Notes On Classification Of Elements In s,p, d, f Blocks - CBSE Class 11 Chemistry
The position of an element in the periodic table is based on its valence shell electronic configuration. Based on the type of atomic orbitals which receives the last electron or differentiating electron the elements are categorized as s-, p-, d- and f-block elements. There are two exceptions to this classification. The first exception is Helium. Based on its electronic configuration, it should be placed in the s-block in the periodic table. But it is placed in the p-block with the other Group 18 elements. This is because, like other Group 18 elements, helium has a completely filled valence shell and as a result exhibits characteristic properties of other noble gases. The other exception is Hydrogen which has only one electron in the s-orbital; and therefore can be placed in Group 1 with other alkali metals. However, hydrogen can also gain an electron to achieve the nearest noble gas configuration of Helium. And hence can be placed along with halogens in group 17. Therefore, it is considered as a special case and is placed separately on the top of the periodic table. The s-block of the periodic table consists of elements that have the last electrons entered the ns orbital. The alkali metals in Group 1 and the alkaline earth metals in Group 2 of the long form of the periodic table belong to this category. These are soft metals with low melting and boiling points with a valence shell electronic configuration of ns¹ and ns² respectively, these elements have extremely low ionization enthalpies, they easily lose their outermost electrons to form uni-positive, ions in the case of alkali metals or dipositive ions in the case of alkaline earth metals. Therefore, the compounds formed by the s-block elements with the exception of those formed by lithium and beryllium are ionic in nature. Due to low ionization enthalpies, s-block elements are highly reactive and therefore, are never found in pure form in nature. They always occur in the combined form in nature. Ex: Magnesium occurs in the combined state as carbonates, sulphates, chlorides and silicates but never as pure magnesium.                                            Minerals of Magnesium   Carbonate Magnesite MgCO₃ Dolomite MgCO₃. CaCO₃     Sulphate Epsam salt MgSO₄.7H₂O Kieserite MgSO₄. H₂O     Chloride Carnalite MgCl₂. KCl. 6H₂O     Silicate Asbestos Mg₃(Si₄O₁₀)(OH)₄ Talc Mg₃(Si₄O₁₀)(OH)₂ The metallic character and chemical reactivity of s-block elements increase as we go down the group. The elements of Group 13 to 18, situated at the right side in the long form of periodic table belong to p- block. The p-block elements together with the s-block elements are called the representative elements or main group elements. The p-block elements have their last electron entering the p orbital; this result in an outermost electronic configuration varying from ns² np¹ to ns² np⁶. At the end of each period is a noble gas element with a stable configuration of ns² np⁶. The p- block consists of metals, non-metals and metalloids. The nonmetallic character increases across a period and the metallic character increases as we go down the group. The p-block elements form compounds by loss or gain of electrons, or by sharing the valence electrons. So they can form both ionic as well as covalent bonds. Two chemically important groups of p-block elements are Group-16 (chalcogens) and Group -17 (halogens). The elements of these two groups have high negative gain enthalpies and they readily gain one or two electrons to attain the stable noble gas electronic configuration which is ns2 np6. Ex: (i) Chlorine, belonging to the halogen family, readily accepts one electron to attain the nearest inert gas configuration of argon. (ii) Oxygen belonging to the chalcogens family needs to accept two electrons to attain the inert gas configuration of neon.. The group 18 elements having the electronic configuration ns² np⁶ are called the noble or inert gases. All of them exist as monatomic gases. Noble gases exhibit very low chemical reactivity because of their completely filled outermost shell. The elements belonging to Group 3 to 12 in the periodic table are called the d—block elements. In these elements the differentiating electron or the last electron enters into the penultimate or second outermost shell (n-1) d orbital. These elements have an outermost electronic configuration of (n-1) d¹⁻¹⁰ ns¹⁻². The d-block elements are placed between the s- and p- block elements. They exhibit the transitional or intermediate behavior between the more active metals in s-block and less active elements in the p-block and are hence, called transition elements. Small size of atom, high nuclear charge and unpaired electrons in d orbitals impart characteristic properties to the transition elements. These properties are, 'all the transition elements are metals' and 'they exhibit variable oxidation states'. Ex: Iron can form compounds in +2 (Ferrous ion) and +3 (Ferric ion). Due to their ability to change oxidation states, most of these transition metals and many of their compounds are used as catalysts in many industrial processes. Ex: Iron is used as a catalyst in the Haber process for the synthesis of ammonia. Most transition metals form colored compounds. Ex: MnCl₂ is a pink colored compound and ferrous sulphate (FeSO₄) is a green colored compound. These elements and their ions are usually paramagnetic (attracted to external magnetic fields) because of the presence of unpaired electrons in their d orbitals. Ex: The elements Iron and Nickel and their ions Fe⁺³ and Ni⁺² exhibit paramagnetism. At the bottom of the long form of periodic table, there are two rows of elements that belong to the f-block. These elements have their last electron entering the f- orbitals and exhibit an outer electronic configuration of (n-2) f¹⁻¹⁴ (n-1) d⁰⁻¹ns². Since f-orbitals are present in the ante-penultimate shell or the shell inner to the penultimate (n-2) shell, these elements are also called inner-transition elements. There are two series of f-block elements with 14 elements each. The elements in which the differentiating or the last electron enters into the 4f-orbital are called lanthanides and the elements in which the last electron enters into 5f -orbital are called actinides. These are all metals with high melting and boiling points. Similar to transition metals the inner transition elements also forms colored compounds and exhibits variable oxidation states. The most common oxidation state in these elements is 3+. Actinides can exhibit a much larger number of oxidation states as compared to lanthanides. Actinides are radioactive in nature. Very few of them like thorium, uranium occur in nature and most of them are artificially made by nuclear reactions. The elements after uranium are called trans-uranium elements. Apart from classifying elements into s-, p-, d- and f-blocks, the periodic table also shows the classification of elements into metals, non-metals, and metalloids. Metals have been listed on the left-side of the table and form 78% of the known elements. Metals generally exist as solids at room temperature. An exception is mercury, which is liquid at room temperature. Metals usually have high melting and boiling points. Gallium and Cesium are exception, as they have very low melting points of 303K and 302K. Metals are good conductors of heat and electricity. Metals are malleable (can be flattened into thin sheets) and are ductile (can be drawn into wires). Non-metals are usually solids or gases at room temperature with low melting and boiling points. Boron and Carbon are exceptions which have high melting point of 2300 and 3930 degree Celsius respectively. And Bromine is the only non-metal that exists in liquid state at room temperature. Non-metals are poor conductors of heat and electricity. Solid non-metals are brittle, non-malleable, and non-ductile. Elements like Silicon, Germanium, Arsenic, Antimony and Tellurium in the p-block of the periodic table exhibit properties of both metals and non-metals and therefore, are called semimetals or metalloids.

#### Summary

The position of an element in the periodic table is based on its valence shell electronic configuration. Based on the type of atomic orbitals which receives the last electron or differentiating electron the elements are categorized as s-, p-, d- and f-block elements. There are two exceptions to this classification. The first exception is Helium. Based on its electronic configuration, it should be placed in the s-block in the periodic table. But it is placed in the p-block with the other Group 18 elements. This is because, like other Group 18 elements, helium has a completely filled valence shell and as a result exhibits characteristic properties of other noble gases. The other exception is Hydrogen which has only one electron in the s-orbital; and therefore can be placed in Group 1 with other alkali metals. However, hydrogen can also gain an electron to achieve the nearest noble gas configuration of Helium. And hence can be placed along with halogens in group 17. Therefore, it is considered as a special case and is placed separately on the top of the periodic table. The s-block of the periodic table consists of elements that have the last electrons entered the ns orbital. The alkali metals in Group 1 and the alkaline earth metals in Group 2 of the long form of the periodic table belong to this category. These are soft metals with low melting and boiling points with a valence shell electronic configuration of ns¹ and ns² respectively, these elements have extremely low ionization enthalpies, they easily lose their outermost electrons to form uni-positive, ions in the case of alkali metals or dipositive ions in the case of alkaline earth metals. Therefore, the compounds formed by the s-block elements with the exception of those formed by lithium and beryllium are ionic in nature. Due to low ionization enthalpies, s-block elements are highly reactive and therefore, are never found in pure form in nature. They always occur in the combined form in nature. Ex: Magnesium occurs in the combined state as carbonates, sulphates, chlorides and silicates but never as pure magnesium.                                            Minerals of Magnesium   Carbonate Magnesite MgCO₃ Dolomite MgCO₃. CaCO₃     Sulphate Epsam salt MgSO₄.7H₂O Kieserite MgSO₄. H₂O     Chloride Carnalite MgCl₂. KCl. 6H₂O     Silicate Asbestos Mg₃(Si₄O₁₀)(OH)₄ Talc Mg₃(Si₄O₁₀)(OH)₂ The metallic character and chemical reactivity of s-block elements increase as we go down the group. The elements of Group 13 to 18, situated at the right side in the long form of periodic table belong to p- block. The p-block elements together with the s-block elements are called the representative elements or main group elements. The p-block elements have their last electron entering the p orbital; this result in an outermost electronic configuration varying from ns² np¹ to ns² np⁶. At the end of each period is a noble gas element with a stable configuration of ns² np⁶. The p- block consists of metals, non-metals and metalloids. The nonmetallic character increases across a period and the metallic character increases as we go down the group. The p-block elements form compounds by loss or gain of electrons, or by sharing the valence electrons. So they can form both ionic as well as covalent bonds. Two chemically important groups of p-block elements are Group-16 (chalcogens) and Group -17 (halogens). The elements of these two groups have high negative gain enthalpies and they readily gain one or two electrons to attain the stable noble gas electronic configuration which is ns2 np6. Ex: (i) Chlorine, belonging to the halogen family, readily accepts one electron to attain the nearest inert gas configuration of argon. (ii) Oxygen belonging to the chalcogens family needs to accept two electrons to attain the inert gas configuration of neon.. The group 18 elements having the electronic configuration ns² np⁶ are called the noble or inert gases. All of them exist as monatomic gases. Noble gases exhibit very low chemical reactivity because of their completely filled outermost shell. The elements belonging to Group 3 to 12 in the periodic table are called the d—block elements. In these elements the differentiating electron or the last electron enters into the penultimate or second outermost shell (n-1) d orbital. These elements have an outermost electronic configuration of (n-1) d¹⁻¹⁰ ns¹⁻². The d-block elements are placed between the s- and p- block elements. They exhibit the transitional or intermediate behavior between the more active metals in s-block and less active elements in the p-block and are hence, called transition elements. Small size of atom, high nuclear charge and unpaired electrons in d orbitals impart characteristic properties to the transition elements. These properties are, 'all the transition elements are metals' and 'they exhibit variable oxidation states'. Ex: Iron can form compounds in +2 (Ferrous ion) and +3 (Ferric ion). Due to their ability to change oxidation states, most of these transition metals and many of their compounds are used as catalysts in many industrial processes. Ex: Iron is used as a catalyst in the Haber process for the synthesis of ammonia. Most transition metals form colored compounds. Ex: MnCl₂ is a pink colored compound and ferrous sulphate (FeSO₄) is a green colored compound. These elements and their ions are usually paramagnetic (attracted to external magnetic fields) because of the presence of unpaired electrons in their d orbitals. Ex: The elements Iron and Nickel and their ions Fe⁺³ and Ni⁺² exhibit paramagnetism. At the bottom of the long form of periodic table, there are two rows of elements that belong to the f-block. These elements have their last electron entering the f- orbitals and exhibit an outer electronic configuration of (n-2) f¹⁻¹⁴ (n-1) d⁰⁻¹ns². Since f-orbitals are present in the ante-penultimate shell or the shell inner to the penultimate (n-2) shell, these elements are also called inner-transition elements. There are two series of f-block elements with 14 elements each. The elements in which the differentiating or the last electron enters into the 4f-orbital are called lanthanides and the elements in which the last electron enters into 5f -orbital are called actinides. These are all metals with high melting and boiling points. Similar to transition metals the inner transition elements also forms colored compounds and exhibits variable oxidation states. The most common oxidation state in these elements is 3+. Actinides can exhibit a much larger number of oxidation states as compared to lanthanides. Actinides are radioactive in nature. Very few of them like thorium, uranium occur in nature and most of them are artificially made by nuclear reactions. The elements after uranium are called trans-uranium elements. Apart from classifying elements into s-, p-, d- and f-blocks, the periodic table also shows the classification of elements into metals, non-metals, and metalloids. Metals have been listed on the left-side of the table and form 78% of the known elements. Metals generally exist as solids at room temperature. An exception is mercury, which is liquid at room temperature. Metals usually have high melting and boiling points. Gallium and Cesium are exception, as they have very low melting points of 303K and 302K. Metals are good conductors of heat and electricity. Metals are malleable (can be flattened into thin sheets) and are ductile (can be drawn into wires). Non-metals are usually solids or gases at room temperature with low melting and boiling points. Boron and Carbon are exceptions which have high melting point of 2300 and 3930 degree Celsius respectively. And Bromine is the only non-metal that exists in liquid state at room temperature. Non-metals are poor conductors of heat and electricity. Solid non-metals are brittle, non-malleable, and non-ductile. Elements like Silicon, Germanium, Arsenic, Antimony and Tellurium in the p-block of the periodic table exhibit properties of both metals and non-metals and therefore, are called semimetals or metalloids.

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