B.Sc. 2nd Year Chemistry – Detailed Notes
StudyWithMe MSN | Standard textbook-style explanation for theory preparation
The periodic table contains elements arranged according to their increasing atomic number. On the basis of the subshell into which the differentiating electron enters, the elements are broadly divided into s-block, p-block, d-block and f-block elements. The s-block and p-block elements together constitute the representative or main-group elements.
The s-block elements are those elements in which the differentiating electron enters the s-subshell of the outermost shell. They mainly occupy Groups 1 and 2 of the modern periodic table. Group 1 elements are known as alkali metals, while Group 2 elements are known as alkaline earth metals.
The p-block elements are those elements in which the differentiating electron enters the p-subshell of the valence shell. They occupy Groups 13 to 18 of the periodic table. The p-block contains metals, metalloids, non-metals and noble gases and therefore shows a very wide range of chemical properties.
The s-block consists mainly of Groups 1 and 2. Their general outer electronic configurations are:
Because of their low number of valence electrons, these elements generally lose electrons easily and form positive ions. Group 1 elements commonly form M+ ions, while Group 2 elements commonly form M2+ ions.
The p-block includes Groups 13 to 18. The general valence-shell configuration is:
The number of valence electrons increases from one in Group 13 to eight in Group 18. Consequently, the p-block displays gradual changes from metallic to non-metallic character across a period.
| Block | Groups | General Configuration | Major Characteristics |
|---|---|---|---|
| s-block | 1–2 | ns1–2 | Highly electropositive, mainly metallic |
| p-block | 13–18 | ns2np1–6 | Contains metals, metalloids, non-metals and noble gases |
The s-block elements possess one or two electrons in their outermost s-orbital. They generally have large atomic size, relatively low ionization energy and high electropositive character.
Most s-block elements are metals. They readily lose their valence electrons and form cations. Their metallic character generally increases down a group because atomic size increases and ionization energy decreases.
s-block elements generally possess comparatively low ionization energies. The outermost electron is relatively weakly held by the nucleus and can therefore be removed easily.
Group 1 elements generally show an oxidation state of +1, while Group 2 elements generally show an oxidation state of +2.
Compounds of s-block metals are generally ionic. Their oxides and hydroxides are usually basic, although some exceptions and trends are observed.
Several s-block metal ions produce characteristic colours in a flame because electrons absorb energy and are excited to higher energy levels. When they return to lower energy levels, radiation of characteristic wavelength is emitted.
| Ion | Characteristic Flame Colour |
|---|---|
| Li+ | Carmine red |
| Na+ | Golden yellow |
| K+ | Violet/lilac |
| Ca2+ | Brick red |
| Sr2+ | Crimson red |
| Ba2+ | Apple green |
Group 1 contains lithium, sodium, potassium, rubidium, caesium and francium. They are called alkali metals because their hydroxides are soluble in water and form strongly alkaline solutions.
The general electronic configuration of alkali metals is ns1. They have only one electron in the valence shell and therefore readily lose it to form M+ ions.
Alkali metals react vigorously with water to form corresponding hydroxides and hydrogen.
The reactivity generally increases down the group. Lithium reacts relatively slowly, whereas potassium and heavier alkali metals react much more vigorously.
Alkali metals react with oxygen to form oxides, peroxides or superoxides depending on the metal.
Alkali metals react readily with halogens to form ionic halides.
Sodium hydroxide is a strong alkali and is commonly called caustic soda. It is widely used in soap manufacture, paper production, petroleum refining and chemical industries.
Sodium carbonate, Na2CO3, is commonly known as washing soda in its hydrated form. It is used in glass manufacture, water softening and chemical industries.
Sodium bicarbonate, NaHCO3, is commonly known as baking soda. It is used in baking, preparation of effervescent mixtures and certain fire extinguishers.
Group 2 consists of beryllium, magnesium, calcium, strontium, barium and radium. These elements are called alkaline earth metals because their oxides and hydroxides are alkaline and many of their compounds occur naturally as minerals.
They generally lose two electrons to form M2+ ions.
The reactivity toward water varies considerably. Beryllium does not react readily with water, magnesium reacts slowly with cold water but more rapidly with hot water, while calcium and heavier members react more readily.
Group 2 metals react with oxygen to form oxides.
Many alkaline earth metals react with dilute acids to produce salts and hydrogen gas.
Magnesium oxide is a basic oxide. It is used in refractory materials because of its high melting point and resistance to heat.
Calcium oxide, CaO, is commonly called quicklime. It is produced by heating limestone.
Calcium oxide reacts with water to form calcium hydroxide.
Calcium hydroxide is commonly called slaked lime. Its aqueous suspension is known as lime water when dilute and clear, while the suspension is commonly referred to as milk of lime.
Calcium carbonate occurs naturally in limestone, marble and chalk. It is an important raw material in cement and lime production.
| Property | Group 1 | Group 2 |
|---|---|---|
| Valence configuration | ns1 | ns2 |
| Common oxidation state | +1 | +2 |
| Atomic size | Generally larger | Generally smaller in the same period |
| Ionization energy | Lower | Higher |
| Hydroxides | Generally strongly basic and soluble | Basicity and solubility vary down the group |
| Nature of compounds | Predominantly ionic | Predominantly ionic, with greater covalent character in smaller cations |
A diagonal relationship refers to the similarity in properties between certain elements belonging to adjacent periods and different groups of the periodic table. The relationship is particularly important for lithium and magnesium, beryllium and aluminium, and boron and silicon.
The similarity arises mainly because movement diagonally down and to the right of the periodic table produces compensating changes in atomic size, charge density, electronegativity and polarizing power.
Hydrogen is the lightest element and has atomic number 1. Its electronic configuration is 1s1. Although it is placed above Group 1 in many periodic tables, hydrogen has properties that distinguish it from both alkali metals and halogens.
| Isotope | Symbol | Important Feature |
|---|---|---|
| Protium | 1H | Most abundant isotope; no neutron |
| Deuterium | 2H or D | Contains one neutron |
| Tritium | 3H or T | Radioactive; contains two neutrons |
Hydrogen can be prepared in the laboratory by the reaction of certain metals with dilute acids.
Hydrogen is a combustible gas. It reacts with oxygen to form water.
Hydrogen also combines with several elements and compounds under suitable conditions. Its reducing property is important in many chemical processes.
The p-block contains elements of Groups 13 to 18. The differentiating electron enters a p-orbital. It is the most diverse block of the periodic table because it contains metals, metalloids, non-metals and noble gases.
Many p-block elements show more than one oxidation state. The difference between the possible oxidation states becomes especially important in heavier elements because of the inert pair effect.
The tendency of the two electrons in the outermost s-orbital to remain non-bonding or unshared in compounds of heavier p-block elements is known as the inert pair effect. It becomes more significant down a group.
Catenation is the ability of an element to form covalent bonds with atoms of the same element, producing chains or rings. Carbon exhibits the most pronounced catenation because of the strong C–C bond and small atomic size.
Smaller p-block elements can form strong multiple bonds such as C=C, C=O, C=N and N=N. The ability to form strong multiple bonds generally decreases for heavier members because effective orbital overlap becomes less efficient.
Group 13 consists of boron, aluminium, gallium, indium and thallium. Their general valence-shell configuration is ns2np1.
Boron is a hard, relatively high-melting metalloid. It differs considerably from the other members of its group because of its small size, relatively high ionization energy and strong covalent bonding.
Aluminium is an important lightweight metal. It has good electrical conductivity, relatively low density and forms a protective oxide layer on exposure to air.
Aluminium oxide, Al2O3, is amphoteric. It can react with both acids and strong bases.
Group 14 contains carbon, silicon, germanium, tin and lead. Their general valence configuration is ns2np2.
The common oxidation states are +4 and +2. The stability of the +2 state increases down the group because of the inert pair effect.
Carbon is a non-metal of exceptional importance. Its ability to form strong covalent bonds with itself and with many other elements gives rise to an enormous number of organic compounds.
Carbon exists in several allotropes. Diamond and graphite are two important crystalline allotropes with very different physical properties.
| Property | Diamond | Graphite |
|---|---|---|
| Structure | Three-dimensional network | Layered structure |
| Hardness | Very hard | Soft and slippery |
| Electrical conductivity | Poor conductor | Good conductor |
| Major use | Cutting tools and jewellery | Electrodes, lubricants and industrial applications |
Silicon is a metalloid and an important semiconductor. It is widely used in electronics, integrated circuits and solar cells.
Tin and lead are metallic members of Group 14. Their compounds show both +2 and +4 oxidation states, although the +2 state is particularly important for the heavier members.
Group 15 consists of nitrogen, phosphorus, arsenic, antimony and bismuth. The general valence configuration is ns2np3.
Common oxidation states include −3, +3 and +5. The stability of the +3 oxidation state increases down the group, especially among the heavier elements.
Nitrogen is a colourless, odourless and relatively unreactive gas under ordinary conditions. Molecular nitrogen contains a strong triple bond.
Ammonia is a colourless gas with a characteristic pungent odour. It is highly soluble in water and forms an alkaline solution.
Phosphorus exists in several allotropes, including white phosphorus and red phosphorus. These allotropes differ in structure, reactivity and physical properties.
Group 16 contains oxygen, sulfur, selenium, tellurium and polonium. Their general valence-shell configuration is ns2np4.
Oxygen is a colourless, odourless gas that supports combustion and is essential for aerobic respiration. Molecular oxygen exists predominantly as O2.
Ozone, O3, is an allotrope of oxygen. It is a strong oxidizing agent and is present in the stratosphere where it absorbs a significant portion of harmful ultraviolet radiation.
Sulfur is a yellow non-metal and occurs in several allotropes. It forms a wide variety of compounds, including hydrogen sulfide, sulfur dioxide and sulfuric acid.
Sulfur dioxide is an important industrial chemical and an acidic oxide. It can act as both an oxidizing and reducing agent depending on the reaction conditions.
Sulfuric acid is one of the most important industrial chemicals. It is widely used in fertilizer manufacture, petroleum refining, chemical synthesis and many other industrial processes.
Group 17 elements are fluorine, chlorine, bromine, iodine and astatine. They are known as halogens because they readily form salts with metals.
Halogens require one additional electron to achieve a stable noble-gas configuration. Therefore, they commonly form X− ions.
| Element | Physical State at Room Temperature |
|---|---|
| Fluorine | Gas |
| Chlorine | Gas |
| Bromine | Liquid |
| Iodine | Solid |
Halogens are strong oxidizing agents because they readily accept electrons. Oxidizing power generally decreases down the group.
Halogens react with hydrogen to form hydrogen halides.
Chlorine is a greenish-yellow gas with a characteristic irritating odour. It is a strong oxidizing agent and has important applications in water treatment and chemical manufacture.
Group 18 contains helium, neon, argon, krypton, xenon and radon. These elements possess stable valence-shell configurations and are characterized by very low chemical reactivity compared with most other elements.
Except helium, which has a stable 1s2 configuration, the noble gases have the general outer configuration:
Noble gases were historically considered completely inert. However, compounds of some heavier noble gases, particularly xenon, are known. Their formation is associated with the ability of heavier noble gases to participate in bonding under suitable conditions.
| Noble Gas | Important Uses |
|---|---|
| Helium | Cooling systems, balloons and specialized applications |
| Neon | Advertising signs and lighting |
| Argon | Inert atmosphere and electric lamps |
| Krypton | Specialized lighting |
| Xenon | Special lamps and specialized technological applications |
Atomic radius generally increases down a group because additional electron shells are added. Across a period, atomic radius generally decreases from left to right because effective nuclear charge increases.
Ionization energy generally decreases down a group because the outer electron becomes farther from the nucleus and is more effectively shielded. Across a period, ionization energy generally increases.
Electronegativity generally decreases down a group and increases across a period. Fluorine is the most electronegative element.
Metallic character increases down a group and generally decreases from left to right across a period.
Across a period, oxides generally change from basic through amphoteric to acidic. This trend reflects the gradual change from metallic to non-metallic character.
| Property | Across a Period | Down a Group |
|---|---|---|
| Atomic radius | Generally decreases | Increases |
| Ionization energy | Generally increases | Decreases |
| Electronegativity | Generally increases | Decreases |
| Metallic character | Generally decreases | Increases |
s- and p-block elements are extremely important in biological systems, agriculture, medicine, industry, environmental science and everyday life. Their compounds form essential materials and participate in numerous chemical and biochemical processes.
Nitrogen, phosphorus and potassium are major plant nutrients. Fertilizers containing these elements are essential for maintaining soil fertility and supporting plant growth.
The s- and p-block elements represent a major portion of the periodic table and include elements with highly diverse physical and chemical properties. The s-block mainly contains electropositive metals that readily form positive ions, while the p-block includes metals, metalloids, non-metals and noble gases.
Understanding electronic configuration is essential for explaining the periodic trends and chemical behaviour of these elements. Properties such as atomic radius, ionization energy, electronegativity, metallic character, oxidation state and reactivity show systematic variations across periods and down groups.
Important concepts such as diagonal relationship, inert pair effect, catenation, allotropy and variable oxidation states provide a deeper understanding of the chemistry of main-group elements. Because these elements are involved in biological processes, agriculture, industry and environmental systems, their chemistry has both theoretical and practical significance.