CBSE Class 12 Chemistry Revision Notes Chapter 7
Class 12 Chemistry Chapter 7 Notes – The p-block elements
The p-Block elements are present on the right-hand side of the periodic table and include elements like Nitrogen, boron, Fluorine, Oxygen, nitrogen families, etc., along with noble gases, except helium. Group 15 elements called the Nitrogen family include nitrogen phosphorus, arsenic, antimony and bismuth elements. The p-block elements are also the Representative Elements, which are placed on the right side of the primary periodic table.
This chapter p-block elements of Class 12 form an essential topic of the CBSE curriculum, and students must thoroughly prepare to score high in exams.
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The p block elements Chapter 7 Chemistry Class 12 Notes will help students understand the core chapter concepts of the p- block elements like Nitrogen and Oxygen in detail. Furthermore, it discusses the importance of these elements in our lives. Moreover, you will also get a hold of other elements present in the p-block. Other than that, it explains their importance and their properties. Thus, p block elements Class 12 Chemistry Chapter 7 Notes will come in handy while learning the same.
Key Topics Covered in Class 12 Chemistry Chapter 7 Notes
The modern periodic table, as conceived by scientist Dimitri Mendeleev, arranges all the elements known to man based on their atomic number, which is unique to every aspect. The conclusion of such an arrangement was the modern periodic table. The elements with the same properties were arranged into a column called a group.
P block elements in the periodic table, p block elements are placed in Groups 13 to 18, with a general valence shell electronic configuration of ns2 np1-6. Properties of p block elements are influenced by variation in atomic size, ionisation energy, electron gain enthalpy and electronegativity. The first element of every group shows anomalous behaviour from the other group members.
Group 15 elements
Nitrogen phosphorus, arsenic, antimony, bismuth and moscovium are all members of Group 15. As we move down the group, we transition from non-metallic to metallic through a metalloid property. Non-metal elements are Nitrogen and phosphorus, metalloids are arsenic, antimony, and bismuth and moscovium are typical metals.
The standard valence shell electronic configuration of group fifteen elements is ns2, np1-6. The electronic structure of helium is 1s2. It has no orbitals. It is a p-block element; hence it takes the physical and chemical properties after that of different p-block elements of the eighteenth group. The p-block elements are generally non-metals, while the remaining are metalloids and metals.
The Group 15 elements include Nitrogen, phosphorus, arsenic, antimony and bismuth. Nitrogen is the actual constituent of the atmosphere and records 78% of it by volume.
It is the ordinary member of this group and happens in a free state as a diatomic Nitrogen gas. It occurs as sodium nitrate, NaNO3 (chile saltpetre) and potassium nitrate(Indian saltpetre).
It occurs in the form of proteins in plants and animals.
Phosphorus is present in minerals of the apatite family, Ca9(PO4)6, CaX2(X= F, Cl or OH) (fluorapatite Ca9(PO4)6, CaF2), which are the essential components of phosphate rocks. Phosphorus is the main constituent of animal and plant matter. Phosphate groups are constituents of nucleic acids, DNA & RNA. Approximately 60% of bones and teeth are made out of phosphates. Phosphoproteins are present in egg yolk, bone marrow and milk. The rest of the elements in the group, which are arsenic, antimony, and bismuth, mostly happen as sulphide minerals. For example, stibnite, arsenopyrite, and bismuth glance.
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Periodic Trends are shown in Group 15 Elements.
In Group 15 elements, as you would go down a group, starting with the lightest element and ending with the heavy ones, you’d notice a general flow in properties as you move down the order. E.g., Nitrogen is a gas and non-metal, but as you go down the group, we encounter metalloids and then, at the down, metal, i.e. bismuth. These trends shown in the periodic table help us better understand the behaviour of atoms and also allow us to predict new elements. Students may refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of group 15 elements.
Some modern periodic table trends regarding group 15 elements of the properties of the p-Block aspects are discussed below.
Image name: Atomic and Physical properties of the group 15 elements
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Electronic Configuration: The valence shell electron configuration plays a significant role in how an element reacts. The valence electron shell electronic configuration of group 15 elements is ns2np3. All the group fifteen elements have the same arrangement, which is why they’re similar. The s-orbital in this group is filled, and the p-orbitals are half-filled, making their electronic configuration extra stable.
Atomic and Ionic Radii: When you see the electronic configuration of elements in the periodic table above, you will observe that with every step you move downwards, new orbitals are added to the atom. This addition of new orbitals increases the Atomic and Ionic radii of group 15 elements. There is a considerable enhancement in covalent radius from Nitrogen to phosphorus. However, from Arsenic to bismuth, only a tiny increase in ionic radius is observed; this is the reason to fill d and f orbitals in heavier members.
Ionisation Enthalpy: Ionization Energy is the amount of energy required to remove an electron from the outermost orbit of the atom. Ionisation enthalpy is usually measured by how hard the nucleus holds on to the electron. The closer an electron is to the nucleus, the stronger it binds; thus, the higher energy is required. As we go down the group, the radius of the atom enhances; therefore, the Ionization energy decreases due to the weaker bind of the nucleus.
Ionisation enthalpy decreases as we lower the group due to a gradual increase in atomic size, Because of the different stable half-filled p orbitals electronic configuration and smaller size. Also, the ionisation enthalpy of group 15 elements is much higher than that of group 14 elements in a particular period.
Electronegativity: The electronegativity value decreases as the group increases. Atomic size goes down because the enhancing distance between the nucleus and the valence shell goes down the group. However, among the heavier elements, the difference is not that pronounced.
Physical Properties: All the group elements exist in a polyatomic state. First, Nitrogen is a diatomic gas, but as you go down, there is a significant increase in the metallic character of the elements. Nitrogen and phosphorus elements are non-metals, arsenic and antimony are metalloids, and bismuth is a metal. These physical and chemical changes can be attributed to the decrease in Ionisation enthalpy and increase in atomic size. The boiling points also, in general, show an increasing trend as you move down in the group, but the melting point increases up to arsenic and then decreases up to bismuth. Except for nitrogen, all the other elements show allotropy. Students may refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of this topic.
Allotropy: All group fifteen elements, aside from bismuth, indicate allotropy. Nitrogen is found in two allotropic structures, that is, alpha nitrogen and beta nitrogen. Phosphorus exists in numerous allotropic structures. The two critical allotropic structures are red phosphorus and white phosphorus.
Arsenic exists in three essential allotropic structures – black, grey, and yellow. Antimony has three basic allotropic structures: yellow, metallic and explosive.
The topic of allotropy has been further elaborated in our Class 12 Chemistry Chapter 7 Notes. To enjoy the maximum benefit of these resources, students may register themselves on Extramarks website and get access to these study resources.
Chemical Properties: The valence shells of the p block elements have a general electronic configuration of ns2 np3. So the elements here can either release 5 electrons or gain 3. The standard oxidation states of these elements are -3, +3 and +5. With a decrease in the Ionisation enthalpy and electronegativity because of the increasing atomic radius, the tendency to accept three electrons to create a -3 oxidation state decreases down the group. bismuth hardly makes any compounds with a -3 oxidation state as we go down the group; the stability of the oxidation state +5 also decreases, and that of +3 increases due to the inert pair effect.
Nitrogen also exhibits a +1, +2, +4 oxidation state when it interacts with Oxygen. The phosphorus element also shows +1 and +4 oxidation states in some oxoacids.
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Covalency: Nitrogen shows higher covalency of four, as only four orbitals (one in s orbital and 3 in p orbital) are available for bonding. The heavier element with a vacant d orbital can show a covalency of more than four. Many group 15 elements form hydrides of the type EH3. All group 15 elements include oxides of type E2O3 and E2O5. The oxides in the higher oxidation state are more acidic than in the lower oxidation state. Details of group 15 react to form halides of type EX3 and EX5. All these elements react with metals to create binary compounds exhibiting a -3 oxidation state, such as Ca3N2 (calcium nitride).
Note: Nitrogen does not form compounds in +5 oxidation states because It has no vacant d orbital; therefore, no excitation can occur. So, the maximum covalence shown by it is 4.
Class 12 Chemistry Chapter 7 also defines the apatite family. It can be described as a group of similar isomorphous hexagonal phosphate minerals. The leading apatite group is composed of Fluorapatite, Chlorapatite, and Hydroxylapatite. The teeth and bones of several animals, including humans, consist of Calcium phosphate, a similar material shown in the apatite family.
The primary Apatite family group incorporates Fluorapatite, Chlorapatite, and Hydroxylapatite. The extended apatite supergroup describes additional minerals like Pyromorphite, Mimetite, and Vanadinite.
Apatite is the primary source of phosphorus, an essential nutrient required by plants. As such, apatite is the critical ingredient in phosphate fertilisers. Most of the phosphorus used in fertiliser comes from phosphate rock, which is mined almost exclusively for this application.
Anomalous behaviour of Nitrogen element
Nitrogen is distinguished from the additional members of this group 15 by its smaller size, high electronegativity, maximum ionisation enthalpy, and lack of d orbitals. Nitrogen can create p𝝅–p𝝅 multiple bonds with itself and other elements having a more diminutive size, electronegativity-rich elements (e.g. C, O). Because the atomic orbitals of more prominent elements in this group are too large and diffuse that they cannot have overlapping effectively, they do not form p𝝅 – p𝝅 bonds. Hence, Nitrogen is a diatomic molecule having a triple bond between the two atoms (one s and two p orbital). In conclusion, the bond enthalpy (941.1 KJ mol−1) is very high. Due to the significant interelectronic repulsion of the non-bonding electrons and the short bond length, the single N – N bond is a weaker bond than the single P𝝅 – P𝝅 bond connection.
In conclusion, Nitrogen, the catenation tendency, is delicate. The absence of d orbitals in Nitrogen’s electron valence shell is another aspect that impacts its modern Chemistry. Since its covalency is limited to four, it cannot form d𝝅 – p𝝅 bonds like the heavier elements, e.g., R3P = 0 or R3P = CH2(R= alkyl group). When their compounds like triethyl phosphorus P(C2H5)3 and triphenyl arsenic As(C6H5)3 behave as ligands, phosphorus and arsenic can create d𝝅 – p𝝅 bonds with transition metals extremely high. On the other hand, phosphorus, arsenic, and antimony elements make metallic connections in their elementals.
The Class 12 Chemistry Chapter 7 Notes compiled by subject experts at Extramarks will help students understand this topic in more detail.It ensures superior knowledge and understanding of the subject because it provides all the information and is complete in every way. Students need not look for answers elsewhere.That itself reduces stress and anxiety to a great extent.
- a) Reactivity Towards Hydrogen: All the elements in Group 15 produce EH3 hydrides, where E = N, P, As, Sb or Bi.
Image Name: Properties of Hydrides of Group 15 elements
The table exhibits some of the features of these hydrides. The characteristics of the hydrides appear to be a steady progression.
The bond dissociation enthalpy of hydrides reduces from NH3 to BiH3, showing that their stability decreases. In conclusion, the hydrides’ reducing nature improves. Ammonia is only a mild reducing agent, but BiH3 is the most vital reducing agent among all the hydrides. The order of basicity also declines. Extramarks Class 12 Chemistry Chapter 7 study material is accessible on the Extramarks website for students who need to clarify their concept of any topic, formulas, exercises etc. This will help them to brush up their knowledge and strengthen their basics to stay ahead of the competition. .
NH3 > PH3 > AsH3 > SbH3 ≥ BiH3
- b) Reactivity Towards Oxygen
E2O3 and E2O5 are the two ways these elements form oxides. The oxide in the element’s maximum oxidation state is more acidic than the oxide in the lower oxidation state. Their acidity character lowers as they progress through the group. Nitrogen and phosphorus oxides are completely acidic, arsenic and antimony oxides are amphoteric, and bismuth oxides are primarily basic.
- c) Reactivity Towards Halogens
These elements combine to create two halide series, EX3 and EX5. Because the d orbitals in Nitrogen’s valence electron shell are not available, it does not form pentahalide. Pentahalides have a maximum covalent bonding strength than trihalides. Excluding Nitrogen’s trihalides, all these elements’ trihalides are stable. Only NF3 is known to be stable in the situation of Nitrogen. Except for BiF3, trihalides are generally covalent.
- d) Reactivity Towards Metals
All these elements combine with metals to generate binary compounds having an oxidation state of –3, like Ca3N2(Calcium nitride), Ca3P2(calcium phosphide), Na3As2 (sodium arsenide), Zn3Sb2 (zinc antimonide), and Mg3Bi2 (magnesium bismuthide). Students may refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of this topic. As reiterated earlier, our subject matter experts have given a step-by-step explanation for each concept and elaborated with examples, exercises and solutions, additional information, practice exercises etc., to make learning easier and engaging for students.
What is Dinitrogen?
Dinitrogen exhibits almost 78% of the earth’s atmosphere. It is the most valuable element present in the atmosphere.
Dinitrogen is the seventh most abundant uncombined element present in the world. Scientist physician Daniel Rutherford discovered dinitrogen back in the year 1772. The symbol of this chemical element, Nitrogen, is N, and its atomic number is 7.
Preparation of Dinitrogen- We can obtain Nitrogen commercially by liquefaction and fractional distillation of air. This method mainly involves two ways:
- i) Dinitrogen is produced commercially by the liquefaction and purification of air. The liquid dinitrogen molecule (b.p. 77.2 K) distils out first, leaving behind liquid Oxygen (boiling point is 90
We must reduce air to liquid by applying high pressure between 100 to 200 atm. After this, we pass this compressed air via a fine jet where it undergoes expansion. We repeat this process several times, which results in the formation of liquid air.
- ii) The liquid formed undergoes fractional distillation. The boiling point of the dinitrogen molecule is lower than that of the liquid Oxygen, and thus it distils out, leaving behind liquid Oxygen. We obtain Nitrogen from the contaminated liquid. In the laboratory, dinitrogen is manufactured by treating an aqueous ammonium chloride solution with sodium nitrite.
NH4Cl aq) + NaNO2(aq) → N2(g)+ 2H2O + NaCl(aq)
The product obtained consists of impurities such as NO and HNO3 nitric acid, which can be separated by the thermal decomposition of ammonium dichromate. Another process in which we can remove the impurities is passing the gaseous mixture through sulphuric acid with potassium dichromate. The thermal decomposition of ammonium dichromate can also obtain it.
(NH4)2Cr2O7 → N2+ 4H2O+ Cr2O3
Thermal Decomposition of sodium or barium azide in the presence of high temperature also results in the formation of pure Nitrogen. Extramarks Class 12 Chemistry Chapter 7 notes are accessible on the Extramarks website for students who want to learn more. Notes are provided here in accordance with the NCERT textbook and CBSE guidelines and curriculum
Ba(N3)2 Ba + 3N2
Physical Properties of Dinitrogen
As we look at the physical properties of dinitrogen.
Dinitrogen is colourless, odourless, tasteless and diamagnetic. It is a non-toxic gas. It is partially soluble in water. Nitrogen undergoes condensation to create a colourless liquid. Solidification results in the formation of snow-like mass. It has two fixed isotopes, N14 and N15. It has very less solubility in water (23.2cm3 per litre of water at 273 K & 1 bar P) and a very low freezing and boiling point. Dinitrogen molecule is inert at room temperature because of the high bond enthalpy of the NN bond. With the increase in temperature, reactivity increases. Students may refer to Extramarks Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of the dinitrogen molecule.
Chemical Properties of Dinitrogen
Dinitrogen molecule has a high bond enthalpy due to the N N bond. Due to this, it is inert at average temperature—however, the reactivity increases as the temperature increases. At high temperatures, nitrogen molecules react with some metals. This chemical reaction results in the preparation of respective ionic nitrides. The molecules interact with non-metals to form covalent nitrides.
6Li +N2 → 2Li3N
3Mg + N2 → Mg3N2
At about 773 K, it reacts with hydrogen in the presence of a catalyst to form ammonia in Haber’s Process.
N2(g) + 3H2(g) ↔ 2NH3(g) at 773 K
Nitric oxide gas is formed by the reaction of a nitrogen molecule with the oxygen molecule at a temperature of 2000 K.
N2(g) + O2(g) ↔ 2NO(g)
Uses of dinitrogen: The main benefit of dinitrogen is manufacturing ammonia and other commercial chemicals containing Nitrogen (e.g., calcium cyanamide CaCN). It also finds use where an inert atm is required (Eg., inert diluent for reactive chemicals in the iron and steel industry). Liquid dinitrogen is utilised as a refrigerant to preserve biological materials, food items and cryosurgery.
Liquid Nitrogen is used as a preservative for specimens. It is used to manufacture compounds like nitric acid etc.
It is used to occur in an inert atmosphere. Students may refer to various study materials such as NCERT solutions, CBSE revision notes, and Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of the uses of dinitrogen that make it convenient for students to remember everything clearly to grasp everything and score well.
Ammonia (NH3) is a vital compound of Nitrogen and hydrogen. It is developed by the regular decaying of vegetable and animal bodies. The demise and decay of animals and plants cause their nitrogen compounds to deteriorate, producing ammonia. Ammonia, likewise, is present in the soil and acts as ammonium salt derivatives.
We can synthesise the compound by the following methods:
Lab preparations: Ammonia is present in little quantities in air and soil, created by the rot of nitrogenous organic matter, e.g., urea.
Hydrolysis of urea: The reaction involved is given below that
NH2CONH2 + H2O (NH4)2CO3 2NH3 + H2O + CO2
By reacting ammonium chloride with calcium hydroxide: on a small scale, ammonia is obtained from ammonium salts which decompose when reacted with caustic soda or lime.
2NH4Cl + Ca(OH)2 2NH3 +CaCl2 + 2H2O
By reacting ammonium sulphate with sodium hydroxide:
(NH4)2SO4 + NaOH Na2SO4 + NH3 + H2O
From Ammonium Chloride
We can create ammonia gas in the research centre by slowly heating ammonium chloride (NH4Cl) and slaked lime [Ca(OH)2].
2NH4Cl + Ca(OH)2 CaCl2 + NH4OH → 2NH3 + 2H2O
Ammonia is a lighter gas than air, requiring its gathering by descending air displacement. Since it is a solvent in water, it can’t accumulate over it.
Advancing ammonia gas through quicklime calcium oxide (CaO) dries it. We can’t dry an essential gas by moving it through concentrated sulphuric acid or phosphorus pentoxide (P2O5). That is because it reacts with them to frame ammonium sulphate or ammonium phosphate separately.
2NH3 + H2SO4 → (NH4)2SO4
P2O5 + 4NH3 + 3H2O → 2 (NH4)3PO4
We can’t use calcium chloride for drying ammonia gas as it forms ammoniates with CaCl2.
CaCl2 + 8NH3 → CaCl2.8NH3
By the Hydrolysis of Metal Nitrides
Hydrolysing metal nitrides like magnesium and aluminium nitrides, with water or alkalis, can likewise deliver ammonia gas.
Mg3N2 + 6H2O → 3Mg(OH2) + 2NH3
AlN + NaOH + H2O → NaAlO2 + NH3
In this, Nitrogen and hydrogen are taken in 1:3 ratios, and catalysts Fe and MoO are used. This reaction can occur at low temperatures (approximately 700 K) and high pressure(200 x 105 Pa or about 200 atm ). The structure of these hydrides is pyramidal, and hybridisation is sp3.
N2(g) + H2(g) <—> 2NH3(g) fH0= -46.1kJmol-1
nitrogen hydrogen ammonia
The reaction occurs at low temperature, high pressure, catalyst Fe, and catalytic promoter MoO (molybdenum oxide). Ammonia is prepared on a large scale by Haber’s process.
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Ammonia is a covalent atom. It is seen as a dot structure. The element is shaped because of the overlap of 3 H atoms and three sp3 hybrid N orbitals in the structure as the central atom.
A lone pair involves the fourth sp³ hybrid orbital.
It provides a trigonal pyramidal shape to the compound. The H-N-H bond edge is 107.3°, somewhat not precisely the tetrahedral edge of 109°28.
The bond pair-lone pair repulsions deviate the N-H bonds somewhat inwards. In solid and liquid states, ammonia compound is related to hydrogen bonds. The ammonia molecule is a trigonal pyramidal structure with the N atom at the apex. Ammonia has three bond pairs & one lone pair of electrons, as shown in the structure below.
Ammonia is a gas and has no colour. It has a sharp pungent smell with a soapy taste. When inhaled directly, it attacks the eyes, bringing tears. It is lighter than air and more soluble in water. It effortlessly melts at room temperature at a pressure of around 8-10 atm.
Liquid ammonia effervesce at 239.6 K (- 33.5°C) under one-atm pressure. It has a maximum value of the latent heat of vaporisation (1370 J for each gram) and solidifies at 195.3 K (- 77.8°C) to give a solid which is white crystalline in appearance. It has a higher melting and boiling point than expected based on its molecular mass.
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The ammonia molecule is exceptionally inert. In any case, It can disintegrate it into hydrogen and Nitrogen by advancing over metallic impetuses that have been heated.
2NH3 2N2(g) + 3H2
Combustibility: It is flammable in the air.
4NH3 + 3O2 2N2(g) + 6H2O
Basic Character- The compound has a natural propensity to give its lone pair of electrons of Nitrogen to other atoms. Consequently, it behaves like a strong Lewis base.
Ammonia gas is more soluble in water. Its aqueous solution is weakly basic because of the formation of hydroxyl ions.
2NH3 + H2O 2NH4+(aq) + OH_
It forms ammonium salts with acids, for example, ammonium chloride, ammonium sulphate etc. As a weak lewis base, its precipitates the hydroxide of many metals from the solution of their salts, e.g.,
Both the reactions above are neutralisation reactions. It behaves as a weak base, which means it precipitates hydroxides.
ZnSO4 + 2NH4OH Zn(OH)2 (s)+ (NH4)2SO4 white ppt
2FeCl3 + 3NH4OH Fe2O3.x.H2O(s) + 3NH4Cl brown ppt
The lone pair present helps form complexes, so we can say it is a complexion agent.
Cu2+(aq) + 4NH3(aq) [Cu(NH3)4]2+ deep blue
Ag+ + Cl– AgCl white ppt
AgCl + 2NH3 [Ag(NH3)2]Cl(aq) colourless diamine silver ion
This complexing nature of ammonia is due to its lone pair so that it can act as an electron donor.
- We use it in the development of urea and rayon.
- We also use it to produce composts(nitrogenous fertilisers), for example, ammonium nitrate, urea diammonium phosphate, ammonium sulphate and so on.
- We frequently use liquid ammonia as a refrigerant in ice plants.
- It finds its utility in the furniture industry as a purging operator for furniture and glass surfaces.
- Ostwald’s procedure is essential to use in the production of nitric acid.
- We also use it in the production of sodium carbonate by Solvay’s procedure.
- Ammonia is also utilised in the preparation of some inorganic nitrogen compounds.
The undermentioned tests of any specimen affirm the presence of the compound:
- The ammoniacal odour of the compound is effectively perceivable, having a trademark pungent smell.
- It turns wet red litmus blue and moist turmeric paper brown.
- When conveyed near ammonia, a glass bar dunked in concentrated HCl causes thick white exhaust.
- When combined with a solution of copper sulphate, ammonia turns the solution deep blue.
NH3 + CuSO4 + nH2O→ [Cu(NH3)4(H2O)n]SO4
- When mixed with Nessler’s reagent (basic arrangement of K2[HgI4], ammonia gives a precipitate brown in colour.
- i) NH4+ + 2[HgI4]2− + 4OH− → HgOHg(NH2)I(ppt)↓ + 7I− + 3H2O
Oxides of Nitrogen
Nitrogen reacts with Oxygen and results in several nitrogen oxides. The oxidation states of all these nitrogen oxides are quite different.
They are in the different range of +1 to +5 oxidation states. We will look at some of the essential oxides below.
i) Dinitrogen Oxide, N2O
It is a colourless and non-flammable gaseous compound which has neutral properties. We know it by the general name, laughing gas. We can prepare it by decomposing ammonium nitrate at a high temperature. The oxidation state of Nitrogen is +1.
NH4NO3 → N2O + 2H2O
ii) Nitrogen monoxide (NO)
It is a colourless and gaseous compound. We can generally prepare it by reducing dilute nitric acid with copper. NO combination shows neutral properties. The oxidation state of Nitrogen is two.
- ii) 2NaNO2 + 2FeSO4 + 3H2SO4 → Fe2(SO4)3 + 2NaHSO4 + 2H2O + 2NO (nitrogen oxide)
iii) Dinitrogen Trioxide, (N2O3)
Dinitrogen trioxide is a deep blue solid and has acidic properties. It is synthesised by mixing equal parts of nitric oxide and nitrogen dioxide and cooling the mixture below −21 °C (−6 °F). The oxidation state of Nitrogen is +3.
2NO + N2O4 → 2 N2O3
iv) Nitrogen Dioxide (NO2)
Nitrogen dioxide is a typical oxide of Nitrogen, a reddish-brown toxic, dangerous gas. You can know its presence with a sharp smell. The oxidation state of Nitrogen is +4.
2Pb(NO3)2 4NO2 + 2PbO
v) Dinitrogen Tetroxide, N2O4
Generally, Dinitrogen tetroxide is a colourless solid that we can find in equilibrium with nitrogen dioxide. It is a powerful oxidiser and a standard reagent producing many chemical compounds. The oxidation state of Nitrogen is +4.
2NO2 ⇌ N2O4
vi) Dinitrogen Pentoxide, N2O5
Normally Dinitrogen pentoxide is a colourless solid, and its characteristic property is that it sublimes slightly above room temperature. It is unstable. It is a potentially dangerous oxidiser. The oxidation state of Nitrogen is +5. We can synthesise it by dehydrating nitric acid (HNO3) with phosphorus (V) oxide:
P4O10 + 4HNO3 → 4HPO3 + 2N2O5
Generally, Lewis dot main resonance structures and bond parameters of all oxides are given in the table below. Students can refer to Class 12 Chemistry Chapter 7 Notes.
Nitrogen forms oxyacids such as H2N2O2 (hyponitrous acid), HNO2 (nitrous acid) and (HNO3)nitric acid. Amongst them, HNO3 is the most essential.
Preparation: In the laboratory, nitric acid is formed by heating potassium nitrate, sodium nitrate, and concentrated sulphuric acid in a glass retort.
NaNO3 + H2SO4 NaHSO4 + HNO3
Industrially: It is prepared from Ostwald’s process. This process is based upon catalytic oxidation of ammonia by atmospheric Oxygen.
The reaction involves various steps:
- i) Oxidation of ammonia
4NH3 (g) + 5O2(g) 4NO(g) + 6H2O (Pt/Rh gauge catalyst, 500 K,9bar)
ii)NO is further oxidised
2NO + O2 2NO2
iii) Then hydrolysis of NO2
3NO2 + H2O 2HNO3 + NO
The nitric acid so prepared is 98% pure, the rest of 2% can be purified by concentrated sulphuric acid. Nitric acid is the most popular and supportive oxoacid of Nitrogen. It has a molecular mass of HNO3, and its molar weight is 63.01 g mol-1. Now, let us see the several properties of nitric acid.
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Physical properties of Nitric acid
- Pure nitric acid is the king of acids but a colourless fuming fluid. It has a unique and pungent odour.
- Upon standing, it creates a yellow shading. It is because of the presence of various oxides of Nitrogen that are dissolved. The primary oxide is NO2.
- The acid is purely soluble in water and the thickness of the pure acid is 1.54 g/mL.
We can see that anhydrous nitric acid effervesces at 355.6 K (83.6°C). It can develop a white solid at 231.4 K (- 41.7°C).
- It can corrode the skin & causes yellow rashes. HNO3 is a planar molecule in a gaseous state with the structure shown below.
Chemical Properties of Nitric Acid
In this topic, we will cover a few of the most important chemical properties of Nitric Acid.
In its pure state, Nitric acid is not stable. Even at ordinary temperatures, it slightly decays when we expose it to daylight. Upon solid heating, it breaks down and gives nitrogen dioxide, Oxygen and water.
4HNO3(aq) 2H2O(l) + 4NO2(g) + O2(g)
Nitrogen dioxide has a reddish brown colour. So, it might further break down in the undecomposed acid and produce its yellowish chestnut shade.
Nitric acid ionises readily in water and is a solid monobasic acid giving hydronium and nitrate ions.
HNO3(l) + H2O(l) H3O+ (aq) + NO3–(aq)
Reactions with Basic Oxides– Nitric acid reacts with basic oxides to give nitrate salt and water. For ex., it combines with calcium oxide to give calcium nitrate and a water molecule.
Reaction with Bases (Hydroxides)– Nitric acid reacts with hydroxides to give nitrate salt and water. For ex., it interacts with sodium hydroxide NaOH to give sodium nitrate and a water molecule.
Reaction with Carbonates and Hydrogen Carbonates– Nitric acid interacts with carbonates and hydrogen carbonates to give nitrate salt, carbon dioxide and water. For ex., it reacts with sodium carbonate to give sodium nitrate, water and carbon dioxide.
Reaction with Metals– Nitric acid does not behave like an acid with metals to form a salt and free hydrogen. However, Magnesium and manganese are the two metals that react with cold and extremely dilute (1%) nitric acid to give out hydrogen.
Test for nitrates: The test is done by adding dilute iron sulphate to an aqueous solution containing NO3–and then adding concentrated sulphuric acid along the side of the test tube. The Brown ring at the interface between the solution and sulphuric acid layer indicates the presence of nitrate in the solution. It is also called the brown ring test. In this, we take a test tube and sample (salt containing nitrate) to it. Then slowly pour concentrated sulphuric acid into it. Then we put a solution of ferrous sulphate and allow it to stand. Then, a brown ring is formed on the interface, showing nitrates’ presence. The ring that is created is chemically [Fe(H2O)5(NO)]2+(Nitroso ferrous sulphate ).
so NO3– + 3Fe2+ + 4H+ NO + 3Fe3+ + 2H2O
[Fe(H2O)6]2+(aq) + NO [Fe(H2O)5(NO)]2+ (brown ring formed) + H2O
Uses of nitric acid: In manufacturing ammonium nitrate and other fertilisers. In the manufacturing of explosives like TNT (trinitrotoluene). In the purification of gold and silver by using aqua regia. It is used as an oxidiser in rocket fuels. Other significant uses are picking stainless steel, etching metals, and as an oxidiser in rocket fuels.
The various uses of nitric acid are given in our Class 12 Chemistry Chapter 7 Notes. These notes are prepared by Chemistry faculty with decades of experience. We are aware of the importance and seriousness of our role that may impact students’ academic performance to help them maximise their potential and ensure 100% results in the exams. .
Allotropic forms of Phosphorus
In this topic, we will discuss the allotropes of phosphorus. This element obtains in a few allotropic forms. The primary vital allotropes of phosphorus include white phosphorus, red phosphorus and black phosphorus. In addition to these, there also exists violet phosphorus. Although, that is not a significant allotrope. So, let us start with the several allotropes of phosphorus.
It is a general allotrope of phosphorus. This White phosphorus is a waxy and translucent solid.
This phosphorus is very soft and needs proper handling. It is poisonous and insoluble in water. However, it dissolves in carbon disulphide and glows in dark or carbon tetrachloride. It breaks down in boiling caustic soda solution in an inert atmosphere and produces sodium hypophosphite and phosphine.
P4 + 3NaOH + 3H2O PH3 + 3NaH2PO2
White phosphorus phosphine sodium hypophosphite
Structure of White Phosphorus: </h3
In the above figure, we discuss the structure of white phosphorus. As we see, it has a tetrahedral shape geometry. Each phosphorus element has a covalent bond with three various atoms of phosphorus. These elements exhibit weak Van der Waals forces of attraction.. We can remember that this element is very active and too harmful. The melting point is relatively low at 44°C.
As we can see, the bond angle in a P4 particle is 60°. It is comparatively significantly less as compared to a regular bond angle or a hypothetical bond angle. Therefore, it has a strain in itself. That is why white phosphorus is highly unstable and reactive. Students are advised to refer to our Class 12 Chemistry Chapter 7 Notes for a more granular and in-depth explanation of the phosphorus varieties.
White phosphorus quickly catches fire suddenly in the air at approx 35°C. As you can observe , this temperature is marginally higher than the standard room temperature. It is the reason why it is kept in water. After combustion, it produces phosphorus pentoxide.
P4 + 5O2 → 2P2O5 or P4O10
White phosphorus undergoes an oxidation reaction if it comes in contact with moist air. This chemical reaction leads to a sparkling discharge of light. As an outcome, it sparkles obliviously. White phosphorus displays chemiluminescence(glows in the dark). Extramarks Class 12 Chemistry Chapter 7 notes are more accessible on the Extramarks website for students who want to learn more. Students can entirely rely upon these notes as these are made following all the latest guidelines and curricula laid by CBSE.
We can create red phosphorus by heating white phosphorus to around 250°Celsius within sight of daylight.
White phosphorus Red phosphorus at 573K in an inert atmosphere
Red phosphorus is iron-grey in colour and a radiant and bright crystalline solid. This red phosphorus is non-poisonous and does not have any odour. Red Phosphorus does not become soluble in water or carbon tetrachloride solvent.
It doesn’t split up in boiling caustic soda-like white phosphorus. It disintegrates in alcoholic potash.
You can find it in the state of a polymeric solid. It is steady under normal conditions and doesn’t catch fire in the air. Although, it experiences burning when we warm it to around 400°C. Red phosphorus doesn’t show chemiluminescence.
We can synthesise black phosphorus from white phosphorus by heating it to 470K at an inert temperature.
White phosphorus Black phosphorus at 470 K and inert atmosphere
Of all phosphorus, Black phosphorus is the most stable allotrope of phosphorus. This black phosphorus has a layered structure and is a very highly polymerised form of the element.
We can find black phosphorus in two primary forms. These are alpha black phosphorus and beta black phosphorus. While beta black phosphorus develops when white phosphorus is heated at 473K under high pressure, alpha black phosphorus forms when we warm red phosphorus at 803K. It can be sublimed crystals. It does not oxidise in the air. Beta black phosphorus conducts electricity, but alpha black phosphorus doesn’t conduct electricity. Extramarks Class 12 Chemistry Chapter 7 notes are accessible on the Extramarks website for students who want to learn more. To enjoy the maximum benefit of these resources, students just need to register themselves at Extramarks’ official website and stay ahead of the competition.
Uses of Phosphorus
Phosphorus compounds assume an important part of life forms. It forms an essential constituent in animal and plant matter.
We find it in the blood, bones and brain of all the animals and living cells. Some of its compounds find applications in industries. The most important of these chemicals are orthophosphoric acid and phosphatic composts.
(PH3) is a chemical compound prepared by heating phosphorous acid or reacting calcium phosphide with water or diluting HCl. Phosphine observes its place in the group of organophosphorus compounds with the PH3 chemical formula. Scientist Philippe Gengembre discovered it in 1783. It can be detected in living tissues, blood, urine, saliva, etc.. We will discuss the chemical phosphine, its physical properties, chemical properties and uses in greater detail in our Class 12 Chemistry Chapter 7 Notes
Phosphine is a chemical that discovers its place in the group of organophosphorus compounds. Philippe Gengembre, the scientist, discovered or acquired this chemical in the year 1783. He was the one who developed phosphine by heating phosphorus in an aqueous solution of potassium carbonate.
It carries the molecular formula of PH3. The concentration of this phosphine compound constantly alters in our environment. As we already mentioned, this phosphine chemical plays an essential role in the biochemical phosphorus cycle.
Preparation of Phosphine
a) From Phosphide
We can find the compound by interacting phosphides with water. For example, calcium phosphide interacts with water to give calcium hydroxide and phosphine.
Ca3P2 + 6H2O 3Ca(OH)2 + 2PH3
b)From Phosphorus Acid
We can also find the compound by heating phosphorus acid. It breaks to give a pure sample of phosphine.
4H3PO3 H3PO4 + PH3
Structure of Phosphine
We can see that the electronic configuration of phosphine resembles ammonia. It has the structure of pyramid symmetry. The bond angle H-P-H is = 93°. On the other hand, ammonia has pyramidal geometry with a bond angle of 107.80. Hence, we see that both these chemicals have a relative bond angle. Phosphorus is less electronegative than Nitrogen. The structure of phosphine has been provided in detail in our Class 12 Chemistry Chapter 7 Notes and students can refer to our study notes for a clear understanding of this topic.
In the cloud of electrons around the central atom, phosphorus is less concentrated than the Nitrogen present in ammonia. Although, the lone pair of electrons causes significantly more contortion in PH3. Hence, we note the decline in the bond angle in PH3 to 93.5°.
Image name: Phosphine>
Properties of Phosphine
It is colourless effervescence gas, & phosphine has a particular odour like rotten fish.
It is a highly poisonous gas. Phosphine PH3 is partially dissolved in water. Although, it can dissolve in natural solvents. PH3 phosphine acts as a Lewis base by giving away its lone pair of electrons by reacting with hydrogen iodide. In a typical way, it is a non-ignitable gas. But, if you heat it, it bursts into flames, forming phosphoric acid. It explodes violently when we expose it to oxidising agents.
Effects of PH3
Phosphine is a hazardous gas. Exposure to even little quantities of the gas can lead to dizziness, loose bowels, cough, cerebral pain, and chest tightness, to name a few. Upon greater exposure, you may risk suffering from convulsions, coma, damage to the kidney and liver and irregular heartbeat.
Properties of PH3 Ligands
Tertiary phosphines, or PR3, are an essential class of ligands since we can change their electronic and steric properties in a very orderly path over a wide range by shifting the R group(s).
These can stabilise a more comprehensive array of metal complexes that might interest organometallic scientific experts as their phosphine complexes (R, P) nM−L. Phosphines are usually spectator ligands and not performer ligands. Like NR3, phosphines also possess a lone pair on the focal particle, giving them to metal. For alkyl phosphines, the π acidity is very weak.
The solution of phosphine in water decomposes in the presence of light, giving red phosphorus and Hydrogen molecules.
3CuSO4 + 2PH3 Cu3P2 + 3H2SO4
Uses of phosphine:
Spontaneous Combustion of PH3 is technically used in Holmes’s signals for ships. Containers containing calcium carbide and phosphide are pierced and thrown into the sea when the gases evolve to burn and serve as a signal. It is also used as a smoke screen in warfare. Students may refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of phosphorus derivative phosphine.
What is a Phosphorus Halide?
A phosphorus halide is a compound that phosphorus forms with a halogen.
A phosphorus halide is of two types. They are PX3 and PX5. Here, we refer to X as a halogen. It should be anything from fluorine, chlorine, bromine or iodine. However, the most common phosphorus halide is chloride. These chlorides are usually covalent in nature.
1) Phosphorus Trichloride
This phosphorus trichloride is an oily and sleek fluid. It is very lethal, and The shape of this compound is that of a triangular pyramid. The atom of phosphorus exhibits sp3 hybridisation.
As shown in the above diagram, phosphorus has its SP3 orbitals. It has only one electron, giving that electron to a p orbital electron from 3 chlorine atoms. The fourth sp3 orbital is full. It is a solitary lone pair. Thus, it cannot form a bond. However, it repels alternate bonds. This phosphorus trichloride creates a state of the shape of trigonal pyramidal.
Extramarks Class 12 Chemistry Chapter 7 notes are accessible on the Extramarks website for students who want to learn more. To enjoy the maximum benefit of these resources, students just need to register themselves at Extramarks’ official website and stay ahead of the competition.
We find phosphorus trichloride by passing dry chlorine heated with white phosphorus. The chemical reaction which takes place is given below.
P4 + 6Cl2 → 4PCl3
So we can also synthesise the PCl3 phosphorus trichloride compound by reacting thionyl chloride with white phosphorus. The reaction is given below:
P + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2
phosphorus trichloride is a colourless oily liquid and hydrolysis in the presence of moisture. Phosphorus trichloride hydrolysis when we dampen it.
PCl3 + 3H2O → H3PO3 + 3HCl
It interacts with natural compounds having a –OH hydroxyl group and gives their ‘chloro’ subsidiaries as products.
3C2H5OH + PCl3 → 3C2H5Cl + H3PO3
Structure of PCl3
The phosphorus element in the centre of PCl3 shows sp3 hybridisation. It has three bond sets and one lone pair of electrons. Because of that reason, it has a pyramidal shape geometry. It behaves as a Lewis base because it can donate its lone pair of electrons to other electron-deficient elements or atoms.
2) Phosphorus Pentachloride
It is a yellowish-white powder. The solid Phosphorus Pentachloride is water soluble. It is soluble in organic solvents, i.e. carbon tetrachloride, benzene, carbon disulphide, and diethyl ether.
Its geometry is that of a trigonal bi-pyramidal. We find this structure primarily in both vaporous and liquid states. In the solid type, we can find it as an ionic solid, [PCl4]+[PCl6]–. Here, the cation, [PCl4]+, is a tetrahedral shape, and the anion, [PCl6]– is an octahedral shape.
We should know that the molecule shows three tropical P-Cl bonds and two pivotal P-Cl bonds. Because of the more prominent repulsion at hub positions compared to the central places, we see that the two axial bonds are more extended than equatorial ones.
We can produce phosphorus pentachloride by reacting to an excess of dry chlorine.
P4 + 10Cl2 → 4PCl5
We can also prepare it by the chemical reaction of SO2Cl2 thionyl chloride and phosphorus.
P4 + 10SO2Cl2 → 4PCl5 + 10SO2
In the presence of moist air, phosphorus pentachloride hydrolyses to POCl3. This compound POCl3 changes over to phosphoric acid over some time.
PCl5 + H2O → POCl3 + 2HCl
POCl3 + 3H2O → H3PO4 + 3HCl
When it is strongly heated, it sublimes and further disintegrates into phosphorus trichloride.
PCl5 → PCl3 + Cl2
It interacts with finely partitioned metals under the influence of heat to create metal chlorides.
2Ag + PCl5 → 2AgCl + PCl3
It reacts with natural compounds containing –OH groups and produces their ‘chloro’ subordinates.
C2H5OH + PCl5 → C2H5Cl + POCl3 + HCl
The topic of phosphorus pentachloride has been covered in more detail in our Class 12 Chemistry Chapter 7 Notes. Our study notes are prepared by Chemistry subject matter experts with years of experience revising the content regularly while adhering to the latest CBSE guidelines and curriculum. So students can confidently rely on our solutions and study notes. The same is true for other subjects as well.
Structure of PCl5
As discussed in our Class 12 Chemistry Chapter 7 Notes, the central phosphorus atom in phosphorus pentachloride experiences SP3d hybridisation. All five electrons are mixed with these hybrid orbitals as bond sets. The molecular shape of the element is trigonal bipyramidal.
After the five electrons are hybridised, we get five of equivalent size and shape. Three frames of a triangle symmetry(120° partition) in the centre. One bond is above, and one is under those three.
However, you must remember that we only notice trigonal bipyramidal geometry in phosphorus pentachloride in its fluid and vaporous state. In its solid form, it exists as a salt.
Oxoacids of Phosphorus
Generally, Oxoacids of Phosphorus are Hypophosphoric acid(H3PO4), Metaphosphoric acid (HPO2), Pyrophosphoric acid (H4P2O7), Hypophosphorous acid(H3PO2), Peroxophosphoric acid (H3PO5), Phosphorous acid (H3PO3)and Orthophosphoric acid (H3PO4). Oxoacids are acids containing Oxygen. The Phosphorus-H bonds in oxoacids are not ionisable to contribute H+ ions. Conversely, the H atoms attached to Oxygen in P-OH form are ionisable. Thus, we can say that basicity is the property exhibited by the H atoms that are attached to Oxygen.
In conclusion, orthophosphoric acid, H3PO3, is dibasic as it has two P-OH bonds. Like phosphoric acid, H3PO4 is tribasic as it has three P-OH bonds. The oxoacids of phosphorus which have P-H bonds show strong reducing properties. For ex., hypophosphorous acid containing two P-H bonds is an excellent reducing agent.
4AgNO3 + 2H2O + H3PO2 → 4Ag + 4HNO3 + H3PO4
Few Popular Oxoacids of Phosphorus
In this section, phosphorus forms several oxoacids. We will explain several of the most essential and popular phosphorus oxoacids. And students can refer to our Class 12 Chemistry Chapter 7 Notes to understand every concept and answer any question easily. This will help the students to master the topic and get good grades..
- i) Phosphorus acid, H3PO3
Phosphorous acid is a diprotic acid. This means that it ionises two protons.
We can explain it better by the structural formula HPO(OH)2. We can prepare phosphorus acid by hydrolysis of phosphorus trichloride with acid or steam.
PCl3 + 3H2O → HPO(OH)2 + 3HCl
ii) Phosphoric acid, H3PO4
Phosphoric acid is a triprotic acid, and This means that it can ionise three protons. It is a non-toxic acid when pure. Phosphoric acid is solid at standard temperature and pressure. We can prepare phosphoric acid by adding sulphuric acid to tricalcium phosphate rock-
Ca5(PO4)3X + 5H2SO4 + 10H2O → 3H3PO4 + 5CaSO4.2H2O + HX
(Halogen X is Fluorine, Chlorine, Bromine, & hydroxyl OH).
iii) Polymetaphosphoric Acid (HPO3)n
We can observe it by warming orthophosphoric acid to around 850 K. It does not exist as a monomer. It exists as a cyclic trimer, cyclic tetramer or polymer.
The oxoacids are interrelated regarding the loss or gain of water molecules or oxygen atoms. The structure of some essential oxoacids is given below:
Students may refer to study notes prepared by Extramarks subject experts. Our Class 12 Chemistry Chapter 7 Notes has a more detailed explanation of Group 15. It is a thoroughly researched material made as per the CBSE examination guidelines.
Group 16 elements
The group sixteen elements of the modern periodic table combination of 5 elements oxygen, sulphur, selenium, tellurium and polonium. The elements in this group are also called chalcogens as well as ore-forming elements due to many elements that can be extracted from sulphide or oxide ores.
The initial four particles of the group are called chalcogens or ore-forming elements. This is because many metal ores are found in the earth’s crust as sulphides or oxides. Oxygen is the most abundant element which is accessible. It shapes 20.946% of Oxygen by volume and 46.6% of the world’s mass as silicates and various compounds like carbonates, oxides, and sulphates.
The essential majority of the Oxygen in the air is delivered by photosynthesis in plants. It additionally occurs as ozone. Sulphur is the sixteenth most inexhaustible element. Sulphur in its combined state is found in ores. Electronic Configuration of Group 16 Elements
Concerning Class 12 Chemistry Chapter 7 note, Group sixteen elements have six electrons in their valence shell, and their typical electronic configuration is ns2np4.
|Element||Symbol||Atomic number||Electronic Configuration|
|Oxygen||O||8||[He] 2s2 2p4|
|Sulphur||S||16||[Ne] 3s2 3p4|
|Selenium||Se||34||[Ar] 3d10 4s2 4p4|
|Tellurium||Te||52||[Kr] 4d10 5s2 5p4|
|Polonium||Po||84||[Xe] 4f14 5d10 6s2 6p4|
The chemical symbol for Oxygen is given O. It is a colourless and odourless gas utilised in the respiration method by humans, which is converted into carbon dioxide. Oxygen shown as a diatomic molecule (O2). Oxygen is also created as a triatomic molecule (O3) in traces. It is called ozone. Oxygen combines readily with many elements. Heat energy evolution occurs during the combination of some elements; this process is known as combustion.
Sulphur is designated by the symbol S. It is a non-metal that ranks ninth based on cosmic abundance. About one atom in every 20,000-30,000 atoms is a sulphur atom. Sulphur is created in the combined state as well as in the free state. About 0.09 % of sulphur is present in seawater as sulphates. The meteorite consists of 12 % of sulphur; a large amount of sulphur is seen from the underground deposits of pure sulphur in dome-like structures. Here the sulphur atom is formed by the action of anaerobic bacteria on the sulphate minerals such as gypsum.
Selenium is rarer than Oxygen or sulphur. It is found in the free state and combined with heavy metals (such as lead, silver, or mercury) in some minerals. The grey form of metallic selenium is the most abundant element under normal conditions.
This Tellurium is a chemical element with atomic number 52, and the properties show of metals and nonmetals both. It is one of the rarest steady elements present in the universe.
It is generally found in a free state and contains elements such as copper, lead, silver or gold.
It is the rarest element among the group of 16 elements. It is a radioactive element. Polonium is rarely used in scientific applications for alpha radiation.
Atomic & Physical Properties and Trends of Group sixteen elements
- Atomic and Ionic Radii– The atomic and ionic radius increases as we move from Oxygen to Polonium. The size of the oxygen atom is, though, exceptionally tiny.
- Ionisation Enthalpy: Ionisation enthalpy decreases with an enhancement in the size of the central atom. Therefore, it decreases as we go down from Oxygen to Polonium since the size of the atom increases as we move down. This ionisation enthalpy is due to the fact that group 15 elements have extra stable half-filled p orbitals electronic configuration.
- Electron Gain Enthalpy– The electron gain enthalpy decreases with an increase in the size of the central atom moving down the group. Oxygen molecules have a less negative electron gain enthalpy than sulphur. Oxygen encounters more repulsion between the electrons effectively present and the approaching electron because of its compressed nature.
- Electronegativity– Next to fluorine, Oxygen has the highest electronegativity value among the elements. The electronegativity value decreases as we go down the group or the atomic number increases. Though, it decreases as we move from Oxygen to polonium due to an increase in nuclear size..
- Nature of the Group sixteen Elements: Oxygen and Sulphur are non-metals, Selenium and Tellurium are metalloids, and Polonium is metal under typical conditions. Polonium is a radioactive and short-lived(half-life 13.8 days) element.
- Allotropy– Each one of the elements of group 16 displays allotropy. Oxygen has two allotropes: Oxygen and Ozone. Sulphur exists in many allotropic forms. Only two allotropes of them are stable, which are shown Rhombic Sulphur and Monoclinic Sulphur. Selenium and Tellurium are observed in both amorphous and crystalline states.
- The Melting and Boiling Points– As the atomic size increases from Oxygen to tellurium, the melting and boiling points also increase. The significant difference between the melting and boiling points of Oxygen and sulphur might be clarified on the premise that Oxygen is observed as a diatomic molecule. In contrast, sulphur exists as a polyatomic particle.
- Oxidation States and trends in chemical reactivity– The group of sixteen elements have an electronic configuration of ns2 np4 in their outer shell orbital; they may accomplish inert gas configuration either by accepting two electrons, framing M-2 or by sharing two electrons; in this manner shaping two covalent bonds. Sulphur selenium and tellurium usually show a +4 oxidation state in their compounds with Oxygen and +6 with fluorine.
Hence, these elements indicate both negative and positive oxidation states. The normal oxidation states are shown by the elements of group 16 incorporating -2, +2, +4 and + 6. Bonding in +4 and +6 states are primarily covalent. Extramarks Class 12 Chemistry Chapter 7 notes are accessible on the Extramarks website for students who want to learn more. It is curated by the subject matter experts and it is a thoroughly researched material made as per the CBSE examination guidelines.
The group of sixteen elements react with hydrogen to form hydrides of the sort H2E, where E could be any element- Oxygen, sulphur, selenium, tellurium or polonium.
H2 + E H2E
Hydrogen Group 16 element Hydride
The Physical States of Hydrides of Group sixteen Elements
Water H2O is an odourless and colourless liquid, but the hydrides of the several elements of this group are poisonous gases that are colourless with an unpleasant odour.
The boiling point of these hydrides extraordinarily decreases from water to hydrogen sulphide and, after that, increases. Water has an anomalously high boiling point since the hydrogen bonds bond its particles in liquid and solid states.
Anomalous Behaviour of Oxygen
The anomalous behaviour of Oxygen, Because of its smaller size and high electronegativity, Oxygen, like different members of the p-block element present in the second period, shows unusual behaviour. Strong hydrogen bonding in H2O, absent in H2S, is an example of the impact of tiny size and maximum electronegativity value.
Due to oxygen shortages, d orbitals, its covalency is limited to four, and it rarely exceeds two in practice. In the case of different group members, however, the valence shell can be enlarged, and covalent approaches four. Students may refer to Extramarks study resources including Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of the anomalous behaviour of Oxygen.
Important note: H2S is less acidic than H2Te because the acidic character increases due to the decrease in (E-H) bond dissociation enthalpy down the group.
Oxygen is 21% by volume of air. It was prepared by Karl Wilhelm Scheele, and Priestly did the other reactions. Isotopes of Oxygen: O16, O17, O18.
Similar to Groups 14 and 15, the lightest member of group sixteen has the best inclination to shape numerous bonds. In this way, elemental Oxygen is observed in nature as a diatomic gas which contains a net twofold bond O=O. Dioxygen molecule means the normal allotrope of Oxygen having two atoms of Oxygen in the molecule.
Likewise, with Nitrogen, electrostatic repulsion between lone pairs of electrons on adjacent atoms keeps Oxygen from framing stable catenated compounds. Ozone (O3), a standout among the most intense oxidants known, is utilised to cleanse drinking water since it doesn’t deliver the characteristic taste of chlorinated water solution.
2H2O2 (l) → 2H2O(l) + O2 (g)
so ΔG° = −119 kJ / mol
Nevertheless, of the quality of the O = O bond (DO2 = 494 kJ/mol), O2 is amazingly reactive and reacts straightforwardly with almost all elements aside from the noble gases. Some properties of O2 and related species, for example, the peroxide and superoxide particles, are mentioned in the accompanying table. The table has been further elaborated in our Class 12 Chemistry Chapter 7 Notes to make it easy for students to comprehend various concepts about this topic.
|Species||Bond Order||Number of unpaired e–||O – Oxygen bond distance in pm|
The most advantageous process for making Oxygen in the lab includes the catalytic decay of potassium chlorate molecules in the solid form, where manganese dioxide works as a catalyst.
2KClO3 → 2KCl + 3O2 (presence of Δ and MnO2)
Another laboratory technique includes the thermal decomposition of metal oxides from the lower part of the electrochemical arrangement. For ex., the thermal decomposition of mercuric oxide and silver oxide produces Oxygen molecules.
2HgO → 2Hg + O2
Mercuric oxide Mercury Dioxygen molecule
2Ag2O → 4Ag + O2
Silver oxide Silver Dioxygen molecule
We can also synthesise Oxygen in the laboratory by thermal treatment of the higher oxides of some metals like lead, manganese and barium.
2PbO2 → 2PbO + O2
Lead (IV) oxide Lead (II) oxide Dioxygen
2MnO2 + 2H2SO4 → 2MnSO4 + 2H2O + O2
Manganese(IV) oxide Sulphuric acid Manganese(II) sulphate Dioxygen
2BaO2 → 2BaO + O2
Barium peroxide Barium oxide Dioxygen
When decomposed thermally, salts rich in Oxygen, such as permanganates and nitrates, also yield Oxygen.
2KNO3 → 2KNO2 + O2
Potassium nitrate Potassium nitrite Dioxygen
2KMnO4 → K2MnO4 + MnO2 + O2
Potassium permanganate Potassium manganate Manganese(IV) oxide Dioxygen
2NaNO3 → 2NaNO2 + O2
Sodium nitrate Sodium nitrite Dioxygen
Physical Properties –
Oxygen is a tasteless, colourless & odourless gas. It is slightly heavier than air.
It is slightly soluble in water. This little amount of Oxygen dissolved is adequate to support aquatic and marine water life. Below pressure, we can condense it to a light blue liquid by compacting the gas at 90 Kelvin. It can likewise be solidified into a bluish-white solid at 55K. It has three isotopes, O16, O17 and O18, which are stable. Dioxygen is paramagnetic, and this can be explained based on MO theory.
Oxygen is an outstandingly reactive element and reacts straightforwardly with at least all metals and non-metals. It doesn’t respond straightforwardly with a few metals like gold and platinum and some noble gases like helium, argon, and neon. i) The Reaction reacts with metals
All metals blaze in oxygen and frame oxides which are, for the most part, basic.
Metal Dioxygen Metal-oxide
2M + O2 → 2MO
4M + O2 → 2M2O
4M + 3O2 → 2M2O3
A large portion of non-metal burns within the presence of oxygen structures acidic oxides. Example: Sulphur blazes within the presence of oxygen and gives sulphur dioxide.
S + O2 → SO2
Reactions With Some Compounds
Sulphur dioxide experiences catalytic oxidation within the presence of vanadium pentoxide (V2O5) to frame sulphur trioxide. This sulphur dioxide is a critical stride in developing sulphuric acid by the contact method.
2SO2 + O2 → 2SO3
Oxygen interacts with many organic compounds such as carbohydrates and hydrocarbons at hoisted temperatures or on start, resulting in carbon dioxide and water formation.
CH4 + 2O2 → CO2 + 2H2O
Methane Dioxygen Carbon dioxide Water
C6H12O6 + 6O2 → 6CO2 + 6H2O
Glucose Dioxygen Carbon dioxide Water
Dioxygen Difluoride denoted with the formula O2F2 is a compound of fluorine and Oxygen. It can be observed that as an orange-coloured solid which liquefies into a red fluid at −163 °C (110 K). It is a great degree solid oxidant and breaks down into oxygen and fluorine even at −160 °C (113 K) at a rate of 4% every day: its lifetime at room temperature is, in this manner, amazingly short.
Uses of dioxygen
It supports life and its importance in normal respiration and combustion processes. It is used in oxy-acetylene flame.
Liquid Oxygen is used as rocket fuel.
Its radioactive isotope is used as a tracer for many chemical reactions. It is used to prepare synthesis gas. Oxygen is used to create oxy-hydrogen or oxy-acetylene flames, which are used for cutting, welding, and manufacturing many metals, particularly steel. Oxygen cylinders are widely used in hospitals, high-altitude flying and mountaineering. The combustion of fuels that are hydrazines in liquid Oxygen provides tremendous thrust in rockets.
An oxide is a binary compound we obtain upon the reaction of Oxygen interacting with other elements. Depending on the oxygen element, we can extensively fix them into mixed and simple oxides. An oxide of a nonmetal commonly tends to be acidic..
On the other hand, an oxide of metal shows an essential tendency. The oxides of particles in or close to the corner band of semi-metals are by and large amphoteric molecules. Let us consider the various types of simple oxides that can be classified as:
A binary compound of Oxygen with another element is called oxide. Oxygen interacts with most of the periodic table elements to develop oxides. In every case, one element forms two or more oxides. A simple oxide is one carrying several oxygen atoms that the general valency of its metal allows. Examples. of this include H2O, MgO, and Al2O3 aluminium oxide.
We get a mixed oxide upon the addition of two simple oxides. Ex., of these include: Lead dioxide (PbO2), and lead monoxide (PbO) combined to develop the mixed oxide Red lead (Pb3O4). In other examples, we can see that Ferric oxide (Fe2O3) and ferrous oxide (FeO) combine and form the mixed oxide Ferro-ferric oxide (Fe3O4). The concepts of oxides are explained in a simple, easy to understand manner. Class 12 Chemistry Chapter 7 Notes are prepared by Extramarks subject matter experts with years of experience have restructured the information into different formats to enable a smooth and deep learning experience so that students need not look elsewhere to supplement their studies to better their performance., They know the challenges faced by students and our solutions are prepared in order to simplify complex topics for students. Needless to say, they completely understand what is legitimate as per the board’s standards.
Classification of Simple Oxides
Based on their chemical behaviour, there are acidic, basic, amphoteric and neutral oxides.
An acidic oxide interacts with water and produces acid. Typically, it is the oxide of non-metals. Examples include SO2, SO3, CO2, Cl2O7, P2O5, and N2O5. It could also be the oxide of some metals with high oxidation states, such as CrO3, Mn2O7, and V2O5.
For ex., Sulphur dioxide interacts with water and creates sulphurous acid.
SO2 + H2O → H2SO3
Chromic anhydride interacts with water and results in chromic acid.
Cr2O3 + H2O → H2Cr2O4
In general metallic oxides are basic. A basic oxide is combined with water to give a base. Examples include the oxide of several metals, such as Na2O, CaO, and BaO. These are basic in nature.
For instance, Calcium oxide reacts with water and produces calcium hydroxide, a base.
CaO + H2O → Ca(OH)2
These are neither acidic nor basic in nature, e.g., CO, NO and N2O
An amphoteric oxide is a metallic oxide showing a dual behaviour. It exhibits the characteristics of both an acid as well as a basic nature. It is combined with both alkalis as well as acids, e.g. Al2O3
For example, zinc oxide acts acidic when it interacts with concentrated sodium hydroxide solution. Though, it acts as a basic oxide while reacting with hydrochloric acid.
ZnO + 2H2O + 2NaOH → Na3Zn[OH]4 + H2
ZnO + 2HCl → ZnCl2 + H2O
Aluminium oxide is another example that reacts with alkalis as well as acids.
Al2O3(s) + 6NaOH(aq) + 3H2O(l) → 2Na3[Al(OH)6](aq)
Al2O3(s) + 6HCl(aq) + 9H2O → 2[Al(H2O)6]3+(aq) + 6Cl–(aq)
Students are recommended to visit the Extramarks website and refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation on this topic of amphoteric oxide.
It is represented as O3. It is found in the upper atmosphere. It is prepared when using ultraviolet rays. They react with Oxygen from the upper atmosphere and split oxygen molecules into oxygen atoms. Then this oxygen atom combines with an oxygen molecule to form ozone:
3O2 —> 2O3
The reaction is endothermic approx. 142.7kJ of heat is needed. For preparing pure ozone, we use ozoniser, where the electric spark is passed through oxygen gas, and we get ozone.
Ozone is an allotrope form of oxygen. It is highly unstable in nature. We can find its traces about 20 kilometres above sea level. So, how is ozone present at enormous heights? It is developed by the reaction of Oxygen to the sun’s ultraviolet rays. The significant role of this gas is preventing the earth’s surface from the harmful ultraviolet radiations.
You know of the depletion of the ozone layer. Do you know why that is? We also use chlorofluorocarbons in refrigerators and other aerosols that release harmful things into the air, leading to gaps in the coating. Due to this, we get UV light, which causes many skin problems and malignancy diseases.
Apart from this, the several oxides of Nitrogen, particularly nitric oxide, react very rapidly with ozone to produce Oxygen and nitrogen dioxide. Hence, the mixed nitrogen oxides of Nitrogen that appear from the fumes frameworks of supersonic fly planes lead to depleting the layer.
As defined in Class 12 Chemistry Chapter 7 Notes, it has an angular structure, and two Oxygen -oxygen bond lengths in ozone molecules are identical (128pm), and the molecule is angular as expected with a bond angle of about 1170. It is a resonance hybrid of two primary forms:
In an extraordinary device, we can prepare ozone by passing a silent electric discharge through dry, unadulterated, and cold Oxygen. This device is what we understand as the ozoniser. In this process, we obtain gas of up to 10% concentration. It is an endothermic process, and we must take care to complete it at maximum temperatures. That is why we use a mute electric discharge. Students may refer to Extramarks Class 12 Chemistry Chapter 7 Notes for a detailed explanation of ozone structure.To enjoy the maximum benefit of these resources, students just need to register themselves at Extramarks official website and stay ahead of the competition.
If we want to produce higher ozone concentrations, we can do so by fractional liquefaction of an oxygen and ozone mixture.
Ozone is a harmful gas with a light blueish colour and has a fishy smell. It condenses at – 120°C to develop a dull blue fluid. On further cooling, it hardens to provide dark violet crystals. Thermodynamically, it is precarious and disintegrates to Oxygen. It is an exothermic process and is catalysed by numerous materials. However, we must know that a high concentration of gas can be hazardous.
Ozone is taken as a powerful oxidising agent. It is mainly because of the ease with which it gives out nascent oxygen atoms. It helps in oxidising:
a)Lead sulphide to lead sulphate:
4O3 + PbS → 4O2 + PbSO4
b)Iodide ions to iodine-
2KI + H2O + O3 → 2KOH + I2 + O2
We use this particular chemical reaction in the process of quantitative estimation of the gas. It liberates iodine when we bring ozone in contact with potassium iodide solution with a borate buffer (pH 9.2). We can titrate this released iodine against a standard solution of sodium thiosulphate with starch as an indicator. The chemical reactions that involve in this process are:
2I– + H2O + O3 2OH– + I2 + O2
I2 + 2Na2S2O3 Na2S4O6 + 2NaI
It also helps in oxidising nitrogen dioxide to dinitrogen pentoxide.
2NO2 + O3 → N2O5 + O2
It is used as an antiseptic, disinfectant and for sterilising water. It is also utilised for bleaching oils, ivory, flour, starch, etc. it acts as an oxidising agent in the manufacture of potassium permanganate.
Allotropes of Sulphur
Sulphur is a chemical element having the atomic number 16. It is easily accessible at room temperature.
It is a splendid yellow crystalline solid. Sulphur is a non-metal, as we know! The position of Sulphur in the modern periodic table is VI A.
Properties of Sulphur-
Following are some of the common characteristics of Sulphur. . More details about the chemical and physical properties will be covered separately.
Sulphur is yellow. It is insoluble in water. However, it is highly soluble in toluene (methylbenzene) and carbon disulphide solvent. It is a non-metal and, though, a poor conductor of electricity and heat. At this point, when we consolidate Sulphur vapour, we get a fine powder, which shapes a pattern resembling a flower.
Most metals and nonmetals are combined with Sulphur under critical conditions. Sulphur burns over the air, including a bright blue fire and develops Sulphur (IV) oxide & a small amount of Sulphur (VI) oxide.
It reacts with Hydrogen at high temperatures and creates hydrogen sulphide. Sulphur vapour combines with hot coke to produce a carbon disulphide fluid. Students may refer to our Class 12 Chemistry Chapter 7 Notes for more details on this topic.
Allotropic Forms of Sulphur</h3
We can find sulphur in several structures in the same physical state. However, the essential crystalline structures are rhombic or octahedral (α – sulphur) and monoclinic sulphur (β – sulphur). We observe the shape of rhombic sulphur at a temperature beneath 96oC. On the other hand, monoclinic sulphur is created at a temperature over 96oCentigrade.
This temperature of 96oCentigrade is the transitional temperature connecting the two crystalline structures. There is another allotrope of sulphur, polymeric sulphur (S8). It is an eight-part ring element.
This is insoluble in organic solvents and synthetic and natural rubbers. It also does not soluble in carbon disulphide solvent.
Class 12 Chemistry Chapter 7 Notes cover the properties of the two primary allotropes of sulphur: rhombic and monoclinic sulphur. A brief explanation of both is given below.
Rhombic Sulphur(⍺- sulphur)
This allotrope is yellow & translucent crystal structures. Alpha Rhombic Sulphur has a melting point of 114o C, and the density of rhombic Sulphur is 2.08 g/cm3. It is stable at temperatures below 96oC. Rhombic sulphur crystals are formed by evaporating the liquid of roll sulphur in CS2. It is partially soluble in water but dissolves in benzene, alcohol, and ether to some extent. It is readily soluble in carbon disulphide.
Monoclinic Sulphur(β sulphur)
These are transparent & amber crystals. This beta sulphur has a melting point of 119oC, and The density of monoclinic sulphur is 1.98 gcm3. It is not stable at temperatures below 96oC and changes into rhombic sulphur form. We must observe that at a temperature of 96oC or above, rhombic sulphur converts to kaleidoscopic and prismatic sulphur. At temperature 96oC or beneath, kaleidoscopic or prismatic sulphur converts to rhombic sulphur. Enantiotropic Allotropes alter their configuration from one form to another by a difference in the temperature. To separate the crust, colourless needle-shaped crystals of beta sulphur are created. It is stable above 369 K and converts into alpha sulphur below..
Colloidal Sulphur – We can produce this sulphur by passing hydrogen sulphide through a saturated and cooled solution of sulphur dioxide in water. Another method includes an explanation of alcohol and sulphur in the water solution.
It behaves as a solvent in carbon disulfide. We use it as a part of medicine.
Milk of Sulphur-
We can develop this by the action of weak hydrochloric acid on ammonium sulphide. Similarly, this milk of sulphur is created by the boiling of sulphur combined with calcium hydroxide (aqueous solution). We separate this mixture and add weak hydrochloric acid to get the milk of sulphur. This compound is non-crystalline and white. It is soluble in carbon disulphide solvent. When we heat it, it transforms to the conventional yellow colour of sulphur that we use as a part of medicines.
Note: Which form of sulphur shows paramagnetic behaviour? -In the vapour state, sulphur partly exists as an S2 molecule with two unpaired electrons in the antibonding pi orbitals(π*) like O2; hence, it shows paramagnetism. A repository of study materials is available on the Extramarks website and students can refer to Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of relevant topics from this chapter. To enjoy a better learning experience students need to sign up at Extramarks website and begin their preparation without any further delay.
Compounds of sulphur
What is Sulphur Dioxide
Sulphur dioxide is a standard gas that has a pungent smell. This chapter will look at this gas’s significance and preparation methods.
Methods of Preparation of Sulphur Dioxide Gas
As explained in Class 12 Chemistry Chapter 7 Notes, we can create sulphur dioxide in the laboratory by using dilute sulphuric acid on sulphites-
Na2SO3 + H2SO4 → Na2SO4 + H2O + SO2↑
Sodium sulphite Sulphuric acid Sodium sulphate Sulphur dioxide
Commercially, chemists produce maximum volumes of sulphur dioxide by heating a sulphide ore, for example, iron sulphide. They then liquefy this gas after drying under 25 atm pressure. It is generally stored in steel barrels. The method of roasting involves the following reaction-
Fe2S3 + 4O2 → FeO + 3SO2
Liquefaction at 25 atm pressure gives place as given below-
SO2(g) → SO2(l)
We can develop Sulphur dioxide likewise, on an extensive scale by burning sulphur(6 to 8%) in the air or Oxygen.
S(g) + O2 (g) → SO2(g)
Industrially, it is manufactured as a by-product of the roasting of sulphide ores.
4FeS2 + 11 O2 2Fe2O3 + 8 SO2
The gas, after drying, is liquefied under pressure and stored in steel cylinders.
Students will get an edge over their peers with superior knowledge and understanding of the subject. We recommend students also register on reliable online learning platforms such as Extramarks which strictly follows books and provides solved exercises and practice questions. It brings clarity to concepts which in turn becomes useful when it comes to answering the most difficult questions. Students will learn several other essential concepts from our Class 12 Chemistry Chapter 7 Notes. These study notes are based on the latest NCERT syllabus and have proven helpful, making it easier for students to revise the whole chapter quickly. Therefore, it is necessary to take guided practice and help to be aware of the mistakes and maximise your potential by learning how to frame the right answer and develop confidence.
Thus it will give students a significant edge when preparing for competitive examinations such as IIT JEE Mains and Advanced.
Structure of Sulphur dioxide-
Sulphur dioxide has an angular shape with an O-S-O bond edge of 119.50. We must indicate that sulphur dioxide has two unique pi bonds, i.e. p pi-p pi & d pi-p pi bonds. The two sulphur-oxygen bond lengths are the same or equivalent. This sulphur dioxide signifies that sulphur dioxide is a resonance hybrid of two canonical structures.
Physical Properties of Sulphur Dioxide
Sulphur dioxide is a dull or colourless gas. It has a very pungent odour and is highly soluble in water. It liquefies at room temperature under a pressure of two atm and a boiling point of 263 K. Its smell resembles smoulder sulphur. It is one of the unique gases to melt quickly. It is because it consolidates at room temperature under a pressure of 2 atmospheres.
Chemical Properties of Sulphur Dioxide
It is an acidic oxide. It is readily dissolvable in water. Sulphur dioxide breaks up in the water molecule and gives out sulphurous acid.
SO2 + H2O → H2SO3
Sulphur dioxide Water Sulphurous acid
It interacts vigorously with sodium hydroxide solution and creates sodium sulphite solution.
SO2 + 2NaOH → Na2SO3 + H2O
In these cases where we pass more sulphur dioxide into this arrangement, we observe sodium hydrogen sulphite.
SO2 + Na2SO3 + H2O → 2NaHSO3
The sulphur element in a sulphur dioxide atom is tetravalent. Subsequently, it can improve its covalency to six by specifically reacting with chemical elements like O2 and Cl2 molecules to shape the comparing addition compounds. For ex., it reacts with chlorine under the influence of charcoal as a catalyst to give sulphuryl chloride (SO2Cl2).
SO2 + Cl2 → SO2Cl2
Within sight of vanadium pentoxide(V2O5) as an impetus, it gives sulphur trioxide.
2SO2 + O2 → 2SO3
In a moist atmosphere, it can start giving nascent oxygen and, along these lines, go about as a reducing agent. For example, it reduces ferric salts to ferrous salt & halogens to halogen acids-
2Fe3+ + SO2 + 2H2O → 2Fe2+ + SO42- + 4H+
X2 + SO2 + 2H2O → SO42- + 2X – + 4H+
Identifying Tests for Sulphur Dioxide Gas
How do we identify and test for the presence of sulphur dioxide gas? Well, it’s elementary! We can perform a few simple steps to identify this gas. We will briefly discuss them here. . Students may refer to our Class 12 Chemistry Chapter 7 Notes for a full explanation about sulphur dioxide.
It decolourises acidified potassium permanganate KMnO4 solution.
It combines with potassium permanganate to give potassium sulphate, manganese sulphate and sulphuric acid.
5SO2 + 2KMnO4 + 2H2O → K2SO4 + 2 MnSO4 + 2H2SO4
It changes a filter paper moistened with acidified potassium dichromate K2Cr2O7 solution green
It reacts with potassium dichromate and sulphuric acid to give potassium sulphate and chromium sulphate-
3SO2 + K2Cr2O7 + H2SO4 → K2SO4 + 2Cr2(SO4)3 (ppt) + H2O
It turns starch iodide paper blue
It reacts with potassium iodate to give potassium hydrogen sulphate, iodine, and sulphuric acid.
5SO2 + 2KIO3 + 4H2O → 2KHSO4 + 3H2SO4 + I2
Sulphur dioxide is utilised in refining petroleum and sugar in bleaching wool and silk and as an anti-chor, disinfectant and preservative and used in the production of H2SO4 and paper made from wood pulp. Bleaching agent for soft materials such as wool, silk, and straw. They are used in petroleum and sugar refining. Sulphuric acid, sodium hydrogen sulphite & calcium hydrogen sulphite(industrial chemicals) are manufactured from sulphur dioxide. Liquid sulphur dioxide is utilised as a solvent to dissolve several organic and inorganic chemicals. We recommend students also register on reliable online learning platforms such as Extramarks which strictly follows books and provides solved exercises and practice questions. It brings clarity to concepts which in turn becomes useful when it comes to answering the most difficult questions. A repository of study materials is available on the Extramarks website and students can refer to Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of group 16 element sulphur and the relevant derivatives.
Oxoacids of Sulphur
Oxoacids are the acids which contain oxygen. Various experiments have shown sulphur to form several oxoacids. For example, these oxoacids are H2SO3,H2SO4, H2SO5, H2SxO6 (x= 2 to 5), H2SO4, H2S2O7,H2SO5,H2S2O8 etc. Some of these acids are not stable and cannot be isolated. They are known in aqueous solutions or the synthesis of their salts.
In the case of oxoacids, sulphur exhibits a tetrahedral structure concerning Oxygen. Usually, these oxoacids have a minimum of one S=O bond and one S-OH bond. We also observe terminal peroxide groups, terminal S=S, and terminal and bridging oxygen atoms in these oxoacids. Now, look at some of the most popular oxoacids and their properties. Structures of some essential oxoacids are shown below.
Students will learn about several essential concepts from our Class 12 Chemistry Chapter 7 Notes. These chapter notes have been proven helpful, making it easier for students to revise the whole chapter quickly. Thus giving them an upper edge in preparation for competitive examinations such as JEE mains and advanced.
Sulphuric acid is also shown as sulfuric acid or H2SO4. This sulphuric acid is an odourless, colourless, oily liquid and is very corrosive. A different name for it is the Oil of Vitriol. Because of its wide applications, it has been called the ‘King of Chemicals.’ Furthermore, we can form it in both combined and free states.
Manufacturing Process of Sulphuric Acid
In general, there are two techniques for the industrial production of sulphuric acid. They are:
- Lead chamber process
- Contact process
Now, let us discuss these processes in detail.
Sulphuric acid is prepared by the contact method, which involves three significant steps. The steps are briefly described below. However, for an in-depth understanding of these three steps, we recommend students to register on Extramarks website and get access to our Class 12 Chemistry Chapter 7 Notes.
- i) Production of Sulphur Dioxide: Sulphur dioxide is produced by heating sulphur or sulphide ores. For example, iron pyrites in excess of air. The critical step in the production of H2SO4is the catalytic oxidation of SO2 with oxygen to give sulphur trioxide in the presence of vanadium pentaoxide.
S (Sulphur) + O2(Oxygen) + Δ(Heating) → SO2(Sulphur dioxide) (in the presence of V2O5)
4FeS(Iron pyrites) + 7O2(Oxygen) + Δ(heating) → 2Fe2O3(Ferric Oxide) + 4SO2(Sulphur dioxide)
- ii) Formation of Sulphur Trioxide: Then sulphur dioxide is oxidised with atmospheric oxygen to sulphur trioxide by using V2O5 as a catalyst.
2SO2(Sulphur dioxide) + O2(Oxygen) → SO3(Sulphur trioxide) (V2O5is used as a catalyst)
iii) Conversion of Sulphur Trioxide into Sulphuric Acid to give oleum: The sulphur trioxide is broken into 98% sulphuric acid oleum. Another name for oleum is pyrosulphuric acid. Then Oleum is diluted with water to give sulphuric acid the desired concentration.
SO3(Sulphur trioxide) + H2SO4(Sulfuric acid-98%) → H2S2O7(Pyrosulfuric acid/Oleum)
H2S2O7(Pyrosulfuric acid/Oleum) + H2O(Dilution) → 2H2SO4(Sulfuric acid)
The H2SO4 obtained by the contact process is 96 to 98% pure. The contact process has been explained in detail in our Class 12 Chemistry Chapter 7 Notes.
Lead Chamber Process
The lead Chamber method is one of the most common manufacturing techniques. It results in around 50-60 B-grade acids. In this method, we use wet SO2 in the presence of nitrogenous oxides (dynamic impetus). As a result, it gets oxidised with the oxygen in the air and forms sulphur trioxide. This reaction is shown as-
2SO2 + O2 → 2SO3
Then sulphur trioxide is made to combine with water to get H2SO4. This reaction is expressed as
SO3 + H2O → H2SO4
Physical Properties of Sulphuric Acid
As explained in our Class 12 Chemistry Chapter 7 Notes, sulphuric acid is a thick, colourless, and syrupy oily fluid with corrosive nature. This sulphuric acid has a specific gravity of 1.84 at 298 K. The boiling point of the acid is 611 K. This chemical’s higher boiling point and thickness is due to hydrogen bonding. This strong chemical reacts with water vigorously, releasing quite a lot of heat. Thus, you must never add water to H2SO4. Instead, you should slowly add the acid to the water with proper stirring.
Chemical Properties of Sulphuric Acid
Sulphuric acid is a strong dibasic acid. Also, it is diprotic and ionises in two stages in the aqueous solution. This chemical is highly corrosive, reactive and soluble in water. It has a very high oxidising power and thus, acts as a strong oxidising agent. It has very low volatility.
For this reason, it plays a part in preparing more volatile acids from their compared salts. Concentrated sulphuric acid is a powerful dehydrating reagent. Thus, this chemical is used for drying several wet gases which do not interact with the acid. It furthermore expels water from natural mixes like starches. As it is the best oxidising agent, it can oxidise non-metals and metals. Moreover, it reduces sulphur dioxide. Students may refer to our Class 12 Chemistry Chapter 7 Notes where our subject matter experts have explained the topic of chemical properties of sulphuric acid in more detail.
Some Common Reactions of Sulphuric Acid
Hot concentrated sulphuric acid oxidises copper to copper sulphate.
Cu + 2H2SO4 → CuSO4 + SO2 + H2O
Concentrated sulphuric acid gives out hydrogen chloride HCl from sodium chloride NaCl. Also, it gives out hydrogen fluoride from calcium fluoride.
CaF2 + H2SO4 → CaSO4 + 2HF
It burns glucose, sugar, and starch molecules to carbon.
C12H22O11 + H2SO4(conc) → 12C + 11H2O
Uses of sulphuric acid:
Sulphuric acid is an essential industrial chemical. It is a common chemical in the preparation of fertilisers. Examples are ammonium sulphate & superphosphate. We use it in the development of dyes, shades, and paints. It is a common chemical in the manufacture of explosives. For example, TNT. Other imperative chemicals need the presence of sulphuric acid. Without sulphuric acid, we cannot get these chemicals. For example, hydrochloric, phosphoric and nitric acid. It is also needed for sodium carbonate. We use it as a part of petroleum refining and as a pickling agent. This chemical is common as a laboratory dehydrating and oxidising agent. It can be used to manufacture nitrocellulose products and the detergent industry. Metallurgical application, e.g., cleaning metals before enamelling, electroplating, galvanising, and storage batteries.
Group 17 Elements
The group seventeen elements include fluorine(F), chlorine(Cl), bromine(Br), iodine(I) and astatine(At) tennessine(Ts) from the top to the bottom. They are collectively called “halogens” because they give salts when they react with metals. The halogens are highly reactive nonmetallic elements like groups 1 and 2, and the element of group 17 show great similarity amongst themselves. Astatine and tennessine are radioactive elements. So, now we know what halogens are! Let’s now discuss the electronic configuration of these elements.
Fluorine and Chlorine are relatively abundant, while Bromine and Iodine are less so. Fluorine is present mainly as insoluble fluorides(fluorosparCaF2, cryolite Na3AlF6 and fluorapatite 3Ca3(PO4)2. CaF2 ), and small quantities are present in the soil, river water plants and bones and teeth of animals.
Electronic Configuration of Group 17 Elements
The valence shell electronic configuration of these electrons is ns2np5. Therefore, seven electrons are in the outermost orbital shell of these elements. The element skips out on the octet configuration by only one electron. Therefore, these elements look to either lose one electron and develop a covalent bond or gain one electron and create an ionic bond. Therefore, these are very reactive non-metals. Many of these elements have seven electrons in their outermost shell, one electron short of the next noble gas.
Fluorine (F) [He]2s2,2p5
Chlorine (Cl) [Ne]3s2,3p5
Astatine (At) [Xe]4f14,5d10,6s2,6p5
Let us now discuss the several nuclear properties of the group seventeen elements. We will talk about the ionic and atomic radii, ionisation enthalpy and more.
i)Ionic and Atomic Radii
In a respective period, The nuclear and atomic radii of these elements keep increasing as we move down the group. It happens because of the addition of an extra energy level. They have minimal atomic radii compared to the other elements in the corresponding periods. These ionic and nuclear radii can be attributed to their atomic charge being quite powerful.
These elements have higher ionisation enthalpy. This value keeps on diminishing as we go down the group. This ionisation enthalpy happens due to the enhancement in the size of the nucleus. Fluorine has the highest ionisation enthalpy than any other halogen, thanks to its minute size.
iii) Electron Gain Enthalpy
The electron gain enthalpy of all these elements becomes lower negative upon moving to go down the group. Fluorine has lower enthalpy than Chlorine. We can attribute it to the tiny size and the smaller 2p subshell of the atom of Fluorine. As a result, there is interelectronic solid repulsion in the relatively small 2p orbitals of Fluorine. Hence, the incoming electron does not experience much attraction.
The halogens show high electro-negativity values. Though, it diminishes slowly on moving down the group from Fluorine to Iodine. This electro negativity can be attributed to the increase in nuclear radii upon moving down the group. Fluorine is the best electronegative element in the modern periodic table.
v) Physical Properties
Let us now discuss the various physical properties of these halogens. Briefly, these properties are described below. All physical properties of halogens are explained by our Chemistry experts in our Class 12 Chemistry Chapter 7 Notes. Students can register on Extramarks website and get access to our study resources.
Physical state: The seventeen elements are found in diverse physical conditions. For example, Fluorine and Chlorine are gases. On the other hand, Bromine is a liquid & Iodine is a solid condition.
Colour: These halogen elements have a variety of colours. For example, while Fluorine is pale yellow, Iodine is dark violet.
Fluorine: Dull yellow. Chlorine: Greenish yellow Bromine: Reddish brown. Iodine: dark violet.
Solubility: Fluorine and Chlorine are water-soluble. Conversely, Bromine and Iodine are significantly less soluble in water. Still, they are soluble in various organic solvents such as chloroform, carbon tetrachloride, carbon disulphide and hydrocarbons to give coloured solutions.
Melting and boiling points: Melting & boiling points of these halogen elements increase, and van der Waal forces also increase as we move down the group from Fluorine to Iodine. Fluorine has the lowest boiling and melting points.
Oxidation states: They show variable oxidation states as-
Chlorine: -1, +1, +3, +7
Bromine : -1, +1, +3, +5, +7
Iodine: -1, +1, +3, +5, +7
We will now discuss some of the chemical properties of these elements below. The same has been covered in detail in our Class 12 Chemistry Chapter 7 Notes.
i) Oxidising Power
All halogens are great oxidising agents. In the list, Fluorine is the most powerful oxidising agent. It is capable of oxidising all the halide particles to halogen. The oxidising power reduces as we move to go down the group. The halide particles also behave as reducing agents. Although, their reducing capacity decreases the group as well.
ii) Reaction with Hydrogen
All halogens combine with hydrogen and form acidic hydrogen halides. The acidity of these all hydrogen halides reduces from HF to HI. Fluorine reacts violently, and Chlorine requires sunlight rays. Conversely, Bromine reacts upon heating, and Iodine needs a catalyst.
iii) Reaction with Oxygen
Halogens interact with oxygen to create oxides. Though, it has been found that the oxides are not steady. Besides oxides, halogens also develop several halogen oxoacids and oxoanions.
iv) Reaction with Metals
As halogens are very reactive, they combine with most metals instantly and produce the resulting metal halides. For ex., sodium reacts with chlorine gas and comprises sodium chloride. This process is exothermic and gives out a bright yellow light and a lot of heat energy.
2Na(s) + Cl2(g) → 2NaCl(s)
Metal halides are ionic. It is because of the halogens’ high electronegative nature and the metals’ maximum electropositivity. These ionic characteristics of the halides reduce from Fluorine to Iodine.
To learn more about s block elements, students are recommended to refer to Extramarks Class 12 Chemistry Chapter 7 Notes. These study resources are prepared and reviewed by Chemistry experts and are authentic and reliable. It will give students a better understanding of the concepts based on NCERT books and in accordance with the latest CBSE guidelines and syllabus.
Anomalous Behaviour of Fluorine
Fluorine’s abnormal behaviour is caused by its tiny size, strong electro-negativity, lower F-F bond dissociation enthalpy, & lack of d orbitals in the valence electron shell. Fluorine’s reactions are primarily exothermic (owing to the short and robust bonds it forms with other elements). It only creates one oxoacid, whereas other halogens have many oxoacids.
Because of strong hydrogen bonding, hydrogen fluoride is liquid (b.p. 293 K). Different hydrogen halides exist in the form of gases.
Uses of halogens
Fluorine compounds constitute an essential ingredient in toothpaste. It is because fluoride compounds react with the enamel of the teeth and take care of teeth rotting. Chlorine is mainly used as a bleach. It is also applicable in the metallurgy of elements like platinum and gold. Iodine is an antiseptic because it kills the germs on the skin. More real-life examples of usage of halogens are covered in Class 12 Chemistry Notes. .
Chlorine was discovered in 1774 by Scheele, who manufactured gas by the activity of HCl hydrochloric acid on MnO2 manganese dioxide. In 1810, Scientist Davy built up its elementary nature and recommended the name Chlorine by its colour. It is greenish-yellow, & it has a pungent odour.
- i) We can prepare the gas by heating manganese dioxide and concentrated hydrochloric acid. We can also synthesise the gas by the activity of hydrochloric acid on bleaching powder (or) lead dioxide potassium (or) permanganate. Various study materials are available on the Extramarks website including Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of group 16 and 17 elements and its related topics.
MnO2 + 4HCl MnCl2 +Cl2 + 2H2O
16 HCl + 2KMnO4 MnCl2 +2KCl + 8H2O + 5Cl2
4HCl + PbO2 PbCl2 +Cl2 + 2H2O
Manufacture of Chlorine
We can acquire the gas by electrolysis salt water in a Nelson cell. It is the least expensive technique and gives the purest form of gas.
Chlorine is liberated at the anode. It is also developed as a by-product in several chemical industries. Students should refer to Extramarks Class 12 Chemistry Chapter 7 Notes for a more detailed and elaborate explanation of chlorine and other such relevant topics.
In this procedure, we can develop the gas by oxidation of hydrochloric acid in the presence of a cuprous chloride catalyst at 723K and a pressure of 1 atmospheric oxygen.
4HCl + O2 2Cl2 + 2H2O
Chlorine is a greenish-yellowish gas and has a pungent and suffocating odour. It has a boiling point of 239.1 K and a melting point of 171.6K. The chlorine gas is harmful to nature. It is 2-5 times denser than air.
It can be effectively condensed. The gas is marginally dissolvable in water. Its valency is seven, and it has an excellent affinity for hydrogen. It interacts with compounds containing hydrogen to form HCl hydrochloric acid.
H2 + Cl2 2HCl
H2S + Cl2 2HCl + S
Hydrolysis in water gives an explosively smelling, yellow arrangement- chlorine water. It drains its yellow colour when it remains in the sunlight. It is because of the structure of a hydrochloric and hypochlorous acid blend.
Hypochlorous acid, being unsteady, dissolves and releases nascent oxygen. The oxygen shape is in charge of Chlorine’s bleaching and oxidising properties.
Cl2 + H2O HCl + HOCl (presence of sunlight)
2HOCl + 2HCl 2O nascent oxygen (decompose)
It combines particularly with all non-metals except Nitrogen, carbon, and oxygen. It reacts very fast with most metals. This chemical reaction concludes in the formation of chlorides.
It has a special attraction for hydrogen molecules. It interacts with hydrogen in the presence of light with a blast to develop hydrochloric acid. H2 + Cl2 2HCl (light)
Chlorine breaks down a few hydrogen compounds to shape hydrochloric acid. It is a decent oxidising agent; it oxidises ferrous to ferric, sulphur dioxide to sulphuric acid, sulfites to sulphates, and Iodine to iodic acid. Because of the release of nascent oxygen, Moist Chlorine is a powerful bleaching agent. It fades organic matter or vegetables. With slaked lime, it frames bleaching powder.
2Ca(OH)2 + 2Cl2 Ca(OCl)2 + CaCl2 + 2H2O
It reacts with unsaturated hydrocarbons to give additional products and substitution products with saturated hydrocarbons. Students can learn more about the chemical properties of chlorine from our Class 12 Chemistry Chapter 7 Notes.
Concerning Class 12 Chemistry, Chapter 7 Notes, Chlorine is a chemical that counteracts bacterial development in stationary water. It’s used to purify sewage and commercial waste. It’s additionally an active ingredient in some cleaning items.
Chlorine poisoning generally happens when you take in or breathe in the chemical. It interacts with water — incorporating the water into your digestive tract. It, hence, forms hydrochloric acid and hydrochlorous acid. Both of these substances are very dangerous to our bodies. Students can refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of chlorine molecules.
You might be most acquainted with Chlorine utilised as a part of pools. Be that as it may, most chlorine poisoning occurs because of ingesting household cleaners, not swimming pool water. A couple of everyday household items and substances containing Chlorine include:
Tablets utilised as a part of swimming pools Swimming pool water Mild household cleaners Bleaching products. We can utilise it as a bleaching agent. It is used in the wood pulp, cotton and textile industry businesses. We can usually use it to clean drinking water.
Chlorine is an antiseptic and disinfectant in swimming pools. It can be utilised in the extraction of gold and platinum. In the arrangement of harmful gases, for example, phosgene, mustard gas and tear gas, we use Chlorine. It is used in developing dyes, drugs and organic compounds such as carbon tetrachloride, chloroform, DDT, refrigerants, etc., in sterilising drinking water. Students may refer to NCERT and CBSE Solutions in addition to Class 12 Chemistry Chapter 7 Notes for a more detailed explanation.
Concerning CBSE Class 12 Chemistry Chapter 7 Notes, hydrochloric acid was discovered by Glauber, and Davy was the one who said that it consists of Hydrogen and Chlorine.
Manufacture of Hydrogen Chloride
We manufacture Hydrogen Chloride in the laboratory by treating sodium chloride with concentrated sulphuric acid. We then heat this mixture to 420K.
NaCl + H2SO4 → NaHSO4 + HCl at 420K
We get Sodium bisulphate as an insoluble by-product. Though, we further mix it with more NaCl sodium chloride. This mixture has to be further heated to a higher temperature of around 823K. It gives dissolvable sodium sulphate and HCl gas.
NaHSO4 + NaCl → Na2SO4 + HCl at 823K
We dry this HCl hydrochloric acid by treating it with concentrated sulphuric acid. HCl is not dehydrated over phosphorus pentoxide or quick lime. This hydrochloric acid is because it reacts with both of these compounds.
Properties of Hydrogen Chloride
Hydrochloric acid is a colourless gas with a very sharp and pungent-smelling odour. It can melt quickly to a colourless fluid at 189K.
HCl hydrogen chloride forms a white solid at 159K upon freezing. It is highly soluble in water. An aqueous solution of HCl hydrogen chloride is what we say of hydrochloric acid. It has a higher dissociation constant and is a strong acid. It interacts with metals and salts to give various chlorides. For example, it interacts with zinc to form zinc chloride. It is highly soluble in water and ionises as follows.
HCl + H2O H3O+ + Cl– ka =107
Its aqueous solution is called hydrochloric acid. The high value of the dissociation constant (Ka) indicates that it is a strong acid in water.
When we use three parts of concentrated HCl and one part of focused HNO3 mixed, aqua regia is formed for dissolving noble metals, e.g., gold and platinum.
Hydrochloric acid combines with iron to form ferrous chloride.
Fe + 2HCl → FeCl2 + H2
It is used to manufacture Chlorine, ammonium chloride and glucose(from starch). It is utilised for extracting glue from bones and purifying bone black. It is used in medicines and as a laboratory reagent.
Group 17 Elements: Halogens
Group seventeen elements are Fluorine, Chlorine, Bromine, Iodine, and astatine from top to bottom. We understand them by ‘halogens’ because they all are salt producers. The members of this seventeen group are highly similar to each other. They show a regular pattern in their physical and chemical properties. Did you know that the astatine element is the only radioactive element in the seventeenth group? These halogen elements have seven electrons in their outer electron shell. Their electronic configuration is ns2 np5.
We can discuss that they are one electron less from the nearest inert gas or octet configuration. These elements have a tiny size because of their effective nuclear charge. Though, these do not possess the tendency to lose electrons. They gain an electron quickly and complete their octet configuration.
Halogens form various oxoacids. These are nothing but the acids consisting of oxygen in the acidic group. The topic has been explained in further detail in our Class 12 Chemistry Chapter 7 Notes. Our study notes are prepared by Chemistry subject experts and are considered to be the most reliable, accurate and authentic learning materials.
Oxoacids of Halogens
An oxoacid is a compound containing hydrogen, oxygen, and not less than one other element. These oxoacids do not have any less than one hydrogen molecule holding to oxygen, and hydrogen can separate into the H+ cation and the acid’s anion.
The fluorine atom is tiny, and thus, it is highly electronegative. Hence we can develop into a single oxoacid, HOF(hypofluorous acid), fluoric(I) acid or hypofluorous acid. The other elements of the halogen family create several oxoacids.
We cannot isolate them in their pure state, which is stable in an aqueous solution. It is also very stable in its salt states. Halogens commonly form four types of oxoacids, namely hypohalous acids (+1 oxidation state), halous acids (+3 oxidation state), halic acids (+5 oxidation state) & perhalic acids (+7 oxidation state).
Structures of Oxoacids of Halogens
We can describe that the focal halogen molecule is sp3 hybridised in these oxoacids. We can present an X-OH bond in each oxoacid. In most of these oxoacids, “X = O” bonds are available. Hypohalous acids incorporate hypofluorous, hypochlorous, hypobromous, and hypoiodous acids. The halogen has the oxidation state of +1 in hypohalous acids. Students can refer to our Class 12 Chemistry Chapter 7 Notes for a more detailed explanation about the structure of oxoacids of halogens.
Some More Examples
As covered in our Class 12 Chemistry Chapter 7 Notes, chlorine can form four different types of oxoacids. These are HOCl (hypochlorous acid), HOClO (chlorous acid), HOClO2(chloric acid) and finally, HOClO3 (perchloric acid). Bromine creates HOBr (hypobromous acid), HOBrO2(bromic acid) and HOBrO3 (perbromic acid).
Iodine develops HOI (hypoiodous acid), HOIO2 (iodic acid) and HOIO3 (periodic acid).
The central nuclear atom in the oxoacids is sp3 hybridised. Every oxoacid has essentially one X-OH bond. At the same time, most oxoacids have X=O bonds present in them. Interhalogen Compounds
We can describe Interhalogen Compounds as the subordinates of halogens. These are the compounds containing two unique sorts of halogens. For example, the common interhalogen compounds include Chlorine monofluoride, bromine trifluoride, iodine pentafluoride, iodine heptafluoride, etc.
Types of Interhalogen Compounds as covered in Class 12 Chemistry Chapter 7 Notes
We can ultimately divide interhalogen compounds into four steps, depending on the number of atoms in the elements. They are as follows:
- XY Compounds
- Compounds XY3
- Compounds XY5
- XY7 Compounds
We must understand that “X” is the more significant (or) less electronegative halogen in these notations. On the other hand, “Y” represents the more miniature (or) more electronegative halogen. We can calculate the number of elements in the atom by the concept of the radius ratio. The general formula for the same is given as follows-
so Radius Ratio = Radius of bigger Halogen Particle / Radius of smaller Halogen Molecule.
With an enhancement in the radius proportion, we consider that the number of atoms per molecule also increases. However, we can determine that Iodine heptafluoride possesses the most significant number of particles per atom. This halogen is because it has a magnificent radius proportion.
Preparation of Interhalogen Compounds
As covered in our Class 12 Chemistry Chapter 7 Notes, we can manufacture these interhalogen compounds using two main methods. One includes the direct mixing of halogens, and the other consists of a reaction of halogens with the lesser Interhalogen compounds.
The halogen atoms interact to form an interhalogen compound. One example is when a volume of Chlorine reacts with equal importance of Fluorine at 473K. The resultant product is chlorine monofluoride.
Cl2 + F2 2ClF at 437 K
In another case, a halogen atom acts with another lower interhalogen to create an interhalogen compound. For ex., Fluorine reacts with iodine pentafluoride at 543K. These interhalogen compounds give rise to the combination of Iodine Heptafluoride.
Properties of Interhalogen Compounds
We can observe Interhalogen compounds in vapour, solid or fluid states. A lot of these compounds are not stable solids or fluids at 298K. Some other combinations are gases as well. As an example, chlorine monofluoride is a gas. Conversely, bromine trifluoride and iodine trifluoride are solid and liquid, respectively. These compounds are covalent and diamagnetic, attributing to the lesser electro-negativity between the bonded molecules. Examples include Chlorine monofluoride, Bromine trifluoride & Iodine heptafluoride.
These compounds are covalent, and these interhalogen compounds are diamagnetic. These interhalogen compounds are because they have bond pairs and lone pairs. Interhalogen compounds are very reactive. One exception to this is Fluorine. The A-X bond in interhalogens is much feeble than the X-X bond in halogens, except for the F-F bond. We can utilise the VSEPR theory to discuss the particular structure of these interhalogens. In chlorine trifluoride, the primary central atom is Chlorine. It has seven electrons in its outermost electron valence shell. Three of these electrons develop three bond pairs with three fluorine molecules leaving four electrons. Students may refer to our Class 12 Chemistry Chapter 7 Notes to get a better understanding about the topic of the properties of interhalogen compounds.
Common Shapes of these Compounds
Applying the VSEPR theory, we consider that it forms a trigonal bipyramidal structure. The lone pairs take up the tropical state. On the other side, bond pairs take up the three different positions. The axial bond pairs bind towards the tropical position. This minimises the repulsions that happen due to lone pair-lone pair bonds. Though, it has the shape of a bowed ‘T’.
Image Name: Chlorine Pentafluoride
Let us now discuss the case of Iodine pentafluoride. The central nuclear atom in Iodine pentafluoride is the iodine atom. They have one lone pair and five bond pairs. This iodine pentafluoride is the reason it has a square pyramidal shape. Similarly, let us consider another case of Iodine heptafluoride. It has seven bond pairs & has the shape of a pentagonal bipyramidal structure.
Image Name: Iodine pentafluoride and Iodine heptafluoride
Students may refer to Extramarks study material Class 12 Chemistry Chapter 7 Notes for a more detailed explanation on this topic.
Uses of Interhalogen Compounds
We use interhalogen compounds as non-watery (nonaqueous)solvents. Also, we use these compounds as fluorinating catalysts in several reactions. We use UF6 in the enrichment of 235U. We can produce this by using ClF3 and BrF3.
U(s) + 3ClF3(l) → UF6(g) + 3ClF(g)
We use these compounds as fluorinating compounds.
Group 18 Elements
Concerning Class 12 Chemistry Chapter 7 , The Group eighteen elements having Helium(He), Neon(Ne), Argon(Ar), Krypton(Kr), Xenon(Xe), & Radon(Rn)elements. All these elements are referred to as noble gases or inert gases. These rare gases mean that these elements are chemically inert and do not take part in any reaction.
Helium, the lightest of the noble or inert gases, had been detected, and Helium is the only element in the periodic table discovered by an astronomer. Helium is the element you can find on the upper right side of the periodic table with atomic number 2. It comes first in the family of noble gases. It held one nuclear atomic orbital and was named by Frankland and Lockyer. Its name is obtained from the Greek word “Helios”, meaning Sun. Scientists understood there was an enormous amount of Helium in the Sun before it was discovered. Helium falls under noble gas since its outermost electron orbital is full of two electrons. It can also be found in lasers, compressed air tanks and coolants in nuclear reactors. It binds the lowest boiling and melting points among all other groups of eighteen elements. The nuclear fusion of hydrogen in stars creates a significant amount of Helium. Helium has two well-known stable isotopes – 3He and 4He. The abundance of helium-3 & helium-4 corresponds to 0.0002% & 99.9998%, respectively. This difference in abundances can be formed in the Earth’s atm, where the ratio of 4He atoms to 3He atoms is approx 1000000:1.The normal use of Helium goes in altitude research and meteorological balloons. It is utilised as an inert protective gas in autogenic welding. There is the only cooler capable of declining temperatures lower than 15K (-434ºF). Helium gas is also used in the production of germanium crystals and silicon crystals. Since it can diffuse through solids much faster than air, Helium is utilised industrially for pipeline leak detection. This inert gas element is also used in gas chromatography as a carrier gas. Liquid Helium has various applications in cryogenics, magnetic resonance imaging (MRI), and superconducting magnets due to its lesser melting point. Students can study more about rare gases from our Class 12 Chemistry Chapter 7 Notes to gain a better and comprehensive understanding about all topics covered in Class 12 Chemistry Chapter 7.
Neon element is reddish-orange coloured in neon lamps and vacuum discharge tubes and is the second-lightest noble gas.
It is not as expensive a refrigerant as Helium in various applications. Its freezing capacity is 40 times more than liquid Helium and three times liquid hydrogen per unit volume basis. It is a rare gas whose molecules consist of a single Neon atom. It is a chemically inert gas and non-toxic. There is no harm to the environment, and it has no impact since it’s non-reactive and does not form compounds. This element causes no ecological damage. It can create exotic combinations with Fluorine in laboratories, an inert element. It also includes an unstable hydrate. The reddish-orange coloured neon lights are utilised in making advertising signs. It’s also generally used in these lights when many other gases are needed to create rays of different colours. Different uses of neon include lightning arrestors, high-voltage indicators, television tubes and metre tubes. Gas lasers are formed with the help of Neon and Helium elements. The electronics industries use neon singly or in complexes with other gases in many types of gas-filled electron tubes. A variety of Helium and neon elements is used for respiration by marine divers in the sea or water bodies since Helium is less soluble in blood than Nitrogen at high pressure.
Argon is a chemical particle in the eighteen groups of the modern periodic table. It is a noble gas, the third most abundant gas in Earth’s atm.
Argon is the most normal gas in the atmosphere besides Nitrogen and Oxygen elements. Argon is a noble gas (like Helium) that is inert. It is an odourless, colourless gas that is inert to other substances. Under extreme conditions, argon can develop certain compounds even though it is a gas. The same solubility level characterises it in water as that of oxygen. It has low thermal conductivity.
They are used in metal industries, and it is used in the production of titanium. It is used in double-dazzled windows to fill the space between the panels. Various study materials are available on the Extramarks website including Class 12 Chemistry Chapter 7 Notes for a more detailed explanation of group 18 elements and other topics.
Specific Facts About Argon: Argon was suspected to be present in the air by scientist Henry Cavendish in the year 1785. According to Chimcool, the bulk of argon is the isotope argon-40 that emerges from the radioactive decay of potassium-40.
Krypton is a chemical element and a noble gas of Group 18 with atomic number 36 and the symbol Kr in the modern periodic table. This gas is approx three times heavier than the air. It is colourless, tasteless, monoatomic, and odourless. This gas is more plentiful in Earth’s atmosphere as its trace is in minerals and meteorites. Earth’s atmosphere contains 1.14 parts per million of the volume of Krypton. It is used in various electric and fluorescent lamps and a flash lamp employed in high-speed photography. They could even react with the very reactive gas fluorine. Xenon is a chemical element with the symbol Xe and atomic number 54 in the periodic table. William Ramsay discovered it in the year 1898.
Xenon is a rare, colourless, odourless, tasteless, chemically not reactive gas. Xenon is a trace gas ( i.e., that makes up less than 1 % by volume of Earth’s atm. It is a component of gases released from some mineral springs. It is also released as a by-product if the air is separated into Nitrogen and oxygen. The element has a molecular atomic number 54 as its nucleus consists of 54 protons. It is present in all forms, i.e., solids, liquids, and gases.
Optics and illumination: Flash lamps are used in Xenon flash lamps. It is also used in Stroboscopic lights and photographic flashes. Lasers are formed with the help of Xe gas.
Medicine, element Xe, behaves as a natural anaesthetic. Inhaling the mixture of oxygen and Xenon formed a hormone that helps to increase Red Blood Cell (RBC) formation. It is used to calculate blood flow and to image the Brain, Heart, Lungs etc. Also, the element is used in the NMR spectroscopy method.
Interesting Facts about Xenon: The name Xenon is derived from the Greek word Xenos which means “stranger”. It is the most expensive and dense of all the gases and produces a bluish purple colour when electrified. It forms very promising compounds with fluorine.
XeF4 and XeF6 are expected to be
XeF4 oxidises potassium iodide
XeF4 + 4I– → 2I2 + 4F– + Xe
XeF4 and XeF6 are expected to oxidise because both reactions, XeF4 and XeF6, caused the oxidation. XeF6 oxidises hydrogen-like other xenon fluorides.
Electronic Configuration of Group 18 Elements
The general configuration of the valence electron shell is ns2np6. The exception is Helium, which has the electronic design of 1s2. They are chemically inert as they already have the octet configuration in their valence shells. They have zero valencies.
All noble gases have a general electronic configuration for them are ns2np6
Helium 1s 22s2
Krypton 1s2,2s2,2p6,3s2,3p6 4s2 3d10 4p6
The Occurrence of These Elements
All of these noble gases form in a free state in the atmosphere. Apart from Radon, every other noble gas is present in the atm. Argon only constitutes 0.93% of the total atm. We can synthesise this element by the fractional distillation of liquid air, and We can observe neon, Helium and argon in specific water springs as disintegrated gases.
Xenon and Radon are the inert elements of the group. Also, we can form Radon by the decay of radium and thorium minerals. The complete set of group 18 elements can be studied in detail from Extramarks Class 12 Chemistry Chapter 7 Notes. Our study materials are prepared by Chemistry subject experts with years of experience and are based on the latest CBSE curriculum.
Students may refer to reliable study resources such as Extramarks Class 12 Chemistry Chapter 7 Notes for a more detailed explanation. .
Trends in the Atomic Properties
Atomic Radius: The nuclear radii increases on moving to go down the group with an increasing nuclear number. This atomic radius results from expanding another shell at each progressive element’s moving go down the group.
Electron Gain Enthalpy: Group eighteen elements exhibit very stable electron configurations. They do not tend to accept an electron and, therefore, have significant positive values of electron gain enthalpy.
Ionisation Potential: They have maximum ionisation potentials, thanks to their closed electronic configurations. This value decreases on moving to go down the group due to an increase in the atomic size.
Physical Properties: Due to their stable nature, we find these elements as monatomic gases in a free state. They are colourless, tasteless and odourless gases. Sparingly soluble in water. The particles of these rare elements have weak Van der Waals interaction forces. This force enhances moving down the group.
It is due to an expansion in the polarising capacity of the molecules. They exhibit low melting and boiling points, which we can attribute to the present weak dispersion forces. The melting & boiling points increase as we move to go down the group. We can condense these elements at shallow temperatures. As the atoms’ size increases, the group decreases, and the ease of liquefaction also increases.
Helium has the lowest boiling point of 4.2 K of any known substance and has the unusual property of diffusing through commonly used laboratory materials like rubber, glass or plastics. Noble gases are the least reactive. Their chemical inertness is due to a filled valence shell. The topic has been covered in more detail in our Class 12 Chemistry Chapter 7 Notes with further explanations about p block elements.
- These rare elements are chemically latent because of their stable electronic configuration.
- Group eighteen elements have high positive electron gain enthalpy and high ionisation enthalpy.
- In 1962, Scientist Neil Bartlett anticipated that Xenon should interact with platinum hexafluoride. A scientist was the first to set up a compound of Xenon known as Xenon hexafluoroplatinate(V). Later, many combinations of Xenon were integrated, including fluorides, oxyfluorides, & oxides.
Xe + PtF6 → Xe[PtF6]
Xenon Platinum Hexafluoride Xenon Hexafluoroplatinate(V)
- The chemical movement of group eighteen elements increases with a diminishment in the ionisation enthalpy on moving down the group.
- The ionisation enthalpies of Helium, argon, and neon are too high for them to shape compounds.
- Krypton only forms krypton difluoride since its ionisation enthalpy is marginally higher than Xenon.
- Although Radon has less ionisation enthalpy than Xenon elements, it shapes just a few compounds like radon difluoride and some complexes; hence Radon has no steady isotopes. At any point, Xenon shapes an especially more prominent number of compounds.
The cages created by the crystal structure of water molecules trap rare gas molecules. Xe atoms will be trapped in the cavities (or shell) produced by the water molecules in the crystal structure of ice during its creation. Clathrate chemicals are the result of this process.
Chemical bonding does not exist. Clathrates compounds don’t have a precise chemical formula, although it’s roughly six water molecules to 1 inert gas molecule. The hollow is just a hair smaller than an atom of a noble gas.
Other organic liquids, dihydroxybenzene, can produce such molecules (for example, quinol solvent).
Compounds of Xenon & Fluorine
Xenon easily combines with fluorine to create xenon fluorides:
Preparation: XeF2 is prepared by heating an excess of Xenon with Fluorine at 673K and 1 bar pressure in a sealed nickel tube. On rapid cooling, a colourless solid of XeF2 is formed.
Xe + F2 —> XeF2
(Xenon) Fluoride XenonFluoride
2: 1 (a catalyst is a nickel, 673k and 1 bar )
Structure: As explained by our experienced Chemistry faculty of Class 12 Chemistry Chapter 7 Notes, the central xenon atom is encircled by five electron regions (Steric Number =5). The five electron regions involve three lone pairs and two Xe–F bond pairs. The lone pairs hold the equatorial positions, while the Xe–F sigma bonds occupy the axial positions. Xenon in XeF2 is sp3d hybridised and shows a trigonal bipyramidal symmetry structure. In this symmetry, the three lone pairs are aligned at an angle of 120o. Because of this symmetrical arrangement of lone pairs, the net repulsions on the Xe–F bond pairs are zero 180o from each other. Thus XeF2 has linear geometry.
Image name : structure of XeF2 and XeF4
The reaction of it with water: XeF2, XeF4 and XeF6 are colourless crystalline solids and sublime readily at 298 K. They are powerful fluorinating agents.
(b) XeF4 xenon tetrafluoride:
Preparation: XeF4 is prepared by heating a mixture of Xenon and Fluorine in the molecular Ratio 1:5 at 873K and 7 bar pressure in a sealed nickel tube. Various study materials are available on the Extramarks website such as Class 12 Chemistry Chapter 7 Notes which will help students to understand the topics of group 18 elements with conceptual clarity and provide deep learning experience so that students need not look elsewhere to supplement their studies to better their performance.
Xe + 4F2 XeF4
1:5 ratio 873 K and 7 bar
The structure of xenon tetrafluoride is square planar shown above.
Image name: Xenon tetrafluoride(XeF4)
The central xenon atom is encircled by six electron regions (Steric Number =6), and the six electron regions consist of two lone pairs and the four Xe–F bond pairs. The lone pairs settled the axial positions, while the Xe–F sigma bonds occupied the equatorial positions. Xenon in XeF4 is sp3d2 hybridisation and holds an octahedral electron pair geometry & square planar molecular geometry structure.
Reaction with water: At normal temperatures, they are explosive.
XeF4 + H2O Xe + HF + XeO3 + O2
XeF4+ H2O HF + XeOF2 at -80 degree celsius
(c) XeF6 xenon hexafluoride:
Preparation: XeF6 is prepared by heating a mixture of Xenon and Fluorine in the Ratio of 1:20, 573K and 60−70 bar in a nickel vessel.
Xe + 3F2 XeF6
1:20 (catalyst nickel, temperature 573 K and 60-70 bar)
Structure: The central xenon element is surrounded by seven electron pairs, of which six Xe–F bonding pairs & one lone pair of electrons are present. With seven electron pairs,
Xe is sp3d3 hybridised, and XeF6 adopts a distorted octahedral geometry. The fluorine atoms settle the vertices of the octahedron while the lone pairs move in the space and are shown at the centre of one of the triangular faces to avoid repulsion.
Because of one lone pair of electrons, the molecule will have lone-pair –bond-pair repulsion, and the molecular geometry will be distorted by octahedral structure.
Image name: Xenon hexafluoride(XeF6)
Reaction with water: partial hydrolysis of XeF6 gives oxyfluorides, XeO4 and XeO2F2.
XeF6 + H2O → XeO4 + 2HF
XeF6 + 2H2O → XeO2F2 + 4HF
(d) XeO3 xenon trioxide
Preparation: XeO3 is prepared by the slow hydrolysis of XeF6 and XeO4
XeF6 + 3H2OXeO3 + 6HF
The hybridisation is sp3.
The geometry is pyramidal.
The structure is: The central Xe atom in XeO3 has three bonding pairs and one lone pair of electrons. Hence, the Xe atom is Sp3 hybridised. The electron geometry of XeO3 is tetrahedral, and the molecular geometry is pyramidal.
(e) XeO4 (xenon tetroxide)
The partial hydrolysis of xenon hexafluoride prepares XeOF4:
XeF6 + H2O→XeOF4 +2HF
The hybridisation is sp3. The geometry is tetrahedral. In XeOF4, the Xe atom forms the central metal atom. The main Xe element has one lone pair of electrons and five bonding pairs. Hence, it undergoes
sp3d2 hybridisation. The electron pair geometry is octahedral structure, and the molecular geometry is square pyramidal.
The structure is given below:
Image name: xenon tetraoxide (XeO4)
(f) XeOF2 (xenon oxy fluoride)
Preparation: XeOF2 is prepared by the slow and partial hydrolysis of XeF4 at low temperatures.
XeF4 +H2O→XeOF2 +2HF
Structure: The hybridisation is sp3d. In XeOF2, Xenon has two fluorine atoms and one oxygen atom. The hybridisation of the molecule is sp3d. As per this hybridisation, the molecule is assumed to have a geometry trigonal pyramidal shape structure. But the molecule’s shape becomes T-shaped due to lone pair-bond pair repulsion. The image is given below:
Image name: XeOF2 xenon oxy fluoride
Preparation: Partial hydrolysis of XeF6 gives xenon di oxy difluoride.
XeF3 +2H2O → XeO2F2 + 4HF
Structure: In XeO2F2, the Xe atom forms the central metal atom. The central Xe atom has1 lone pair of electrons and four bonding pairs. Hence, it undergoes sp3d hybridisation. The molecular geometry of XeO2F2 is trigonal bipyramidal, but the actual shape will be see-sawed due to the presence of lone pairs in the equatorial position. The repulsion between the bond pair & lone pair of electrons will be added. Fluorine will be axial atoms, & oxygen will also be tropical atoms. The hybridisation is sp3d2. The geometry is given below:
Image name :Xenon dioxyfluoride(XeO2F2)
Although a noble gas, Xenon forms many compounds with elements of higher electro-negativity, such as Fluorine and oxygen. Although, no other noble gas forms such compounds. The reason is that Xenon is a bigger atomic nuclear size. A weak force of attraction occurs between the outer electron and the protons in the nucleus, due to which the outermost electrons are readily available to form a compound. Students may refer to NCERT and CBSE Solutions in addition to Class 12 Chemistry Chapter 7 Notes for a more detailed explanation.
Further our Class 12 Chemistry Chapter 7 Notes, explains the reason behind the existence of xenon compounds. It also describes the preparation, structure and hydrolysis of some xenon compounds.
Why do Xenon compounds exist?
Out of all rare gases, only Xenon forms several chemical compounds. It is because Xenon is large and has a higher atomic mass. Due to the larger nuclear radius, a weak force of attraction exists between the outer electron and the protons in the nucleus. Hence the outermost electrons are readily available to form a compound.
Uses of Noble gases
Helium obtains its use in filing air balloons and aircraft. This noble gas is because it is flammable and has a very low density. Because it has a shallow boiling point of only 4.2 K, we can use fluid Helium as a cryogenic agent to accomplish tests at significantly lesser temperatures. Liquid Helium is ordinary in the method of cryoscopy that we need for superconductivity. Fluid helium is a general ingredient to cool the superconducting magnets in atomic magnetic resonance spectrometers. We also utilise it as the cooling gas in gas-cooled nuclear reactors and as a streaming gas in gas-fluid chromatography. This noble gas is also essential to our oxygen cylinders in deep sea diving. Argon is mainly utilised to provide an inert atmosphere in high-temperature metallurgical processes (such as arc welding metals or alloys) and to fill electric bulbs. It’s also used in the lab to handle compounds that are sensitive to air. In light bulbs tailored for specific tasks, Xenon and Krypton are employed.
Class 7 Chemistry Chapter 7 – Chapter Summary Notes
Inert pair effect: the inability of electronic valence shell ns2 electrons to take part in bond formation in p block is known as the inert pair effect.
Disproportionation: The same element gets oxidised and reduced in a chemical reaction.
p-pi- p-pi bond: The chemical bond (pi)formed by lateral (sidewise) overlapping of p orbitals of two atoms in a compound or ion.
dpi – dpi bond– the chemical bond (pi) developed by lateral overlapping of p & d orbitals of two atoms in a compound or ions.
Allotropy: Phenomenon of the existence of an element in different physical forms.
Catenation: Tendency of self-linking elements
Chalcogens: Ore forming. Given to group 16.
Halogens molecules: salt producers. They are given to group 17.
Interhalogens compound: compounds created between two different halogens.
Noble gases: The elements in group 18 have their valence shell orbitals wholly filled and, therefore, react with only a few elements under certain conditions. So they are now known as Noble gases.
Anomalous behaviour: The abnormal behaviour of the first member of a particular group from other members of the same group.
Stranger Gas: The term Xenon is derived from the Greek word ” Xenos”, which means foreign or strange. Xenon belongs to a rare gas group where elements are not reactive, but it combines with elements like Fluorine and oxygen to form new compounds. Therefore, it is known as stranger gas.
Uses of Boron: Metal borides are a standard part of nuclear reactors. Here, we are used as defensive shields and control rods. This boron is because of the high capacity of the 10B isotope to ingest neutrons.
Boron is an essential component in the steel industry. It observes its usage in the expansion of the hardness of steel.
Every time, we use boron as a semiconductor for forming electronic gadgets. Boron compounds are turning out to be successively important as the rocket fills. It is primarily because of their high energy/weight proportion. The fibres from boron are common in the process of making light composite materials for airships. Boron is an essential element in the plant digestion system. Boron carbide filaments are rigid, even though they are incredibly lightweight. That is the reason they are primarily used for making bulletproof vests.
Ethanol (C2H5OH): Ethanol goes by the common name ethyl alcohol. We generally know it as alcohol or spirit. It mainly constitutes alcoholic drinks. This organic compound is used in manufacturing medicines, like tincture iodine, cough syrup, etc. Though, one must never consume pure ethanol as it could be fatal.
Ethanoic Acid (CH3COOH): Acetic acid is the common name for ethanoic acid. Acetic acid has a melting point of 290 K. This acidic compound freezes in winter; hence, we term it glacial acetic acid. The vinegar we consume is a 5% to 8% solution of acetic acid in water. It is also used to preserve pickles.
Now, we will look at a fascinating concept. Do you consider what your soaps are made of? Well, organic Chemistry has a vital role to play in this!
Saponification: Concerning Class 12 Chemistry, Chapter 7 notes that Ester of higher fatty acids gives sodium salt of a higher fatty acid. This chemical reaction happens when we heat it with glycerol and sodium hydroxide. Sodium salts of higher fatty acids are shown as soaps. This reaction is known as saponification or the method of soap making.
What are Soaps & Detergents?
Soapy Soaps are esters of higher fatty acids. We manufacture these by reacting esters of a higher fatty acid with sodium hydroxide. The sodium salt, thus produced, has cleansing properties.
Detergent: Soap cannot form a lather in hard water. To overcome this situation, chemists initiate detergents. These detergents are soapless soaps. It is the sodium salt of benzene sulphonic acid or sodium salt of long chain alkyl hydrogen sulphate.
Cleansing Action of Soap: A soap molecule has two ends. One end shows hydrophilic, and another end shows hydrophobic. One end is lipophobic (hydrophilic), and another is lipophilic (hydrophobic).
When we dissolve a soap bar in water and soak the clothes in the soapy solution, soap molecules converge systematically to make a structure. This structure is a micelle. The hydrophobic ends of different molecules neighbouring a particle of grease make the micelle. It has a spherical shape.
The hydrophilic site end is outside the sphere, and the hydrophobic end is towards the sphere’s centre. Soapy soap molecules wash away dirt and grease by forming micelles around them.
So, these are the summary notes of The p block elements Class 12 Chemistry Chapter 7. This paper concludes all the information you need to revise this particular topic. Chapter 7 is also essential as it has more theoretical and fewer numerical problems for students to solve. Students can expect two or five mark questions from it, and these marks can be easily secured if you thoroughly prepare this chapter using the course material prepared by Chemistry subject experts at Extramarks.
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CBSE Solutions Class 12 Chemistry Chapter 7: Exercise & Answer Solutions
Class 12 Chemistry Chapter 7 Notes include exercise and answer solutions that are covered in NCERT textbook and NCERT exemplar books. The exercise and questions cover all important chapter topics such as p block elements, halogen and inert gases structures, preparation and uses, etc. to add more value to the chapter notes.
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- Chapter 7: Exercise 7.1 – Question and Solutions
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- Chapter 7: Exercise 7.5 – Question and Solutions
- Chapter 7: Exercise 7.6 – Question and Solutions
- Chapter 7: Exercise 7.7 – Question and Solutions
- Chapter 7: Exercise 7.8 – Question and Solutions
- Chapter 7: Exercise 7.9 – Question and Solutions
- Chapter 7: Exercise 7.10 – Question and Solutions
- Chapter 7: Exercise 7.11 – Question and Solutions
- Chapter 7: Exercise 7.12 – Question and Solutions
- Chapter 7: Exercise 7.13 – Question and Solutions
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Reference Books Class 12 Chemistry Chapter 7
|i) Modern’s ABC of Chemistry Class by SP Jauhar|
- ii) Concise inorganic Chemistry by J.D.Lee
iii) NCERT textbook Class 12 Chemistry
- iv) Pradeep’s new course of Chemistry by SC Khaterpal
- v) Modern Approach to Chemical Calculation by RC Mukherjee
NCERT Exemplar for Class 12 Chemistry Chapter 7
NCERT Exemplar books are an important element for Mathematics and Science preparation for secondary and higher secondary grades. The exemplar covers questions of different formats with varying degrees of difficulty levels that help students to clear their concepts and prepare well for the exam. It provides a deep learning experience which brings clarity of concepts which in turn becomes useful when it comes to answering the most difficult questions.
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Key Features of Class 12 Chemistry Chapter 7 Notes
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The Chemistry Class 12 Chapter 7 Notes help students understand all the concepts and give a detailed explanation of every synthesis and formula that is covered in this chapter. It allows students to study independently for the examinations without having to look for any further assistance elsewhere. . Students may refer to Class 12 Chemistry Chapter 7 Notes by registering on Extramarks website.
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Yes, Class 12 Chemistry Chapter 7 Notes covers all topics and provides students with full information to understand the chapter quickly. The notes are prepared by Chemistry teachers who have years of teaching experience. They refer to the latest NCERT textbook and other reference materials to prepare detailed notes for students.
4. How many elements exist in the p-Block element?
There are 35 p-block elements, all in p orbital with valence electrons.
5. What is a ‘Halogen’?
Any of the five elements, fluorine, chlorine, bromine, iodine and astatine, form part of group VII A of the modern periodic table and exist in the free state, usually as diatomic molecules.
6. Define 'Oxo-acids' in one word?
An oxo-acid (oxyacid) is an acid that contains oxygen.
7. How vital are Class 12 Chemistry Chapter 7 Notes?
The Class 12 Chapter 7 Chemistry Notes help students with a thorough understanding of all the basic concepts of solutions. It gives a detailed explanation of various p block elements preparation, properties, uses and ideas mentioned in the chapter. So along with the NCERT textbook, Class 12 Chemistry Chapter 7 Notes will help students prepare for this chapter comprehensively. It gives students a joyful learning experience. It guides students and encourages them to be curious and look for answers themselves. It’s good to follow the NCERT book first and then strengthen your knowledge and learning with Extramarks reference materials to create your own milestone in Chemistry.
8. Where can I find notes on Chapter 7, the p block elements of Class 12 Chemistry?
You can find Class 12 Chemistry Chapter 7 Notes prepared by experienced faculty on Extramarks website and their mobile app. These chapter notes are prepared by experienced Chemistry subject experts with simple language, illustrations and -labelled diagrams and verified facts. These notes can be used to understand the topic and also helps students in last-minute revisions ahead of the competition. .
9. Justify the position of O, S, Se, Te and Po in the same periodic table group regarding electronic configuration, oxidation state and hydride formation.
The elements of group 16 are collectively called chalcogens. Elements of group sixteen have six valence electrons each. These elements’ general electronic configuration valence shell is
ns2np4, where n varies from 2 to 6.
Oxidation state: These eighteen elements have six valence electrons (ns2np4), so they should show an oxidation state of −2. Only oxygen predominantly shows the oxidation state of −2 owing to its high electro-negativity. It also presents the oxidation state of −1 (H2O2), zero(O2), and +2(OF2). The stability is −2 oxidation state decreases on moving down a group because of the lower electro-negativity of the p block elements. The heavier elements of the group exhibit an oxidation state of +2, +4, and +6 due to the availability of d-orbitals.
Formation of hydrides:
These elements form hydrides of formula H2E, where E = O, S, Se, Te, PO Oxygen and sulphur also form hydrides of type H2E2. These hydrides are pretty volatile.
10. Why is dioxygen a gas but sulphur a solid?
Oxygen is smaller in size when compared to sulphur. Since its small size, it can form pπ−pπ bonds and create O2 (O=O) molecules. Also, the intermolecular interaction force in oxygen is weak in Van der Wall’s, which causes it to exist as a gas. Conversely, sulphur does not form an M2 molecule but exists as a puckered structure held together by strong covalent bonds. Hence, it is solid.
11. Knowing the electron gain enthalpy values for O → O -1 and O → O 2- as −141 and 702 kJ mol −1 .How can you account for the formation of many oxides having O2−species and not O−? (Hint: Consider lattice energy factor in the construction of compounds).
The more the lattice energy of a compound, the more stable it will be. The stability of an ionic compound depends on its lattice energy.
12. Which aerosols deplete ozone?
The aerosol which is responsible for the depletion of ozone is Freons or chlorofluorocarbons (CFCs). The molecules of CFS break down when there is the presence of ultraviolet radiation and form Chlorine free radicals, which then combine with ozone to create oxygen.
13. Give some practical uses of carbon.
Carbon is mostly obtained from coal deposits and it’s processed before being used for commercial purposes. Some of the most important uses of carbon are:
- We use impure carbon in the form of charcoal (from wood) and coke (from coal) in metal smelting.
- Graphite is used in pencils as a lead. We also use graphite to make brush for electric motors and furnace linings.
- Carbon in the solid form is known as dry ice.Carbon monoxide is also useful for the reduction of a variety of metallurgical processes.
- Activated charcoal finds its usage in purification and filtration in respirators and kitchen extractor hoods.
- Industrial diamonds are a standard tool for cutting rocks and drilling.
14. What is the geometry of xenon tetrafluoride XeF4?
The shape of (xenon tetrafluoride) XeF4 is square planar. The central Xe atom has four bond pairs of electrons and two lone pairs of electrons. It undergoes sp3d2 hybridisation, which concludes in octahedral electron geometry and square planar molecular geometry. The two lone pairs present are at opposite corners of an octahedron.
Question 15. What is XeF4 used for?
Answer 15: XeF4 is a compound that usually occurs as a colourless/white, crystalline solid. This compound comprises Xenon (an inert gas) and fluoride (a naturally occurring mineral). XeF4 can help find and analyse trace metals that contaminate silicone rubber.
15. Is Xenon element a rare metal?
Xenon (Xe) is present as a colourless, odourless gas and is chemically inert. It has the nuclear atomic number 54 in the periodic table and belongs in Group eighteen, the Noble Gases. It is a non-metal with the denoted symbol Xe. Xenon was discovered in 1898 by Scientists William Ramsay and Morris Travers.
16. What are some characteristics of Xenon?
Xenon is one of the inert or noble gases and is odourless, colourless, tasteless and chemically non-reactive. While not toxic, its compounds are strong oxidising agents that are highly toxic
17. Name two poisonous gases prepared from Chlorine.
The two poisonous gases prepared from Chlorine are mustard gas, phosgene, and tear gas.
18. Why is the I-Cl bond more polar reactive than I2?
I-Cl bond is polar and weaker than the I-I bond, so I-Cl is more reactive than I2.
19. The bleaching effect of Chlorine is permanent. Why?
It is due to the process of oxidation of coloured substances to colourless substances by nascent oxygen.
20. Compare the oxidising power of Fluorine and Iodine molecules.
The electrode potential of F2 is maximum, but that of I2 is minimum. F2 can be reduced very quickly, but I2 is reduced less readily. Thus, F2 is the most vital oxidising agent, while I2 is the weakest oxidising agent.
21. Does the hydrolysis of Xenon hexafluoride lead to a redox reaction?
No, the products of hydrolysis are XeOF4 & XeO2F2, where the oxidation states of all the elements remain the same as it was in the reacting form.