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First year Chemistry Lecture Notes

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CHEMISTRY HIGHER SECONDARY - FIRST YEAR VOLUME - II REVISED BASED ON THE RECOMMENDATIONS OF THE TEXT BOOK DEVELOPMENT COMMITTEE A Publication Under Government of Tamilnadu Distribution of Free Textbook Programme (NOT FOR SALE) Untouchability is a sin Untouchability is a crime Untouchability is inhuman TAMILNADU TEXTBOOK CORPORATION College Road, Chennai - 600 006Syllabus : Higher Secondary - First Year Chemistry INORGANIC CHEMISTRY Unit I - Chemical Calculations Significant figures - SI units - Dimensions - Writing number in scientific notation - Conversion of scientific notation to decimal notation - Factor label method - Calculations using densities and specific gravities - Calculation of formula weight - Understanding Avogadro’s number - Mole concept-mole fraction of the solvent and solute - Conversion of grams into moles and moles into grams - Calculation of empirical formula from quantitative analysis and percentage composition - Calculation of molecular formula from empirical formula - Laws of chemical combination and Dalton’s atomic theory - Laws of multiple proportion and law of reciprocal proportion - Postulates of Dalton’s atomic theory and limitations - Stoichiometric equations - Balancing chemical equation in its molecular form - Oxidation reduction-Oxidation number - Balancing Redox equation using oxidation number - Calculations based on equations. - Mass/Mass relationship - Methods of expressing concentration of solution - Calculations on principle of volumetric analysis - Determination of equivalent mass of an element - Determination of equivalent mass by oxide, chloride and hydrogen displacement method - Calculation of equivalent mass of an element and compounds - Determination of molar mass of a volatile solute using Avogadro’s hypothesis. 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 - I 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. - Quantum designation of electron - Pauli’s exclusion principle - Hund’s rule of maximum multiplicity - Aufbau principle - Stability of orbitals - Classification of elements based on electronic configuration. Unit 4 - Periodic Classification - I Brief history of periodic classification - IUPAC periodic table and IUPAC nomenclature of elements with atomic number greater than 100 - Electronic configuration and periodic table - Periodicity of properties Anomalous periodic properties of elements. Unit 5 - Group-1s 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-13. Boron Group - Important ores of Boron - Isolation of Born-Properties - Compounds of Boron- Borax, Boranes, diboranes, Borazole-preparation. properties - Uses of Boron and its compounds - Carbon group - Group -14 - Allotropes of carbon - Structural difference of graphite and diamond - General physical and chemical properties of oxides, carbides, halides and sulphides of carbon group - Nitrogen - Group-15 - Fixation of nitrogen - natural and industrial - HNO -Ostwald process 3 - Uses of nitrogen and its compounds - Oxygen - Group-16 - Importance of molecular oxygen-cell fuel - Difference between nascent oxygen and molecular oxygen - Oxides classification, acidic basic, amphoteric, neutral and peroxide - Ozone preparation, property and structure - Factors affecting ozone layer. (vi)Physical Chemistry Unit 8 - Solid State - I Classification of solids-amorphous, crystalline - Unit cell - Miller indices - Types of lattices belong to cubic system. 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 - I 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 - I 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 - I 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. ORGANIC CHEMISTRY Unit 15 - Basic Concepts of Organic Chemistry Catenation - Classification of organic compounds - Functional groups - Nomenclature - Isomerism - Types of organic reactions - Fission of bonds - Electrophiles and nucleophiles - Carbonium ion Carbanion - Free radicals - Electron displacement in covalent bond. 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 18 - Hydrocarbons Classification of Hydrocarbons - IUPAC nomenclature - Sources of alkanes - General methods of preparation of alkanes - Physical properties - (viii)Chemical properties - Conformations of alkanes - Alkenes - IUPAC nomenclature of alkenes - General methods of preparation - Physical properties - Chemical properties - Uses - Alkynes - IUPAC Nomenclature of alkynes - General methods of preparation - Physical properties - Chemical properties - Uses. 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. (ix)CHEMISTRY PRACTICALS FOR STD XI I. Knowledge of using Burette, Pipette and use of logarithms is to be demonstrated. II. Preparation of Compounds. 1. Copper Sulphate Crystals from amorphous copper sulphate solutions 2. Preparation of Mohr’s Salt 3. Preparation of Aspirin 4. Preparation of Iodoform 5. Preparation of tetrammine copper (II) sulphate III. Identification of one cation and one anion from the following. (Insoluble salt should not be given) ++ ++ ++ 2+ ++ ++ ++ ++ + Cation : Pb , Cu , Al , Mn , Zn , Ca , Ba , Mg , NH . 4 Anions : Borate, Sulphide, Sulphate, Carbonate, Nitrate, Chloride, Bromide. IV. Determination of Melting point of a low melting solid. V. Acidimetry Vs Alkalimetry 1. Preparation of Standard solution of Oxalic acid and Sodium Carbonate solution. 2. Titration of HCl Vs NaOH 3. Titration of HCl Vs Na CO 2 3 4. Titration of Oxalic acid Vs NaOH (x) 10. CHEMICAL BONDING OBJECTIVES To know about bonding as binding forces between atoms to form • molecules. To learn about Kossel-Lewis approach to chemical bonding, the octet • rule, its limitations and Lewis representations of simple molecules. To know about ionic bond, lattice energy and Born-Haber cycle. • To understand covalent bond, directional character. • To learn about VSEPR model and predict the geometry of simple • molecules. To understand the concepts of hybridisation, 1DQGŒERQGVUHVRQDQFH • and coordinate covalent bonds. 10.1 Elementary theories on Chemical Bonding The study on the "nature of forces that hold or bind atoms together to form a molecule" is required to gain knowledge of the following- i) to know about how atoms of same element form different compounds combining with different elements. ii) to know why particular shapes are adopted by molecules. iii) to understand the specific properties of molecules or ions and the relation between the specific type of bonding in the molecules. 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. 1 Classification of molecules Molecules having two identical atoms like H , O , Cl , N etc. are 2 2 2 2 called as homonuclear diatomic molecules. Molecules containing two different atoms like CO, HCl, NO, HBr etc., are called as heteronuclear diatomic molecules. Molecules containing identical but many atoms bonded together such as P, S etc., are called as homonuclear 4 8 polyatomics. In most of the molecules, more than two atoms of different kinds are bonded such as in molecules like NH , CH COOH, SO , HCHO 3 3 2 and they are called as heteronuclear polyatomics. 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 1916, W.Kossel and G.N.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. 10.1.1 Kossel-Lewis approach to Chemical Bonding W.Kossel laid down the following postulates to the understanding of ionic bonding: In the periodic table, the highly electronegative halogens and the • highly electropositive alkali metals are separated by the noble gases. Therefore one or small number of electrons are easily gained and transferred to attain the stable noble gas configuration. The formation of a negative ion from a halogen atom and a positive • ion from an alkali metal atom is associated with the gain and loss of an electron by the respective atoms. 2 The negative and positive ions so formed attains stable noble gas • electronic configurations. 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 2 6 electrons) with a general representation ns np . The negative and positive ions are bonded and stabilised by force of • electrostatic attraction. 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: lossofe + ──→ ─ ─ (i) Na Na + e 1 Ne 3s Ne where Ne = electronic configuration of Neon 2 6 = 2s 2p gain ofe - ─ ─→ ─ ─ (ii) Cl + e Cl 2 5 Ne3s 3p Ar Ar = electronic configuration of Argon + - + - electrostatic (iii) Na+Cl NaCl(or)Na Cl ───→ ─ ─ attraction NaCl is an electrovalent or ionic compound made up of sodium ions and chloride ions. 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. The + - coulombic or electrostatic attraction between Na and Cl ions result in NaCl formation. Similarly formation of MgO may be shown to occur by the transfer of two electrons as: 3 - loss ofe 2+ ──→ ─ ─ (i) Mg Mg + 2e 2 Ne3s Ne - gain ofe 2- ───→ ─ ─ (ii) O + 2e O 2 4 2 6 He2s 2p He2s 2p (or) Ne 2+ 2- 2+ 2- electrostatic (iii) Mg +O─ ──→ ─ ─ MgO(or)Mg O attraction The bonding in MgO is also electrovalent or ionic and the electrostatic 2+ 2- forces of attraction binds Mg ions with O ions. 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 H , O , Cl etc., and in other organic molecules that have 2 2 2 ions. G.N.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 : "a bond is formed between two atoms by mutual sharing of pairs of electrons to attain a stable outer-octet of electrons for each atom involved in bonding". 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 F , Cl , H etc., the bond is formed by the 2 2 2 sharing of a pair of electrons between the atoms. For example, consider the formation of a fluorine molecule (F). The atom has electronic 2 2 2 5 configuration. He2s 3s 3p which is having one electron less than the electronic configuration of Neon. In the fluorine molecule, each atom 4 contributes one electron to the shared pair of the bond of the F molecule. 2 In this process, both the fluorine atoms attain the outershell octet of a noble gas (Argon) (Fig. 10.1(a)). Dots ( ) represent electrons. 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 : Each bond is the result of sharing of an electron pair between the • atoms comprising the bond. Each combining atom contributes one electron to the shared pair. • The combining atoms attain the outer filled shells of the noble gas • configuration. 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. x • (or) F-F 8e 8e Fig. 10.1(a) F molecule 2 In carbon dioxide (CO ) two double bonds are seen at the centre carbon 2 atom which is linked to each oxygen atom by a double bond. The carbon and the two oxygen atoms attain the Neon electronic configuration. Fig. 10.1 (b) CO molecule 2 5 When the two combining atoms share three electron pairs as in N 2 molecule, a triple bond is said to be formed. Each of the Nitrogen atom shares 3 pairs of electrons to attain neon gas electronic configuration. Fig. 10.1 (c) N molecule 2 10.2 Types of Bond There are more than one type of chemical bonding possible between atoms which makes the molecules to show different characteristic properties. 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. 10.3 Ionic (or) Electrovalent bond The electrostatic attraction force existing between the cation and the anion produced by the electron transfer from one atom to the other is known as the ionic (or) electrovalent bond. 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. 6 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 0.9 while Cl has 3.0, thus Na and Cl atoms when brought together will form an ionic bond. For example, NaCl is formed by the electron ionisation of sodium atom + to Na ion due to its low ionisation potential value and chlorine atom to chloride ion by capturing the odd electron due to high electron affinity. Thus, NaCl (ionic compound) is formed. In NaCl, both the atoms possess unit charges. ionisation + - ──→ ─ ─ i) Na(g) Na + e (g) 2 6 1 2 6 2s2p3s 2s sp sodium cation affinity - - ─ ─→ ─ ─ ii) Cl(g) + e Cl 2 5 2 6 3s3p 3s , 3p chloride anion + - electrostatic iii) Na + Cl ───→ ─ ─ NaCl attraction Sodium Chloride ionic/crystalline ion ion compound is formed Fig. 10.2 Electron transfer between Na and Cl atoms during ionic bond formation in NaCl In CaO, which is an ionic compound, the formation of the ionic bond involves two electron transfers from Ca to O atoms. Thus, doubly charged positive and negative ions are formed. ionisation 2+ - ──→ ─ ─ Ca Ca +2e (Calcium Cation) 6 2 3p 4s electron - 2 (Oxide anion) + O 2e─ ─→ ─ ─ O 2 4 2 6 aaffinity 2s 2p 2s 2p 7 electrostatic 2+ 2- Ca + O───→ ─ ─ CaO attraction ionic compound Ionic bond may be also formed between a doubly charged positive ion with single negatively charged ion and vice versa. The molecule as a whole remains electrically neutral. For example in MgF , Mg has two positive 2 2+ charges and each fluorine atom has a single negative charge. Hence, Mg - binds with two fluoride (F ) ions to form MgF which is electrically neutral. 2 2+ - Mg Mg + 2e → 2 6 2 2 6 (2s 2p 3s (2s 2p ) - 2e + 2F 2F → 2 5 2 6 (2s 2p) (2s 2p ) 2+ - i.e:- Mg + 2F MgF → 2 Magnesium - fluoride (an ionic compound). Similarly in Aluminium bromide (AlBr ), Aluminium ion has three 3 positive charges and therefore it bonds with three Bromide ions to form AlBr which is a neutral ionic molecule. 3 3+ - Al Al + 3e → 6 2 1 2 6 2p 3s 3p (2s , 2p ) - - 3 Br + 3e 3 Br → 2 5 2 6 (4s 4p) (4s 4p ) 3+ - Al + 3Br AlBr (ionic bond) → 3 10.3.1 Lattice energy and Born - Haber's cycle Ionic compounds in the crystalline state exist as three dimensionally ordered arrangement of cations and anions which are held together by columbic interaction energies. 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 8 solid into its respective ions at infinity separation is taken the lattice enthalpy. + - MX M + X (s) → (g) (g) H = L.E Δr° Lattice enthalpy is a positive value. -1 For example, the lattice enthaply of NaCl is 788 kJ.mol . This means that 788 kJ of energy is required to separate 1 mole of solid NaCl into 1 + - mole of Na and 1 mole of Cl (g) to an infinite distance. (g) 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. 0 0 H H (1) (3) e M M M (g) (s) (g) 0 0 H H 2 4 1/2X X X 2(g) (g) (g) +e 0 H f MX (s) o û+ = enthalpy change for sublimation of M to M 1 (s) (g) 9 o û+ = enthalpy change for dissociation of 1/2 X to X 2 2(g) (g) o + û+ = ionization energy of M to M 3 (g) (g) o û+ = electronic affinity or electron gain energy for conversion of X 4 (g) - to X (g) o û+ = the lattice enthalpy for formation of solid MX (1 mole). 5 o û H = enthalpy change for formation of MX solid directly from the f respective elements such as 1 mole of solid M and 0.5 moles of X . 2(g) According to Hess's law, o o o o o o û+ = û+ û+ û+ û+ û+ f 1 2 3 4 5 Some important features of lattice enthalpy are: i. The greater the lattice enthalpy the more stabler the ionic bond formed. ii. The lattice enthalpy is greater for ions of higher charge and smaller radii. iii. The lattice enthalpies affect the solubilities of ionic compounds. Calculation of lattice enthalpy of NaCl Let us use the Born - Haber cycle for determining the lattice enthalpy of NaCl as follows : o 7KHVWDQGDUGHQWKDOS\FKDQJHû H overall for the reaction, f -1 Na + 1/2 Cl NaCl is - 411.3 kJmol (s) 2(g) → (s) o H Δf Na + ½ Cl NaCl (s) 2(g) (s) o 7KHYDOXHRIû+ calculated using the equation of Born - Haber cycle should be 5 reversed in sign 10 Sublimation ½ Dissociation o o H H 1 2 Δ Δ H 5 Δ Na Cl (g) (g) Ionization Energy Electron Affinity o o H H Δ 3 Δ4 + - Na Cl (g) (g) Fig. 10.3 Born-Haber cycle for Lattice enthalpy determination involving various stepwise enthalpic processes for NaCl solid formation Since the reaction is carried out with reactants in elemental forms and products in their standard states, at 1 bar, the overall enthalpy change of the reaction is also the enthalpy of formation for NaCl. 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. Atomisation : 1 û+ for Na(s) Na(g) is + 108.70 (kJ mol ) °1 → Dissociation: û+ for ½ Cl (g) Cl(g) is + 122.0 °2 2 → Ionisation : + û+ for Na(g) Na (g) + e is + 495.0 °3 → Electron affinity : - û+ for e + Cl(g) Cl (g) is - 349.0 °4 → Lattice enthalpy : + - û+ for Na (g) + Cl (g) NaCl(g) is ? 5 ° → 11 û H û+ û+ û+ û+ û+ f ° °1 °2 °3 °4 °5 û+ -411.3 = 108.70 + 122.0 + 495 - °5 -1 û+ = -788.0 kJ mol ∴ °5 But the lattice enthalpy of NaCl is defined by the reaction + - NaCl(g) Na (g) + Cl (g) only. → /DWWLFHHQWKDOS\YDOXHIURPû+ is written with a reversed sign. 5 ∴ ° -1 Lattice enthalpy of NaCl = +788.0 kJ mol . ∴ Problem 1 Calculation of lattice enthalpy of MgBr from the given data. 2 Solution The enthalpy of formation of MgBr according to the reaction 2 Mg(s) + Br (l) MgBr V û H = -524 kJ/mol 2 → 2 f ° -1 û+ for Mg(s) Mg(g) = + 148 kJ mol °1 → 2+ - -1 û+ for Mg(g) Mg (g) +2e = +2187kJ mol °2 → -1 û+ for Br (l) Br (g) = 31 KJ mol 3 2 2 ° → -1 û+ for Br (g) 2Br(g) = 193 KJ mol 4 2 ° → - - -1 û+ for Br(g) + e (g) Br = -331 KJ mol 5 ° → 2+ - û+ for Mg (g)+2Br (g) Mg Br(s) = ? °6 → 2 ûH  û+ û+ û+ û+ û+ û+ f ° °1 °2 °3 °4 °5 °6 -1 -524 kJ mol = (+148 + 2187 + 31 0 -1 + 193 - 2(331) + H ) kJ mol Δ 6 -1 0 = -2421 KJ mol= H Δ 6 -1 Hence, lattice enthalpy of Mg Br  û+ = 2421 kJ mol 2 °6 10.3.2 Properties of electrovalent (or) ionic compounds Ionic compounds possess characteristic properties of their own like physical state, solubility, melting point, boiling point and conductivity. The nature of these properties are discussed as follows. i. 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. ii. 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. iii. 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. iv. 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. v. Electrovalent compounds having the same electronic configuration exhibit isomorphism. 10.4 Covalent bond A covalent bond is a chemical bond formed when two atoms mutually share a pair of electron. 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 (H ) a covalent bond results by the overlap of the two s 2 orbitals each containing an electron from each of the two H atoms of the 2 ' molecule. Each H atom attains '1s filled K shell. 13