Arihant Chemistry Handbook PDF

hand book KEY NOTES TERMS DEFINITIONS FORMULAE Chemistry Highly Useful for Class XI & XII Students, Engineering & Medical Entrances and Other Competitions hand book KEY NOTES TERMS DEFINITIONS FORMULAE Chemistry Highly Useful for Class XI & XII Students, Engineering & Medical Entrances and Othe...

hand book KEY NOTES TERMS DEFINITIONS FORMULAE Chemistry Highly Useful for Class XI & XII Students, Engineering & Medical Entrances and Other Competitions hand book KEY NOTES TERMS DEFINITIONS FORMULAE Chemistry Highly Useful for Class XI & XII Students, Engineering & Medical Entrances and Other Competitions Preeti Gupta Supported by Saleha Khan Shahana Ansari ARIHANT PRAKASHAN, (SERIES) MEERUT  Arihant Prakashan (Series), Meerut All Rights Reserved © Publisher No part of this publication may be re-produced, stored in a retrieval system or distributed in any form or by any means, electronic, mechanical, photocopying, recording, scanning, web or otherwise without the written permission of the publisher. Arihant has obtained all the information in this book from the sources believed to be reliable and true. However, Arihant or its editors or authors or illustrators don’t take any responsibility for the absolute accuracy of any information published and the damages or loss suffered there upon. All disputes subject to Meerut (UP) jurisdiction only. Administrative & Production Offices Regd. Office ‘Ramchhaya’ 4577/15, Agarwal Road, Darya Ganj, New Delhi -110002 Tele: 011- 47630600, 43518550; Fax: 011- 23280316 Head Office Kalindi, TP Nagar, Meerut (UP) - 250002 Tele: 0121-2401479, 2512970, 4004199; Fax: 0121-2401648 Sales & Support Offices Agra, Ahmedabad, Bengaluru, Bareilly, Chennai, Delhi, Guwahati, Hyderabad, Jaipur, Jhansi, Kolkata, Lucknow, Meerut, Nagpur & Pune ISBN : 978-93-13196-49-5 Published by Arihant Publications (India) Ltd. For further information about the books published by Arihant log on to www.arihantbooks.com or email to [email protected] /arihantpub /@arihantpub Arihant Publications /arihantpub PREFACE Handbook means reference book listing brief facts on a subject. So, to facilitate the students in this we have released this Handbook of Chemistry this book has been prepared to serve the special purpose of the students, to rectify any query or any concern point of a particular subject. This book will be of highly use whether students are looking for a quick revision before the board exams or just before other examinations like Engineering Entrances, Medical Entrances or any similar examination, they will find that this handbook will answer their needs admirably. This handbook can even be used for revision of a subject in the time between two shift of the exams, even this handbook can be used while travelling to Examination Centre or whenever you have time, less sufficient or more. The format of this handbook has been developed particularly so that it can be carried around by the students conveniently. The objectives of publishing this handbook are : — To support students in their revision of a subject just before an examination. — To provide a focus to students to clear up their doubts about particular concepts which were not clear to them earlier. — To give confidence to the students just before they attempt important examinations. However, we have put our best efforts in preparing this book, but if any error or what so ever has been skipped out, we will by heart welcome your suggestions. A part from all those who helped in the compilation of this book a special note of thanks goes to Ms. Shivani of Arihant Publications. Author  CONTENTS 1. Basic Concepts of Chemistry 1-14 — Chemistry — Dalton's Atomic Theory — Matter — Mole Concept — Atoms and Molecules — Atomic Mass — Physical Quantities and — Molecular Mass Their Measurement Units — Equivalent Mass — Dimensional Analysis — Stoichiometry — Scientific Notation — Per cent Yield — Precision and Accuracy — Empirical and Molecular — Laws of Chemical Formulae Combinations 2. Atomic Structure 15-29 — Atom — Planck's Quantum Theory — Electron — Bohr's Model — Proton — Sommerfeld Extension to — Neutron Bohr's Model — Thomson's Atomic Model — de-Broglie Principle — Rutherford's Nuclear — Heisenberg's Uncertainty Model of Atom Principle — Atomic Number — Quantum Mechanical Model — Mass Number of Atom — Electromagnetic Wave — Quantum Numbers Theory (Maxwell) — Electronic Configuration 3. Classification of Elements and Periodicity 30-42 in Properties — Classification of Elements — Mendeleev's Periodic Table — Earlier Attempts of — Modern Periodic Table Classify Elements — Periodic Properties 4. Chemical Bonding and Molecular Structure 43-59 — Chemical Bond — Resonance — Ionic Bond — VSEPR Theory — Born Haber Cycle — VBT Theory — Covalent Bond — Hybridisation — Octet Rule — MO Theory — Bond Characteristics — Hydrogen Bond — Dipole Moment — Metallic Bond — Fajan's Rule 5. States of Matter 60-72 — Factors Deciding Physical — Graham's Law Diffusion State of a Substance — Dalton's Law — The Gaseous state — Kinetic Theory of Gases — Boyle's Law — Van der Waals' Equation — Charles' Law — Liquefaction of Gases and — Gay Lussac's Law Critical Points — Avogadro's Law — Liquid State — Ideal Gas Equation 6. The Solid State 73-86 — Solids — Structure of Ionic Crystals — Bragg's Equation — Imperfections Defects in Solids — Unit Cell — Point Defects — Seven Crystal Systems — Classification of Solids on the — Packing Fraction Basis of Electrical — Coordination Number Conductivity — Density of Unit Cell — Magnetic Properties of Solids 7. Thermodynamics 87-100 — Thermodynamic — Various forms of Enthalpy of Properties Reaction — Laws of Thermochemistry — Thermodynamic Process — Bond Enthalpy — Internal — Zeroth Law of (E or U) Energy — Entropy (S) Thermodynamics — Spontaneous Process — First Law of — Second Law of Thermodynamics Thermodynamics — Enthalpy (H) — Joule Thomson Effect  — Carnot Cycle — Third Law of — Gibbs Free Energy Thermodynamics 8. Chemical Equilibrium 101-107 — Physical and Chemical — Law of Mass Action Processes — Relation Between Kc and Kp — Types of Chemical — Types of Equilibrium Reactions — Reaction Quotient — Equilibrium State — Le-Chatelier's Principle 9. Ionic Equilibrium 108-120 — Electrolytes Weak Acid and Weak Base — Calculation of the Degree — Buffer Solutions of Dissociation (a) — Salts — Ostwald's Dilution Law — Common Ion Effect — Acids and Bases — Solubility Product — The pH Scale — Acid Base Indicator — Dissociation Constant of 10. Solutions 121-135 — Solubility — Azeotropic Mixture — Henry's Law — Colligative Properties — Concentration of — Osmotic Properties Solutions — Abnormal Molecular Masses — Raoult's Law — van't Hoff Factor (i) 11. Redox Reactions 136-143 — Oxidation Number — Balancing of Redox Chemical Equations 12. Electrochemistry 144-159 — Conductors — Electrochemical Series — Electrochemical Cell and — Nernst Equation Electrolytic Cell — Concentration Cell — Electrode Potential — Conductance (G) — Reference Electrode — Specific Conductivity — Electromotive Force (emf) — Molar Conductivity of a Cell — Kohlrausch's Law  — Electrolysis — Batteries — Faraday's Laws of — Fuel Cells Electrolysis — Corrosion 13. Chemical Kinetics 160-169 — Rate of Reaction — Methods to Determine Order — Rate Law Expressions of Reaction — Rate Constant — Arrhenius Equation — Order and Molecularity of — Activated Complex a Reaction — Role of Catalyst in a Chemical — Zero Order Reactions Reaction — First Order Reactions — Theory of Reaction Rates — Pseudo First Order — Photochemical Reactions Reaction 14. Surface Chemistry 170-179 — Adsorbtion — Enzyme Catalysis — Catalysis 15. Colloidal State 180-187 — Classification of Colloids Solution — Preparation of Colloids — Protective Colloids — Purification of Colloidal — Emulsion Solutions — Gels — Properties of Colloidal — Applications of Colloids 16. Principles & Processes of Isolation 188-203 of Elements — Elements in Nature — Purification of Crude Metals — Minerals and Ores — Occurance and Extraction of — Metallurgy Some Metals — Thermodynamic Principle in Extraction of Metals 17. Hydrogen 204-216 — Position of Hydrogen in the — Water Periodic Table — Heavy Water — Dihydrogen — Soft and Hard Water — Different Forms of — Hydrogen Peroxide Hydrogen 18. The s-Block Elements 217-236 — Alkali Metals — Anomalous Behaviour of Be — Anomalous Behaviour Li — Compounds of Calcium — Compounds of Sodium — Cement — Alkaline Earth Metals 19. The p-Block Elements 237-283 — Elements of Group-13 — LPG — Anomalous Behaviour of — Compounds of Silicon Boron — Compounds of Lead — Boron and Its Compounds — Elements of Group-15 — Compounds of — Nitrogen and Its Compounds Aluminium — Phosphorus and Its — Elements of Group-14 Compounds — Carbon and Its — Elements of Group-16 Compounds — Oxygen and Its Compounds — Coal Gas — Compounds of Sulphur — Natural Gas — Elements of Group-17 — Oil Gas — Chlorine and Its Compounds — Wood Gas — Elements of Group-18 20. The d-and f-Block Elements 284-296 — Transition Elements — Silver Nitrate — Potassium Dichromate — Inner-Transition Elements — Potassium Permanganate — Lanthanides — Copper Sulphate — Actinoids 21. Coordination Compounds 297-310 — Terms Related to — Bonding in Coordination Coordination Compounds Compounds — Types of Complexes — Werner's Theory — Effective Atomic Number — VBT (EAN) — CFT — IUPAC Naming of — Importance and Applications Complex Compounds of Coordination Compounds — Isomerism in — Organometallic Compounds Coordination Compounds 22. Environmental Chemistry 311-323 — Classification of Global Warming Environment — Acid Rain — Environmental Pollution — Stratospheric Pollution — Pollutants — Water Pollution — Tropospheric Pollution — Soil or Land Pollution — Air Pollution — Radioactive Pollution — Smog — Bhopal Gas Tragedy — Green House Effect and — Green Chemistry 23. Purification and Characterisation of 324-333 Organic Compounds — Purification of Organic — Quantitative Estimation of Compounds Elements — Qualitative Analysis of — Determination of Empirical Organic Compounds and Molecular Formula 24. General Organic Chemistry 334-360 — Organic Chemistry Compounds — Classification of Organic — Fission of a Covalent Bond Compounds — Attacking Reagents — Classification of Carbon — Reaction Intermediate and Hydrogen Atoms — Inductive Effect — Functional Group — Electromeric Effect — Homologous Series — Hyperconjugation — Representation of — Resonance Different Formulae — Isomerism — Nomenclature of Organic — Types of Organic Reactions 25. Hydrocarbons 361-383 — Alkanes — Benzene — Conformations of Alkanes — Petroleum — Alkenes — Octane Number — Conjugated Diene — Cetane Number — Alkynes 26. Haloalkanes and Haloarenes 384-397 — General Methods of Halides Preparation of — Dihalogen, Trihalogen, Haloalkanes and Aryl Polyhalogen Derivatives 27. Alcohols, Phenols and Ethers 398-419 — Alcohols and Phenols and Phenols — Classification — Dihydric Alcohols — Structure — Trihydric Alcohols — Nomenclature — Ethers — Preparation of Alcohols 28. Aldehydes, Ketones and Carboxylic Acids 420-442 — Aldehydes and Ketones — Nomenclature — Nomenclature — Preparation — Classification — Properties — Preparation — Derivatives of Carboxylic — Carboxylic Acids Acids — Classification — Properties 29. Amines 443-457 — Structure Chloride — Preparation — Alkyl Cyanides — Properties — Alkyl Isocyanides — Benzene Diazonium — Nitro Compounds 30. Polymers 458-474 — Polymerisation — Natural Rubber — Classification — Neoprene — Types of Polymerisation — Buna-N — Molecular Mass of — Polyesters Polymers — Biopolymers and — Polyolefins Biodegradable Polymers — Resin 31. Biomolecules 475-494 — Carbohydrates — Lipids — Amino Acids — Acid Value — Proteins — Blood — Enzymes — Hormones — Nucleic Acids — Vitamins 32. Chemistry in Everyday Life 495-509 — Medicines or Drugs — Chemistry in Colouring — Chemicals in Food Matter — Food Preservatives — Chemistry in Cosmetics — Cleansing Agents — Rocket Propellants 33. Nuclear Chemistry 510-515 — Nucleons and Nuclear — Artificial Radioactivity Forces — Artificial Transmutation — Parameter of Nucleus — Nuclear Reactions — Factors Affecting Stability — Nuclear Fission Nucleus — Nuclear Fusion — Group Displacement Law — Applications of Radioactivity — Disintegration Series 34. Analytical Chemistry 516-539 — Qualitative Analysis of Organic Compounds Inorganic Compounds — Titrimetric Exercises — Qualitative Analysis of Appendix 540-560  1 Basic Concepts of Chemistry Chemistry It is the branch of science which deals with the composition, structure and properties of matter. Antoine Laurent Lavoisier is called the father of chemistry. Branches of Chemistry Inorganic chemistry is concerned with the study of elements (other than carbon) and their compounds. Organic chemistry is the branch of chemistry which is concerned with organic compounds or substances produced by living organisms. Chemistry Physical chemistry is concerned with the explanation of fundamental principles. Analytical chemistry is the branch of chemistry which is concerned with qualitative and quantitative analysis of chemical substances. In addition to these, biochemistry, war chemistry, nuclear chemistry, forensic chemistry, earth chemistry etc., are other branches of chemistry. 2 Handbook of Chemistry Matter Anything which occupies some space and has some mass is called matter. It is made up of small particles which have space between them. The matter particles attract each other and are in a state of continuous motion. Classification of Matter Matter Physical classification Chemical classification Homogeneous Solid Liquid Gas Pure substances Mixtures (For physical classification Heterogeneous see chapter 4) Elements Compounds Metals Non-metals Metalloids Inorganic compounds Organic compounds Pure Substances They have characteristics different from the mixtures. They have fixed composition, whereas mixtures may contain the components in any ratio and their composition is variable. Elements It is the simplest form of pure substance, which can neither be decomposed nor be built from simpler substances by ordinary physical and chemical methods. It contains only one kind of atoms. The number of elements known till date is 118. An element can be a metal, a non-metal or a metalloid. Hydrogen is the most abundant element in the universe. Oxygen (46.6%), a non-metal, is the most abundant element in the earth crust. Al is the most abundant metal in the earth crust.  Basic Concepts of Chemistry 3 Compounds It is also the form of matter which can be formed by combining two or more elements in a definite ratio by mass. It can be decomposed into its constituent elements by suitable chemical methods, e.g. water (H 2O) is made of hydrogen and oxygen in the ratio 1 : 8 by mass. Compounds can be of two types : (i) Inorganic compounds Previously, it was believed that these compounds are derived from non-living sources, like rocks and minerals. But these are infact the compounds of all the elements except hydrides of carbon (hydrocarbons) and their derivatives. (ii) Organic compounds According to earlier scientists, these compounds are derived from living sources like plants and animals, or these remain buried under the earth; (e.g. petroleum). According to modern concept, these are the hydrides of carbon and their derivatives. Mixtures These are made up of two or more pure substances. They can possess variable composition and can be separated into their components by some physical methods. Mixtures may be homogeneous (when composition is uniform throughout) or heterogeneous (when composition is not uniform throughout). Mixture Separation Methods Common methods for the separation of mixtures are: (a) Filtration Filtration is the process of separating solids that are suspended in liquids by pouring the mixture into a filter funnel. As the liquid passes through the filter, the solid particles are held on the filter. (b) Distillation Distillation is the process of heating a liquid to form vapours and then cooling the vapours to get back the liquid. This is a method by which a mixture containing volatile substances can be separated into its components. (c) Sublimation This is the process of conversion of a solid directly into vapours on heating. Substances showing this property are called sublimate, e.g. iodine, naphthalene, camphor. This method is used to separate a sublimate from non-sublimate substances. 4 Handbook of Chemistry (d) Crystallisation It is a process of separating solids having different solubilities in a particular solvent. (e) Magnetic separation This process is based upon the fact that a magnet attracts magnetic components of a mixture of magnetic and non-magnetic substances. The non-magnetic substance remains unaffected. Thus, it can be used to separate magnetic components from non-magnetic components. (f) Atmolysis This method is based upon rates of diffusion of gases and used for their separation from a gaseous mixture. Atoms and Molecules Atom is the smallest particle of an element which can take part in a chemical reaction. It may or may not be capable of independent existence. Molecule is the simplest particle of matter that has independent existence. It may be homoatomic, e.g. H2 , Cl2 , N2 (diatomic), O3 (triatomic) or heteroatomic, e.g. HCl, NH3 , CH4 etc. Physical Quantities and Their Measurements Physical quantity is a physical property of a material that can be quantified by measurement and their measurement does not involve any chemical reaction. To express the measurement of any physical quantity, two things are considered: (i) Its unit, (ii) The numerical value. Magnitude of a physical quantity = numerical value × unit Unit It is defined as ‘‘some fixed standard against which the comparison of a physical quantity can be done during measurement.’’ Units are of two types: (i) Basic units (ii) Derived units (i) The basic or fundamental units are length (m), mass (kg), time (s), electric current (A), thermodynamic temperature (K), amount of substance (mol) and luminous intensity (Cd). (ii) Derived units are basically derived from the fundamental units, e.g. unit of density is derived from units of mass and volume.  Basic Concepts of Chemistry 5 Different systems used for describing measurements of various physical quantities are: (a) CGS system It is based on centimetre, gram and second as the units of length, mass and time respectively. (b) FPS system A British system which used foot (ft), pound (lb) and second (s) as the fundamental units of length, mass and time respectively. (c) MKS system It is the system which uses metre (m), kilogram (kg) and second (s) respectively for length, mass and time; ampere (A) was added later on for electric current. (d) SI system (1960) International system of units or SI units contains following seven basic and two supplementary units: Basic Physical Quantities and Their Corresponding SI Units Physical quantity Name of SI unit Symbol for SI unit Length (l ) metre m Mass (m) kilogram kg Time (t ) second s Electric current (I) ampere A Thermodynamic temperature (T ) kelvin K Amount of substance (n) mole mol Luminous intensity (Iv ) candela Cd Supplementary units It includes plane angle in radian and solid angle in steradian. Prefixes The SI units of some physical quantities are either too small or too large. To change the order of magnitude, these are expressed by using prefixes before the name of base units. The various prefixes are listed as: 6 Handbook of Chemistry Multiple Prefix Symbol Multiple Prefix Symbol 24 –1 10 yotta Y 10 deci d 21 –2 10 zeta Z 10 centi c 1018 exa E 10–3 milli m 1015 peta P 10–6 micro µ 12 10 tera T 10–9 nano n 109 giga G 10–12 pico p 6 –15 10 mega M 10 femto f 103 kilo K 10−18 atto a 102 hecto h 10−21 zepto z 10 deca da 10−24 yocto y Some Physical Quantities (i) Mass It is the amount of matter present in a substance. It remains constant for a substance at all the places. Its unit is kg but in laboratories usually gram is used. (ii) Weight It is the force exerted by gravity on an object. It varies from place to place due to change in gravity. Its unit is Newton (N) (iii) Temperature There are three common scale to measure temperature °C (degree celsius), °F (degree fahrenheit) and K (kelvin). K is the SI unit. The temperature on two scales (°C and °F) are related to each other by the following relationship: 9 °F = ( ° C) + 32 5 The kelvin scale is related to celsius scale as follows: K = ° C + 273.15 (iv) Volume The space occupied by matter (usually by liquid or a gas) is called its volume. Its unit is m3. (v) Density It is defined as the amount or mass per unit volume and has units kg m −3 or g cm −3. Scientific Notation In such notation, all measurements (how so ever large or small) are expressed as a number between 1.000 and 9.999 multiplied or divided by 10. In general it can be given as = N × 10n  Basic Concepts of Chemistry 7 Here, N is called digit term (1.000–9.999) and n is known as exponent. e.g. 138.42 cm can be written as 1.3842 × 102 and 0.0002 can be written as 2.0 × 10−4. Precision and Accuracy Precision refers to the closeness of the set of values obtained from identical measurements of a quantity. Precision is simply a measure of reproducibility of an experiment. Precision = individual value – arithmetic mean value Accuracy is a measure of the difference between the experimental value or the mean value of a set of measurements and the true value. Accuracy = mean value – true value In physical measurements, accurate results are generally precise but precise results need not be accurate. Significant Figures Significant figures are the meaningful digits in a measured or calculated quantity. It includes all those digits that are known with certainty plus one more which is uncertain or estimated. Greater the number of significant figures in a measurement, smaller the uncertainty. Rules for determining the number of significant figures are: 1. All digits are significant except zeros in the beginning of a number. 2. Zeros to the right of the decimal point are significant. e.g. 0.132, 0.0132 and 15.0, all have three significant figures. 3. Exact numbers have infinite significant figures. Calculations Involving Significant Figures 1. In addition or subtraction, the final result should be reported to the same number of decimal places as that of the term with the least number of decimal places, e.g. 2.512 (4 significant figures) 2.2 (2 significant figures) 5.23 (3 significant figures) 9.942 ⇒ 9.9 (Reported sum should have only one decimal point.) 8 Handbook of Chemistry 2. In multiplication and division, the result is reported to the same number of significant figures as least precise term or the term with least number of significant figures, e.g. 15.724 ÷ 0.41 = 38.3512195121(38.35) Rounding Off the Numerical Results When a number is rounded off, the number of significant figures is reduced, the last digit retained is increased by 1 only if the following digit is ≥ 5 and is left as such if the following digit is ≤ 4, e.g. 12.696 can be written as 12.7 18.35 can be written as 18.4 13.93 can be written as 13.9 Dimensional Analysis Often while calculating, there is a need to convert units from one system to other. The method used to accomplish this is called factor label method or unit factor method or dimensional analysis. In this, Information sought = Information given × Conversion factor Important Conversion Factors −5 1dyne = 10 N 1L = 1000 mL –2 1atm = 101325 Nm = 1000 cm3 = 101325 Pa (pascal) = 10−3 m3 1bar = 1 × 105 Nm–2 = 1 dm3 5 = 1 × 10 (pascal) 1 L atm = 101.325 J = 24.21 cal 1 gallon = 3.7854 L 19 1cal = 4.184 J = 2.613 × 10 eV 1 eV/atom = 96.485 kJ mol −1 1eV = 1.602189 × 10–19 J 1amu or u = 1.66 × 10−27 kg 1 J = 10 7 erg = 931.5 MeV −10 1 Å = 10 m 1esu = 3.3356 × 10−10 C Laws of Chemical Combinations The combination of elements to form compounds is governed by the following six basic laws:  Basic Concepts of Chemistry 9 Law of conservation of mass (Lavoisier, 1789) This law states that during any physical or chemical change, the total mass of the products is equal to the total mass of reactants. It does not hold good for nuclear reactions. Law of definite proportions (Proust, 1799) According to this law, a chemical compound obtained by different sources always contains same percentage of each constituent element. Law of multiple proportions (Dalton, 1803) According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers, e.g. in NH3 and N 2H 4, fixed mass of nitrogen requires hydrogen in the ratio 3 : 2. Law of reciprocal proportions (Richter, 1792) According to this law, when two elements (say A and B ) combine separately with the same weight of a third element (say C), the ratio in which they do so is the same or simple multiple of the ratio in which they ( A and B) combine with each other. Law of definite proportions, law of multiple proportions and law of reciprocal proportions do not hold good when same compound is obtained by using different isotopes of the same element, e.g. H 2O andD2O. Gay Lussac’s law of gaseous volumes (In 1808) It states that under similar conditions of temperature and pressure, whenever gases react together, the volumes of the reacting gases as well as products (if gases) bear a simple whole number ratio. Avogadro’s hypothesis It states that equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules. Dalton’s Atomic Theory (1803) This theory was based on laws of chemical combinations. It’s basic postulates are : 1. All substances are made up of tiny, indivisible particles, called atoms. 2. In each element, the atoms are all alike and have the same mass. The atoms of different elements differ in mass. 10 Handbook of Chemistry 3. Atoms can neither be created nor destroyed during any physical or chemical change. 4. Compounds or molecules result from combination of atoms in some simple numerical ratio. Limitations (i) It failed to explain how atoms combine to form molecules. (ii) It does not explain the difference in masses, sizes and valencies of the atoms of different elements. Atomic Mass It is the average relative atomic mass of an atom. It indicates that how 1 many times an atom of that element is heavier as compared with th 12 part of the mass of one atom of carbon-12. average mass of an atom Average atomic mass = 1 × mass of an atom of C12 12 The word average has been used in the above definition and is very significant because elements occur in nature as mixture of several isotopes. So, atomic mass can be computed as RA (1) × at. mass (1) + RA (2) × at. mass (2) l Average atomic mass = RA(1) + RA(2) Here, RA is relative abundance of different isotopes. l In case of volatile chlorides, the atomic weight is calculated as At. wt. = Eq. wt. × valency 2 × vapour density of chloride and valency = eq. wt. of metal + 35.5 l According to Dulong and Petit’s rule, Atomic weight × specific heat = 6.4 Gram Atomic Mass (GAM) Atomic mass of an element expressed in gram is called its gram atomic mass or gram-atom or mole-atom.  Basic Concepts of Chemistry 11 Molecular Mass It is the mass of a molecule, i.e. number of times a molecule is heavier 1 than th mass of C-12 atom. Molecular mass of a substance is an 12 additive property and can be calculated by taking algebraic sum of atomic masses of all the atoms of different elements present in one molecule. average relative mass of one molecule Molecular mass = 1 × mass of C-12 atom 12 Gram molecular mass or molar mass is molecular mass of a substance expressed in gram. Molecular mass = 2 × VD (Vapour density) Formula Mass Some substances such as sodium chloride do not contain discrete molecules as their constituent units. The formula such as NaCl is used to calculate the formula mass instead of molecular mass as in the solid state sodium chloride does not exist as a single entity. e.g. formula mass of sodium chloride is 58.5 u. Equivalent Mass It is the mass of an element or a compound which would combine with or displaces (by weight) 1 part of hydrogen or 8 parts of oxygen or 35.5 parts of chlorine. wt. of metal Eq. wt. of metal = × 1.008 wt. of H 2 displaced wt. of metal or = ×8 wt. of oxygen combined wt. of metal or = × 35.5 wt. of chlorine combined wt. of metal Eq. wt. of metal = × 11200 volume of H 2 (in mL) displaced at STP In general, Wt. of substance A Eq. wt. of substance A = Wt. of substance B Eq. wt. of substance B 12 Handbook of Chemistry or for a compound (I) being converted into another compound (II) of same metal, Wt. of compound I Wt. of compound II eq. wt. of metal + eq. wt. of anion of compound I = eq. wt. of metal + eq. wt. of anion of compound II formula mass Eq. mass of a salt = total positive or negative charge atomic mass or molecular mass Equivalent mass = n factor n factor for various compounds can be obtained as (i) n factor for acids i.e. basicity (Number of ionisable H + per molecule is the basicity of acid.) Acid HCl H2SO 4 H 3 PO 3 H 3 PO 4 H2C 2O 4 Basicity 1 2 2 3 2 (ii) n factor for bases, i.e. acidity. (Number of ionisable OH − per molecule is the acidity of a base.) Base NaOH Mg(OH)2 Al(OH) 3 Acidity 1 2 3 (iii) In case of ions, n factor is equal to charge of that ion. (iv) In redox titrations, n factor is equal to change in oxidation number. Cr2O72– + 6e– + 14H + → 2Cr3 + + 2H 2O n factor = 6 MnO–4 + 8H + + 5e− → Mn2+ + 4H 2O n factor = 5 Equivalent mass of organic acid (RCOOH) is calculated by the following formula Eq. wt. of silver salt of acid ( RCOOAg) Wt. of silver salt = Eq. wt. of Ag (or 108) Wt. of silver  Basic Concepts of Chemistry 13 Mole Concept Term mole was suggested by Ostwald (Latin word mole = heap) A mole is defined as the amount of substance which contains same number of elementary particles (atoms, molecules or ions) as the number of atoms present in 12 g of carbon (C-12). 1 mol = 6.023 × 1023 atoms = one gram-atom = gram atomic mass 1 mol = 6.023 × 1023 molecules = gram molecular mass In gaseous state at STP (T = 273 K, p = 1 atm) Gram molecular mass = 1 mol = 22.4 L = 6.022 × 1023 molecules Standard number 6.023 × 1023 is called Avogadro number in honour of Avogadro (he did not give this number) and is denoted by N A. The volume occupied by one mole molecules of a gaseous substance is called molar volume or gram molecular volume. amount of substance (in gram) Number of moles = molar mass Multipli ed lied by Amount of a by Multip Number NA substance Molar mass Mol (6.023 × 10 ) 23 of (in gram) entities Divided by by Divided by D d l ie iv id tip ed u l 22.4 L by M Volume of gas (in L) at STP Number of molecules = number of moles × N A NA Number of molecules in 1g compound = g-molar mass Number of molecules in 1 cm3 (1 mL) of an ideal gas at STP is called Loschmidt number (2.69 × 1019 ). 1 One amu or u (unified mass) is equal to exactly the th of the mass 12 1 of 12C atom, i.e. 1 amu or u = × mass of one carbon (C12 ) atom 12 1 1 amu = = 1 Avogram = 1 Aston NA = 1 Dalton = 1.66 × 10 −24 g One mole of electrons weighs 0.55 mg (5.5 × 10 −4 g). 14 Handbook of Chemistry Empirical and Molecular Formulae Empirical formula is the simplest formula of a compound giving simplest whole number ratio of atoms present in one molecule, e.g. CH is empirical formula of benzene (C6H 6 ). Molecular formula is the actual formula of a compound showing the total number of atoms of constituent elements present in a molecule of compound, e.g. C6H 6 is molecular formula of benzene. Molecular formula = (Empirical formula)n where, n is simple whole number having values 1, 2, 3, …, etc., and can be calculated as molecular formula mass n= empirical formula mass Stoichiometry The relative proportions in which the reactants react and the products are formed, is called stoichiometry (from the Greek word meaning ‘to measure an element’.) Limiting reagent It is the reactant which is completely consumed during the reaction. Excess reagent It is the reactant which is not completely consumed and remains unreacted during the reaction. In a irreversible chemical reaction, the extent of product can be computed on the basis of limiting reagent in the chemical reaction. Per cent Yield The actual yield of a product in any reaction is usually less than the theoretical yield because of the occurrence of certain side reactions. actual yield Per cent yield = × 100 theoretical yield  2 Atomic Structure Atom John Dalton proposed (in 1808) that atom is the smallest indivisible particle of matter. Atomic radii are of the order of 10−8 cm. It contains three subatomic particles namely electrons, protons and neutrons. Electron Electron was discovered as a result of study of cathode rays by JJ Thomson. It was named by Stony. It carries a unit negative charge ( −1.6 × 10−19 C). Mass of electron is 9.11 × 10−31 kg and mass of one mole of electron is 0.55 mg. Some of the characteristics of cathode rays are: (i) These travel in straight line away from cathode and produce fluorescence when strike the glass wall of discharge tube. (ii) These cause mechanical motion in a small pin wheel placed in their path. (iii) These produce X-rays when strike with metal and are deflected by electric and magnetic field. Charge to Mass Ratio of Electron In 1897, British physicist JJ Thomson measured the ratio of electrical charge (e) to the mass of electron ( me ) by using cathode ray tube and applying electrical and magnetic field perpendicular to each other as well as to the path of electrons. Thomson argued that the amount of deviation of the particles from their path in the presence of electrical or magnetic field may vary as follows: (i) If greater the magnitude of the charge on the particles, greater is the deflection. (ii) The mass of the particle, lighter the particle, greater the deflection. 16 Handbook of Chemistry (iii) The deflection of electrons from its original path increase with the increase in the voltage. By this Thomson determined the value e/ me as 1.758820 × 1011 C kg −1. Proton Rutherford discovered proton on the basis of anode ray experiment. It carries a unit positive charge (+1.6 × 10−19 C). The mass of proton is 1.007276 u. e e The ratio of proton is 9.58 × 10−4 C /g. ( ratio is maximum for m m hydrogen gas.) Some of the characteristics of anode rays are: (i) These travel in straight line and possess mass many times heavier than the mass of an electron. (ii) These are not originated from anode but are produced in the space between the anode and the cathode. (iii) These also cause mechanical motion and are deflected by electric and magnetic field.  e (iv) Specific charge   for these rays depends upon the nature of  m the gas taken and is maximum for H 2. Neutron Neutrons are neutral particles. It was discovered by Chadwick (1932). The mass of neutron is 1.675 × 10−24 g or 1.008665 amu or u. 9 4 Be + 42He → 12 6C + 1 0n (α ′ − particles) ( Neutron) Some Other Subatomic Particles (a) Positron Positive electron ( +10 e), discovered by Dirac (1930) and Anderson (1932). (b) Neutrino and antineutrino Particles of small mass and no charge as stated by Fermi (1934). (c) Meson Discovered by Yukawa (1935) and Kemmer. They are unstable particles and include pi ions [π + , π − or π 0]. (d) Anti-proton It is negative proton produced by Segre and Weigland (1955).  Atomic Structure 17 Thomson’s Atomic Model Atom is a positive sphere with a number of electrons distributed within the sphere. It is also known as plum pudding model. It explains the neutrality of an atom. This model could not explain the results of Rutherford scattering experiment. Rutherford’s Nuclear Model of Atom It is based upon α-particle scattering experiment. Rutherford presented that (i) most part of the atom is empty. (ii) atom possesses a highly dense, positively charged centre, called nucleus of the order 10−13 cm. (iii) entire mass of the atom is concentrated inside the nucleus. (iv) electrons revolve around the nucleus in circular orbits. (v) electrons and the nucleus are held together by electrostatic forces of attraction. Drawbacks of Rutherford’s Model (i) According to electromagnetic theory, when charged particles are accelerated, they emit electromagnetic radiations, which comes by electronic motion and thus orbit continue to shrink, so atom is unstable. It doesn’t explain the stability of atom. (ii) It doesn’t say anything about the electronic distribution around nucleus. Atomic Number (Z) Atomic number of an element corresponds to the total number of protons present in the nucleus or total number of electrons present in the neutral atom. Mass Number (A) The mass of the nucleus is due to protons and neutrons, thus they are collectively called nucleons. The total number of nucleons is termed as mass number of the atom. Mass number of an element = number of protons + number of neutrons Representation of an Atom Mass number A Symbol of the element Atomic number Z 18 Handbook of Chemistry Different Types of Atomic Species (a) Isotopes Species with same atomic number but different mass number are called isotopes, e.g. 1H1 , 1H 2. (b) Isobars Species with same mass number but different atomic number are called isobars, e.g. 18 Ar40 , 19K 40. (c) Isotones Species having same number of neutrons are called isotones, e.g. 1H3 and 2He4 are isotones. (d) Isodiaphers Species with same isotopic number are called isodiaphers, e.g. 19K39 , 9F19. Isotopic number = mass number − [2 × atomic number] (e) Isoelectronic Species with same number of electrons are called isoelectronic speices, e.g. Na + , Mg2+. (f) Isosters Species having same number of atoms and same number of electrons, are called isosters, e.g. N 2 and CO. Developments Leading to the Bohr’s Model of Atom Two developments played a major role in the formulation of Bohr’s model: (i) Dual character of the electromagnetic radiation which means that radiation possess wave like and particle like properties. (ii) Atomic spectra explained by electronic energy level in atoms. Electromagnetic Wave Theory (Maxwell) The energy is emitted from source continuously in the form of radiations and magnetic fields. All electromagnetic waves travel with the velocity of light (3 × 108 m/ s) and do not require any medium for their propagation. An electromagnetic wave has the following characteristics: (i) Wavelength It is the distance between two successive crests or troughs of a wave. It is denoted by the Greek letter λ (lambda). (ii) Frequency It represents the number of waves which pass through a given point in one second. It is denoted by ν (nu). (iii) Velocity (v) It is defined as the distance covered in one second by the waves. Velocity of light is 3 × 1010 cms− 1. (iv) Wave number It is the reciprocal of wavelength and has units cm − 1. It is denoted by ν (nu bar). (v) Amplitude (a) It is the height of the crest or depth of the trough of a wave.  Atomic Structure 19 Wavelength ( λ ), frequency ( ν ) and velocity ( v ) of any electromagnetic radiations are related to each other as v = νλ. Electromagnetic wave theory was successful in explaining the properties of light such as interference, diffraction etc., but it could not explain the 1. Black body radiation 2. Photoelectric effect These phenomena could be explained only if electromagnetic waves are supposed to have particle nature. Max Planck provided an explanation for the behaviour of black body and photoelectric effect. Particle Nature of Electromagnetic Radiation : Planck’s Quantum Theory Planck explain the distribution of intensity of the radiation from black body as a function of frequency or wavelength at different temperatures. hc E = hν = (Q c = νλ ) λ where, h = Planck’s constant = 6.63 × 10−34 J-s E = energy of photon or quantum ν = frequency of emitted radiation If n is the number of quanta of a particular frequency and ET be total energy then ET = nhν Black Body Radiation If the substance being heated is a black body, the radiation emitted is called black body radiation. Photoelectric Effect It is the phenomenon in which beam of light of certain frequency falls on the surface of metal and electrons are ejected from it. This phenomenon is known as photoelectric effect. It was first observed by Hertz. W 0 = hν 0 hν 1 mv 2 hc W0 = 2 λ max Metal hν0 [work function] 20 Handbook of Chemistry Threshold frequency ( ν 0 ) = minimum frequency of the radiation Work function (W 0 ) = required minimum energy of the radiation E = KE + W 0 1 ∴ mv 2 = h( ν − ν 0 ) [Kinetic energy of ejected electron = h( ν − ν 0 )] 2 where, ν = frequency of incident radiation ν 0 = threshold frequency Electromagnetic Spectrum The different types of electromagnetic radiations differ only in their wavelengths and hence, frequencies. When these electromagnetic radiations are arranged in order of their increasing wavelengths or decreasing frequencies, the complete spectrum obtained is called electromagnetic spectrum. Different Types of Radiations and Their Sources Type of radiation Wavelength (in Å) Generation source Gamma rays 0.01 to 0.1 Radioactive disintegration X-rays 0.1 to 150 From metal when an electron strikes on it UV-rays 150 to 3800 Sun rays Visible rays 3800 to 7600 Stars, arc lamps Infrared rays 7600 to 6 × 10 6 Incandescent objects Micro waves 6 × 106 to 3 × 109 Klystron tube Radio waves 3 × 1014 From an alternating current of high frequency Electromagnetic spectra may be emission or absorption spectrum on the basis of energy absorbed or emitted. An emission spectrum is obtained when a substance emits radiation after absorbing energy. An absorption spectra is obtained when a substance absorbs certain wavelengths and leave dark spaces in bright continuous spectrum. A spectrum can be further classified into two categories such as (i) Continuous or band spectrum A spectrum in which there is no sharp boundary between two different radiations. (ii) Discontinuous or line spectrum A spectrum in which radiations of a particular wavelength are separated from each other through sharp boundaries.  Atomic Structure 21 Bohr’s Model Neils Bohr proposed his model in 1931. Bohr’s model is applicable only for one electron system like H, He+ , Li2+ etc. Assumptions of Bohr’s model are 1. Electrons keep revolving around the nucleus in certain fixed permissible orbits where it doesn’t gain or lose energy. These orbits are known as stationary orbits. circumference of orbit Number of waves in an orbit = wavelength 2. The electrons can move only in those orbits for which the angular h momentum is an integral multiple of , i.e. 2π nh mvr = ( n = 1, 2, 3..... ) 2π where, m = mass of electron; v = velocity of electron; r = radius of orbit n = number of orbit in which electrons are present 3. Energy is emitted or absorbed only when an electron jumps from higher energy level to lower energy level and vice-versa. hc ∆E = E2 − E1 = hν = λ 4. The most stable state of an atom is its ground state or normal state. From Bohr’s model, energy, velocity and radius of an electron in nth Bohr orbit are (i) Velocity of an electron in nth Bohr orbit Z ( vn ) = 2.165 × 106 m/s n (ii) Radius of nth Bohr orbit n2 n2 (rn ) = 0.53 × 10−10 m = 0.53 Å Z Z Z2 (iii) En = − 2.178 × 10−18 2 J/atom n Z2 = − 1312 2 kJ/ mol n Z2 = − 13.6 2 eV/atom n 22 Handbook of Chemistry  1 1 ∆E = − 2.178 × 10−18  2 − 2  Z 2 J/atom  n1 n 2  where, n = number of shell; Z = atomic number As we go away from the nucleus, the energy levels come closer, i.e. with the increase in the value of n, the difference of energy between successive orbits decreases. Thus, E2 − E1 > E3 − E2 > E4 − E3 > E5 − E4, etc. Emission Spectrum of Hydrogen According to Bohr’s theory, when an electron jumps from ground state to excited state, it emits a radiation of definite frequency (or wavelength). Corresponding to the wavelength of each photon of light emitted, a bright line appears in the spectrum. The number of spectral lines in the spectrum when the electron comes n( n − 1) from nth level to the ground level = 2 Hydrogen spectrum consist of line spectrum. Series Region n1 n2 (i) Lyman UV 1 2, 3, 4, … (ii) Balmer Visible 2 3, 4, 5, … (iii) Paschen IR 3 4, 5, 6, … (iv) Brackett IR 4 5, 6, 7, … (v) Pfund far IR 5 6, 7, … (vi) Humphery far IR 6 7, 8, 9, … Wave number ( ν ) is defined as reciprocal of the wavelength. 1  1 1 ν= ⇒ ν = RZ 2  2 − 2  λ  n1 n 2  where, n1 = 1, 2...... n 2 = n1 + 1, n1 + 2...... Here, λ = wavelength R = Rydberg constant = 109677.8 cm –1 First line of a series is called line of longest wavelength (shortest energy) and last line of a series is the line of shortest wavelength (highest energy, n 2 = ∞).  Atomic Structure 23 Sommerfeld Extension to Bohr’s Model According to this theory, the angular momentum of revolving electron h in an elliptical orbit is an integral multiple of , i.e. 2π kh mvr = 2π nh From Bohr model, mvr = 2π For K shell, n = 1, k = 1 Circular shape L shell, n = 2, k = 1, 2 Circular M shell, n = 3, k = 1, 2, 3 Elliptical N shell, n = 4, k = 1, 2, 3, 4 Elliptical Limitations of Bohr’s Theory (i) It is unable to explain the spectrum of atom other than hydrogen like doublets or multielectron atoms. (ii) It could not explain the ability of atom to form molecules by chemical bonds. Hence, it could not predict the shape of molecules. (iii) It is not in accordance with the Heisenberg uncertainty principle and could not explain the concept of dual character of matter. (iv) It is unable to explain the splitting of spectral lines in the presence of magnetic field (Zeeman effect) and electric field (Stark effect). Towards Quantum Mechanical Model of the Atom Two important developments which contributed significantly in the formulation of such a model were given below 1. de-Broglie Principle (Dual Nature) de-Broglie explains the dual nature of electron, i.e. both particle as well as wave nature. h h λ= or =λ [ p = mv (momentum)] mv p where, λ = wavelength; v = velocity of particle; m = mass of particle h λ= 2m × K E where, KE = kinetic energy. 24 Handbook of Chemistry 2. Heisenberg’s Uncertainty Principle According to this principle, ‘‘it is impossible to specify at any given instant both the momentum and the position of subatomic particles simultaneously like electron.’’ h ∆x ⋅ ∆p ≥ 4π where, ∆x = uncertainty in position; ∆p = uncertainty in momentum Quantum Mechanical Model of Atom It is the branch of chemistry which deals with dual behaviour of matter. It is given by Werner Heisenberg and Erwin Schrodinger. Schrodinger wave equation is ∂ 2ψ ∂ 2ψ ∂ 2ψ 8π 2m + + + (E − U ) ψ = 0 ∂x 2 ∂y 2 ∂z 2 h2 where, x , y , z = cartesian coordinates m = mass of electron, E = total energy of electron U = potential energy of electron, h = Planck’s constant ψ (Psi) = wave function which gives the amplitude of wave ψ 2 = probability function For H-atom, the equation is solved as H$ ψ = Eψ where, H$ is the total energy operator, called Hamiltonian. If the sum of kinetic energy operator (T ) and potential energy operator (U ) is the total energy, E of the system, H = T +U (T + U )ψ = Eψ The atomic orbitals can be represented by the product of two wave functions (i) radial wave function (ii) angular wave function. The orbital wave function, ψ has no significance, but ψ 2 has significance, it measures the electron probability density at a point in an atom. ψ can be positive or negative but ψ 2 is always positive.  Atomic Structure 25 Difference between Orbit and Orbital Orbit Orbital 1. An orbit is a well defined circular path An orbital is the three dimensional space around the nucleus in which the around the nucleus within which the electron revolves. probability of finding an electron is maximum. 2. The maximum number of electrons in The maximum number of electrons present any orbit is given by 2 n2 where n is the in any orbital is two. number of the orbit. Shapes of Atomic Orbitals The shapes of the orbitals are s-spherical, p-dumb bell, d-double-dumb-bell, f-Diffused These orbitals combine to form subshell. (i) s-subshell will have only one spherical orbital. Y Z X (ii) p-subshell has three orbitals ( px , py , pz ). px py pz z z z x x x y y y (iii) d-subshell has five orbitals ( dxy , d yz , dzx , dx 2 − y 2 and dz 2 ). dxy dxz z z x x y y 26 Handbook of Chemistry dyz dx2_ y2 z z y x x y dz2 z y Wave function distribution The orbital wave function ( ψ ) for an electron in an atom has no physical meaning. It is a mathematical function of the coordinates of the electron. ψ ψ (r ) 1s (r ) 2s + + Node r – r Probability Diagrams The graph plotted between ψ 2 and distance from nucleus is called probability diagram. ψ2 ψ2 2s (r) 1s (r) r (nm) Node r (nm) Variation of ψ 2 with distance from the nucleus for 1s and 2s orbitals.  Atomic Structure 27 Node A region or space, where probability of finding an electron is maximum, is called a peak, while zero probability space is called node. Nodes are of two types : (a) Radial nodes (b) Angular nodes (i) ( n − l − 1) = radial node (ii) ( l ) = angular node (iii) ( n − 1) = total nodes Number of Peaks and Nodes for Various Orbitals S. No. Type of orbital Number of peaks Number of nodes 1. s n n −1 2. p n −1 n−2 3. d n−2 n− 3 4. f n− 3 n− 4 Quantum Numbers Each electron in an atom is identified in terms of four quantum numbers. Principal Quantum Number (Neils Bohr) It is denoted by n. It tells us about the main shell in which electron resides. It also gives an idea about the energy of shell and average distance of the electron from the nucleus. Value of n = any integer. Azimuthal Quantum Number (Sommerfeld) It is denoted by l. It tells about the number of subshells ( s, p, d , f ) in any main shell. It also represents the angular momentum of an electron and shapes of subshells. The orbital angular momentum of an h electron = l( l + 1) 2π Value of l = 0 to n − 1. l = 0 for s, l = 2 for d l = 1 for p, l = 3 for f Number of subshells in main energy level = n. Magnetic Quantum Number (Lande) It is denoted by m. It tells about the number of orbitals and orientation of each subshell. Value of m = − l to +l including zero. Number of orbitals in each subshell = ( 2l + 1) 28 Handbook of Chemistry S. No. Subshell Orbital 1. s 1 2. p 3 3. d 5 4. f 7 Number of orbitals in main energy level = n 2. Maximum number of electrons in nth shell = 2n 2 Spin Quantum Number (Ublenbeck and Goldsmith) It is denoted by ms or s. It indicates the direction of spinning of electron, i.e. clockwise or anti-clockwise. Maximum number of electrons in main energy level = 2n 2 Electronic Configuration Arrangement of electrons in various shells, subshells and orbitals in an atom is known as electronic configuration. Filling of Orbitals in Atom Aufbau Principle According to this principle, in the ground state of an atom, the electrons occupy the lowest energy orbitals available to them, i.e. the orbitals are filled in order of increasing value of n + l. For the orbitals having the same value of n + l, the orbtial having lower value of n is filled up first. The general order of increasing energies of the orbital is 1s < 2s < 2 p < 3s < 3 p < 4s < 3d < 4 p < 5s < 4d < 5 p < 6s < 4 f < 5d < 6 p < 7s < 5 f < 6d < 7 p Thus, the filling of electrons in various subshells within the atom can be summerised through following figure. 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 5f 6d 7s 7p The energy of atomic orbitals for H-atom varies as 1s < 2s = 2 p < 3s = 3 p = 3d < 4s = 4 p = 4d = 4 f  Atomic Structure 29 Pauli Exclusion Principle It states, no two electrons in an atom can have identical set of four quantum numbers. The maximum number of electrons in s subshell is 2, p subshell is 6, d subshell is 10 and f subshell is 14. Hund’s Rule of Maximum Multiplicity It states, (i) In an atom no electron pairing takes place in the p, d or f-orbitals until each orbital of the given subshell contains one electron. (ii) The unpaired electrons present in the various orbitals of the same subshell should have parallel spins. Methods of Writing Electronic Configuration (i) Orbital method In this, the electrons present in respective orbitals are denoted. e.g. Cl(17) = 1s2 , 2s2 , 2 p6 , 3s2 , 3 p5. (ii) Shell method In this, the number of electrons in each shell is continuously written. e.g. Cl (17) = 1s2 , 2s2 , 2 p6 , 3s2 , 3 p5 K L M 2, 8, 7 (iii) Box method In this method, each orbital is denoted by a box and electrons are represented by half-headed ( ) or full-headed ( ↑ ) arrows. An orbital can occupy a maximum of two electrons. e.g. Cl(17) = 1s2 2s2 2p6 3s2 3p5 Half-filled and completely filled electronic configurations are more stable. Hence, outer configuration of Cr is 3d5 4s1 and Cu is 3d10 4s1. Electronic Configuration of Ions To write the electronic configuration of ions, first write the electronic configuration of neutral atom and then add (for negative charge) or remove (for positive charge) electrons in outer shell according to the nature and magnitude of charge present on the ion. e.g. O( 8) = 1s2 , 2s2 2 p4, O2− (10) = 1s2 , 2s2 2 p6  3 Classification of Elements and Periodicity in Properties Classification of Elements With the discovery of a large number of elements, it became difficult to study the elements individually, so classification of elements was done to make the study easier. Earlier Attempts to Classify Elements Many attempts were made to classify the known elements from time to time. The earlier attempts are as follows: Prout’s Hypothesis (1815) According to this theory, hydrogen atom was considered as the fundamental unit from which all other atoms were made. It is also known as unitary theory. Dobereiner’s Triads (1829) Dobereiner classified the elements into groups of three elements with similar properties in such a manner so that the atomic weight of the middle element was the arithmetic mean of the other two, e.g. Element Li Na K Atomic weight 7 23 39 7 + 39 Mean of atomic masses = = 23 2 Similarly Cl, Br, I; Ca, Sr, Ba are two more examples of such triads.  Classification of Elements and Periodicity of Properties 31 Limitations Dobereiner could not arrange all the elements known at that time into triads. He could identify only three such triads that have been mentioned. Newland’s Octaves (1864) (Law of Octaves) Newland states that when elements are arranged in order of increasing atomic masses, every eighth element has properties similar to the first just like in the musical note [Every eighth musical note is the same as the first mentioned note]. This can be illustrated as given below sa re ga ma pa dha ni Li Be B C N O F Na Mg Al Si P S Cl Limitations 1. This classification was successful up to the element calcium. 2. When noble gas elements were discovered at a later stage, their inclusion in these octaves disturbed the entire arrangement. Lother Meyer’s Atomic Volume Curve (1869) Meyer presented the classification of elements in the form of a curve between atomic volume and atomic masses and stated that the properties of the elements are the periodic functions of their atomic volumes.  Molecular mass Here, atomic volume =   Density  He concluded that the elements with similar properties occupy similar position in the curve. Mendeleev’s Periodic Table Mendeleev’s periodic table is based upon Mendeleev’s periodic law which states “The physical and chemical properties of the elements are a periodic function of their atomic masses.” At the time of Mendeleev, only 63 elements were known. This periodic table is divided into seven horizontal rows (periods) and eight vertical columns (groups). Zero group was added later on in the modified Mendeleev’s periodic table. 32 Handbook of Chemistry Importance of Mendeleev’s Periodic Table Few important achievements of periodic table are (i) Systematic study of the elements. (ii) Prediction of new elements and their properties, he left space for the elements yet to be discovered, e.g. he left spaces for Ga and Ge and named these elements as EKa-aluminium (Ga) and EKa-silicon (Ge) respectively. (iii) Atomic mass correction of doubtful elements on the basis of their expected positions and properties. Modified Form of Mendeleev’s Periodic Table Group → I II III IV V VI VII VIII 0 Period ↓ A B A B A B A B A B A B A B Zero 1 H He 1.008 4.003 2 Li Be B C N O F Ne 6.94 9.01 10.82 12.01 14.00 16 19 20.183 8 3 Na Mg Al Si P S Cl Ar 22.99 24.32 26.98 28.09 30.975 32.06 35.46 39.944 4 K Ca Sc Ti V Cr Mn Fe Co Ni Kr 39.10 40.08 44.96 47.90 50.95 52.01 54.94 55.85 58.94 58.69 83.80 Cu Zn Ga Ge As Se Br 63.54 65.38 69.72 72.60 74.91 79.91 78.96 5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Xe 85.48 87.63 88.92 91.22 92.91 95.95 99 101.1 102.91 106.7 131.3 Ag Cd In Sn Sb Te I 107.88 112.41 114.76 118.70 121.76 127.61 126.9 6 Cs Ba La Hf Ta W Re Os Ir Pt Rn 132.91 137.36 138.92 178.6 180.92 183.92 186.31 190.2 192.2 195.23 222 Au Hg Tl Pb Bi Po At 197.0 200.61 204.39 207.21 209 210 7 Fr Ra Ac 223 226.05 227 Defects in the Mendeleev’s Periodic Table (i) Position of hydrogen Hydrogen has been placed in group IA (alkali metals), but it also resembles with halogens of group VIIA. Thus, its position in the Mendeleev’s periodic table is controversial. (ii) Position of isotopes As Mendeleev’s classification is based on atomic weight, isotopes would have to be placed in different  Classification of Elements and Periodicity of Properties 33 positions due to their different atomic weights, e.g. 11H , 12H , 31H would occupy different positions. (iii) Anomalous positions of some elements Without any proper justification, in some cases the element with higher atomic mass precedes the element with lower atomic mass. For example, Ar (atomic weight = 39.9) precedes K (atomic weight = 39.1) and similarly Co (atomic weight = 58.9) has been placed ahead of Ni (atomic weight = 58.7). (iv) Position of lanthanoids and actinoids Lanthanoids and actinoids were not placed in the main periodic table. Modern Periodic Table (1913) Moseley modified Mendeleev’s periodic law. He stated “Physical and chemical properties of elements are the periodic function of their atomic numbers.” It is known as modern periodic law and considered as the basis of Modern Periodic Table. When the elements were arranged in increasing order of atomic numbers, it was observed that the properties of elements were repeated after certain regular intervals of 2, 8, 8, 18, 18 and 32. These numbers are called magic numbers and cause of periodicity in properties due to repetition of similar electronic configuration. Structural Features of Long Form of Periodic Table (i) Long form of periodic table is called Bohr’s periodic table. There are 18 groups and seven periods in this periodic table. (ii) The horizontal rows are called periods. First period (1H — 2He) contains 2 elements. It is the shortest period. Second period (3 Li —10Ne) and third period (11Na 18 Ar) contain 8 elements each. These are short periods. Fourth period (19K —36 Kr) and fifth period (37 Rb —54 Xe) contain 18 elements each. These are long periods. Sixth period ( 55 Cs — 86 Rn) consists of 32 elements and is the longest period. Seventh period starting with 87 Fr is incomplete and consists of 19 elements. (iii) The 18 vertical columns are known as groups. Elements of group 1 are called alkali metals. Elements of group 2 are called alkaline earth metals. 34 Handbook of Chemistry Elements of group 16 are called chalcogens [ore forming elements]. Elements of group 17 are called halogens. [sea salt forming elements] Elements of group 18 are called noble gases. Anomalous behaviour of the first element of a group. The first element of a group differs considerably from its congeners (i.e. the rest of the elements of its group). This is due to (i) small size (ii) high electronegativity and (iii) non availability of d-orbitals for bonding. Anomalous behaviour is observed among the second row elements (i.e. Li to F). (iv) The periodic table is divided into four main blocks (s, p, d and f ) depending upon the subshell to which the valence electron enters into. (a) s-block elements Ist and IInd group elements belong to this block and the last electron enters in s-subshell. General electronic configuration = ns1 − 2. (b) p-block elements Group 13th to 18th belong to this block in which last electron enters in p-orbital. Their general electronic configuration is ns2np1 − 6. This is the only block which contains metal, non-metal and metalloids. Examples of metalloids are B, Si, Ge, As, Sb, Te and At. The elements of s-and p-block elements are collectively called representative elements. (c) d-block elements Group 3rd to 12th belong to this block, in which last electron enters in d-orbital. They have inner incomplete shell, so known as transition elements. General electronic configuration is ns1 − 2 ( n − 1)d1 − 10. d-block elements are generally coloured, paramagnetic and exhibit variable valency. (d) f-block elements They constitute two series 4f (lanthanoids) and 5f (actinoids) in which last electron is in 4f and 5f subshell respectively. General electronic configuration  Classification of Elements and Periodicity of Properties 35 ( n − 2) f 1 − 14 ( n − 1) d 0 − 1 ns2 The f-block elements are also called as inner-transition elements. Elements with atomic number greater than 92 (U92 ) are called the transuranic or transuranium elements. All these elements are man-made through artificial nuclear reactions. Very recently, on August 16, 2003, IUPAC approved the name for the element of atomic number 110, as Darmstadtium, with symbol Ds. Limitations of Long Form of Periodic Table In the long form of the periodic table: (i) The position of hydrogen still remains uncertain. (ii) The inner-transition elements do not find a place in the main body of the table. They are placed separately. Predicting the Position of an Element in the Periodic Table First of all write the complete electronic configuration. The principle quantum number of the valence shell represents the period of the element. The subshell in which the last electron is filled corresponds to the block of the element. Group of the element is predicted from the electrons present in the outermost ( n ) or penultimate ( n − 1) shell as follows : For s-block elements, group number = number of ns-electrons (Number of valence electrons) For p-block elements, group number = 10 + number of ns and np electrons For d-block elements, group number = the sum of the number of ( n − 1) d and ns electrons. For f-block elements, group number is always 3. 36 Handbook of Chemistry  Classification of Elements and Periodicity of Properties 37 IUPAC Nomenclature of Elements With Z > 100 The names are derived directly from the atomic numbers using numerical roots for 0 and numbers from 1-9 and adding the suffix ium. Digit 0 1 2 3 4 5 6 7 8 9 Root nil un bi tri quad pent hex sept oct enn Abbreviation n u b t q p h s o e The IUPAC names and symbols of elements with Z > 100 are Z 101 102 103 104 105 106 107 108 109 110 Unnilu Unnilb Unniltr Unnilq Unnilp Unnilh Unnils Unnilo Unnil Ununn IUPAC nium ium ium uadiu entium exium eptium ctium enniu ilium name m m Symbol Unu Unb Unt Unq Unp Unh Uns Uno Une Uun Metals, Non-metals and Metalloids l Metals comprise more than 78% of all known elements and appears on the left side of the periodic table. l In contrast, non-metals are located at the top right handside of the periodic table. l Within the non-metals, some elements show the properties of both metals and non-metals, i.e. metalloids. These elements border the zig-zag line beginning from boron and running diagonally across the p-block. Periodic Properties The properties which are directly or indirectly related to their electronic configuration and show gradual change when we move from left to right in a period or from top to bottom in a group are called periodic properties. Atomic Radius It is the distance from the centre of the nucleus to the outermost shell containing of electrons. It is an hypothetical definition because in a single atom, it is almost impossible to measure this distance. Hence, practically, atomic radius is defined in the following four ways : 38 Handbook of Chemistry Covalent radius If the combining atoms are non-metals (except noble gases) and the bond between them is the single covalent bond then their radius is called the covalent radius. It is measured as the half of their internuclear distance, i.e. For an atom A in a molecule A2. r + rA dA − A rA = A = 2 2 [Distance A − A = Radius of A + Radius of A] For heterodiatomic molecule AB, dA − B = rA + rB + 0.09 ( X A − X B ) Where, X A and X B are electronegativities of A and B. van der Waals’ Radius It is defined as one-half of the distance between the nuclei of two non-bonded isolated atoms or two adjacent atoms belonging to two neighbouring molecules of an element in the solid state. Metallic Radius It is defined as one-half of the internuclear distance between the centres of nuclei of the two adjacent atoms in the metallic crystal. Ionic Radius An atom can be changed to a cation by loss of electrons and to an anion by gain of electrons. A cation is always smaller than the parent atom because during its formation effective nuclear charge increases and sometimes a shell may also decrease. On the other hand, the size of an anion is always larger than the parent atom because during its formation effective nuclear charge decreases. In case of iso-electronic ions, the higher the nuclear charge, smaller is the size, e.g. Al3 + < Mg2+ < Na + < F – < O2– < N3 – The order of radii is : covalent radius < metallic radius < van der Waals’ radius In general, the atomic size decreases on moving from left to right in a period due to increase in effective nuclear charge and increases on moving from top to bottom in a group due to addition of new shells. The concept of effective nuclear charge is discussed below : Effective Nuclear Charge In a multielectron atom, the electron of the inner-shell decrease the force of attraction exerted by the nucleus on the valence electrons. This is called shielding effect. Due to this, the nuclear charge ( Z ) actually present on the nucleus, reduces and is called effective nuclear charge ( Z eff ).  Classification of Elements and Periodicity of Properties 39 It is calculated by using the formula Z eff = Z − σ where, σ = screening constant The magnitude of σ is determined by Slater’s rules. Slater Rules (i) Write the electronic configuration in the following order and groups. (1s) ( 2s, 2 p) ( 3s, 3 p) ( 3d ), ( 4s, 4 p) ( 4d ) ( 4 f ) ( 5s, 5 p) etc. (ii) Electrons of ( n + 1) shell (shell higher than considering electrons) do not contribute in shielding, i.e. σ = 0 (iii) All other electrons in ( ns, np) group contribute σ = 0.35 each. (iv) All electrons of ( n − 1) s and p shell contribute σ = 0.85 each. (v) All electrons of ( n − 2) s and p shell or lower shell contribute σ = 1.00 each (vi) All electrons of nd and nf orbital contribute σ = 0.35 and those of ( n − 1) and f or lower orbital contribute σ = 1.00 each. e.g. Be (4) = 1s2 , 2s2 (for 2s) for 1s σ = 0.35 + 2 × 0.85 = 2.05 Z eff = Z − σ = 4 − 2.05 = 1.95 Ionisation Enthalpy (IE) It is the amount of energy required to remove the loosely bound electron from the isolated gaseous atom. A( g) + IE → A+ ( g) + e− Various factors with which IE varies are : (i) Atomic size : varies inversely (ii) Screening effect : varies inversely (iii) Nuclear charge : varies directly Generally left to right in periods, ionisation enthalpy increases; down the group, it decreases. IE values of inert gases are exceptionally higher due to their stable configurations. Successive ionisation enthalpies IE3 > IE2 > IE1 IE1 of N is exceptionally greater than that of oxygen due to stable half-filled 2 p-orbitals. Among transition elements of 3d-series, 24Cr and 29Cu have higher IE2 due to half-filled and fully-filled stable d-orbitals. 40 Handbook of Chemistry Electron Gain Enthalpy (∆e g ) It is the amount of energy released when an electron is added in an isolated gaseous atom. First electron gain enthalpy is negative while the other successive electron gain enthalpy will be positive due to repulsion between the electrons already present in the anion and the electron being added. O ( g) + e− → O− ( g) ; ∆e g H = − 141 kJ mol−1 O− ( g) + e− → O2− ( g) ; ∆e g H = + 780 kJ mol−1 Various factors with which electron gain enthalpy varies are : (i) Atomic size : varies directly (ii) Nuclear charge : varies directly Along a period, electron gain enthalpy becomes more and more negative while on moving down the group, it becomes less negative. Noble gases have positive electron gain enthalpies. Halogens have maximum value of ∆e g H within a period due to smallest atomic size. F and O atom have small size and high charge density, therefore they have lower values of electron gain enthalpy, than Cl and S respectively. Cl > F; S > O Elements having half-filled and fully-filled orbitals exhibit more stability, therefore, electron gain enthalpy will be low for such elements. Electron gain enthalpy can be measured by Born-Haber cycle and elements with high ∆e g H , are good oxidising agent. Electronegativity (EN) It is defined as the tendency of an atom to attract the shared electron pair towards itself in a polar covalent bond. Various factors with which electronegativity varies are : (i) Atomic size : varies inversely (ii) Charge on the ion : varies directly, e.g. Li < Li+ , Fe2+ < Fe3 + (iii) Hybridisation : (Electronegativity ∝ % age s-character in the hybrid orbital) Electronegativity of carbon atom = C2H 6 < C2H 4 < C2H 2 In periods as we move from left to right electronegativity increases, while in the groups electronegativity decreases down the group. For noble gases, its value is taken as zero.  Classification of Elements and Periodicity of Properties 41 Electronegativity helps to predict the polarity of bonds and dipole moment of molecules. Electronegativity order of some elements (on Pauling scale) is F > O > N ≈ Cl > Br (4.0) (3.5) (3.0) (3.0) (2.8) (i) Mulliken scale IE + ∆e g H Electronegativity ( x ) = 2 (ii) Pauling scale The difference in electronegativity of two atoms A and B is given by the relationship xB − xA = 0.208 ∆ where, ∆ = EA − B − EA − A × EB − B (∆ is known as resonance energy.) EA − B , EA − A and EB − B represent bond dissociation energies of the bonds A − B, A − A and B − B respectively. (iii) Allred and Rochow’s scale 0.359 Z eff Electronegativity = 0.744 + r2 Where, Z eff is the effective nuclear charge = Z − σ Where, σ is screening constant. It’s value can be determined by Slater’s rule. Valency It is defined as the combining capacity of the element. The valency of an element is related to the electronic configuration of its atom and usually determined by electrons present in the valence shell. On moving along a period from left to right, valency increases from 1 to 4 and then decreases to zero (for noble gases) while on moving down a group the valency remains the same. Transition metals exhibit variable valency because they can use electron from outer as well as penultimate shell. Chemical Reactivity Reactivity of metal increases with decrease in IE, electronegativity and increase in atomic size as well as electropositive character. Reactivity of non-metals increases with increase in electronegativity as well as electron gain enthalpy and decrease in atomic radii. 42 Handbook of Chemistry Melting and Boiling Points On moving down the group, the melting point and boiling point for metallic elements go on decreasing due to the decreasing forces of attraction. However, for non-metals, melting point and boiling point generally increase down the group. Along a period from left to right, melting point and boiling point increases and reaches a maximum value in the middle of the period and then start decreasing. Tungsten (W) has highest melting point (3683 K) among metals, carbon (diamond) has the highest melting point among non-metals. Helium has lowest melting point ( −270° C) among all elements, Electropositivity or Metallic Character The tendency of an atom of the element to lose valence electrons and form positive ion is called electropositivity. Greater the electropositive character, greater is the metallic character. Electropositive character decreases on moving across the period and increases on moving down the group. Alkali metals are the most electropositive and halogens are the least electropositive element in their respective period. Basic nature of oxides ∝ metallic character, i.e. it also decreases along a period and increases down the group. Density Li metal has minimum density while osmium (Os) metal has maximum density. Diagonal Relationship Certain elements of 2nd period show similarity in properties with their diagonal elements in the 3rd period as shown below : Group 1 Group 2 Group 13 Group 14 2nd period Li Be B C 3rd period Na Mg Al Si Thus, Li resembles Mg, Be resembles Al and B resembles Si. This is called diagonal relationship and this is due to the reason that these pairs of elements have almost identical ionic radii and polarizing power (i.e. charge/size ratio). Elements of third period, i.e. Mg, Al and Si are known as bridge elements.  4 Chemical Bonding and Molecular Structure Chemical Bond It is defined as the attractive force which hold the various chemical constituents (atoms, ions, etc.) together in different chemical species. Bond forms to get the stability, with a release of energy. Kossel-Lewis Approach to Chemical Bonding According to this theory, atoms take part in the bond formation to complete their octet or to acquire the electronic configuration