halogens, etc.) and revise them regularly. Finally, schedule periodic revisions to reinforce learning and identify areas that need more focus.

Studying s-block, hydrogen, and p-block elements for JEE Mains requires a systematic approach that combines understanding, memorization, and consistent practice. Begin with a strong grasp of the periodic table layout, focusing on how groups and periods define properties. Start with s-block (Group 1 and 2 elements), followed by hydrogen and then p-block (Groups 13 to 18).For s-block and p-block, first understand their general electronic configurations and periodic trends (like atomic size, ionization energy, reactivity, and oxidation states). Don’t just memorize; try to relate trends to electronic structure and nuclear charge. For each group, make concise notes on key properties, important compounds, and anomalous behaviors. For hydrogen, focus on its unique position, isotopes, and reactions with other elements.Use the NCERT textboo k as your primary reference and complement it with well-organized notes and concept maps. After studying theory, practice as many questions as possible including NCERT in-text and exercise problems, previous years’ JEE questions, and online quizzes. Make a list of frequently asked reactions (like preparation of NaOH, properties of hydrogen peroxide, trends in halogens, etc.) and revise them regularly. Finally, schedule periodic revisions to reinforce learning and identify areas that need more focus.

The s-block elements are found in the first two groups (columns) of the periodic table Group 1 (alkali metals) and Group 2 (alkaline earth metals). These elements are named “s-block” because their outermost electrons occupy an s-orbital, which can hold a maximum of two electrons. Examples include lithium, sodium, potassium (Group 1) and magnesium, calcium (Group 2).S-block elements are characterized by their high reactivity, low ionization energies, and tendency to form ionic compounds. Alkali metals have a single electron in their outermost shell (ns¹), making them highly reactive, especially with water. Alkaline earth metals have two valence electrons (ns²), which they readily lose to form +2 ions.These elements display several common trends: they are generally soft, have low melting points (especially alkali metals), and conduct electricity well. Their chemistry is significant in daily life and industry think of table salt (sodium chloride), milk of magnesia (magnesium hydroxide), and calcium in bones. Understanding s-block chemistry is crucial for grasping broader chemical trends and applications in both academic and real-world contexts.

The s-block elements are found in the first two groups (columns) of the periodic table Group 1 (alkali metals) and Group 2 (alkaline earth metals). These elements are named “s-block” because their outermost electrons occupy an s-orbital, which can hold a maximum of two electrons. Examples include lithium, sodium, potassium (Group 1) and magnesium, calcium (Group 2).S-block elements are characterized by their high reactivity, low ionization energies, and tendency to form ionic compounds. Alkali metals have a single electron in their outermost shell (ns¹), making them highly reactive, especially with water. Alkaline earth metals have two valence electrons (ns²), which they readily lose to form +2 ions.These elements display several common trends: they are generally soft, have low melting points (especially alkali metals), and conduct electricity well. Their chemistry is significant in daily life and industry think of table salt (sodium chloride), milk of magnesia (magnesium hydroxide), and calcium in bones. Understanding s-block chemistry is crucial for grasping broader chemical trends and applications in both academic and real-world contexts.

Yes, s-block elements are very much a part of the JEE Advanced syllabus . Both s-block and p-block elements are foundational topics in the study of inorganic chemistry and are explicitly mentioned in the official JEE Advanced syllabus. The exam tests your understanding of their properties, preparation, reactions, trends, and applications.You are expected to be familiar with the electronic configuration, occurrence, extraction, and uses of alkali and alkaline earth metals, as well as important compounds such as sodium carbonate, sodium chloride, calcium carbonate, and magnesium sulfate. Questions often focus on periodic trends, chemical reactivity, anomalous properties, and comparison within the group or with hydrogen.Additionally, JEE Advanced emphasizes the application of concepts rather than rote memorization. For instance, you might be asked to compare the solubility of group 2 carbonates, or predict the outcome of a reaction between group 1 metal and water. Therefore, thorough understanding, coupled with regular problem-solving and revision, is essential to master this part of the syllabus and perform well in the entrance exam .

Yes, s-block elements are very much a part of the JEE Advanced syllabus . Both s-block and p-block elements are foundational topics in the study of inorganic chemistry and are explicitly mentioned in the official JEE Advanced syllabus. The exam tests your understanding of their properties, preparation, reactions, trends, and applications.You are expected to be familiar with the electronic configuration, occurrence, extraction, and uses of alkali and alkaline earth metals, as well as important compounds such as sodium carbonate, sodium chloride, calcium carbonate, and magnesium sulfate. Questions often focus on periodic trends, chemical reactivity, anomalous properties, and comparison within the group or with hydrogen.Additionally, JEE Advanced emphasizes the application of concepts rather than rote memorization. For instance, you might be asked to compare the solubility of group 2 carbonates, or predict the outcome of a reaction between group 1 metal and water. Therefore, thorough understanding, coupled with regular problem-solving and revision, is essential to master this part of the syllabus and perform well in the entrance exam .

Hydrogen’s properties are unique and do not fully match those of either Group 1 (alkali metals) or Group 17 (halogens), despite superficial similarities in electron configuration. Like Group 1 elements, hydrogen has a single electron in its outermost shell (1s¹), which suggests it could lose an electron to form H⁺. However, unlike alkali metals, hydrogen is a non-metal, is diatomic (H₂), and exists as a gas at room temperature.Hydrogen also shares similarities with Group 17 halogens, as both require one more electron to achieve a noble gas configuration. This enables hydrogen to form covalent bonds by sharing electrons (like in HCl). However, hydrogen is much less electronegative than halogens, forms positive as well as negative ions (hydrides), and exhibits distinct physical and chemical behaviors.The reason for these differences lies in hydrogen’s small size, high ionization energy, and absence of inner electrons. Hydrogen’s chemistry is versatile; it can lose, gain, or share electrons, leading to a wide range of compounds. This versatility makes hydrogen unique, often justifying its separate placement in the periodic table and its study as a special case in chemistry.

Hydrogen’s properties are unique and do not fully match those of either Group 1 (alkali metals) or Group 17 (halogens), despite superficial similarities in electron configuration. Like Group 1 elements, hydrogen has a single electron in its outermost shell (1s¹), which suggests it could lose an electron to form H⁺. However, unlike alkali metals, hydrogen is a non-metal, is diatomic (H₂), and exists as a gas at room temperature.Hydrogen also shares similarities with Group 17 halogens, as both require one more electron to achieve a noble gas configuration. This enables hydrogen to form covalent bonds by sharing electrons (like in HCl). However, hydrogen is much less electronegative than halogens, forms positive as well as negative ions (hydrides), and exhibits distinct physical and chemical behaviors.The reason for these differences lies in hydrogen’s small size, high ionization energy, and absence of inner electrons. Hydrogen’s chemistry is versatile; it can lose, gain, or share electrons, leading to a wide range of compounds. This versatility makes hydrogen unique, often justifying its separate placement in the periodic table and its study as a special case in chemistry.

Hydrogen gas (H₂) can be produced through several laboratory and industrial methods, each suited to specific needs and resource availability. The simplest laboratory method involves the reaction of a dilute acid with a reactive metal—for example, when zinc granules are added to dilute hydrochloric acid, hydrogen gas is evolved:Zn + 2HCl → ZnCl₂ + H₂↑This reaction is commonly used in school labs for small-scale hydrogen production. In dustrially, hydrogen is most often made by steam reforming of methane (natural gas), where methan e reacts with steam at high temperatures in the presence of a catalyst:CH₄ + H₂O → CO + 3H₂Another important industrial process is the electrolysis of water, where an electric current is passed through water (usually containing a small amount of acid or base to increase conductivity) to decompose it into hydrogen and oxygen:2H₂O(l) → 2H₂(g) + O₂(g)Each method has its advantages laboratory reactions are quick and simple, while industrial processes are optimized for large-scale, efficient, and cost-effective production. The choice depends on the volume needed, purity requirements, and available resources.

Hydrogen gas (H₂) can be produced through several laboratory and industrial methods, each suited to specific needs and resource availability. The simplest laboratory method involves the reaction of a dilute acid with a reactive metal—for example, when zinc granules are added to dilute hydrochloric acid, hydrogen gas is evolved:Zn + 2HCl → ZnCl₂ + H₂↑This reaction is commonly used in school labs for small-scale hydrogen production. In dustrially, hydrogen is most often made by steam reforming of methane (natural gas), where methan e reacts with steam at high temperatures in the presence of a catalyst:CH₄ + H₂O → CO + 3H₂Another important industrial process is the electrolysis of water, where an electric current is passed through water (usually containing a small amount of acid or base to increase conductivity) to decompose it into hydrogen and oxygen:2H₂O(l) → 2H₂(g) + O₂(g)Each method has its advantages laboratory reactions are quick and simple, while industrial processes are optimized for large-scale, efficient, and cost-effective production. The choice depends on the volume needed, purity requirements, and available resources.

Hydrogen is classified as a non-metal. Despite its position above Group 1 alkali metals in the periodic table and its single valence electron, hydrogen’s physical and chemical properties are distinctly non-metallic. It exists as a colorless, odorless, diatomic gas (H₂) under standard conditions quite different from metals, which are usually solid, lustrous, and good conductors of electricity.Chemically, hydrogen forms covalent bonds by sharing its single electron with other non-metals (as in water, H₂O) and ionic bonds with metals (as in sodium hydride, NaH). It does not display typical metallic characteristics such as malleability, ductility, or high electrical conductivity in its elemental form.Moreover, hydrogen is highly reactive, has high ionization energy, and can form positive (H⁺) or negative (H⁻) ions, which is a property not generally seen in metals or metalloids. These features justify its classification as a non-metal, despite its unique versatility in chemical reactions and placement in the periodic table .

Hydrogen is classified as a non-metal. Despite its position above Group 1 alkali metals in the periodic table and its single valence electron, hydrogen’s physical and chemical properties are distinctly non-metallic. It exists as a colorless, odorless, diatomic gas (H₂) under standard conditions quite different from metals, which are usually solid, lustrous, and good conductors of electricity.Chemically, hydrogen forms covalent bonds by sharing its single electron with other non-metals (as in water, H₂O) and ionic bonds with metals (as in sodium hydride, NaH). It does not display typical metallic characteristics such as malleability, ductility, or high electrical conductivity in its elemental form.Moreover, hydrogen is highly reactive, has high ionization energy, and can form positive (H⁺) or negative (H⁻) ions, which is a property not generally seen in metals or metalloids. These features justify its classification as a non-metal, despite its unique versatility in chemical reactions and placement in the periodic table .

Hydrogen is a colorless, odorless, tasteless, and highly flammable diatomic gas (H₂) at room temperature and standard atmospheric pressure. It is the lightest and most abundant element in the universe, making up about 75% of its elemental mass, predominantly in stars and interstellar space.Hydrogen gas is non-toxic, but because it is extremely light, it rises rapidly in air and can disperse quickly. Its most notable physical property is its high flammability—hydrogen can form explosive mixtures with air or oxygen, igniting with a pale blue, nearly invisible flame. This property requires careful handling and storage.In practical applications, hydrogen is used in the synthesis of ammonia (via the Haber process), as a reducing agent in metallurgy, in hydrogen fuel cells for clean energy, and in various laboratory procedures. Its unique properties make it crucial for energy, industrial, and scientific uses.

Hydrogen is a colorless, odorless, tasteless, and highly flammable diatomic gas (H₂) at room temperature and standard atmospheric pressure. It is the lightest and most abundant element in the universe, making up about 75% of its elemental mass, predominantly in stars and interstellar space.Hydrogen gas is non-toxic, but because it is extremely light, it rises rapidly in air and can disperse quickly. Its most notable physical property is its high flammability—hydrogen can form explosive mixtures with air or oxygen, igniting with a pale blue, nearly invisible flame. This property requires careful handling and storage.In practical applications, hydrogen is used in the synthesis of ammonia (via the Haber process), as a reducing agent in metallurgy, in hydrogen fuel cells for clean energy, and in various laboratory procedures. Its unique properties make it crucial for energy, industrial, and scientific uses.

The chemical symbol H₂ represents a molecule of hydrogen gas, which consists of two hydrogen atoms covalently bonded together. Under standard conditions, hydrogen does not exist as single atoms because the atom is highly reactive and unstable due to its unpaired electron. Instead, two hydrogen atoms share their electrons, achieving a stable configuration similar to the noble gas helium.This diatomic nature is not unique to hydrogen other non-metals like oxygen (O₂), nitrogen (N₂), and halogens (Cl₂, F₂) also exist as diatomic molecules in their elemental forms. The subscript “2” in H₂ indicates the presence of two atoms in the molecule.Understanding this notation is fundamental in chemical equations and reactions. For example, in the combustion of hydrogen:2H₂ + O₂ → 2H₂OHere, H₂ represents molecular hydrogen, not atomic hydrogen. This distinction is crucial for correctly balancing chemical equations and predicting reaction outcomes in both academic and practical chemistry contexts.

The chemical symbol H₂ represents a molecule of hydrogen gas, which consists of two hydrogen atoms covalently bonded together. Under standard conditions, hydrogen does not exist as single atoms because the atom is highly reactive and unstable due to its unpaired electron. Instead, two hydrogen atoms share their electrons, achieving a stable configuration similar to the noble gas helium.This diatomic nature is not unique to hydrogen other non-metals like oxygen (O₂), nitrogen (N₂), and halogens (Cl₂, F₂) also exist as diatomic molecules in their elemental forms. The subscript “2” in H₂ indicates the presence of two atoms in the molecule.Understanding this notation is fundamental in chemical equations and reactions. For example, in the combustion of hydrogen:2H₂ + O₂ → 2H₂OHere, H₂ represents molecular hydrogen, not atomic hydrogen. This distinction is crucial for correctly balancing chemical equations and predicting reaction outcomes in both academic and practical chemistry contexts.

Alkali metals (Group 1 elements: lithium, sodium, potassium, etc.) belong to the s-block of the periodic table because their outermost electron occupies an s-orbital. Specifically, their general electronic configuration is ns¹, where “n” is the period number. For example, lithium is [He] 2s¹, sodium is [Ne] 3s¹, and potassium is [Ar] 4s¹.The s-orbital can hold a maximum of two electrons, and in alkali metals, this single s-electron is easily lost, making these elements highly reactive and prone to forming +1 ions. Their placement in the s-block is determined by this shared electronic feature, which also governs their similar chemical properties.For example, both sodium and potassium react vigorously with water, forming hydroxides and releasing hydrogen gas:2Na + 2H₂O → 2NaOH + H₂↑Because their reactivity and properties stem from their s-orbital electron, their classification in the s-block is logical and consistent with periodic trends.

Alkali metals (Group 1 elements: lithium, sodium, potassium, etc.) belong to the s-block of the periodic table because their outermost electron occupies an s-orbital. Specifically, their general electronic configuration is ns¹, where “n” is the period number. For example, lithium is [He] 2s¹, sodium is [Ne] 3s¹, and potassium is [Ar] 4s¹.The s-orbital can hold a maximum of two electrons, and in alkali metals, this single s-electron is easily lost, making these elements highly reactive and prone to forming +1 ions. Their placement in the s-block is determined by this shared electronic feature, which also governs their similar chemical properties.For example, both sodium and potassium react vigorously with water, forming hydroxides and releasing hydrogen gas:2Na + 2H₂O → 2NaOH + H₂↑Because their reactivity and properties stem from their s-orbital electron, their classification in the s-block is logical and consistent with periodic trends.

S-block elements exhibit several distinctive properties that stem from their electronic configuration: High Reactivity: Both alkali and alkaline earth metals are highly reactive, especially with water and halogens, due to their tendency to lose electrons easily. Softness and Low Density: Alkali metals are particularly soft (can be cut with a knife) and have low densities—lithium, sodium, and potassium float on water. Low Ionization Energies: The energy required to remove the outermost electron is low, resulting in the easy formation of cations (Na⁺, Mg²⁺, etc.). Formation of Ionic Compounds: S-block elements typically form ionic compounds, like NaCl, CaCl₂, and KBr, due to their strong electropositive character. Flame Colors: They impart characteristic colors to flames—sodium gives yellow, potassium gives lilac, and calcium gives orange-red. Good Conductors: Both groups conduct electricity well in solid and molten states due to the presence of free-moving electrons. These properties make s-block elements essential in chemical industries, biological systems (like sodium and potassium in nerve function), and laboratory applications.

S-block elements exhibit several distinctive properties that stem from their electronic configuration: High Reactivity: Both alkali and alkaline earth metals are highly reactive, especially with water and halogens, due to their tendency to lose electrons easily. Softness and Low Density: Alkali metals are particularly soft (can be cut with a knife) and have low densities—lithium, sodium, and potassium float on water. Low Ionization Energies: The energy required to remove the outermost electron is low, resulting in the easy formation of cations (Na⁺, Mg²⁺, etc.). Formation of Ionic Compounds: S-block elements typically form ionic compounds, like NaCl, CaCl₂, and KBr, due to their strong electropositive character. Flame Colors: They impart characteristic colors to flames—sodium gives yellow, potassium gives lilac, and calcium gives orange-red. Good Conductors: Both groups conduct electricity well in solid and molten states due to the presence of free-moving electrons. These properties make s-block elements essential in chemical industries, biological systems (like sodium and potassium in nerve function), and laboratory applications.

No, hydrogen is not formally included in the s-block of the periodic table, despite having a single electron in its 1s orbital. Hydrogen is placed at the top of Group 1 due to its electronic configuration, but its properties differ significantly from alkali metals.While hydrogen can lose its electron like alkali metals, it also readily gains or shares electrons (as seen with halogens), and forms covalent molecules like H₂ and HCl. Due to these unique characteristics, hydrogen is generally considered separately and not as a true s-block element. Its special status is acknowledged in textbooks and the periodic table, often with a distinct placement above the alkali metals but not officially part of their group.

No, hydrogen is not formally included in the s-block of the periodic table, despite having a single electron in its 1s orbital. Hydrogen is placed at the top of Group 1 due to its electronic configuration, but its properties differ significantly from alkali metals.While hydrogen can lose its electron like alkali metals, it also readily gains or shares electrons (as seen with halogens), and forms covalent molecules like H₂ and HCl. Due to these unique characteristics, hydrogen is generally considered separately and not as a true s-block element. Its special status is acknowledged in textbooks and the periodic table, often with a distinct placement above the alkali metals but not officially part of their group.

Hydrogen’s position in the periodic table is unique and somewhat debated among chemists. Technically, hydrogen is not assigned to any specific block (s, p, d, or f) because, while its electronic configuration (1s¹) suggests s-block, its chemical behavior is quite distinct.In most periodic tables, hydrogen is placed above Group 1 (s-block) to reflect its single valence electron, but it is not considered an alkali metal. Some periodic tables show hydrogen above Group 17 (halogens) due to similarities in electron gain behavior, but again, hydrogen is fundamentally different from halogens as well.Because of this, hydrogen is often shown as a special case, outside of strict block classification, and its unique chemistry is treated as an independent topic in academic studies.

Hydrogen’s position in the periodic table is unique and somewhat debated among chemists. Technically, hydrogen is not assigned to any specific block (s, p, d, or f) because, while its electronic configuration (1s¹) suggests s-block, its chemical behavior is quite distinct.In most periodic tables, hydrogen is placed above Group 1 (s-block) to reflect its single valence electron, but it is not considered an alkali metal. Some periodic tables show hydrogen above Group 17 (halogens) due to similarities in electron gain behavior, but again, hydrogen is fundamentally different from halogens as well.Because of this, hydrogen is often shown as a special case, outside of strict block classification, and its unique chemistry is treated as an independent topic in academic studies.

A hydride is a compound formed when hydrogen combines with another element. S-block hydrides specifically refer to compounds where s-block metals (alkali and alkaline earth metals) react with hydrogen.There are two main types: Ionic (saline) hydrides: Formed by Group 1 and Group 2 metals (except Be and Mg), these are white, crystalline solids containing the hydride ion (H⁻). For example, sodium hydride (NaH) and calcium hydride (CaH₂). These hydrides react vigorously with water, releasing hydrogen gas: NaH + H₂O → NaOH + H₂↑ Covalent hydrides: BeH₂ and MgH₂ form more covalent bonds due to the smaller size and higher charge density of Be²⁺ and Mg²⁺ ions. S-block hydrides are important in organic synthesis, reduction reactions, and as hydrogen storage materials.

A hydride is a compound formed when hydrogen combines with another element. S-block hydrides specifically refer to compounds where s-block metals (alkali and alkaline earth metals) react with hydrogen.There are two main types: Ionic (saline) hydrides: Formed by Group 1 and Group 2 metals (except Be and Mg), these are white, crystalline solids containing the hydride ion (H⁻). For example, sodium hydride (NaH) and calcium hydride (CaH₂). These hydrides react vigorously with water, releasing hydrogen gas: NaH + H₂O → NaOH + H₂↑ Covalent hydrides: BeH₂ and MgH₂ form more covalent bonds due to the smaller size and higher charge density of Be²⁺ and Mg²⁺ ions. S-block hydrides are important in organic synthesis, reduction reactions, and as hydrogen storage materials.

No, s-block elements and hydrogen have not been removed from the JEE syllabus or from standard chemistry curricula. Both topics remain fundamental in chemistry education and competitive exams like JEE Mains and Advanced. These chapters cover key concepts in the periodic table, element trends, and chemical reactivity, and are crucial for building a solid foundation in both inorganic and physical chemistry.However, it’s wise to regularly check the latest official syllabus released by exam authorities or educational boards, as minor updates may occur year-to-year. For now, both topics are essential and should be prepared thoroughly.

No, s-block elements and hydrogen have not been removed from the JEE syllabus or from standard chemistry curricula. Both topics remain fundamental in chemistry education and competitive exams like JEE Mains and Advanced. These chapters cover key concepts in the periodic table, element trends, and chemical reactivity, and are crucial for building a solid foundation in both inorganic and physical chemistry.However, it’s wise to regularly check the latest official syllabus released by exam authorities or educational boards, as minor updates may occur year-to-year. For now, both topics are essential and should be prepared thoroughly.

Hydrogen itself is not an s-orbital; rather, the only electron in a hydrogen atom occupies the 1s orbital. The term “s-orbital” refers to a type of atomic orbital with a spherical shape where electrons are most likely to be found. In the hydrogen atom (atomic number 1), the single electron is found in the 1s orbital, the lowest energy state available.Thus, hydrogen’s electron configuration is written as 1s¹, indicating one electron in the s-orbital of the first energy level. This configuration influences many of hydrogen’s unique chemical properties.

Hydrogen itself is not an s-orbital; rather, the only electron in a hydrogen atom occupies the 1s orbital. The term “s-orbital” refers to a type of atomic orbital with a spherical shape where electrons are most likely to be found. In the hydrogen atom (atomic number 1), the single electron is found in the 1s orbital, the lowest energy state available.Thus, hydrogen’s electron configuration is written as 1s¹, indicating one electron in the s-orbital of the first energy level. This configuration influences many of hydrogen’s unique chemical properties.

Although helium’s electrons occupy the s-orbital (1s² configuration), helium is not included in the s-block. This is because helium’s chemical properties are vastly different from those of s-block elements (Group 1 and 2 metals).Helium is a noble gas, chemically inert due to its completely filled valence shell (stable octet/duplet), and does not participate in the typical reactions of s-block elements. In the periodic table, helium is placed with other noble gases in Group 18(far right), reflecting its unreactive nature rather than its electron configuration.

Although helium’s electrons occupy the s-orbital (1s² configuration), helium is not included in the s-block. This is because helium’s chemical properties are vastly different from those of s-block elements (Group 1 and 2 metals).Helium is a noble gas, chemically inert due to its completely filled valence shell (stable octet/duplet), and does not participate in the typical reactions of s-block elements. In the periodic table, helium is placed with other noble gases in Group 18(far right), reflecting its unreactive nature rather than its electron configuration.

Hydrogen can be categorized in several ways, but one common classification is based on isotopes atoms with the same number of protons but different numbers of neutrons. The three isotopes of hydrogen are: Protium (^1H): The most abundant form, with one proton and no neutrons. Deuterium (^2H or D): Contains one proton and one neutron; used in heavy water (D₂O) and nuclear reactors. Tritium (^3H or T): Contains one proton and two neutrons; radioactive and used in fusion research and certain luminous paints. Another way to classify hydrogen is by its allotropic forms ( atomic , diatomic, nascent), but the isotope classification is most widely accepted in textbooks and exams.

Hydrogen can be categorized in several ways, but one common classification is based on isotopes atoms with the same number of protons but different numbers of neutrons. The three isotopes of hydrogen are: Protium (^1H): The most abundant form, with one proton and no neutrons. Deuterium (^2H or D): Contains one proton and one neutron; used in heavy water (D₂O) and nuclear reactors. Tritium (^3H or T): Contains one proton and two neutrons; radioactive and used in fusion research and certain luminous paints. Another way to classify hydrogen is by its allotropic forms ( atomic , diatomic, nascent), but the isotope classification is most widely accepted in textbooks and exams.

Hydrogen does exist, but the elemental hydrogen atom (H) is extremely reactive and rarely exists in isolation under normal conditions. Because a single hydrogen atom has just one electron and is highly unstable, it pairs with another hydrogen atom to form a stable diatomic molecule (H₂), which is what we encounter in laboratories and nature.Atomic hydrogen can be produced in laboratories or stars (as plasma), but it quickly combines to form molecular hydrogen due to energy stability. So, while hydrogen atoms exist, they are almost always found as H₂ molecules, hydrides, or compounds rather than as free atoms.

Hydrogen does exist, but the elemental hydrogen atom (H) is extremely reactive and rarely exists in isolation under normal conditions. Because a single hydrogen atom has just one electron and is highly unstable, it pairs with another hydrogen atom to form a stable diatomic molecule (H₂), which is what we encounter in laboratories and nature.Atomic hydrogen can be produced in laboratories or stars (as plasma), but it quickly combines to form molecular hydrogen due to energy stability. So, while hydrogen atoms exist, they are almost always found as H₂ molecules, hydrides, or compounds rather than as free atoms.

Hydrogen is an s-block element in terms of its electron configuration (1s¹), meaning its lone electron is in an s-orbital. The term “sp” generally refers to hybridized orbitals found in molecules, not to the elemental classification.So, in the periodic table context, hydrogen is s-block (1s¹), but in molecular structures (like methane, CH₄), hydrogen can bond with atoms that use sp³ hybrid orbitals, though hydrogen itself does not hybridize.

Hydrogen is an s-block element in terms of its electron configuration (1s¹), meaning its lone electron is in an s-orbital. The term “sp” generally refers to hybridized orbitals found in molecules, not to the elemental classification.So, in the periodic table context, hydrogen is s-block (1s¹), but in molecular structures (like methane, CH₄), hydrogen can bond with atoms that use sp³ hybrid orbitals, though hydrogen itself does not hybridize.

There is no 1p orbital because the first principal energy level (n=1) only allows the formation of an s-orbital (1s). The p-orbitals start appearing from the second energy level (n=2) onwards. The rules of quantum mechanics restrict the angular momentum quantum number (l) to values from 0 to (n–1). For n=1, l can only be 0 (s-orbital), so 1p does not exist.This quantum restriction explains why the periodic table starts with hydrogen and helium (1s electrons), and why the p-block begins only from the second period.

There is no 1p orbital because the first principal energy level (n=1) only allows the formation of an s-orbital (1s). The p-orbitals start appearing from the second energy level (n=2) onwards. The rules of quantum mechanics restrict the angular momentum quantum number (l) to values from 0 to (n–1). For n=1, l can only be 0 (s-orbital), so 1p does not exist.This quantum restriction explains why the periodic table starts with hydrogen and helium (1s electrons), and why the p-block begins only from the second period.

SPDF refers to the four types of atomic orbitals (subshells) in which electrons reside: S (sharp) P (principal) D (diffuse) F (fundamental) These names originated from early spectroscopic studies and correspond to different values of the angular momentum quantum number (l = 0 for s, 1 for p, 2 for d, 3 for f). Each type of orbital has a distinct shape and capacity—s (2 electrons), p (6), d (10), and f (14). The arrangement of electrons in these orbitals determines the chemical properties and periodic classification of the elements.

SPDF refers to the four types of atomic orbitals (subshells) in which electrons reside: S (sharp) P (principal) D (diffuse) F (fundamental) These names originated from early spectroscopic studies and correspond to different values of the angular momentum quantum number (l = 0 for s, 1 for p, 2 for d, 3 for f). Each type of orbital has a distinct shape and capacity—s (2 electrons), p (6), d (10), and f (14). The arrangement of electrons in these orbitals determines the chemical properties and periodic classification of the elements.

S-block Elements and Hydrogen Class 11 Notes – Properties, Compounds & Trends

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Fundamental Properties: Understand the core physical and chemical properties, including electronic configuration, density, ionization energy, and flame colouration for both Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals).
Key Chemical Reactions: Learn how s-block elements react with water, oxygen, hydrogen, nitrogen, and halogens, with detailed reaction mechanisms and stability trends.
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Introduction

s-block consists of group 1 (alkali metals) and group 2 (alkaline earth metals) elements.

The elements are

Group I:

s-block

Group I I:

s-block

PHYSICAL PROPERTIES

Electronic Configuration

Group I elements contain only one electron in s-orbital (ns¹) while group II elements contain two electrons in s-orbital or valence shell (ns²).

Density

Increases down the group as atomic mass increases. In group II, density first deceases from Be to Ca then increases.

Melting and Boiling Points

Melting and boiling points are quite low which decreases further with increase in atomic number due to weakening of metallic bonds.

Structural Geometry

At room temperature, all alkali metals possess body centred cubic structure with co-ordination number 8. However, at low temperature lithium forms a hexagonal close packed structure with co-ordination number 12.

Flame Colouration

In case of s-block elements, the electrons easily get excited to higher energy level by small energy provided by the Bunsen flame due to low ionization energy. These electrons emit energy during de-excitation to ground state. The energy emitted is in the visible region of electromagnetic spectrum which imparts colour to the flame. Moving down the group I.E. decreases hence frequency of emitted light increases. Hence, the colour imparted to the flame shows gradual shift from red to violet on moving down the group.

Group I

Element	Colour
Na	Golden yellow
K	Pale violet
Rb	Purple
Cs	Sky blue

Group II

Be and Mg have high I.E., so they do not impart characteristic colour to the flame. However, Ca, Sr and Ba imparts characteristic colour to the flame because of their relative low ionization energies.

Element	Colour
Ca	Brick red
Sr	Crimson red
Ba	Apple green
Ra	Crimson

Ionization Energy

Alkali metals have low I.E. and can easily loose their outermost electron, hence shows the following properties.

(i) Photo electric effect

(ii) Electrical conduction: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

(iii) Reducing character: Na < Rb < Cs < Li

Example: Which of the following is stronger reducing agent in aqueous solution?

(A) Li

(B) Na

(C) K

(D) Rb

Solution: (A).

Lithium has smallest ionic radius due to this high hydration energy has synergistic effect to lose electron in aqueous solution in respective of high ionization potential of the same.

Hydration of Ions

Alkali metal in aqueous solution exists as M⁺ (H₂O)_x. Degree of hydration decreases with ionic size. Hence, Li⁺ ion is most hydrated. More the hydration, bigger is the size of the hydrated ion. Hence, it has lowest ionic mobility.

Example: Which of the alkali metal has polarizing power close to that of magnesium?

(A) Li

(B) Na

(C) K

(D) Rb

Solution: (A).

Metals having identical charge-size ratio.

Chemical Properties

Alkali metals are quite reactive and readily form ionic compounds. However, lithium can form covalent compounds like alkyl lithium (R – Li) due to its high ionization enthalpy. Chemical reactivity of alkaline earth metals is lower than alkali metals. In case of alkaline earth metals, Be mostly form covalent compounds while Mg compounds are partly covalent and partly ionic while those of Ca, Sr and Ba are purely ionic.

Some chemical properties of alkali and alkaline earth metals are discussed below.

Reaction with Water

In case of alkali metals, electropositive character increases on moving down the group. Hence, reactivity follows the order:

Li < Na < K < Rb < Cs

Similarly, in case of alkaline earth metals, reactivity increases moving down the group. Be does not react with water. While Mg reacts with boiling water. Ca, Sr and Ba react even with cold water.

Group I:

reaction with water

Group II:

s-block

Reaction with Oxygen

Group I:

Affinity for oxygen increases down the group. Thus, Li forms Li₂O (lithium oxide) while Na forms peroxide (Na₂O₂). K, Rb and Cs form super oxide (KO₂, RbO₂, CsO₂). As superoxides have unpaired electron so they are paramagnetic in nature and coloured.

Na₂O₂ shows yellow colour due to presence of superoxide NaO₂ as impurity whereas KO₂, RbO₂, and CsO₂are orange, brown and orange respectively. Alkali metals are kept under hydrocarbon solution to prevent their reaction with air and moisture.

Group II:

Similarly, for group II elements monoxides and peroxides are formed.

monoxides and peroxides

Reaction with Hydrogen

Both group I and group II elements form hydrides of molecular formula, MH and MH₂ respectively with hydrogen. Stability of hydrides of group I decreases from LiH to CsH. Be does not react with hydrogen, so BeH₂ is obtained by reducing BeCl₂ with LiAlH₄. CaH₂ is known as hydrolith. All hydrides in water evolve hydrogen gas and act as reducing agent.

Reaction with Nitrogen

In group I, only Li reacts with nitrogen to form lithium nitride (Li₃N).

nitrogen reaction

While all group II elements form nitrides with N₂, which evolve NH₃ gas with water.

nitrogen

Reaction with Halogen

All group I and group II elements form halides of molecular formula, MX and MX₂ respectively. Reactivity of group I elements towards a particular halogen is in the order:

Li < Na < K < Rb < Cs.

While for a particular metal reactivity of halogen is

F₂ > Cl₂ > Br₂ > I₂

LiF is ionic while all other lithium halides have covalent character (Fajan’s rule). Hence, they are soluble in organic solvents. All halides of group II elements are hygroscopic in nature and their solubility decreases in the order: MgX₂ > CaX₂ > SrX₂ > BaX₂

On descending down the group, the size of atom increases so lattice energy and hydration energy both decreases. But hydration energy decrease more rapidly. Hence, solubility goes on decreasing. However, with fluorides reverse happens.

Reaction with Non-metals

Alkali metals react with sulphate and phosphorus on heating to form disulphide and phosphides.

M + S₈ → M₂S (Sulphide)
M + P₄ → M₂P (Phosphide)

Except Be, all other group II elements form carbides of general formula MC₂. Be forms BeC₂ called methanide.

Strength of Hydroxides

Group I element hydroxides are strong bases and their strength increases from LiOH to CsOH due to decreasing bond energy between metal and oxygen.

Similarly, for alkaline earth metal hydroxides, Be(OH)₂ is amphoteric while from Be to Ba basic strength increases.

Solubility Trends

1. Solubility in liquid NH₃: All alkali and alkaline earth metals dissolve in liquid ammonia. Alkali metal gives deep blue solution due to presence of ammoniated (solvated) electrons in solution.

M + (x+y)NH₃ → M⁺(NH₃)_x + e^-(NH₃)_y

These electrons absorb photons of energy falling in red region of spectrum hence appear blue. Similarly, alkaline earth metals are soluble in liquid NH₃. In dilute solution, they give blue colour while in concentrated solution they impart bronze colour due to the formation of metal clusters.

2. Solubility of s-block element hydroxides increases going down the group since lattice energy decreases more rapidly than the hydration energy.

Solubility of sulphate decreases in the order:

BeSO₄ > MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄ for group II elements.

As the size of cations increases, heat of hydration decreases while lattice energy remains almost same (due to large size of sulphate). For group I element sulphates, Li₂SO₄ is insoluble while others are soluble.

3. Similar is the case with carbonates. Hence, trend is

BeCO₃ > MgCO₃ > CaCO₃ > SrCO₃ > BaCO₃

Extremely low solubility of alkaline earth metal carbonates in water is used in the precipitation of Ba²⁺, Sr²⁺ and Ca²⁺ as their carbonates in V^th group of qualitative analysis of basic radicals.

Solubilities of carbonates and bicarbonates increases as we move down the group due to lower lattice energies. Thus, order is LiHCO₃ < NaHCO₃ < KHCO₃ < RbHCO₃ < CsHCO₃

THERMAL STABILITY

Alkali Carbonates and Bicarbonates

Li₂CO₃ is much less stable and decomposes on strong heating to give Li₂O and CO₂

Li₂CO₃ ——^{Red Hot}——> Li₂O + CO₂

Order for stability is:

Li₂CO₃ < Na₂CO₃ < K₂CO₃ < Rb₂CO₃ < Cs₂CO₃

All bicarbonates on heating form carbonates with the evolution of CO₂, e.g.

2NaHCO₃ ——^Δ——> Na₂CO₃ + CO₂ + H₂O

LiHCO₃ does not exist as a solid but exists in solution.

Alkaline earth metal carbonates decompose on heating giving metal oxide and CO₂.

The order is:

BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃

Thermal Stabilities of Sulphates and Nitrates

LiNO₃ decomposes to give NO₂ and O₂ while nitrates of all other alkali metals decomposes on heating to form nitrites and O₂.

thermal stabilities of sulphates and nitrates

Thermal stability of sulphates increases down the group, i.e.

BeSO₄ < MgSO₄ < CaSO₄ < SrSO₄

Anomalous behaviour: Lithium in case of alkali metals and beryllium in alkaline earth metals show anomalous behaviour due to (i) higher electronegativity and (ii) smaller atomic and ionic radii.

COMPOUNDS OF s-BLOCK ELEMENTS

Sodium Carbonate (Na₂CO₃)

This is prepared by Solvay’s process. Excess of CO₂ is bubbled through nearly saturated brine solution containing ammonia. Sodium bicarbonate formed is reacted with sodium chloride. Sodium bicarbonate is filtered off and calcined to give sodium carbonate and CO₂. This is “Solvay’s process”.

sodium carbonate na2co3

Properties

(i) Aqueous solution absorbs CO₂ yielding sparingly soluble sodium bicarbonate.

Na₂CO₃ + H₂CO₃ ⟶ NaHCO₃

(ii) With acids, it liberates CO₂ and with lime it produces caustic soda (NaOH).

Na₂CO₃ + 2HCl ⟶ 2NaCl + H₂O + CO₂

Na₂CO₃ + Ca(OH)₂ ⟶ 2NaOH + CaCO₃

(iii) Sodium carbonate forms sodium silicate with silica.

Na₂CO₃ + SiO₂ ⟶ Na₂SiO₃ + CO₂

Sodium Bicarbonate (NaHCO₃)

It is sparingly soluble in water and on heating it is converted to anhydrous sodium carbonate. On the reverse, it can be prepared from sodium carbonate as follows.

Na₂CO₃ + CO₂ + H₂O ⟶ 2NaHCO₃

Sodium Sulphate (Na₂SO₄)

Anhydrous salts prepared from sodium chloride.

NaCl + H₂SO₄ ⟶ NaHSO₄ + HCl ↑

NaCl + NaHSO₄ ⟶ Na₂SO₄ + HCl ↑

Hydrated salt Na₂SO₄·10H₂O is called Glauber’s salt and is prepared from anhydrous salt by crystallization from water below 32°C. It is a water soluble crystalline salt which forms large monoclinic prism type crystals.

Calcium Oxide or Quick Lime (CaO)

It is made by decomposing limestone at high temperature of about 1000ºC.

Reaction: CaCO₃ →^1000ºC CaO + CO₂

It is a white amorphous powder, which emits intense white light when heated in oxy-hydrogen flame.

It reacts with silica to form easily fusible calcium silicate.

Reaction: CaO + SiO₂ → CaSiO₃

With water it evolves high amount of heat and produces slaked lime.

Reaction: CaO + H₂O → Ca(OH)₂

Calcium Carbonate (CaCO₃)

In nature it occurs as limestone, marble, coral, chalk, etc. The white powder is prepared by dissolving limestone or marble in hydrochloric acid and removing iron and aluminum present by precipitating with NH₃ and then adding ammonium carbonate to the solution. The precipitate is filtered and dried.

Reaction: CaCl₂ + (NH₄)₂CO₃ → CaCO₃ + 2NH₄Cl

It dissolves in water forming bicarbonate.
CaCO₃ + H₂O + CO₂ ⟶ Ca(HCO₃)₂

Bleaching Powder (CaOCl₂)

Preparation

It is prepared by passing chlorine over slaked lime.
3Ca(OH)₂ + 2Cl₂ ⟶ CaOCl₂ + CaCl₂·Ca(OH)₂·H₂O
Bleaching powder

Properties

It reacts with dilute acids to give chlorine known as available chlorine.
CaOCl₂ + 2HCl ⟶ CaCl₂ + H₂O + Cl₂ ↑
CaOCl₂ + H₂SO₄ ⟶ CaSO₄ + H₂O + Cl₂ ↑
With water it decomposes into calcium chloride and calcium hypochlorite.
CaOCl₂ + H₂O ⟶ CaCl₂ + Ca(OCl)₂ + H₂O
It reacts with CO₂ from atmosphere and evolves Cl₂ which acts as a bleaching agent and oxidizing agent.
CaOCl₂ + CO₂ ⟶ CaCO₃ + Cl₂
On heating bleaching powder, it gives a mixture of chlorate and chloride.
6CaOCl₂ ⟶Δ⟶ Ca(ClO₃)₂ + 5CaCl₂

Sodium Carbonate or Washing Soda (Na₂CO₃·10H₂O)

Ammonia – soda process (Solvay process):

When carbon dioxide is bubbled through a brine solution saturated with ammonia it results in the formation of sodium hydrogen carbonate.

NH₃ + H₂O + CO₂ ⟶ NH₄HCO₃

NaCl + NH₄HCO₃ ⟶ NaHCO₃ ↓ + NH₄Cl

2NaHCO₃ ———(ignited)——→ Na₂CO₃ + CO₂ + H₂O

The filtrate after removal of NaHCO₃ contains ammonium salts such as NH₄HCO₃ and NH₄Cl.

 NH₄HCO₃ —Δ→ NH₃ + H₂O + CO₂
 2NH₄Cl + Ca(OH)₂ —Δ→ 2NH₃ + 2H₂O + CaCl₂

Recovery of ammonia

Note: Solvay’s process cannot be used for the manufacture of K₂CO₃ because potassium bicarbonate is soluble in water.

Sodium Hydroxide (NaOH)

It is prepared by the following process.

Caustic process (Gossage process):

In this process a 10 – 22% solution of sodium carbonate is reacted with calcium hydroxide.

Na₂CO₃ + Ca(OH)₂ ⟶ CaCO₃ ↓ + 2NaOH

Electrolytic process:

This process involves electrolysis of an aqueous solution of NaCl.
At anode: 2Cl⁻ ⟶ Cl₂ + 2e⁻
At cathode: 2Na⁺ + 2e⁻ ⟶ 2Na

2Na + 2H₂O ⟶ 2NaOH + H₂

Since, chlorine reacts with NaOH even in cold conditions, therefore, electrolysis is carried out in specially designed electrolytic cell, which avoids contact between Cl₂ and NaOH. The two such cells are:

Nelson cell
Castner–Kellner cell

Magnesium Sulphate or Epsom salt (MgSO₄·7H₂O)

Magnesium sulphate can be prepared from magnesite (MgCO₃) or from kieserite (MgSO₄·H₂O).

MgCO₃ + H₂SO₄ ⟶ MgSO₄ + CO₂ + H₂O

The resulting solution on concentration and cooling gives crystals of MgSO₄·7H₂O. MgSO₄·7H₂O is efflorescent, i.e. it loses water of crystallization on exposure to air.

Plaster of Paris (Hemihydrate of Calcium Sulphates)

It is prepared by heating gypsum (CaSO₄·2H₂O) to 390 K.

2CaSO₄·2H₂O ⟶_{390 K} (CaSO₄)₂·H₂O + 3H₂O

The temperature should not be allowed to rise above 390 K because above this temperature, the water of crystallization is lost. The resulting anhydrous CaSO₄ is called dead burnt plaster because it loses the property of setting with water.

ANOMALOUS BEHAVIOUR OF LITHIUM

Amongst the alkali metals, Li has the smallest size and hence the highest polarizing power. Due to this reason, lithium is different from rest of its family members.

(i) Lithium is harder than any other alkali metal.

(ii) Lithium combines with O₂ to form lithium monoxide whereas other alkali metals form peroxides and superoxides.

(iii) Lithium hydride is most stable than the other alkali metal hydrides.

(iv) Lithium carbonate decomposes on heating to evolve CO₂ whereas other alkali metal carbonates do not.

(v) Lithium when heated with ammonia forms lithium imide (Li₂NH) while other alkali metals form amides of the general formula, MNH₂.

ANOMALOUS BEHAVIOUR OF BERYLLIUM

Beryllium is anomalous in many of its properties and shows a diagonal relationship to aluminium in group – 3

Be is very small and has a high charge density so by Fajans’ rules it should have a strong tendency to covalence. Thus the melting point of its compounds is lower (BeF₂ m.pt. 800°C, rest of group about 1300°C) and they are all soluble in organic solvents and hydrolyses in water rather like the compounds of aluminium.
Be forms many complexes – not typical of Group – 1 and 2.
Be like Al is rendered passive by nitric acid.
Be is amphoteric, liberating H₂ with NaOH and forming beryllates. Al forms aluminates.
Be(OH)₂ like Al(OH)₃ is amphoteric.
The salts are extensively hydrolyzed.
The halides are polymeric, with multicentre bonding associated with electron deficiency. BeCl₂ usually forms chains but also exists as the dimmer, whilst AlCl₃ is dimeric.
The hydrides and alkyls are also polymeric.
Be salts are among the most soluble known.
Be₂C, like Al₄C₃, yields methane of hydrolysis.

HYDROGEN

Hydrogen is the first element in the periodic table and is also the lightest element known. Its atomic form exists only at high temperatures. In the normal element form, it exists as a diatomic molecule, i.e..

Position of Hydrogen in the Periodic Table

Hydrogen, the first element in the periodic table, has the simplest atomic structure of all the elements, and consists of a nucleus of charge + 1 and one orbital electron. The alkali metals also have a single outer orbital electron, but they tend to lose this electron in reactions and form positive ions; by contrast, hydrogen has little tendency to lose this electron but a great tendency to pair the electron and form a covalent bond. The halogens 17, like hydrogen, are one electron short of an inert gas structure. In many reactions the halogens gain an electron and so form negative ions, although hydrogen can only do this in reactions with highly electropositive metals. This behaviour is explained by the atomic structure, the extremely small size of hydrogen atoms and the low electronegativity value. These unique properties make it difficult to place hydrogen in the periodic table, since its properties differ from those of both Group 1 and group 17 elements, and although it is included in both groups in the tables in this book, it could well be put in a group on its own.

Isotopes of Hydrogen

It has been found by mass spectrograph that hydrogen has three isotopes namely; protium, deuterium and tritium. The relative abundance of three isotopes of hydrogen is as under

isotopes of hydrogen

(a) Protium or hydrogen

It is represented by the symbol D or ₁²H. Its atomic number is 1 and mass number is also 1. It has one proton (but no neutron) in its nucleus and one electron in its 1s orbital. Naturally occurring hydrogen contains 99.985% of this isotopes.

(b) Deuterium or heavy hydrogen

It is represented by the symbol D or. It’s atomic number is 1 and mass number is 2. It has one proton and one neutron in its nucleus and one electron in its 1s orbital. Naturally occurring hydrogen has 0.015% of this isotopes mostly in the form of HD.

It is represented by the symbol T or ₁³H. Its nucleus has one proton and 2 neutron and there is one electron in its 1s orbital. It is an extremely rare isotope. Out of 10¹⁷ molecules of ordinary hydrogen there is just one molecule of tritium. This isotope of hydrogen is radioactive in nature and emits low energy $β$ - particles (t_1/2 = 12.33y)

tritium

It may be noted that three isotopes of hydrogen have similar chemical properties because of the same electronic configuration,1s¹. However due to different mass numbers they have different rates of chemical reactions. For example, reaction between protium and chlorine is 13.4 times faster than that between deuterium and chlorine. Similarly electrolysis of ordinary water (H₂O) occurs more rapidly than of heavy water (D₂O).Difference in properties arising due to the difference in mass number is referred to as isotopic effect.

(d) Ortho and para hydrogen

When the spins of the nuclei are in the same direction (parallel spins), dihydrogen is called ortho hydrogen and when the spins are in the opposite direction (anti parallel spins), dihydrogen is called para hydrogen.

Example: The property of hydrogen which distinguishes it from other alkali metals is

(A) its electropositive character

(B) its affinity for non-metals

(C) its reducing character

(D) its non-metallic character

Solution: (D).

When fused NaH is electrolyte hydrogen deposited on anode while sodium on

cathode.

HYDRIDES

Dihydrogen combines with a number of elements to form binary compounds called hydrides. Their general formula being where M represents the element and x the number of hydrogen atoms.

Depending upon the physical and chemical properties, the hydrides have been divided into the following three broad categories:

Ionic or salt-like or saline hydrides
Metallic or Interstitial hydrides
Molecules or Covalent hydrides

Saline hydrides or ionic hydrides

These are binary compounds of hydrogen and elements which are more electropositive than hydrogen such as alkali metals, alkaline earth metals (except Be), etc. Saline hydrides are formed by the transference of electron from metal to hydrogen. Some common examples of this category are: etc.

Covalent hydrides or molecular hydrides

These are binary compounds of hydrogen and elements of comparatively high electronegativity such as p- block elements. In these hydrides, H atoms are bonded to the other atoms by covalent bonds. Some examples of covalent hydrides are, HCl, H₂O, PH₃, NH₃ etc.

Interstitial hydride or metallic hydrides

These are binary compounds of hydrogen and transition elements.

These hydrides are generally formed by the

(a) transition metals of group 3, 4, 5 of d- block;

(b) Cr metal of group 6 and

It may be noted that elements of group 7, 8, 9 of d – block do not form hydrides at all. This inability of metal, of group 7, 8, 9 of periodic table to form hydrides is referred to as hydride gap of d – block.

In these compounds H atoms are supposed to occupy interstitial position in the metal lattices. Some scientists consider these compounds as simply solid solutions of hydrogen. The composition of these hydrides may not correspond to simple whole number ratio and therefore, they are also called non-stoichimotric hydrides. Their composition is also found to vary with the conditions of temperature and pressure. Some examples of interestial hydrides of elements of group 3 to 5 are ScH₂, YH₂, YH₃, LiH₃, CrH, TiH₂, ZrH₂, HfH₂, VH, NbH, NbH₂, TaH etc.

HYDROGEN PEROXIDE (H₂O₂)

Preparation

(i) Lab Method: It is prepared by the action of cold, dilute sulphuric acid on sodium or barium peroxide.

preparatin

Anhydrous barium oxide is not used because the precipitated BaSO₄ forms a protective layer on the unreacted barium peroxide and thus prevent its further participation in the reaction. However, it can be overcome by using phosphoric acid.

(ii) By Electrolysis: It can also be prepared by the hydrolysis of peroxydisulphuric acid which is obtained by the electrolytic oxidation of sulphuric acid.

electrolysis

(iii) It can be prepared by auto-oxidation of 2-ethyl anthraquinol. The net reaction is a catalytic union of H₂and O₂ to yield hydrogen peroxide.

auto oxidation

Structure of hydrogen peroxide

As established by X-ray studies, hydrogen peroxide molecules has a non-planar structure. The molecules dimensions in gas phase and that in solid phase have given in figure. (a) and (b) respectively. In the crystal, the dihedral angle (111.5^o) reduces to on account of hydrogen bonding. The two oxygen atoms are joined by a single electron-pair bond.

The O–O Linkage is also called peroxide linkage.

structure of hydrogen peroxide

Chemical properties

(a) Decomposition

It is a unstable liquid readily decomposes on heating or on long standing to give water and dioxygen. It is an example of disproportionation decomposition is suppressed by addition of glycerol, acetanilide or phosphoric acid.

(b) Acidic behaviour

Pure hydrogen peroxide is a weak acid (at 298 K).

It ionizes in water as:

H₂O₂ → H⁺ + H₂O^- (hydroperoxide)

HO₂^- → H⁺+ O₂^2-(peroxide ion)

Its acidic character can be shown its ability to neutralize bases such as NaOH, Na₂CO₃ etc, to form corresponding peroxides.

2NaOH + H₂O₂ ⟶ Na₂O₂ + 2H₂O

2Na₂CO₃ + H₂O₂ ⟶ Na₂O₂ + H₂O + CO₂

Hydrogen peroxide has an interesting chemistry because of its ability to act as oxidizing as well as reducing agent both in acidic and basic solutions.

OXIDIZING NATURE OF H₂O₂

H₂O₂ can act as oxidizing agents in acidic as well as basic medium as described below.

In acidic medium

H₂O₂ + 2H⁺ + 2e^- ⟶ 2H₂O

In basic medium

H₂O₂ + 2e^- ⟶ 2OH^-

Bleaching action

The bleaching action of hydrogen peroxide is due to the nascent oxygen which it liberates on decomposition.

H₂O₂ ⟶ H₂O + [O]

The nascent oxygen combines with colouring matter which, in turn, gets oxidised. Thus, the bleaching action of H₂O₂ is due to the oxidation of colouring matter by nascent oxygen. It is used for the bleaching of delicate materials like ivory, feather silk, wool etc.

Colouring matter + [O] ⟶ Colourless matter + [O]

Addition reactions

Hydrogen peroxide reacts with alkenes to form glycols.

Strength of H₂O₂ solution:

Strength of H₂O₂ is expressed in terms of volume of oxygen at NTP that one volume of hydrogen peroxide gives on heating.

The commercial sample are marked as 10 volume, 15 volume etc. 10 volume means that one volume of hydrogen peroxide gives 10 volume of oxygen at NTP.

Let, the volume strength is “V”.

2H₂O₂ → 2H₂O + O₂

2(2 + 32) = 68 g 22400 ml at NTP

Then at NTP, “V” litre oxygen will be given by 1 litre of H₂O₂.

g/litre strength of H₂O₂ = (68 / 22.4) × V

N = strength / EM = (68 / 22.4) × (V / 17) = V / 5.6

i.e. volume strength = 5.6 × normality
Similarly, volume strength = 11.2 × molarity

Points to Remember

1. On moving down the group from Li to Cs, electropositivity, atomic radii, atomic volume, reactivity, reducing power, conductivity, solubility of salts with small anion and density show increasing trend.

2. Stability and solubility orders of carbonates, nitrates of alkali metals and bicarbonates are:

Li₂CO₃ < Na₂CO₃ < K₂CO₃ < Rb₂CO₃ < Cs₂CO₃

Li₂CO₃ < NaNO₃ < KNO₃ < RbNO₃ < CsNO₃

LiHCO₃ < NaHCO₃ < KHCO₃ < RbHCO₃ < CsHCO₃

3. Stability of alkali metal peroxides and superoxides are in the order:

N₂O₂ < K₂O₂ < Rb₂O₂ < Cs₂O₂

NaO₂ < KO₂ < RbO₂ < CsO₂

4. Solubility, stability and basic strength of alkali and alkaline earth metal hydroxides increases in the following order:

LiOH < NaOH < RbOH < CsOH

Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂

5. Degree of hydration of alkali metal ions decreases in the order:

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

The relative ionic radii in water also decreases in the same order.

6. Only lithium combines directly with carbon to form lithium carbide, Li₂C₂. Other alkali metals react with ethylene to form corresponding metal carbides.

7. Magnesium ions are present in chlorophyll- a, which is a green pigment in plants and which absorbs light and is essential for photosynthesis.

8. Aqueous Ba(OH)₂ is known as baryta water.

9. Be and Mg crystallize in hcp, Ca and Sr in ccp and Ba in bcc structure.

10. K₂CO₃ cannot be prepared by Solvay process as KHCO₃ is more soluble in water than NaHCO₃.

11. Isotopes of hydrogen: Protium (¹H), Deuterium (²H or D), Tritium (³H or T).

12. Allotropes of hydrogen: orthohydrogen (o-H₂), parahydrogen (p-H₂).

13. Molecular forms: H₂.2H₂ (D₂), 3H₂ (T₂), HD, HT etc.

14. Ortho-hydrogen have parallel nuclear spins and para-hydrogen have antiparallel nuclear spins.

15. Hydrogen under very high pressure is expected to behave like a metal.

16. H₂O₂ is a colourless some what dense, less volatile liquid than water.

17. H₂O₂ (gas phase) dihedral angle 111.5^o.

18. H₂O₂ (solid pahse) dihedral angle 110K is 90.2^o.

19. H₂O₂ is weak acid K_a = 1.55 ´ 10^-¹² at 298 K.

20. H₂O₂ can used as oxidizing and reducing agent in both medium acidic and basic medium.

Frequently Asked Questions

Yes, it is highly recommended to study chemical bonding before delving into s-block elements and hydrogen for IIT JEE preparation. Chemical bonding forms the foundation of understanding how atoms combine to form compounds, and this knowledge is directly applicable when studying the properties and reactions of s-block elements (like sodium, magnesium, etc.) and hydrogen.

When you first learn chemical bonding, you get a strong grasp of electronegativity, bond formation, molecular shapes, hybridization, and electron configuration. These concepts are vital for understanding why s-block elements behave the way they do for instance, why they readily lose electrons to form positive ions or why hydrogen can form both ionic and covalent bonds. Without this base, students often find it challenging to comprehend trends in reactivity, atomic size, ionization energies, and the unique chemistry of hydrogen.

Moreover, chemical bonding acts as a conceptual bridge between atomic structure and periodic classification, making subsequent chapters less overwhelming. Many IIT JEE problems involve the application of chemical bonding principles to real-life chemical reactions or periodic trends. So, start with chemical bonding, master its theories, then move to s-block and hydrogen for smoother, deeper, and more effective learning this strategy significantly boosts retention and exam performance.

Inorganic chemistry can seem daunting due to the volume of facts, trends, and exceptions. However, a smart strategy makes it manageable and even enjoyable for JEE Advanced aspirants. Start by breaking the syllabus into logical sections—periodic table fundamentals, s-block, p-block, d-block, f-block, coordination compounds, and qualitative analysis.

First, ensure you understand periodic trends, chemical bonding, and the logic behind electronic configurations. This helps you see patterns in the properties of elements, rather than rote memorization. Use NCERT textbooks as your primary resource; many questions are either directly from or inspired by NCERT content. Supplement your reading with summary tables and handwritten notes—these help consolidate learning and quick revisions.

Next, focus on solving a variety of problems, including previous years’ JEE questions and mock tests. Practice helps identify frequently tested concepts and sharpens your application skills. Make flashcards for common reactions, color changes, and unique properties (especially in p-block and d-block). Group study or peer discussions can also be useful for clarifying doubts and remembering tricky points.

Finally, regular revision is key. Set aside weekly time slots to revise summary notes and revisit incorrect answers from your practice sets. This continuous reinforcement turns information into long-term knowledge and maximizes your inorganic chemistry scores.

S.no	Formulas List
1.	Atomic Structure
2.	Some Basic Concepts of chemistry
3.	Gaseous State
4.	Periodic Properties
5.	Chemical Bonding
6.	Thermodynamics & Thermochemistry
7.	Chemical Equilibrium
8.	Ionic Equilibrium
9.	Surface Chemistry
10.	S-block elements and Hydrogen

S-block elements and Hydrogen

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