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Tamil Nadu 11th Chemistry Book

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Text book published by Government of Tamil Nadu. This book has been prepared by the Directorate of School Education on behalf of the Government of Tamilnadu. This book has been printed on 60 G.S.M paper. 11th Standard Diamond Chemistry TM / வேதியியல்: bestthing.info: Mrs. Chitra Malayaman, TamilNadu State Board Syllabus: Books.

The results can be checked by the students on the official sites at tnresults. It appoints through exams all Indian Services and group A and B of central services. Tamil Nadu 11th Result announced today at tnresults. Students can check Tamil Nadu Plus one Results name wise marks list and without date of birth school wise results can be download from below. We will also update on this page after the official declaration 12th Result Date Tamilnadu: Tamil Nadu School Education Board has published the exam time table for HSC students on the internet at dge.

Also, you can see below for Tamilnadu 10th Time Table The students who are studying on this board must appear in the exam. All students this time, they are eagerly waiting their TN Board 11th Class Exam Date Sheet very long time, now waiting time is over because examination authority are planning to announce the 11th Class Public Exam Time Table on the official web portal in upcoming month.

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Tamil Nadu Board of School Education conducts the public exam for 10 th and 12 th students, every year. Candidates who are in 12th standard they can check the timetable which is given below, there are timetable in two patterns this year are advised to go through the timetable of the previous year given below till the date sheet for is announced. Tamil Nadu board 12th Exam Admit card is expected to be released in February, Read all news including political news, current affairs and news headlines online on 11th Public Exam Result Date today.

After releasing the Tamilnadu 11th Results we will update on this page. Read on to find out. Bangladesh has only on Madrasah board. Read to know the SSC exam result details. In this table, we have also provided exam date, exam day, the name of the exam very clearly, students can start your preparation with the help of this Kerala Plus One Exam Time Table Students can check results through SMS alert as well.

Tamilnadu 11th Result Announced. The results for plus two or class 12th results would be released on official websites tnresults. Have a wonderful and awesome result which gives full satisfaction. The number of vacancies in different services and posts to be filled up on results of examination will be announced in due time. Official result publish date and website info.

Tamil Nadu 12th class exams will start on 1st March, and ends on 6th April, In this article, we will provide you with all the information regarding Tamil Nadu 11th Result An international delegation of 11 members from Australia, South Korea, Bosnia, Bhutan and Bangladesh visited India to analyse the nitty-gritty of the election management system.

Many candidates also appeared the Class 11 Exam and this time state board going to release the 11th Result Tamilnadu. Once the result of the same will declare, the candidates are requested to check the results from the official web portal of Karnataka Public service Commission i.

It will be really helpful to you to prepare and get high marks in your exams. The result of 11th class will soon going to be uploaded here. Bangladesh educational sectaries Sohrab Hossain announce the result to date. Here is the full details of 11th public exam result date, Plus 1 public result fastest server, 11th public result date with mark update android mobile, 11th Result TN — 11th Public Exam Result Date The 11th examination question papers and answer key is available after the completion of the examination and the answer key will be an unofficial key.

SSC exam was held on 2nd February Monday to 10th March Tuesday for written exam and practical exam held on 11th February for general board. Through the branch of science we call chemistry we have gained an understanding of the matter which makes up our world and of the interactions between particles on which it depends. The ancient Greek philosophers had their own ideas of the nature of matter, proposing atoms as the smallest indivisible particles. However, although these ideas seems to fit with modern models of matter, so many other Ancient Greek ideas were wrong that chemistry cannot truly be said to have started there.

Alchemy was a mixture of scientific investigation and mystical quest, with strands of philosophy from Greece, China, Egypt and Arabia mixed in. The main aims of alchemy that emerged with time were the quest for the elixir of life the drinking of which would endue the alchemist with immortality , and the search for the philosopher's stone, which would turn base metals into gold.

Improbable as these ideas might seem today, the alchemists continued their quests for around years and achieved some remarkable successes, even if the elixir of life and the philosopher's stone never appeared.

Towards the end of the eighteenth century, pioneering work by Antoine and Marie Lavoisier and by John Dalton on the chemistry of air and the atomic nature of matter paved the way for modern chemistry. During the nineteenth century chemists worked steadily towards an understanding of the relationships between the different chemical elements and the way they react together. A great body of work was built up from careful observation and experimentation until the relationship which we now represent as the periodic table emerged.

This brought order to the chemical world, and from then on chemists have never looked back. Modem society looks to chemists to produce, amongst many things, healing drugs, pesticides and fertilisers to ensure better crops and chemicals for the many synthetic materials produced in the twenty-first century. It also looks for an academic understanding of how matter works and how the environment might be protected from the source of pollutants.

Fortunately, chemistry holds many of the answers! Hi Following the progressing trend in chemistry, it enters into other branches of chemistry and answers for all those miracles that are found in all living organisms. The questions that are given in each and every chapter can be taken only as model questions.

A lot of self evaluation questions, like, choose the best answer, fill up the blanks and very short answer type questions are given in all chapters. They must be prepared to answer the questions and problems from the entire text. Learning objectives may create an awareness to understand each and every chapter.

Sufficient reference books are suggested so as to enable the students to acquire more informations about the concepts of chemistry. Chemical Bonding 1 Colligative Properties 36 Thermodynamics - 1 64 Chemical Equihbrium - 1 88 Chemical Kinetics - 1 Organic Chemistry Basic Concepts of Organic Chemistry Purification of Organic compounds Detection and Estimation of Elements Hydrocarbons Aromatic Hydrocarbons Organic Halogen Compounds v Syllabus: Unit 2 - General Introduction to Metallurgy Ores and minerals - Sources from earth, living system and in sea - Purification of ores-Oxide ores sulphide ores magnetic and non magnetic ores - Metallurgical process - Roasting-oxidation - Smelting-reduction - Bessemerisation - Purification of metals-electrolytic and vapour phase refining - Mineral wealth of India.

Unit 3 - Atomic Structure - 1 Brief introduction of history of structure of atom - Defects of Rutherford's model and Niels Bohr's model of an atom - Sommerfeld's extension of atomic structure - Electronic configuration and quantum numbers - Orbitals-shapes of s, v p and d orbitals. Unit 4 - Periodic Classification - 1 Brief history of periodic classification - IUPAC periodic table and IUPAC nomenclature of elements with atomic number greater than - Electronic configuration and periodic table - Periodicity of properties Anomalous periodic properties of elements.

Unit 5 - Group-Is Block elements Isotopes of hydrogen - Nature and application - Ortho and para hydrogen - Heavy water - Hydrogen peroxide - Liquid hydrogen as a fuel - Alkali metals - General characteristics - Chemical properties - Basic nature of oxides and hydroxides - Extraction of lithium and sodium - Properties and uses.

Unit 6 - Group - 2s - Block elements General characteristics - Magnesium - Compounds of alkaline earth metals. Unit 7 -p- Block elements General characteristics of p-block elements - Group- Unit 9 - Gaseous State Four important measurable properties of gases - Gas laws and ideal gas equation - Calculation of gas constant "R" - Dalton's law of partial pressure - Graham's law of diffusion - Causes for deviation of real gases from ideal behaviour - Vanderwaal's equation of state - Critical phenomena - Joule-Thomson effect and inversion temperature - Liquefaction of gases - Methods of Liquefaction of gases.

Unit 10 - Chemical Bonding Elementary theories on chemical bonding - Kossel-Lewis approach - Octet rule - Types of bonds - Ionic bond - Lattice energy and calculation of lattice energy using Born-Haber cycle - Properties of electrovalent compounds - Covalent bond - Lewis structure of Covalent bond - Properties of covalent compounds - Fajan's rules - Polarity of Covalent bonds - VSEPR Model - Covalent bond through valence bond approach - Concept of resonance - Coordinate covalent bond.

Unit 11 - Colligative Properties Concept of colligative properties and its scope - Lowering of vapour pressure - Raoul's law - Ostwald - Walker method - Depression of freezing point of dilute solution - Beckmann method - Elevation of boiling point of dilute solution - Cotrell's method - Osmotic pressure - Laws of Osmotic pressure - Berkley-Hartley's method - Abnormal colligative properties Van' t Hoff factor and degree of dissociation.

Unit 12 - Thermodynamics - 1 Thermodynamics - Scope - Terminology used in thermodynamics - Thermodynamic properties - nature - Zeroth law of thermodynamics - Internal energy - Enthalpy - Relation between "H and "E - Mathematical form of First law - Enthalpy of transition - Enthalpy of formation - Enthalpy of combustion - vii Enthalpy of neutralisation - Various sources of energy-Non-conventional energy resources.

Unit 13 - Chemical Equilibrium - 1 Scope of chemical equilibrium - Reversible and irreversible reactions - Nature of chemical equilibrium - Equilibrium in physical process - Equilibrium in chemical process - Law of chemical equilibrium and equilibrium constant - Homogeneous equilibria - Heterogeneous equilibria. Unit 14 - Chemical Kinetics - 1 Scope - Rate of chemical reactions - Rate law and rate determining step - Calculation of reaction rate from the rate law - Order and molecularity of the reactions - Calculation of exponents of a rate law - Classification of rates based on order of the reactions.

Unit 16 - Purification of Organic compounds Characteristics of organic compounds - Crystallisation - Fractional Crystallisation - Sublimation - Distillation - Fractional distillation - Steam distillation - Chromotography.

Unit 17 - Detection and Estimation of Elements Detection of carbon and hydrogen - Detection of Nitrogen - Detection of halogens - Detection of sulphur - Estimation of carbon and hydrogen - Estimation of Nitrogen - Estimation of sulphur - Estimation of halogens.

Unit 19 - Aromatic Hydrocarbons Aromatic Hydrocarbons - IUPAC nomenclature of aromatic hydrocarbons - Structure of Benzene - Orientation of substituents on the benzene ring - Commercial preparation of benzene - General methods of preparation of Benzene and its homologues - Physical properties - Chemical properties - Uses - Carcinogenic and toxic nature.

Unit 20 - Organic Halogen Compounds Classification of organic hydrogen compounds - IUPAC nomenclature of alkyl halides - General methods of preparation - Properties - Nucleophilic substitution reactions - Elimination reactions - Uses - Aryl halide - General methods of preparation - Properties - Uses - Aralkyl halides - Comparison arylhalides and aralkyl halides - Grignard reagents - Preparation - Synthetic uses.

Knowledge of using Burette, Pipette and use of logarithms is to be demonstrated. Preparation of Compounds. Copper Sulphate Crystals from amorphous copper sulphate solutions 2.

Preparation of Mohr's Salt 3. Preparation of Aspirin 4. Preparation of Iodoform 5. Identification of one cation and one anion from the following. Insoluble salt should not be given Cation: Determination of Melting point of a low melting solid. Acidimetry Vs Alkalimetry 1. Preparation of Standard solution of Oxalic acid and Sodium Carbonate solution.

Chemical bond Existence of a strong force of binding between two or many atoms is referred to as a Chemical Bond and it results in the formation of a stable compound with properties of its own. The bonding is permanent until it is acted upon by external factors like chemicals, temperature, energy etc. It is known that, a molecule is made up of two or many atoms having its own characteristic properties which depend on the types of bonding present.

Molecules containing identical but many atoms bonded together such as P4, Ss etc. Chemical bonds are basically classified into three types consisting of i ionic or electrovalent bond ii covalent bond and iii coordinate - covalent bond.

Mostly, valence electrons in the outer energy level of an atom take part in the chemical bonding. In , W. Kossel and G. Lewis, separately developed theories of chemical bonding inorder to understand why atoms combined to form molecules. According to the electronic theory of valence, a chemical bond is said to be formed when atoms interact by losing, gaining or sharing of valence electrons and in doing so, a stable noble gas electronic configuration is achieved by the atoms.

Except Helium, each noble gas has a stable valence shell of eight electrons. The tendency for atoms to have eight electrons in their outershell by interacting with other atoms through electron sharing or electron-transfer is known as the octet rule of chemical bonding.

Kossel laid down the following postulates to the understanding of ionic bonding: Therefore one or small number of electrons are easily gained and transferred to attain the stable noble gas configuration. The noble gases with the exception of helium which has two electrons in the outermost shell have filled outer shell electronic configuration of eight electrons octet of electrons with a general representation ns np.

Kossel's postulates provide the basis for the modern concepts on electron transfer between atoms which results in ionic or electrovalent bonding. For example, formation of NaCl molecule from sodium and chlorine atoms can be considered to take place according to Kossel's theory by an electron transfer as: The bonding in NaCl is termed as electrovalent or ionic bonding. Sodium atom loses an electron to attain Neon configuration and also attains a positive charge.

Chlorine atom receives the electron to attain the Argon configuration and also becomes a negatively charged ion. Similarly formation of MgO may be shown to occur by the transfer of two electrons as: Thus, "the binding forces existing as a result of electrostatic attraction between the positive and negative ions", is termed as electrovalent or ionic bond.

The electrovalency is considered as equal to the number of charges on an ion. Thus magnesium has positive electrovalency of two while chlorine has negative electrovalency of one. The valence electron transfer theory could not explain the bonding in molecules like H2, O2, CI2 etc.

Lewis, proposed the octet rule to explain the valence electron sharing between atoms that resulted in a bonding type with the atoms attaining noble gas electronic configuration. The statement is: This type of valence electron sharing between atoms is termed as covalent bonding. Generally homonuclear diatomics possess covalent bonds. It is assumed that the atom consists of a "Kernel' which is made up of a nucleus plus the inner shell electrons.

The Kernel is enveloped by the outer shells that could accommodate a maximum of eight electrons. The eight outershell electrons are termed as octet of electrons and represents a stable electronic configuration.

Atoms achieve the stable outer octet when they are involved in chemical bonding. In case of molecules like F2, Ch, H2 etc. For example, consider the formation of a fluorine molecule F2. The atom has electronic 2 2 5 configuration. In the fluorine molecule, each atom contributes one electron to the shared pair of the bond of the F2 molecule.

In this process, both the fluorine atoms attain the outershell octet of a noble gas Argon Fig. Such structures are called as Lewis dot structures. Lewis dot structures can be written for combining of like or different atoms following the conditions mentioned below: If the two atoms share a pair of electrons, a single bond is said to be formed and if two pairs of electrons are shared a double bond is said to be formed etc.

All the bonds formed from sharing of electrons are called as covalent bonds. The carbon and the two oxygen atoms attain the Neon electronic configuration. Each of the Nitrogen atom shares 3 pairs of electrons to attain neon gas electronic configuration. The different types of chemical bonding that are considered to exist in molecules are i ionic or electrovalent bond which is formed as a result of complete electron transfer from one atom to the other that constitutes the bond; ii covalent bond which is formed as a result of mutual electron pair sharing with an electron being contributed by each atom of the bond and iii coordinate - covalent bond which is formed as a result of electron pair sharing with the pair of electrons being donated by only one atom of the bond.

The formation and properties of these types of bonds are discussed in detail in the following sections. The compounds containing such a bond are referred to as ionic or electrovalent compounds. Ionic bond is non directional and extends in all directions. Therefore, in solid state single ionic molecules do not exist as such. Only a network of cations and anions which are tightly held together by electro-static forces exist in the ionic solids.

To form a stable ionic compound there must be a net lowering of energy. That is, energy is released as a result of electovalent bond formation between positive and negative ions. When the electronegativity difference between the interacting atoms are greatly different they will form an ionic bond.

In fact, a difference of 2 or more is necessary for the formation of an ionic bond. Na has electronegativity 0. Thus, NaCl ionic compound is formed. In NaCl, both the atoms possess unit charges. Thus, doubly charged positive and negative ions are formed. The molecule as a whole remains electrically neutral.

For example in MgF2, Mg has two positive charges and each fluorine atom has a single negative charge. The three dimensional network of points that represents the basic repetitive arrangement of atoms in a crystal is known as lattice or a space lattice. Thus a qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation. Lattice enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

That is, the enthalpy change of dissociation of MX ionic solid into its respective ions at infinity separation is taken the lattice enthalpy. E Lattice enthalpy is a positive value. For example, the lattice enthaply of NaCl is kJ. In ionic solids, the sum of the electron gain enthaply and the ionisation enthalpy may be positive but due to the high energy released in the formation of crystal lattice, the crystal structure gets stabilised.

Born Haber's Cycle Determination of Lattice enthalpy It is not possible to calculate the lattice enthalpy directly from the forces of attraction and repulsion between ions but factors associated with crystal geometry must also be included.

The solid crystal is a three-dimensional entity. The lattice enthalpy is indirectly determined by the use of Born - Haber Cycle. The procedure is based on Hess's law, which states that the enthalpy change of a reaction is the same at constant volume and pressure whether it takes place in a single or multiple steps long as the initial reactants and the final products remain the same.

Also it is assumed that the formation of an ionic compound may occur either by direct combination of elements or by a step wise process involving vaporisation of elements, conversion of gaseous atoms into ions and the combination of the gaseous ions to form the ionic solid.

For example consider the formation of a simple ionic solid such as an alkali metal halide MX, the following steps are considered.

The greater the lattice enthalpy the more stabler the ionic bond formed. The lattice enthalpy is greater for ions of higher charge and smaller radii. The lattice enthalpies affect the solubilities of ionic compounds. Also, the formation of NaCl can be considered in 5 steps. The sum of the enthalpy changes of these steps is considered equal to the enthalpy change for the overall reaction from which the lattice enthalpy of NaCl is calculated.

Problem 1 Calculation of lattice enthalpy of MgBr2 from the given data. The nature of these properties are discussed as follows. Due to strong coulombic forces of attraction between the oppositely charged ions, electrovalent compounds exist mostly as hard 12 crystalline solids. Due to the hardness and high lattice enthalpy, low volatility, high melting and boiling points are seen. Because of the strong electrostatic forces, the ions in the solid are not free to move and act as poor conductor of electricity in the solid state.

However, in the molten state, or in solution, due to the mobility of the ions electrovalent compounds become good conductor of electricity. Ionic compounds possess characteristic lattice enthalpies since they exist only as ions packed in a definite three dimensional manner. They do not exist as single neutral molecule or ion. Ionic compounds are considered as polar and are therefore, soluble in high dielectric constant solvents like water.

In solution, due to solvation of ions by the solvent molecules, the strong interionic attractions are weakened and exist as separated ions. Electrovalent compounds having the same electronic configuration exhibit isomorphism. By doing so, the atoms attain stable octet electronic configuration. In covalent bonding, overlapping of the atomic orbitals having an electron from each of the two atoms of the bond takes place resulting in equal sharing of the pair of electrons.

Also the interatomic bond thus formed due to the overlap of atomic orbitals of electrons is known as a covalent bond. Generally the orbitals of the electrons in the valency shell of the atoms are used for electron sharing. The shared pair of electrons lie in the middle of the covalent bond. Including the shared pair of electrons the atoms of the covalent bond attain the stable octet configuration. Thus in hydrogen molecule H2 a covalent bond results by the overlap of the two s orbitals each containing an electron from each of the two H atoms of the molecule.

Each H atom attains 'Is filled K shell. Consider Ck molecule. The outer shell electronic configuration of atom is 3 s 2p x 2p y 2p z. When each chlorine atom mutually share the 2p z unpaired electron contributed from each CI atom of the molecule, a covalent bond is formed.

By doing so, each chlorine atom attains argon electron configuration. The Lewis dot structure will be: More than one e" can also be mutually shared to result in two covalent bonds between the atoms of a bond. The two unpaired electrons in 2p y and 2p z orbitals of each of O atom is mutually shared so that after the double bond formation stable octet electronic configuration is attained by each oxygen atom of the molecule. In the phosphine PH3 molecule, three hydrogen atoms combine with one phosphorous atom.

Each hydrogen atom shares its Is electron with phosphorous. So that three covalent bonds are formed in PH3. The lewis dot structure is in Fig. In case of ethane molecule, the six C-H bonds and a C-C bond are covalent in nature. They are formed by mutual sharing of a pair of electrons between the two atoms of a bond.

Each carbon atom completes its stable octet and each H atom has completed K shell. The electronic configuration of O atom is Is 2s 2p. By sharing two more electrons from the other O atom, each O atom attains 2s 2p , filled configuration. Similar to oxygen molecule in ethylene which is an organic molecule, a double covalent bond exists between the two carbon atoms due to the mutual sharing of two pairs of electrons.

Each carbon atom attains the stable octet electron configuration. Covalent compounds are formed by the mutual sharing of electrons. There is no transfer of electrons from one atom to another and therefore no charges are created on the atom. No ions are formed. These compounds exist as neutral molecules and not as ions. Although some of the covalent molecules exist as solids, they do not conduct electricity 15 in fused or molten or dissolved state.

They possess low melting and boiling points. This is because of the weak intermolecular forces existing between the covalent molecules. Since, no strong coulombic forces are seen, some of covalent molecules are volatile in nature.

Mostly covalent compounds possess low melting and boiling points. Covalent bonds are rigid and directional therefore different shapes of covalent molecules are seen. Most of the covalent molecules are non polar and are soluble in nonpolar low dielectric constant solvents like benzene, ether etc and insoluble in polar solvents like water. Carbon tetrachloride CCU is a covalent nonpolar molecule and is soluble in benzene. Dashed lines represent hypothetical unpolarized ions When cations and anions approach each other, the valence shell of anions are pulled towards cation nucleus due to the coulombic attraction and thus shape of the anion is deformed.

This phenomenon of deformation of anion by a cation is known as polarization and the ability of cation to polarize a nearby anion is called as polarizing power of cation. Fajan points out that greater is the polarization of anion in a molecule, more is covalent character in it.

This is Fajan's rule.

Fajan also pointed out the influence of various factors on cations for polarization of anion. The general trend in the polarizing power of cations: Therefore more effectively the cation pulls the valence electrons towards its nucleus. This results in more polarization effect. That is, for the same charge of the anion, larger sized anion is more polarized than a smaller sized anion. The trend in the polarization of anions: Thus more will be covalent nature in the bonding of the molecule.

Thus polarizing power: Covalent character: This prevents polarization of anion by the cation. Therefore AlCb behaves as an ionic molecule in water, while it is a covalent molecule in the free state. In the covalently bonded molecules like H2, Ch, F2 homonuclear diatomics , the bond is a pure covalent bond.

In case of heteronuclear molecules like, HF, HC1, CO, NO etc, the shared electron pair gets displaced more towards the atom possessing higher electronegativity value than the other one.

In HF, the shared electron pair is displaced more towards fluorine because the electronegativity of Fluorine is far greater than that of Hydrogen. This results in partial ionic character induced in the covalent bond and is represented as: Thus the extent of ionic character in a covalent bond will depend on the relative attraction of electrons of the bonded atoms which depends on the electro negativity differences between the two atoms constituting the covalent bond.

This causes a dipolar molecule formation. Some dipolar molecules are ". In a triatomic molecule like water two covalent bonds exist between the oxygen atom and the two H atoms. Oxygen with higher electronegativity attracts the shared pair of electrons to itself and thus oxygen becomes the negative end of the dipole while the two hydrogen atoms form the positive end. Thus the two covalent bonds in the water molecule possess partial ionic character.

Generally larger the electronegativity difference between the atoms consisting the bond, greater will be the ionic character. For H atom electronegativity is 2. Thus H-Cl covalent bond is polarised and it has more ionic character.

Many of the physical and chemical properties of molecules arise due to different shapes of the molecules. Some of the common geometrical shapes found among the molecules are: The theory was originally proposed by Sigdwick and Powell in It was further developed and modified by Nyholm and Gillespie It is convenient to divide molecule into two categories i molecules in which the central atom has no lone pairs of electrons and i molecules in which the central atom has one or more lone pairs.

Table In compounds of AB2, AB3, AB4, AB5, AB6, types the arrangement of electron pairs bonded pairs as well as the B atoms around the central atom A are, linear, trigonal planar, tetrahedral, trigonal- bipyramidal and octahedral respectively. In these type of molecules, both lone 21 pairs and bond pairs of electrons are present. The lone pairs are localised on the central atom, and bonded pairs are shared between two atoms. Consequently, the lone pair electrons in a molecule occupy more space as compared to the bonding pair electrons.

This causes greater repulsions between lone pairs of electrons as compared to the lone pairs of electrons to the lone pair lp - bonding pair and bonding pair - bonding pair repulsions bp. In sulphur dioxide molecule there are three electron pairs on the S atom. The overall arrangement is trigonal planar.

However, because one of the three electron pairs is a lone pair, the SO2 molecule has a "bent' shape and due to the lp - lp repulsive interactions the bond angle is reduced 22 to 1 In the ammonia NH3 molecule, there are three bonding pairs and one lone pair of electrons.

The overall arrangement of four electron pairs is tetrahedral. In NH3, one of the electron pairs, on nitrogen atom is a lone pair, so the geometry of NH3 is pyramidal with the N atom at the apex of the pyramid.

The water H2O molecule, oxygen atom contains two bonding pairs and two lone pairs of electrons. The overall arrangement for four electron pairs is tetrahedral, but the lp - lp repulsions being greater than lp-bp repulsions in H 2 0. The HOH angle is reduced to The molecule has a bent shape. The geometrical arrangement will be a regular octahedral.

In a o bond, maximum extent overlap of orbitals are possible and the bond formed is also stronger. For e. H-H bond is a o bond. Consider the valence bond description of O2 molecule: It is conventional to take z axis as the inter nuclear axis or molecular axis.

Along the molecular axis, overlap of 2P Z orbital of two O atoms occur with cylindrical symmetry thus forming a o bond. The remaining two 2p y orbitals of two O atoms cannot overlap to the full extent like a a bond as they do not have cylindrical symmetry around the internuclear axis. Instead, 2p y , orbitals overlap laterally sideways above and below the axis and share the pair of electrons.

The bond formed by lateral overlap of p orbitals above and below the axis together is called a 7i Pi bond. Since 2p y orbitals are perpendicular to 2p z orbitals, 71 bond formed is perpendicular to the a bond. Thus bonding in oxygen molecule is represented as in fig. There are two bonds in O2 molecule.

One of which is a o bond and another is n bond. Similarly, in N2 molecule, 3 bonds are present between 2N atoms. The nature of orbital overlaps in the 3 bonds can be considered as in fig. Thus bonding in oxygen molecule is represented as in Fig. The nature of orbital overlaps in the 3 bonds can be considered as in Fig. Based on the valence bond orbital overlap theory, the H2O molecule is viewed to be formed by the overlap Is orbital of a H atom with 2p y orbital of O atom containing one electron each forming a o bond.

Another a bond is also formed by the overlap of Is orbital of another H atom with 2p x orbital of O atom each containing an unpaired electron. Therefore based on VB theory, pure orbital overlaps does not explain the geometry in H2O molecule. Since 2p x , 2p y and 2p z orbitals of N are mutually perpendicular.

There are three major processes that are considered to occur in hybridisation of orbitals. These are: In case of Boron VB theory predicts univalency due to the presence of one unpaired electron but in practice Boron is trivalent since compounds as BCb, BH3 etc. The stable state Ground State electronic configuration of C is 2 11 2s 2p x 2p y. Electronic configuration of C suggests only bivalency.

But carbon forms over a million compounds in all of which carbon is tetravalent. This suggests only tetravalency. This deficiency is overcome by allowing for promotion or the excitation of an electron to an orbital of higher energy.

Although for electron promotion energy is needed, if that energy is recovered back during a covalent bond formation, or by a bond with a greater strength or by many number of bonds formation, then the electron promotion becomes energetically allowed and assumed to take place initially. In carbon, promotion of an electron to an orbital which is close to itself with an empty orbital of only slightly higher energy which is the 2p z orbital can take place.

Then the electron pair is unpaired itself by absorbing the required energy available by the atom from its surrounding and one of the electrons in the original orbital 2s or 2p shifts to the empty higher energy orbital. Each electron can now be utilised to form a covalent bond by sharing an electron coming from the combining atom.

Thus four a covalent bonds are possible, each with equivalent strength and overlapping tendency. Further, chemical and physical evidences reveal the four bonds of carbon to be equivalent and that they are tetrahedrally oriented. The promotion of an electron from 2s to 2p orbital leads to four half filled orbitals which can form four bonds leading to greater energy lowering.

This energy is more than the initial energy required for the promotion of 2s electron to 2p orbital. Hybridisation mixing of orbitals After an electron promotion the 4 electrons are not equivalent, since one of them involves with an s orbital while the other three involve with p orbitals. To explain the equivalence of the four bonds, the concept of hybridisation is introduced.

Dissimilar orbitals like s,p,d with one or many numbers, with nearly the same energy on the same atom may combine or mix completely to form an equal number of equivalent energy new orbitals with properties of their own. This is called as hybridisation of orbitals. The new orbitals formed are known as hybrid orbitals and these orbitals possess the properties of the pure orbitals that are mixed to form them. The hybrid orbitals of an atom are symmetrically distributed around it in space. Essentially, mixing up of orbitals to form new orbitals explains the different geometries of many compounds like CH4, SF6 etc.

Each such structure is called as canonical structure. A resonance hybrid consists of many canonical structures. All the canonical structures are equally possible to represent the structure of the molecule. For example, in ozone O3 molecule, the two canonical structures as shown below and their hybrid represents the structure of O3 more 29 accurately. Resonance is represented by a double headed arrow placed between the canonical structures.

There are two canonical forms of O3. The resonance structures are possible for molecular ions also. For example, consider resonance in CO3 ion: According to experimental findings all carbon to oxygen bonds in CO3 are equivalent.

Therefore the carbonate ion is best described as a resonance hybrid of the canonical forms as shown in Fig. Therefore, a single lewis structure cannot depict the structure of CO2 as a whole and it is best described as a resonance hybrid of the canonical forms given in Fig.

In N2O molecule which is a linear molecule, structures with charges on atoms can be written similar to CO2. Therefore N2O exists as a hybrid structure of the two canonical forms with a linear geometry.

In each shared pair of electrons one electron is contributed from each atom of the bond. However in some bond formation, the whole of the shared pair of electrons comes from only one of the combining atoms of the bond, which is to referred as the donor atom. The other atom which does not contribute the electron to the shared pair but tries to pull the pair of electron towards itself is called as the acceptor atom.

The bond thus formed is between the donor and acceptor atoms is called as the co-ordinate or co-ordinate - covalent or dative bond. A coordinate bond is showed as an arrow which points from the donor to the acceptor atom. In some cases, the donated pair of electron comes from a molecule as a whole which is already formed to an already formed acceptor molecule as a whole. For Example, coordination bond between H3N: The molecule, ammonia donor which gives a pair of electron lone pair to BF3 molecule which is electron deficient acceptor which has an empty orbital to accommodate the pair of electrons.

When Proton is added to ammonia, a pair of electron is donated by nitrogen to proton and then proton shares the electron pair to form coordinate covalent bond. Few examples of covalent - coordinate bond: The two chlorine atoms act as bridge to link the two Aluminium atoms.

In some complex ion formations, if the central transition metal-ion has empty v d' orbitals then lone pair of electrons from neutral molecules or anions are donated resulting in the formation of coordination bonds. In Nickel tetracarbonyl, the four bonds between central Ni atom and the carbonyl ligands are mainly covalent -coordinate type. This complex exists in square planar geometry.

Questions A. Choose the correct answer 1. The crystal lattice of electrovalent compounds is composed of a Atoms b Molecules c Oppositely charged ions d Both molecules and ions 2. Fill in the blanks 3. Linear overlap of two atomic p-orbitals leads to. Born-Haber cycle is related with. Two atoms of similar electronegativity are expected to form compounds.

Repulsion between bond pair-bond pair is than in between lonepair- lone pair. Match the following 1. Electrovalent bonding a. Benzene 2. Covalent bonding b. Heitler and London 3. Valence Bond theory c. Electron transfer 4. Polarised Bond d.

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Electron sharing 5. Resonance e. Fajan's theory f. Aluminium chloride 33 D. Write in one or two sentence 1. Find o and 71 bonds in the following: What is the structure of BeCl 2? Write the differences between electrovalent and covalent bonds. Give reason: Why are the bond angles different in three cases? What is octet rule? Explain with an example. What are the different types of bonds? What is meant by electrovalent bond.

Explain the bond formation in AlBr 3 and CaO. Give the electron dot representation for PH3 and ethane.

Write the Lewis dot structures for the following. What are the important features of valence bond theory? What is meant by hybridisation? Define resonance. Give the various resonance structures of CO2 and CO3 2 " ion.

Explain briefly on the following 1. Discuss the important properties of electrovalent compounds. Calculate the lattice energy of NaCl using Born-Haber cycle. Explain the important properties of covalent compounds. Discuss the partial covalent character in ionic compounds using Fajan's rule. Explain the polarity of covalent bonds in H2O and HC1.

Discuss the shapes of following molecules: Discuss VSEPR model applied for linear, trigonal planar, tetrahedral and octahedral geometries of molecules. Explain the formation and difference between a sigma bond and a pi- bond. Which has more bond strength? Calculate the lattice enthalpy of CaCb given that the enthalpy of: Except helium, atoms share or transfer valence electrons to attain the stable octet shell as the electronic configuration.

Electrostatic force of attraction between ions describe the ionic bonding.

Mutual sharing of electrons between the two atoms result in covalent bonding. The directional character, partial ionic character by the pure orbital overlaps are also studied with suitable examples. Tetrahedral; BC1 3: Resonance in benzene, carbonate ion, molecules are understood. Ab Ch is covalent but in water, it is ionic. Coordinate-covalent bonding in Ni CO 4 is also understood. It is said to be "binary' if two substances are present and "ternary' if three substances are present and "quaternary' if four substances are being present etc.

In a binary solution, the component present in larger amount is called as solvent and the component in smaller amounts is called as solute. Solvent and solute together make a solution. In dilute solutions, very small amount of the solute is present. A colligative property of a solution depends purely on the number of particles dissolved in it, rather than on the chemical nature of the particles.

The colligative properties can be regarded as the properties of the solvent in a given solution. Generally, the solute is considered as non-volatile. The various colligative properties are as below: Lowering of vapour pressure of the solvent Ap 36 ii. Elevation of boiling point of the solvent ATb iii.

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Depression of freezing point of the solvent ATf iv. Osmotic pressure The important scope of the measurement of colligative properties lies on its use to determine the molar mass of the non-volatile solute dissolved in the dilute solution. The vapour exerts a pressure on the walls of the container and exists in equilibrium with the liquid.

This pressure is referred as the vapour pressure of the liquid. When a non-volatile solute is dissolved in the solvent so that a dilute and homogeneous solution results, then again the vapour pressure of the solution will be made up of entirely from the solvent since the solute does not evaporate. This vapour pressure of the dilute solution is found to be lower than the vapour pressure of the pure solvent.

From Fig. Consequently the vapour pressure of the solvent molecules gets lowered on the surface of the solution. According to Raoult's law, at constant temperature the vapour pressure of the solution p is directly proportional to the molefraction of 37 the solvent Xi present in the solution.

The value of k is known as follows: Mi and M2 are the molar masses of solvent and solute respectively. The relative lowering of the vapour pressure, is defined as the ratio of the lowering of vapour pressure to the vapour pressure of the pure solvent.

Thus, the statement of Raoult's law for dilute solutions containing non-volatile 38 non-electrolyte solute is: Relative lowering of vapour pressure is equal to the mole fraction of the solute.

Thus, relative lowering of vapour pressure is a colligative property. Ap n2 W2. Knowing Mi,Wi and W2 and from the measurement of lowering of vapour pressure, M2 the molar mass of the solute can be determined using equation 1 1. What will be the vapour pressure of the solution. A weighed amount of anhydrous and dry calcium chloride is taken in the U-tube c connected at the end. The chambers and the U-tube are connected by a series of delivery tubes d through which air is passed.

The dry air is first allowed to pass through the solution chamber until the air is saturated with the solvent vapour to maintain the vapour pressure of the solution v p'. Consequently, a loss in weight of the solution results in the solution chamber since some amount solvent molecules have evaporated. Problem 2 Dry air was passed successively through a solution of 5 gm of solute dissolved in The loss in weight of the solution was 2. What is the molecular weight of the solute?

The molefraction of the solute is 0. The vapour pressure of the pure solvent is 0. According to Raoult's law, addition of a non- volatile solute to solvent lowers the vapour pressure of the solvent and hence, the vapour pressure of pure solvent is greater than the vapour pressure of solution.

Thus the temperature at which the solution and its solid form existing in equilibrium and possessing the equal vapour pressures, is lowered. That is, the freezing point of solution is lowered. The lowering of the freezing point of the solution from that of the freezing point of the pure solvent is known as depression in freezing point of the solution.

Generally when 42 the temperature of a solid substance that is used as the solvent is raised, the vapour pressure also raises.

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AB curve depicts this. Similarly curve BC represents the increase in vapour pressure of the liquid solvent with increase in temperature. Since the vapour pressure of the solution is always lower than that of its pure solvent, the vapour pressure curve of the solution DE always lie below that of the pure solvent. D is the point of intersection of the vapour pressure curves of solution and pure solvent.

The measured depression in freezing point ATf is found to be directly proportional to the molality m of the solute in solution. It is also the depression in freezing point of one molal solution. Freezing point depression of a dilute solution is found to be directly proportional to the number of moles and hence the no. Also ATf is independent of the nature of the solute as long as it is non- volatile. Hence depression in freezing point is considered as a colligative property. W1 Kf. Problem 4 1.

The freezing point depression constant of benzene is 5. Find the molecular mass of the solute. Beckmann thermometer is not used in determining the absolute value of freezing temperature of the solvent or that of the solution.

It is therefore called a differential thermometer. Temperature differences of even 0. As the capillary has fine bore, a small change of temperature causes a considerable change in the height of mercury column level in the capillary.

The whole scale of a Beckmann thermometer covers only about 6K. Initially the level of mercury in the capillary should be on the scale. This is achieved by transferring mercury from the lower bulb to the reservoir and viceversa.

When the Beckmann thermometer is used at high temperatures, some of the mercury from the thermometer bulb is transferred into the upper reservoir. At lower temperature mercury from the reservoir falls down in to the thermometer bulb. Measurement of freezing point depression by Beckmann method A simple Beckmann apparatus is shown in Fig. It consists of a freezing tube a with a side arm c through which a known amount of a solute can be introduced.

A stopper carrying a Beckmann thermometer b and a stirrer d is fitted in to the freezing tube. To prevent rapid cooling of the contents of the freezing tube, A, a guard tube e surrounds the tube so that there is an air space between a and e. It is cooled with gentle and continuous stirring. As a result of super cooling, the temperature of the solvent will fall by about 0. Vigorous stirring is then set in when solid starts separating and the temperature rises to the exact freezing point.

The tube a is taken out, warmed to melt the solid and a known weight of the solute is added through the side arm c.

When the solute is dissolved in to the solvent forming a solution, the tube a is put back in to the original position and the freezing point of the solution T is redetermined in the same manner as before. The difference between the two readings gives the freezing point depression ATf. From this value, the molecular mass of the non-volatile solute can be determined using the expression and known Kf value.

Since the vapour pressure of a solution is always lower than that of the pure solvent, it follows that the boiling point of a solution will always be higher than of the pure solvent. The lower curve represents the vapour pressure - temperature dependance of a dilute solution with known concentration.

It is evident that the vapour pressure of the solution is lower than that of the pure solvent at every temperature. Elevation of boiling point is found directly proportional to the molality of the solution or inturn the number of molecules of solute. Also it is independent of the nature of the solute for a non-volatile solute. Hence, boiling point elevation is a colligative property.

ATb am It is defined as the elevation of boiling point of one molal solution. Wi 49 Table PtK Kb K. The molal elevation constant of benzene is 2. An inverted funnel tube b placed in the boiling tube collects the bubbles rising from a few fragments of a porous pot placed inside the liquid. When the liquid starts boiling, it pumps a stream of a liquid and vapour over the bulb of the Beckmann thermometer f held a little above the liquid surface. In this way, the bulb is covered with a thin layer of boiling liquid which is in equilibrium with the vapour.

This ensures that the temperature reading is exactly that of the boiling liquid and that superheating is minimum. After determining the boiling point of the pure solvent, a weighed amount of the solute is added and procedure is repeated for another reading. The vapours of the boiling liquid is cooled in a condenser C which has circulation of water through d and e.

The cooled liquid drops into the liquid in a. To Condenser C Problem 7 Fig.

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This solution boiled 0. Molal elevation constant for benzene is 2. Calculate the molecular weight of the solute. Molecular weight! The flow of the solvent from its side a to solution side b separated by semipermeable membrane c can be stopped if some definite extra pressure is applied on the solution risen to height h.

This pressure that just stops the flow of solvent is called osmotic pressure of the solution. Osmosis is a process of prime importance in living organisms. The salt concentration in blood plasma due to different species is equivalent to 0. If blood cells are placed in pure water, water molecules rapidly move into the cell. The movement of water molecules into the cell dilutes the salt content. As a result of this transfer of water molecules the blood cells swell and burst.

Hence, care is always taken to ensure that solutions that flow into the blood stream have the same osmotic pressure as that of the blood. Osmosis is also critically involved in the functioning of kidneys. They are known as isotonic solutions. He concluded that, a substance in solution behaves exactly like gas and the osmotic pressure of a dilute solution is equal to the pressure which the solute would exert if it is a gas at the same temperature occupying the same volume as the solution. Thus it is proposed 53 that solutions also obey laws similar to gas laws.

Boyle's -VantHoff law The osmotic pressure 71 of the solution at constant temperature is directly proportional to the concentration C of the solution. Charle's - Vant Hoff law At constant concentration the osmotic pressure n of the solution is directly proportional to the temperature T.

Determination of molecular weight by osmotic pressure measurement The osmotic pressure is a colligative property as it depends, on the number of solute molecules and not on their identity. The apparatus Fig. The inner tube a is made of semipermeable membrane c with two side tubes. The outer tube b is made of gun metal which contains the solution. The solvent is taken in the inner tube. As a result of osmosis, there is fall of level in the capillary indicator d attached to the inner tube.

The external pressure is applied by means of a piston e attached to the outer tube so that the level in the capillary indicator remains stationary at d. The osmotic pressure is recorded directly and the method is quick.

There is no change in the concentration of the solution during the measurement of osmotic pressure. The osmotic pressure is balanced by the external pressure and there is minimum strain on the semipermeable membrane.

Problem 8 lOg of an organic substance when dissolved in two litres of water gave an osmotic pressure of 0. Calculate the molecular weight of the substance. However, in some cases experimental values of colligative properties differ widely from those obtained theoretically. Such experimental values are referred to as abnormal colligative properties. The abnormal behaviour of colligative properties has been explained in terms of dissociation and association of solute molecules.

Dissociation of solute molecules Such solutes which dissociate in solvent water i. This effect results in an increase in colligative properties obtained experimentally. We can calculate the degree of dissociation a using the equation. In such case, the number of effective particles increases and therefore observed colligative property is greater than normal colligative property. Problem 9 A 0. Wi Observed molecular mass 1.

Association of the solute molecules Such solute which associate in a solvent show a decrease in number of particles present in solution. This effect results in a decrease in colligative properties obtained experimentally.

In this case, the number of particles is reduced to half its original value due to dimerisation. In such case, the experimental colligative property is less than normal colligative property.

Kb M,W, 3.Electrovalent bonding a. A liquid in equilibrium with its vapours in a sealed tube represents a closed system since the sealed container may be heated or cooled to add or remove energy from its contents while no matter liquid or vapour can be added or removed.

What is the structure of BeCl 2? In case of Boron VB theory predicts univalency due to the presence of one unpaired electron but in practice Boron is trivalent since compounds as BCb, BH3 etc. For H atom electronegativity is 2.

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