Chemistry 9.pdf

CHEMISTRY 9 Publisher: CARAVAN BOOK HOUSE, LAHORE  All rights (Copy right etc.) are reserved with the publisher. Approved by the Federal Ministry of Education (Curriculum Wing), Islamabad, according to the National Curriculum 2006 under the National Tex...

CHEMISTRY 9 Publisher: CARAVAN BOOK HOUSE, LAHORE  All rights (Copy right etc.) are reserved with the publisher. Approved by the Federal Ministry of Education (Curriculum Wing), Islamabad, according to the National Curriculum 2006 under the National Textbook and Learning Materials Policy 2007. N.O.C. F.2-2/2010-Chem. Dated 2-12-2010. This book has also been published by Punjab Textbook Board under a print licence arrangement for free distribution in all Government School in Punjab. No part of this book can be copied in any form especially guides, help books etc., without the written permission of the publisher. CONTENTS Unit 1 Fundamentals of Chemistry 1 Unit 2 Structure of Atoms 27 Unit 3 Periodic Table and Periodicity of Properties 44 Unit 4 Structure of Molecules 58 Unit 5 Physical States of Matter 75 Unit 6 Solutions 96 Unit 7 Electrochemistry 113 Unit 8 Chemical Reactivity 138 Authors: Dr. Jaleel Tariq Dr. Irshad Ahmad Chatha Designed by: Sakhawat Hussain Prepared by: The Caravan Book House Kachehri Road, Lahore  Chapter1 Fundamentals of Chemistry Major Concepts 1.1 Branches of Chemistry Time allocation 1.2 Basic Definitions Teaching periods 12 1.3 Chemical species Assessment periods 03 1.4 Avogadro' s Number and Mole Weightage 10% 1.5 Chemical Calculations Students Learning Outcomes Students will be able to: Identify and provide examples of different branches of chemistry. Differentiate among branches of chemistry. Distinguish between matter and a substance. Define ions, molecular ions, formula units and free radicals. Define atomic number, atomic mass, atomic mass unit. Differentiate among elements, compounds and mixtures. Define relative atomic mass based on C-12 scale. Differentiate between empirical and molecular formula. Distinguish between atoms and ions. Differentiate between molecules and molecular ions. Distinguish between ion and free radicals. Classify the chemical species from given examples. Identify the representative particles of elements and compounds. Relate gram atomic mass, gram molecular mass and gram formula mass to mole. Describe how Avogadro's number is related to a mole of any substance. Distinguish among the terms gram atomic mass, gram molecular mass and gram formula mass. Change atomic mass, molecular mass and formula mass into gram atomic mass, gram molecular mass and gram formula mass. Chemistry - IX 2 Unit 1: Fundamentals of Chemistry Introduction The knowledge that provides understanding of this world and how it works, is science. The branch of science which deals with the composition, structure, properties and reactions of matter is called chemistry. It deals with every aspect of our life. The development of science and technology has provided us a lot of facilities in daily life. Imagine the role and importance of petrochemical products, medicines and drugs, soap, detergents, paper, plastics, paints and pigments, insecticides, pesticides which all are fruit of the efforts of chemists. The development of chemical industry has also generated toxic wastes, contaminated water and polluted air around us. On the other hand, chemistry also provides knowledge and techniques to improve our health and environment and to explore and to conserve the natural resources. In this chapter, we will study about different branches of chemistry, basic definitions and concepts of chemistry. 1.1 BRANCHES OF CHEMISTRY It is a fact that we live in the world of chemicals. We all depend upon different living organisms which require water, oxygen or carbon dioxide for their survival. Today chemistry has a wide scope in all aspects of life and is serving the humanity day and night. Chemistry is divided into following main branches: physical chemistry, organic chemistry, inorganic chemistry, biochemistry, industrial chemistry, nuclear chemistry, environmental chemistry and analytical chemistry. 1.1.1 Physical Chemistry Physical Chemistry is defined as the branch of chemistry that deals with the relationship between the composition and physical properties of matter along with the changes in them. The properties such as structure of atoms or formation of molecules behavior of gases, liquids and solids and the study of the effect of temperature or radiation on matter are studied under this branch. 1.1.2 Organic Chemistry Organic Chemistry is the study of covalent compounds of carbon and hydrogen (hydrocarbons) and their derivatives. Organic compounds occur naturally and are also synthesized in the laboratories. Organic chemists determine the structure and properties of these naturally occurring as well as synthesized compounds. Scope of this branch covers petroleum, petrochemicals and pharmaceutical industries. 1.1.3 Inorganic Chemistry Inorganic chemistry deals with the study of all elements and their compounds except those of compounds of carbon and hydrogen (hydrocarbons) and their derivatives. It has applications in every aspect of the chemical industry such as glass, cement, ceramics and metallurgy (extraction of metals from ores). Chemistry - IX 3 Unit 1: Fundamentals of Chemistry 1.1.4 Biochemistry It is the branch of chemistry in which we study the structure, composition, and chemical reactions of substances found in living organisms. It covers all chemical processes taking place in living organisms, such as synthesis and metabolism of biomolecules like carbohydrates, proteins and fats. Biochemistry emerged as a separate discipline when scientists began to study how living things obtain energy from food or how the fundamental biological changes occur during a disease. Examples of applications of biochemistry are in the fields of medicine, food science and agriculture, etc. 1.1.5 Industrial Chemistry The branch of chemistry that deals with the manufacturing of chemical compounds on commercial scale, is called industrial chemistry. It deals with the manufacturing of basic chemicals such as oxygen, chlorine, ammonia, caustic soda, nitric acid and sulphuric acid. These chemicals provide the raw materials for many other industries such as fertilizers, soap, textiles, agricultural products, paints and paper, etc. 1.1.6 Nuclear Chemistry Nuclear Chemistry is the branch of chemistry that deals with the radioactivity, nuclear processes and properties. The main concern of this branch is with the atomic energy and its uses in daily life. It also includes the study of the chemical effects resulting from the absorption of radiation within animals, plants and other materials. It has vast applications in medical treatment (radiotherapy), preservation of food and generation of electrical power through nuclear reactors, etc. 1.1.7 Environmental Chemistry It is the branch of chemistry in which we study about components of the environment and the effects of human activities on the environment. Environmental chemistry is related to other branches like biology, geology, ecology, soil and water. The knowledge of chemical processes taking place in environment is necessary for its improvement and protection against pollution. 1.1.8 Analytical Chemistry Analytical chemistry is the branch of chemistry that deals with separation and analysis of a sample to identify its components. The separation is carried out prior to qualitative and quantitative analysis. Qualitative analysis provides the identity of a substance (composition of chemical species). On the other hand, quantitative analysis determines the amount of each component present in the sample. Hence, in this branch different techniques and instruments used for analysis are studied. The scope of this branch covers food, water, environmental and clinical analysis. Chemistry - IX 4 Unit 1: Fundamentals of Chemistry i. In which branch of chemistry behaviour of gases and liquids is studied? ii. Define biochemistry? iii. Which branch of chemistry deals with preparation of paints and paper? iv. In which branch of chemistry the metabolic processes of carbohydrates and proteins are studied? Test yourself v. Which branch of chemistry deals with atomic energy and its uses 1.1 in daily life? vi. Which branch of chemistry deals with the structure and properties of naturally occurring molecules? 1.2 BASIC DEFINITIONS Matter is simply defined as anything that has mass and occupies space. Our bodies as well as all the things around us are examples of matter. In chemistry, we study all types of matters that can exist in any of three physical states: solid, liquid or gas. A piece of matter in pure form is termed as a substance. Every substance has a fixed composition and specific properties or characteristics. Whereas, impure matter is called a mixture; which can be homogeneous or heterogeneous in its composition. We know that every substance has physical as well as chemical properties. The properties those are associated with the physical state of the substance are called physical properties like colour, smell, taste, hardness, shape of crystal, solubility, melting or boiling points, etc. For example, when ice is heated, it melts to form water. When water is further heated, it boils to give steam. In this entire process only the physical states of water change whereas its chemical composition remains the same. The chemical properties depend upon the composition of the substance. When a substance undergoes a chemical change, its composition changes and a new substances are formed. For example, decomposition of water is a chemical change as it produces hydrogen and oxygen gases. All materials are either a substance or a mixture. Figure 1.1 shows simple classification of the matter into different forms. MATTER Mixture Substance Homogeneous Heterogeneous Hetrogeneous Elements Compounds mixture mixture Fig. 1.1. Classification of matter Chemistry - IX 5 Unit 1: Fundamentals of Chemistry 1.2.1 Elements, Compounds and Mixtures 1.2.1.1 Elements In the early ages, only nine elements (carbon, gold, silver, tin, mercury, lead, copper, iron and sulphur) were known. At that time, it was considered that elements were the substances that could not be broken down into simpler units by ordinary chemical processes. Until the end of nineteenth century, sixty-three elements had been discovered. Now 118 elements have been discovered, out of which 92 are naturally occurring elements. Modern definition of element is that it is a substance made up of same type of atoms, having same atomic number and cannot be decomposed into simple substances by ordinary chemical means. It means that each element is made up of unique type of atoms that have very specific properties. Elements occur in nature in free or combined form. All the naturally occurring elements found in the world have different percentages in the earth's crust, oceans and atmosphere. Table 1.1. shows natural occurrence in percentage by weight of some major elements around us. It shows concentrations of these major elements found in the three main systems of our environment. Table 1.1 Natural Occurrences by Weight % of Some Major Elements ‘s crust Elements may be solids, liquids or gases. Majority of the elements exist as solids e.g. sodium, copper, zinc, gold, etc. There are very few elements which occur in liquid state e.g. mercury and bromine. A few elements exist as gases e.g. nitrogen, oxygen, chlorine and hydrogen. On the basis of their properties, elements are divided into metals, non-metals and metalloids. About 80 percent of the elements are metals. Chemistry - IX 6 Unit 1: Fundamentals of Chemistry Major part of a living body is made up of water i.e. 65% to 80% by mass. Six elements constitute about 99% of our body mass; namely: Oxygen 65 %, Carbon 18%, Hydrogen 10 %, Nitrogen 3%, Calcium 1.5% and Phosphorus 1.5%. Potassium, Sulphur, Magnesium and Sodium constitute 0.8% of Do you know? our body mass. Whereas Copper, Zinc, Fluorine, Chlorine, Iron, Cobalt and Manganese constitute only 0.2% of our body mass. Elements are represented by symbols, which are abbreviations for the name of elements. A symbol is taken from the name of that element in English, Latin, Greek or German. If it is one letter, it will be capital as H for Hydrogen, N for Nitrogen and C for Carbon etc. In case of two letters symbol, only first letter is capital e.g. Ca for Calcium, Na for Sodium and Cl for Chlorine. The unique property of an element is valency. It is combining capacity of an element with other elements. It depends upon the number of electrons in the outermost shell. In simple covalent compounds, valency is the number of hydrogen atoms which combine with one atom of that element or the number of bonds formed by one atom of that element e.g. in the following compounds. The valency of chlorine, oxygen, nitrogen and carbon is 1, 2, 3 and 4, respectively. In simple ionic compounds valency is the number of electrons gained or lost by an atom of an element to complete its octet. Elements having less than four electrons in their valence shell; prefer to lose the electrons to complete their octet. For example, atoms of Na, Mg and Al have 1, 2 and 3 electrons in their valence shells respectively. They lose these electrons to have valency of 1, 2 and 3, respectively. On the other hand, elements having five or more than five electrons in their valence shells, gain electrons to complete their octet. For example, N, O and Cl have 5, 6 and 7 electrons in their valence shells respectively. They gain 3, 2 and 1 electrons respectively to complete their octet. Hence, they show valency of 3, 2 and 1, respectively. A radical is a group of atoms that have some charge. Valencies of some common elements and radicals are shown in Table 1.2. Chemistry - IX 7 Unit 1: Fundamentals of Chemistry Table 1.2 Some Elements and Radicals with their Symbols and Common Valencies Some elements show more than one valency, i.e. they have variable valency. For example, in ferrous sulphate (FeSO4) the valency of iron is 2. In ferric sulphate (Fe2(SO4)3), the valency of iron is 3. Generally, the Latin or Greek name for the element (e.g., Ferrum) is modified to end in 'ous' for the lower valency (e.g. Ferrous) and to end in 'ic' for the higher valency (e.g. Ferric). 1.2.1.2 Compound Compound is a substance made up of two or more elements chemically combined together in a fixed ratio by mass. As a result of this combination, elements lose their own properties and produce new substances (compounds) that have entirely different properties. Compounds can't be broken down into its constituent elements by simple physical methods. For example, carbon dioxide is formed when elements of carbon and oxygen combine chemically in a fixed ratio of 12:32 or 3:8 by mass. Similarly, water is a compound formed by a chemical combination between hydrogen and oxygen in a fixed ratio of 1:8 by mass. Chemistry - IX 8 Unit 1: Fundamentals of Chemistry Compounds can be classified as ionic or covalent. Ionic compounds do not exist in independent molecular form. They form a three dimensional crystal lattice, in which each ion is surrounded by oppositely charged ions. These oppositely charged ions attract each other very strongly, as a result ionic compounds have high melting and boiling points. These compounds are represented by formula units e.g. NaCl, KBr, CuSO4. The covalent compounds mostly exist in molecular form. A molecule is a true representative of the covalent compound and its formula is called molecular formula e.g. H2O, HC1, H2SO4, Ch4. Table 1.3 Some Common Compounds with their Formulae Remember Always use: Standard symbols of elements Chemical formulae of compounds Proper abbreviations of scientific terms Standard values and SI units for constants 1.2.1.3 Mixture When two or more elements or compounds mix up physically without any fixed ratio, they form a mixture. On mixing up, the component substances retain their own chemical identities and properties. The mixture can be separated into parent components by physical methods such as distillation, filtration, evaporation, crystallisation or magnetization. Mixtures that have uniform composition throughout are called homogeneous mixtures e.g. air, gasoline, ice cream. Whereas, heterogeneous mixtures are those in which composition is not uniform throughout e.g. soil, rock and wood. Chemistry - IX 9 Unit 1: Fundamentals of Chemistry Air is a mixture of nitrogen, oxygen, carbon dioxide, noble gases and water vapours. Soil is a mixture of sand, clay, mineral salts, water and air. Milk is a mixture of water, sugar, fat, proteins, mineral salts and vitamins. Do you know? Brass is a mixture of copper and zinc metals. Table 1.4 Difference between a Compound and a Mixture Compound Mixture i It is formed by a chemical combination of Mixture is formed by the simple mixing atoms of the elements. up of the substances. ii. The constituents lose their identity and Mixture shows the properties of the form a new substance having entirely constituents. different properties from them. iii. Compounds always have fixed Mixtures do not have fixed composition. composition by mass. iv. The components cannot be separated by The components can be separated by physical means. simple physical methods. v. Every compound is represented by a It consists of two or more components chemical formula. and does not have any chemical formula. vi. Compounds have homogeneous They may be homogeneous or composition. heterogeneous in composition vii. Compounds have sharp and fixed melting Mixtures do not have sharp and fixed points melting points. i. Can you identify mixture, element or compound out of the following: Coca cola, petroleum, sugar, table salt, blood, gun powder, urine, aluminium, silicon, tin, lime and ice cream. ii. How can you justify that air is a homogenous mixture. Identify substances present in it. iii. Name the elements represented by the following symbols: Hg, Au, Fe, Ni, Co, W, Sn, Na, Ba, Br, Bi. Test yourself iv. Name a solid, a liquid and a gaseous element that exists at the 1.2 room temperature. v. Which elements do the following compounds contain? Sugar, common salt, lime water and chalk. 1.2.1 Atomic Number and Mass Number The atomic number of an element is equal to the number of protons present in the nucleus of its atoms. It is represented by symbol ‘Z’. As all atoms of an element have the same number of protons in their nuclei, they have the same atomic number. Chemistry - IX 10 Unit 1: Fundamentals of Chemistry Hence, each element has a specific atomic number termed as its identification number. For example, all hydrogen atoms have 1 proton, their atomic number is Z=l. All atoms in carbon have 6 protons, their atomic number is Z=6. Similarly, in oxygen all atoms have 8 protons having atomic number Z=8 and sulphur having 16 protons shows atomic number Z = 16. The mass number is the sum of number of protons and neutrons present in the nucleus of an atom. It is represented by symbol 'A'. It is calculated as A=Z+n where n is the number of neutrons. Each proton and neutron has lamu mass. For example, hydrogen atom has one proton and no neutron in its nucleus, its mass number A=l+0 =1. Carbon atom has 6 protons and 6 neutrons, hence its mass number A=12. Atomic numbers and mass numbers of a few elements are given in Table 1.5 Table 1.5 Some Elements along with their Atomic and Mass Numbers Example 1.1 How many protons and neutrons are there in an atom having A = 238 and Z = 92. Solution: First of all, develop data from the given statement of the example and then solve it with the help of data. Data A=238 Z=92 Number of protons ? Number of neutrons? Number of protons = Z = 92 Chemistry - IX 11 Unit 1: Fundamentals of Chemistry Number of Neutrons =A- Z = 238 – 92 = 146 1.2.3 Relative Atomic Mass and Atomic Mass Unit As we know that the mass of an atom is too small to be determined practically. However, certain instruments enable us to determine the ratio of the atomic masses of various elements to that of carbon-12 atoms. This ratio is known as the relative atomic mass of the element. The relative atomic mass of an element is the average mass of the atoms of that element as compared to 1/12th (one-twelfth) the mass of an atom of carbon- 12 isotope (an element having different mass number but same atomic number). Based on carbon-12 standard, the mass of an atom of carbon is 12 units and l/2th of it comes to be 1 unit. When we compare atomic masses of other elements with atomic mass of carbon- 12 atom, they are expressed as relative atomic masses of those elements. The unit for relative atomic masses is called atomic mass unit, with symbol 'amu'. One atomic mass unit is 1/12th the mass of one atom of carbon-12th. When this atomic mass unit is expressed in grams, it is: For example: i) How many amu 1 g of a substance has? ii) Is atomic mass unit a SI unit of an atomic mass? iii) What is the relationship between atomic number and atomic mass? iv) Define relative atomic mass. Test yourself v) Why atomic mass of an atom is defined as relative atomic mass? 1.3 1.2.4 How to write a Chemical Formula It represents the name of the Compounds are represented by chemical substance e.g. formulae as elements are represented by H2O (water). It also It tells the represents symbols. Chemical formulae of compounds are name of the one mole of elements the molecules written keeping the following steps in as present Significance in the balanced consideration. in the of chemical compound. chemical equation. i. Symbols of two elements are written formula side by side, in the order of positive ion It indicates It is in fact the mass of the one molecule or first and negative ion later. compound in amus formula unit of or grams. the compound. ii. The valency of each ion is written on the right top corner of its symbol, e.g. Na+, Ca2+, CI and O2. Chemistry - IX 12 Unit 1: Fundamentals of Chemistry iii. This valency of each ion is brought to the lower right corner of other ion by 'cross- exchange' method, e.g. They are written as: iv. If the valencies are same, they are offset and are not written in the chemical formula. But if they are different, they are indicated as such at the same position, e.g. in case of sodium chloride both the valencies are offset and formula is written as NaCl, whereas, calcium chloride is represented by formula CaCl2. v. If an ion is a combination of two or more atoms which is called radical, bearing a net charge on it, e.g. SO42 (sulphate) and PO43 (phosphate), then the net charge represents the valency of the radical. The chemical formula of such compounds is written as explained in (iii) and (iv); writing the negative radical within the parenthesis. For example, chemical formula of aluminium sulphate is written as Al2(SO4)3 and that of calcium phosphate as Ca3(PO4)2. 1.2.4.1 Empirical formula Chemical formulae are of two types. The simplest type of formula is empirical formula. It is the simplest whole number ratio of atoms present in a compound. The empirical formula of a compound is determined by knowing the percentage composition of a compound. However, here we will explain it with simple examples. The covalent compound silica (sand) has simplest ratio of 1:2 of silicon and oxygen respectively. Therefore, its empirical formula is SiO2. Similarly, glucose has simplest ratio 1:2:1 of carbon, hydrogen and oxygen, respectively. Hence, its empirical formula is CH2O. As discussed earlier, the ionic compounds exist in three dimensional network forms. Each ion is surrounded by oppositely charged ions in such a way to form electrically neutral compound. Therefore, the simplest unit taken as a representative of an ionic compound is called formula unit. It is defined as the simplest whole number ratio of ions, as present in the ionic compound. In other words, ionic compounds have only empirical formulae. For example, formula unit of common salt consists of one Na+ and one CI ion and its empirical formula is NaCl. Similarly, formula unit of potassium bromide is KBr, which is also its empirical formula. 1.2.4.2 Molecular Formula Chemistry - IX 13 Unit 1: Fundamentals of Chemistry Molecules are formed by the combination of atoms. These molecules are represented by molecular formulae that show actual number of atoms of each element present in a molecule of that compound. Molecular formula is derived from empirical formula by the following relationship: Molecular formula = (Empirical formula)n Where n is 1,2,3 and so on. For example, molecular formula of benzene is C6H6 which is derived from the empirical formula CH where the value of n is 6. The molecular formula of a compound may be same or a multiple of the empirical formula. A few compounds having different empirical and molecular formulae are shown in Table 1.6. Table 1.6 Some Compounds with their Empirical and Molecular Formulae Some compounds may have same empirical and molecular formula e.g. water (H20), hydrochloric acid (HC1), etc. 1.2.5 Molecular Mass and Formula Mass The sum of atomic masses of all the atoms present in one molecule of a molecular substance, is its molecular mass. For example, molecular mass of chlorine (Cl2) is 71.0 amu, of water (H2O) is 18 amu and that of carbon oxide (CO2) is 44 amu. Example 1.2 Calculate the molecular mass of Nitric acid, HNO3. Solution Atomic mass of H = 1 amu Atomic mass of N = 14 amu Atomic mass of O = 16 amu Molecular formula = HNO3 Molecular mass = 1 (At. mass of H) + 1 (At. mass of N) + 3 (At. mass of O) = 1 + 14 + 3(16) = 1 + 14 + 48 = 63 amu Some ionic compounds that form three dimensional solid crystals, are represented by their formula units. Formula mass in such cases is the sum of atomic masses of all the atoms present in one formula unit of a substance. For example, formula mass of sodium chloride is 58.5 amu and that of CaCO3 is 100 amu. Chemistry - IX 14 Unit 1: Fundamentals of Chemistry Example 1.3 Calculate the formula mass of Potassium sulphate K2SO4 Solution Atomic mass of K = 39 amu Atomic mass of S = 32 amu Atomic mass of O = 16 amu Formula unit = K2SO4 Formula mass of K2SO4 = 2(39) + 1(32) + 4(16) = 78 + 32 + 64 = 174 amu i.What is the relationship between empirical formula and formula unit? ii. How can you differentiate between molecular formula and empirical formula? iii. Identify the following formulae as formulas or unit molecular formulae: H2O2, CH4, C6H12O6, C12H22O1, BaCO3, KBr iv. What is empirical formula of acetic acid (CH3COOH)? Test yourself Find out its molecular mass. 1.3 CHEMICAL 1.4 v. SPECIES Calculate the formula masses of: Na2S04, ZnSO4 and CuCO3. 1.3.1 Ions (Cations and Anions), Molecular Ions and Free Radicals Ion is an atom or group of atoms having a charge on it. The charge may be positive or negative. There are two types of ions i.e. cations and anions. An atom or group of atoms having positive charge on it is called cation. The cations are formed when atoms lose electrons from their outermost shells. For example, Na+, K+ are cations. The following equations show the formation of cations from atoms. An atom or a group of atoms that has a negative charge on it, is called anion. Anion is formed by the gain or addition of electrons to an atom. For example, Cl and O2. Following examples show the formation of an anion by addition of electrons to an atom. Table 1.7 Difference between Atoms and Ions Chemistry - IX 15 Unit 1: Fundamentals of Chemistry 1.3.1.1 Molecular Ion When a molecule loses or gains an electron, it forms a molecular ion. Hence, molecular ion or radical is a species having positive or negative charge on it. Like other ions they can be cationic molecular ions (if they carry positive charge) or anionic molecular ions (if they carry negative charge). Cationic molecular ions are more abundant than anionic molecular ions. For example, CH4+, He+, N2+. When gases are bombarded with high energy electrons in a discharge tube, they ionize to give molecular ions. Table 1.8 shows some differences between molecule and molecular ion. Table 1.8 Difference between Molecule and Molecular Ion 1.3.1.2 Free Radicals Free radicals are atoms or group of atoms possessing odd number of (unpaired) electrons. It is represented by putting a dot over the symbol of an element e.g. H, CI, H3C. Free radicals are generated by the homolytic (equal) breakage of the bond between two atoms when they absorb heat or light energy. A free radical is extremely reactive species as it has the tendency to complete its octet. Table 1.9 shows some of the differences between ions and free radicals. Chemistry - IX 16 Unit 1: Fundamentals of Chemistry Most of the universe exists in the form of plasma, the fourth state of matter. Both the cationic and anionic molecular ions are present Do you know? in it. sunlight Table 1.9 Differencesunlight between Ions and Free Radicals 1.3.2 Types of Molecules A molecule is formed by the chemical combinations of atoms. It is the smallest unit of a substance. It shows all the properties of the substance and can exist independently. There are different types of molecules depending upon the number and types of atoms combining. A few types are discussed here. A molecule consisting of only one atom is called monoatomic molecule. For example, the inert gases helium, neon and argon all exist independently in atomic form and they are called monoatomic molecules. If a molecule consists of two atoms, it is called diatomic molecule. For example: hydrogen (H2), oxygen (O2), chlorine (Cl2) and hydrogen chloride (HCl). If it consists of three atoms, it is called triatomic molecule. For example :H2O and CO2. If a molecule consists of many atoms, it is called polyatomic. For example: methane (CH4), sulphuric acid (H2SO4) and glucose (C6H12O6). A Molecule containing same type of atoms, is called homoatomic molecule. For example: hydrogen (H2), ozone (O3), sulphur (S8) and phosphorus (P4) are the examples of molecules formed by the same type of atoms. When a molecule consists of different kinds of atoms, it is called heteroatomic molecule. For example: CO2, H2O and NH3. Chemistry - IX 18 Unit 1: Fundamentals of Chemistry mass of a substance. Avogadro's Number is a collection of 6.02  1023 particles. It is represented by symbol 'NA'. Hence, the 6.02  1023 number of atoms, molecules or formula units is called Avogadro's number that is equivalent to one 'mole' of respective substance. In simple words, 6.02  1023 particles are equal to one mole as twelve eggs are equal to one dozen. To understand the relationship between the Avogadro's number and the mole of a substance let us consider a few examples. i. 6.02  1023 atoms of carbon are equivalent to one mole of carbon. ii. 6.02  1023 molecules of H2O are equivalent to one mole of water. iii. 6.02  1023 formula units of NaCl are equivalent to one mole of sodium chloride. Amaedo Avogadro (1776-1856) was an Thus, 6.02  1023 atoms of elements or 6.02  1023 molecules Italian scholar. He is of molecular substance or 6.02  1023 formula units of ionic famous for molecular compounds are equivalent to 1 mole. theory commonly known as Avogadro's For further explanation about number of atoms in molecular law. In tribute to him, compounds or number of ions in ionic compounds let us discuss the number of particles two examples: (atoms, molecules, ions) in mole of a i. One molecule of water is made up of 2 atoms of hydrogen substance 6.02  1023 is and 1 atom of oxygen, hence 2  6.02  1023 atoms of k n o w n a s t h e hydrogen and 6.02  1023 atoms of oxygen constitute one Avogadro's constant. mole of water. ii. One formula unit of sodium chloride consists of one sodium ion and one chloride ion. So there are 6.02  1023 number of Na ions and 6.02  1023 CI ions in one mole of sodium chloride. Thus, the total number of ions in 1 mole of NaCl is 12.04l023 or 1.204  1024. 1.5.2 Mole (Chemist secret unit) A mole is defined as the amount(mass) of a substance that contains 6.02  l023 number of particles (atoms, molecules or formula units). It establishes a link between mass of a substance and number of particles as shown in summary of molar calculations. It is abbreviated as 'mol'. You know that a substance may be an element or compound (molecular or ionic). Mass of a substance is either one of the following: atomic mass, molecular mass or formula mass. These masses are expressed in atomic mass units (amu). But when these masses are expressed in grams, they are called as molar masses. Scientists have agreed that Avogadro's number of particles are present in one molar mass of a substance. Thus, quantitative definition of mole is the atomic mass, molecular mass or formula mass of a substance expressed in grams is called mole. Chemistry - IX 19 Unit 1: Fundamentals of Chemistry For example: Atomic mass of carbon expressed as 12 g = 1 mol of carbon Molecular mass of H2O expressed as 18 g = 1 mol of water Molecular mass of H2SO4 expressed as 98 g = 1 mol of H2SO4 Formula mass of NaCl expressed as 58.5 g = 1 mol of NaCl Thus, the relationship between mole and mass can be expressed as: the Or, Mass of substance (g) = number of moles x molar mass A detailed relationship between a substance and a mole through molar mass and number of particles is presented here. Summary showing a relationship between a substance and a mole. SUBSTANCE Compound t en em El Molecular mass (amu) (Expressed in g) 6.02 x 1023 (Contains) gram molecules molecular mass (is equivalent to) Atomic mole Formula mass mass (amu) (amu) (Expressed in g) (Expressed in g) 6.02 x 1023 (Contains) gram 6.02 x 1023 gram (Contains) atoms atomic mass formula units formula mass (is equivalent to) (is equivalent to) mole mole Chemistry - IX 20 Unit 1: Fundamentals of Chemistry i. Which term is used to represent the mass of 1 mole of molecules of a substance? ii. How many atoms are present in one gram atomic mass of a substance ? iii. Explain the relationship between mass and mole of a substance. Test yourself iv. Find out the mass of 3 moles of oxygen atoms. 1.6 v. How many molecules of water are present in half mole of water? Example 1.4 Calculate the gram molecule (number of moles) in 40 g of H3PO4. Solution Therefore, 40 grams will contain 0.408 gram molecule (mol) of H3PO4. 1.6 CHEMICAL CALCULATIONS In chemical calculations, we calculate number of moles and number of particles of a given mass of a substance or vice versa. These calculations are based upon mole concept. Let us have a few examples of these calculations. Calculating the number of moles and number of particles from known mass of a substance. First calculate the number of moles from given mass by using equation Then calculate number of particles from the calculated number of moles with the help of following equation: 1.6.1 Mole-Mass Calculations In these calculations, we calculate the number of moles of a substance from the known mass of the substance with the help of following equation: of the When we rearrange the equation to calculate mass of a substance from the number of moles of a substance we get, Chemistry - IX 21 Unit 1: Fundamentals of Chemistry Example 1.5 You have a piece of coal (carbon) weighing 9.0 gram. Calculate the number of moles of coal in the given mass. Solution The mass is converted to the number of moles by the equation: So, 9.0 g of coal is equivalent to 0.75 mol. 1.6.2 Mole-Particle Calculations In these calculations, we can calculate the number of moles of a substance from the given number of particles. (These particles are the atoms, molecules or formula units). On rearranging above equation we get, Summary of Molar Calculations: mole NA mole molar mass Mass of Substances Mole Number of Particles known mass number of particles molar mass NA Remember Never calculate the number of particles from mass of the substance or vice versa. Always make calculations through moles. For calculations of the number of atoms in molecular compounds and the number of ions in ionic compounds; first calculate the number of molecules or formula units and then calculate the number of atoms or ions. Example 1.6 Calculate the number of moles, number of molecules and number of atoms present in 6 grams of water. Chemistry - IX 22 Unit 1: Fundamentals of Chemistry Solution The number of molecules contained in 6 grams of water are 1.98 x 1023 As we know 1 molecule of water consists of 3 atoms, therefore: Example 1.7 There are 3.01  1023 molecules of CO2 present in a container. Calculate the number of moles and its mass. Solution We can calculate the number of molecules of CO2 by putting the values in equation known number of molecules mol Then by putting this value in this equation we get i. How many atoms of sodium are present in 3 moles of sodium and what is the mass of it? ii. How many atoms are in 1 amu and 1 g of hydrogen (H)? iii. How many atoms are present in 16 g of O and 8g of S? iv. Is the mass of 1 mole of O and 1 mole of S same? v. What do you mean by 1 atom of C and 1 gram atom of C? Test yourself vi. If 16 g of oxygen contains 1 mole of oxygen atoms calculate the mass of 1.7 one atom of oxygen in grams. vii. How many times is 1 mole of oxygen atom heavier than 1 mole of hydrogen atom? viii. Why does 10 g nitrogen gas contain the same number of molecules as 10 g of carbon monoxide? Chemistry - IX 23 Unit 1: Fundamentals of Chemistry THE MOLECULARITY OF THE PHYSICAL WORLD. The nature of the physical world as perceived through men's senses has been investigated in depth. The biggest lesson we learnt in 20th century is that Chemistry has become central science. It leads to the discovery of every chemical reaction in any living and non-living thing based on formation of "molecules". A reaction in the smallest living organism or in the most developed species like man, always takes place through the process of molecule formation. Hence it provides basis of "molecularity" of the physical world. CORPUSCULAR NATURE OF MATTER. In 1924 de Broglie put forward the theory of dual nature of matter i.e. matter has both the properties of particles as well as waves. He explained the background of two ideas. He advocated that these two systems could not remain detached from each other. By mathematical evidences, he proved that every moving object is attached with waves and every wave has corpuscular nature as well. It formulated a basis to understand corpuscular nature of matter. THE WORKS OF DIFFERENT SCIENTISTS AT THE SAME TIME HANDICAP OR PROMOTE THE GROWTH OF SCIENCE. Over the course of human history, people have developed many interconnected and validated ideas about the physical, biological, psychological and social worlds. Those ideas have enabled successive generations to achieve an increasingly comprehensive and reliable understanding of the human species and its environment. The means used to develop these ideas are particular ways of observing, thinking, experimenting and validating. These ways represent a fundamental aspect of the nature of science and reflect how science tends to differ from other modes of knowing. It is the union of science, mathematics and technology that forms the scientific endeavor and that makes it so successful. Although, each of these human enterprises has a character and history of its own, each is dependent on and reinforces the others. MOLE - A QUANTITY A computer counting with a speed of 10 million atoms a second would take 2 billion years to count one mole of atoms. If one mole of marbles were spread over the surface of the Earth, our planet would be covered by a 3 miles thick layer of marbles. A glass of water, which contains about 10 moles of water, has more water molecules than the grains of sand in the Sahara desert. Key Points Chemistry is study of composition and properties of matter. It has different branches. Substances are classified into elements and compounds. Elements consist of only one type of atoms. Compounds are formed by chemical combination of atoms of the elements in a fixed ratio. Mixtures are formed by mixing up elements or compounds in any ratio. They are classified as homogeneous and heterogeneous mixtures. Chemistry - IX 24 Unit 1: Fundamentals of Chemistry Each atom of an element has a specific atomic number (Z) and a mass number or atomic mass (A). Atomic mass of an atom is measured relative to a standard mass of C-12. Relative atomic mass of an element is the mass of an element compared with 1/12 mass of an atom of C-12 isotope. Atomic mass unit is 1/12 of the mass of one atom of C-12, lamu = 1.66  l024g Empirical formula is the simplest type of chemical formula, which shows the relative number of atoms of each element in a compound. Molecular formula gives the actual number of atoms of each element in a molecule. Formula mass is the sum of atomic masses of all the atoms in one formula unit of a substance. An atom or group of atoms having a charge on it is called an ion. If it has positive charge it is called a cation and if it has negative charge it is called an anion. There are different types of molecules: monoatomic, polyatomic, homoatomic and heteroatomic. The number of particles in one mole of a substance is called Avogadro's number. The value of this number is 6.02  1023 It is represented as NA. The amount of a substance having 6.02  1023 particles is called a mole. The quantitative definition of mole is atomic mass, molecular mass or formula mass expressed in grams. EXERCISE Multiple Choice Questions Put a ( ) on the correct answer 1. Industrial chemistry deals with the manufacturing of compounds: (a) in the laboratory (b) on micro scale (c) on commercial scale (d) on economic scale 2. Which one of the following compounds can be separated by physical means? (a) mixture (b) element (c) compound (d) radical 3. The most abundant element occurring in the oceans is: (a) oxygen (b) hydrogen (c) nitrogen (d) silicon 4. Which one of the following elements is found in most abundance in the Earth's crust? (a) oxygen (b) aluminium (c) silicon (d) iron 5. The third abundant gas found in the Earth's atmosphere is: (a) carbon monoxide (b) oxygen. (c) nitrogen (d) argon 6. One amu (atomic mass unit) is equivalent to: Chemistry - IX 25 Unit 1: Fundamentals of Chemistry 7. Which one of the following molecule is not tri-atomic? (a) H2 (b) O3 (c) H2O (d) CO2 8. The mass of one molecule of water is: (a) 18 amu (b) 18 g (c) 18 mg (d) 18 kg 9. The molar mass of H2SO4 is: (a)98g (b) 98 amu (c) 9.8 g (d) 9.8 amu 10. Which one of the following is a molecular mass of O2 in amu? 11. How many number of moles are equivalent to 8 grams of Co2? (a) 0.15 (b) 0.18 (c) 0.21 (d) 0.24 12. In which one of the following pairs has the same number of ions? (a) 1 mole of NaCl and 1 mole of MgCl2 (b) 1/2 mole of NaCl and 1/2 mole of MgCl2 (c) 1/2 mole of NaCl and 1/3 mole of MgCl2 (d) 1/3 mole of NaCl and 1/2 mole of MgCl2 Which one of the following pairs has the same mass? 13. (a) 1 mole of CO and 1 mole of N2 (b) 1 mole of CO and 1 mole of CO2 (c) 1 mole of O2 and 1 mole of N2 (d) 1 mole of O2 and 1 mole of Co2 Short answer questions. 1. Define industrial chemistry and analytical chemistry. 2. How can you differentiate between organic and inorganic chemistry? 3. Give the scope of biochemistry. 4. How does homogeneous mixture differ from heterogeneous mixture? 5. What is the relative atomic mass? How is it related to gram? 6. Define empirical formula with an example. 7. State three reasons why do you think air is a mixture and water a compound? 8. Explain why are hydrogen and oxygen considered elements whereas water as a compound. 9. What is the significance of the symbol of an element? 10. State the reasons: soft drink is a mixture and water is a compound. 11. Classify the following into element, compound or mixture: i. He and H2 ii. CO and Co iii. Water and milk iv. Gold and brass v. Iron and steel 12. Define atomic mass unit. Why is it needed? Chemistry - IX 26 Unit 1: Fundamentals of Chemistry 13. State the nature and name of the substance formed by combining the following: i. Zinc + Copper ii. Water + Sugar iii. Aluminium + Sulphur iv. Iron + Chromium + Nickel 14. Differentiate between molecular mass and formula mass, which of the followings have molecular formula? H2O, NaCl, KI, H2SO4 15. Which one has more atoms: 10 g of Al or 10 g of Fe? 16. Which one has more molecules: 9 g of water or 9 g of sugar (C12H22O11)? 17. Which one has more formula units: 1 g of NaCl or 1 g of KC1? 18. Differentiate between homoatomic and heteroatomic molecules with examples. 19. In which one of the followings the number of hydrogen atoms is more? 2 moles of HC1 or 1 mole of NH3 (Hint: 1 mole of a substance contains as much number of moles of atoms as are in 1 molecule of a substance Long Answer Questions. 1. Define element and classify the elements with examples. 2. List five characteristics by which compounds can be distinguished from mixtures. 3. Differentiate between the following with examples: i. Molecule and gram molecule ii. Atom and gram atom iii. Molecular mass and molar mass iv. Chemical formula and gram formula 4. Mole is SI unit for the amount of a substance. Define it with examples? Numericals 1. Sulphuric acid is the king of chemicals. If you need 5 moles of sulphuric acid for a reaction, how many grams of it will you weigh? 2. Calcium carbonate is insoluble in water. If you have 40 g of it; how many Ca2+ and CO32 ions are present in it? 3. If you have 6.02 x 1023 ions of aluminium; how many sulphate ions will be required to prepare Al2(SO4)3? 4. Calculate the number of molecules in the following compounds: a. 16 g of H2CO3 b. 20 g of HNO3 c. 30 g of C6H12O6 5. Calculate the number of ions in the following compounds: a. 10 g of AlCl3 b. 30 g of BaCl2 c. 58 g of H2SO4(aq) 6. What will be the mass of 2.05l016 molecules of H2SO4 7. How many atoms are required to prepare 60 g of HNO3? 8. How many ions of Na+ and Cl will be present in 30 g of NaCl? 9. How many molecules of HC1 will be required to have 10 grams of it? 10. How many grams of Mg will have the same number of atoms as 6 grams of C have?  Chapter 2 Structure of Atoms Major Concepts 2.1 Theories and Experiments related Time allocation to Atomic Structure Teaching periods 16 2.2 Electronic Configuration Assessment periods 03 2.3 Isotopes Weightage 10% Students Learning Outcomes Students will be able to: Describe the contributions that Rutherford made to the development of the Atomic Theory. Explain how Bohr's atomic theory differed. Describe the structure of atom including the location of the proton, electron and neutron. Define isotopes. Compare istopes of an atom. Discuss the properties of the isotopes of H, C, CI, U. Draw the structure of different isotopes from mass number and atomic number. State the importance and uses of isotopes in various fields of life. Describe the presence of subshells I shell. Distinguish between shells and subshells. Write the electronic configuration of first 18 elements in the Periodic Table. Introduction Ancient Greek philosopher Democritus suggested that matter is composed of tiny indivisible particles called atoms. The name atom was derived from the Latin word 'Atomos' meaning indivisible. In the beginning of 19th century John Dalton put forward Atomic Theory. According to it 'all matter is made up of very small indivisible particles called atoms'. Till the end of 19th century it was considered that atom cannot be subdivided. However, in the beginning of 20th century experiments performed by Goldstein, J. J. Thomson, Rutherford, Bohr and other scientist revealed that atom is made up of subatomic particles like electron, proton and neutron. Properties of these subatomic particles will be discussed in this chapter. 2.1 THEORIES AND EXPERIMENTS RELATED TO STRUCTURE OF Chemistry - IX 28 Unit 2: Structure of Atoms ATOM According to Dalton, an atom is an indivisible, hard, dense sphere. Atoms of the same element are alike. They combine in different ways to form compounds. In the light of Dalton's atomic theory, scientists performed a series of experiments. But in the late 1800's and early 1900's, scientists discovered new subatomic particles. In 1886, Goldstein discovered positively charged particles called protons. In 1897, J.J. Thomson found in an atom, the negatively charged particles known as electrons. It was established that electrons and protons are fundamental J.J. Thomson (1856- particles of matter. Based upon these observations Thomson 1940) was a British put forth his “plum pudding” theory. He postulated that physicist. He was atoms were solid structures of positively charge with tiny awarded the 1906 Noble negative particles stuck inside. It is like plums in the pudding. Prize in Physics for the discovery of electron and for his work on the Cathode rays and Discovery of Electron conduction of electricity in gases In 1895 Sir William Crooks performed experiments by passing electric current through gases in a discharge tube at very low pressure. He took a glass tube fitted with two metallic electrode, which were connected to a high voltage battery. The pressure inside the tube was kept 104 atm. When high voltage current was passed through the gas, shiny rays were emitted from the cathode which travel towards the anode as shown in figure 2.1. These rays were given the name of “cathode rays” as these were originated from the cathode. Beam of electrons Sir William Crooks (1832-1919) was a (+) Anode British chemist and physicist. He was pioneer of vacuum to vacuum pump tubes. He worked on Battery spectroscopy. Chemistry - IX 29 Unit 2: Structure of Atoms The cathode rays were studied in detail and their properties were determined, which are given below: i. These rays travel in straight lines perpendicular to the cathode surface. ii. They can cast a sharp shadow of an opaque object if placed in their path. iii. They are deflected towards positive plate in an electric field showing that they are negatively charged. iv. They raise temperature of the body on which they fall. v. JJ. Thomson discovered their charge/mass (e / m) ratio. vi. Light is produced when these rays hit the walls of the discharge tube. vii. It was found that the same type of rays were emitted no matter which gas and which cathode was used in the discharge tube. All these properties suggested that the nature of cathode rays is independent of the nature of the gas present in the discharge tube or material of the cathode. The fact that they cast the shadow of an opaque object suggested that these are not rays but they are fast moving material particles. They were given the name electrons. Since all the materials produce same type of particles, it means all the materials contain electrons. As we know materials are composed of atoms, hence the electrons are fundamental particles of atoms. Discovery of Proton In 1886 Goldstein observed that in addition to cathode rays, other rays were also present in the discharge tube. These rays were traveling in opposite direction to cathode rays. He used a discharge tube having perforated cathode as shown in figure 2.2. He found that these rays passed through holes present in the cathode and produced a glow on the walls of the discharge tube. He called these rays as "canal rays". Canal Rays Battery + - - + + - () perforated cathode (+) anode to vacuum pump Fig 2.2 Discharge tube used for the production of canal rays. The properties of these rays were as following: i. These rays travel in straight lines in a direction opposite to the cathode rays. Chemistry - IX 30 Unit 2: Structure of Atoms ii. Their deflection in electric and magnetic field proved that these are positively charged. iii. The nature of canal rays depends upon the nature of gas, present in the discharge tube. iv. These rays do not originate from the anode. In fact these rays are produced when the cathode rays or electrons collide with the residual gas molecules present in the discharge tube and ionize them as follows: v. Mass of these particles was found equal to that of a proton or simple multiple of it. The mass of a proton is 1840 times more than that of an electron. Thus, these rays are made up of positively charged particles. The mass and charge of these particles depend upon the nature of the gas in the discharge tube. Hence, different gases produce different types of positive rays having particles of different masses and different charges. Keep in mind that positive particles produced by a gas will be of the same type i.e. positive rays produced by the lightest gas hydrogen contain protons. Discovery of Neutron Rutherford observed that atomic mass of the element could not be explained on the basis of the masses of electron and proton only. He predicted in 1920 that some neutral particle having mass equal to that of proton must be present in an atom. Thus scientists were in search of such a neutral particle. Eventually in 1932 Chadwick discovered neutron, when he bombarded alpha particles on a beryllium target. He observed that highly penetrating radiations were produced. These radiations were called neutron. Properties of neutron are as following: i. Neutrons carry no charge i.e. they are neutral. ii. They are highly penetrating. iii. Mass of these particles was nearly equal to the mass of a proton. i. Do you know any element having no neutrons in its atoms? ii. Who discovered an electron, a proton and a neutron? iii. How does electron differ from a neutron? iv. Explain, how anode rays are formed from the gas present in the Test yourself 2.1 discharge tube? Chemistry - IX 31 Unit 2: Structure of Atoms 2.1.1 Rutherford's Atomic Model Rutherford performed 'Gold Foil' experiment to understand how negative and positive charges could coexist in an atom. He bombarded alpha particles on a 0.00004 cm thick gold foil. Alpha particles are emitted by radioactive elements like radium and polonium. These are actually helium nuclei (He2+). They can penetrate through matter to some extent. He observed the effects of -particles on a photographic plate or a screen coated with zinc sulphide as shown in figure 2.3. He proved that the 'plum-pudding' model of the atom was not correct. small tion deflec majority of particles pass undeflected led repel back large deflection Fig 2.3 Scattering of alpha particles by the atoms of gold foil. Observations made by Rutherford were as follows: i. Almost all the particles passed through the foil un-deflected. ii. Out of 20000 particles, only a few were deflected at fairly large angles and very few bounced back on hitting the gold foil. Results of the experiment Keeping in view the experiment, Rutherford proposed planetary model for an atom and concluded following results: Chemistry - IX 32 Unit 2: Structure of Atoms i. Since most of the particles passed through the foil un- deflected, therefore most of the volume occupied by an atom is empty. ii. The deflection of a few particles proved that there is a 'center of positive charges' in an atom, which is called 'nucleus' of an atom. iii. The complete rebounce of a few particles show that the nucleus is very dense and hard. iv. Since a few particles were deflected, it shows that the Rutherford was a British- size of the nucleus is very small as compared to the total New Zealand chemist. He performed a series of volume of an atom. experiments using a - v. The electrons revolve around the nucleus. particles. He won the 1908 Noble Prize in Chemistry. vi. An atom as a whole is neutral, therefore the number of In 1911, he proposed the electrons in an atom is equal to the number of protons. nuclear model of the atom and performed the first vii. Except electrons, all other fundamental particles that lie experiment to split atom. Because of his great within the nucleus, are known as nucleons. contributions, he is considered the father of nuclear science. Defects in Rutherford's Model Although Rutherford's experiment proved that the 'plum-pudding' model of an atom was not correct, yet it had following defects: i. According to classical theory of radiation, electrons being the charged particles should release or emit energy continuously and they should ultimately fall into the nucleus. ii. If the electrons emit energy continuously, they should form a continuous spectrum but in fact, line spectrum was observed. Although the scientists had objections on the atomic model presented by Rutherford, yet it cultivated thought provoking ideas among them. They initiated the quest to answer the following questions: i. How can an atom collapse or why are atoms stable? ii. Why does an atom give line spectrum? iii. Scientists considered there must be another model of atom. It indicated that Rutherford's model was not perfect. Chemistry - IX 33 Unit 2: Structure of Atoms 2.1.2 Bohr's Atomic Theory Keeping in view the defects in Rutherford's Atomic Model, Neil Bohr presented another model of atom in 1913. The Quantum Theory of Max Planck was used as foundation for this model. According to Bohr's model, revolving electron in an atom does not absorb or emit energy continuously. The energy of a revolving electron is 'quantized' as it revolves only in orbits of fixed energy, called 'energy levels' by him. The Bohr's atomic model is shown in figure 2.4. Neil Bohr was a Danish physicist who joined Rutherford in 1912 for his post doctoral research. In 1913, Bohr presented his atomic model based upon Quantum theory. He won the 1922 Noble Prize for Physics for his work on the structure of an atom. Fig 2.4 Bohr's atomic model showing orbits. The Bohr's atomic model was based upon the following postulates: i. The hydrogen atom consists of a tiny nucleus and electrons are revolving in one of circular orbits of radius ‘r’ around the nucleus. ii. Each orbit has a fixed energy that is quantized. iii. As long as electron remains in a particular orbit, it does not radiate or absorb energy. The energy is emitted or absorbed only when an electron jumps from one orbit to another. iv. When an electron jumps from lower orbit to higher orbit, it absorbs energy and when it jumps from higher orbit to lower orbit it radiates energy. This change in energy, E is given by following Planck's equation Where, h is Planck's constant equal to 6.63  1034 Js, and v is frequency of light. v. Electron can revolve only in orbits of a fixed angular moment mvr, given as: Where 'n' is the quantum number or orbit number having values 1,2,3 and so on. Chemistry - IX 34 Unit 2: Structure of Atoms Quantum means fixed energy. It is the smallest amount of energy that can be emitted or absorbed as electromagnetic radiation. Quanta is plural of quantum. In 1918 Noble prize in physics was awarded to German physicist Do you know? Max Planck (1858-1947) for his work on the quantum theory. Summary of differences between two theories: Rutherford's Atomic Theory Bohr's Atomic Theory i. It was based upon classical theory. It was based upon quantum theory. ii. Electrons revolve around the nucleus. Electrons revolve around the nucleus in orbits of fixed energy. iii. No idea about orbits was introduced. Orbits had angular momentum. iv. Atoms should produce continuous Atoms should produce line spectrum. spectrum. v. Atoms should collapse. Atoms should exist. i. How was it proved that the whole mass of an atom is located at its centre? ii. How was it shown that atomic nuclei are positively charged ? iii. Name the particles which determine the mass of an atom. iv. What is the classical theory of radiation? How does it differ from quantum theory? v. How can you prove that angular momentum is quantized? Hint: Let angular momentum (mvr) of 1st orbit is mvr = nh/2 Test yourself 2.2 By putting the values of h and  2.2 ELECTRONIC CONFIGURATION Before discussing electronic configuration let us first understand the concept of shells and subshells. We have learnt about the structure of atom i.e. it consists of a tiny nucleus lying in the center and electrons revolving around the nucleus. Now we will discuss how the electrons revolve around the nucleus? The electrons revolve around the nucleus in different energy levels or shells according to their respective energies (potential energy). The concept of potential energy of an electron shall be discussed in higher classes. Energy levels are represented by 'n' values 1, 2, 3 and so on. They are designated by the alphabets K, L, M and so on. A shell closer to the nucleus is of minimum energy. Since K shell is closest to the nucleus, the energy of shells increases from K shell onwards. Such as: Chemistry - IX 34 Unit 2: Structure of Atoms Quantum means fixed energy. It is the smallest amount of energy that can be emitted or absorbed as electromagnetic radiation. Quanta is plural of quantum. In 1918 Noble prize in physics was awarded to German physicist Do you know? Max Planck (1858-1947) for his work on the quantum theory. Summary of differences between two theories: Rutherford's Atomic Theory Bohr's Atomic Theory i. It was based upon classical theory. It was based upon quantum theory. ii. Electrons revolve around the nucleus. Electrons revolve around the nucleus in orbits of fixed energy. iii. No idea about orbits was introduced. Orbits had angular momentum. iv. Atoms should produce continuous Atoms should produce line spectrum. spectrum. v. Atoms should collapse. Atoms should exist. i. How was it proved that the whole mass of an atom is located at its centre? ii. How was it shown that atomic nuclei are positively charged ? iii. Name the particles which determine the mass of an atom. iv. What is the classical theory of radiation? How does it differ from quantum theory? v. How can you prove that angular momentum is quantized? Hint: Let angular momentum (mvr) of 1st orbit is mvr = nh/2 Test yourself 2.2 By putting the values of h and  2.2 ELECTRONIC CONFIGURATION Before discussing electronic configuration let us first understand the concept of shells and subshells. We have learnt about the structure of atom i.e. it consists of a tiny nucleus lying in the center and electrons revolving around the nucleus. Now we will discuss how the electrons revolve around the nucleus? The electrons revolve around the nucleus in different energy levels or shells according to their respective energies (potential energy). The concept of potential energy of an electron shall be discussed in higher classes. Energy levels are represented by 'n' values 1, 2, 3 and so on. They are designated by the alphabets K, L, M and so on. A shell closer to the nucleus is of minimum energy. Since K shell is closest to the nucleus, the energy of shells increases from K shell onwards. Such as: Chemistry - IX 36 Unit 2: Structure of Atoms N shell can accommodate 32 electrons. As we know there is a slight difference between the energies of the subshells within a shell, therefore, filling of electrons in subshells of a shell is such as that V subshell is filled first and then its p subshell and then other subshells are filled. The maximum capacity of subshells to accommodate electrons is: ‘s’ subshell can accommodate 2 electrons. ‘p’ subshell can accommodate 6 electrons. Let us write the electronic configuration of the elements and their ions with the help of a few examples. Keep in mind, we should know three things: i. The number of electrons in an atom. ii. The sequence of shells and subshells according to the energy levels. iii. The maximum number of electrons that can be placed in different shells and subshells. Example 2.1 1 Write the electronic configuration of an element having 11 8 M electrons. 2 L K Solution: Keep in mind that all electrons do not have the same energy. Therefore, they are accommodated in different shells according to increasing energy and capacity of the shell. First of all, the electrons will go to K shell which has minimum energy. It can accommodate 2 electrons. After this, electrons will go to L shell that can accommodate 8 electrons. Thus K and L shells have accommodated 10 electrons. The remaining 1 electron will go to M shell, the outermost shell of maximum energy in this case. The electronic configuration will be written as: But it is not necessary to write the subshells. Therefore, it is simply written as 2,8 and 1. Further distribution of electrons in subshells will be: Example 2.2 88 Write down the electronic configuration of Cl ion 2 M K L Solution: We know that chlorine has 17 electrons and chloride ion (Cl) has 17 + 1 = 18 electrons. Its electronic configuration will be 2, 8, 8, which is presented in the figure. The further distribution of 85 electrons in subshells will be 2 M K L Example 2.3 An element has 5 electrons in M shell. Find out its atomic number. Chemistry - IX 37 Unit 2: Structure of Atoms Solution: When there are 5 electrons in M shell, it means K and L shell are completely filled with their maximum capacity of 10 electrons. Hence the electronic configuration of the element is: So the total number of electrons is 2 + 8 + 5 = 15 As we know, the number of electrons in an atom is equal to its atomic number. Therefore, atomic number of this element is 15. 2.2.2 The electronic configuration of first 18 elements The sequence of filling of electrons in different subshells is as following: Where number represents the shell number, while letters (s and p) represent subshells. The superscript shows the number of electrons in a subshell. The sum of superscripts number is the total number of electrons in an atom. i.e. atomic number of an element. The electronic configuration of first 18 elements is shown in the Table 2.1 Table 2.1 Electronic Configuration of First Eighteen Elements Chemistry - IX 38 Unit 2: Structure of Atoms i. How many the maximum number of electrons that can be accommodated in a p-sub shell? ii. How many subshells are there in second shell? iii. Why does an electron first fill 2p orbital and then 3s orbital? iv. If both K and L shells of an atom are completely filled; how many Test yourself total number of electrons are present in them? 2.3 v. How many electrons can be accommodated in M shell? vi. What is the electronic configuration of a hydrogen atom? vii. What is atomic number of phosphorus? Write down its electronic configuration. viii. If an element has atomic number 13 and atomic mass 27; how many electrons are there in each atom of the element? ix. How many electrons will be in M shell of an atom having atomic number 15, x. What is maximum capacity of a shell? 2.3 ISOTOPES 2.3.1 Definition Isotopes are defined as the atoms of an element that have same atomic number but different mass numbers. They have same electronic configuration and number of protons but they differ in the number of neutrons. Isotopes have similar chemical properties because these depend upon electronic configuration. But they have different physical properties because these depend upon mass numbers. Most of the elements have isotopes. Here we will discuss the isotopes of hydrogen, carbon, chlorine and uranium only. 2.3.2 Examples i) Isotopes of Hydrogen The naturally occurring hydrogen is combination of its three isotopes, present in different abundances. The three isotopes of hydrogen are named as protium, deuterium and tritium Each one of them has 1 proton and 1 electron, but number of neutrons are different as shown in Table 2.2 The isotopes are represented as: protium ( 11 H ) deuterium ( 21 H ) tritium ( 31 H ) Chemistry - IX 39 Unit 2: Structure of Atoms ii) Isotopes of Carbon There are two stable isotopes of carbon 12C and 13C and one radioactive isotope 14 C. The isotope 12C is present in abundance of 98.9 %, while 13C and 14C are both present only 1.1 % in nature. All of them have the same number of protons and electrons but differ in number of neutrons. They are represented as follows: p=6 p=6 p=6 n=6 n=7 n=8 carbon (12 6 C) carbon (13 6 C) carbon (14 6 C) iii) Isotopes of Chlorine There are two isotopes of chlorine, iv) Isotopes of Uranium There are three isotopes of uranium i.e. in nature nearly 99%. The difference in their number of electrons, protons and neutrons is shown below: Table 2.2 Atomic Number, Mass Number, Number of Protons and Neutrons of H, C, CI and U APPLICATION OF ISOTOPES In science and many different technological fields isotopes have vast applications. The biggest application is in the field of medicine. They are applied in diagnosis, radiotherapy and treatment of many diseases like cancer. Chemistry - IX 40 Unit 2: Structure of Atoms 2.3.3 Uses With the advancement of the scientific knowledge, the isotopes have found many applications in our lives. Following are the major fields in which isotopes have vast applications: i. Radiotherapy (Treatment of Cancer) For the treatment of skin cancer, isotopes like P-32 and Sr-90 are used because they emit less penetrating beta radiations. For cancer, Co-60, affecting within the body, is used because it emits strongly penetrating gamma rays. ii. Tracer for Diagnosis and Medicine The radioactive isotopes are used as tracers in medicine to diagnose the presence of tumor in the human body. Isotopes of Iodine-131 are used for diagnosis of goiter in thyroid gland. Similarly technetium is used to monitor the bone growth. iii. Archaeological and Geological Uses The radioactive isotopes are used to estimate the age of fossils like dead plants and animals and stones, etc. The age determination of very old objects based on the half- lives of the radioactive isotope is called radioactive-isotope dating. An important method of age determination of old carbon containing objects (fossils) by measuring the radioactivity of C-14 in them is called radio-carbon dating or simply carbon dating. iv. Chemical Reaction and Structure Determination The radioisotopes are used in a chemical reaction to follow a radioactive element during the reaction and ultimately to determine the structure. For example: C-14 is used to label CO2. As CO2 is used by the plants for photosynthesis to form glucose, its movement is detected through the various intermediate steps up to glucose. v. Applications in Power Generation Radioactive isotopes are used to generate electricity by carrying out controlled nuclear fission reactions in nuclear reactors. For example, when U-235 is bombarded with slow moving neutrons, the uranium nucleus breaks up to produce Barium-139 and Krypton-94 and three neutrons. A large amount of energy is released which is used to convert water into steam in boilers. The steam then drives the turbines to generate electricity. This is the peaceful use of atomic energy for development of a nation. Chemistry - IX 41 Unit 2: Structure of Atoms i. Why do the isotopes of an element have different atomic masses? ii. How many neutrons are present in C-12 and C-13? iii. Which of the isotopes of hydrogen contains greater number of neutrons? iv. Give one example each of the use of radioactive isotope in medicine Test yourself and radiotherapy. 2.4 v. How is the goiter in thyroid gland detected? vi. Define nuclear fission reaction. vii. When U-235 breaks up, it produces a large amount of energy. How is this energy used? viii. How many neutrons are produced in the fission reaction of U-235? ix. U-235fission produces two atoms of which elements? TESTING PREVAILING THEORIES BRINGS ABOUT CHANGE IN THEM Science is a process for producing knowledge. The process depends both on making careful observations of phenomenae and inventing theories for making sense out of those observations. Change in knowledge is inevitable because new observations may challenge prevailing theories. No matter how well one theory explains a set of observations, it is possible that another theory may fit just as well or better, or may fit a still wider range of observations. In science, the testing and improving and occasional discarding of theories, whether new or old, go on all the time. Scientists assume that even if there is no way to secure complete and absolute truth, increasingly accurate approximations can be made to account for the world and how it works. Key Points Cathode rays were discovered in last decade of nineteen century. The properties of cathode rays were determined and they led to the discovery of electron. Canal rays were discovered in 1886 by Goldstein. The properties of canal rays resulted in the discovery of proton in the atom. Neutron in the atom was discovered in 1932 by Chadwick. First of all structure of an atom was presented by Rutherford in 1911. He proposed that an atom contains nucleus at the center and electrons revolve around this nucleus. Bohr presented an improved model of an atom in 1913 based upon four postulates. He introduced the concept of circular orbit, in which electrons revolve. As long as electron remains in a particular orbit, it does not radiate energy. Release and gain of energy is because of change of orbit. The concept of shells and subshells is explained. A shell consists of subshells. Isotopes are defined as the atoms of elements that have the same atomic number but different atomic mass. Hydrogen, carbon and uranium have three isotopes each, whereas chlorine has two isotopes. Chemistry - IX 42 Unit 2: Structure of Atoms EXERCISE Multiple Choice Questions Put a ( ) on the correct answer 1. Which one of the following results in the discovery of proton (a) cathode rays (b) canal rays (c) X-rays (d) alpha rays. 2. Which one of the following is the most penetrating. (a) protons (b) electrons (c) neutrons (d) alpha particles 3. The concept of orbit was used by (a) J. J. Thomson (b) Rutherford (c) Bohr (d) Planck 4. Which one of the following shell consists of three subshells. (a) O shell (b) N shell (c) L shell (d) M shell 5. Which radioisotope is used for the diagnosis of tumor in the body? (a) cobalt-60 (b) iodine-131 (c) strontium-90 (d) phosphorus-30 6. When U-235 breaks up, it produces: (a) electrons (b) neutrons (c) protons (d) nothing 7. The p subshell has: (a) one orbital (b) two orbitals (c) three orbitals (d) four orbitals 8. Deuterium is used to make: (a) light water (b) heavy water (c) soft water (d) hard water 9. The isotope C-12 is present in abundance of: (a) 96.9 % (b) 97.6 % (c) 99.7 % (d) none of these 10. Who discovered the proton: (a) Goldstein (b) J. J. Thomson (c) Neil Bohr (d) Rutherford Short Short answer questions. 1. What is the nature of charge on cathode rays? 2. Give five characteristics of cathode rays. 3. The atomic symbol of a phosphorus ion is given as ^P3~ (a) How many protons, electrons and neutrons are there in the ion? (b) What is name of the ion? (c) Draw the electronic configuration of the ion. (d) Name the noble gas which has the same electronic configuration as the phosphorus ion has. 4. Differentiate between shell and subshell with examples of each. 5. An element has an atomic number 17. How many electrons are present in K, L and M shells of the atom? Chemistry - IX 43 Unit 2: Structure of Atoms 6. Write down the electronic configuration of Al3+. How many electrons are present in its outermost shell? 7. Magnesium has electronic configuration 2, 8, 2, (a) How many electrons are in the outermost shell? (b) In which subshell of the outermost shell electrons are present? (c) Why magnesium tends to lose electrons? 8. What will be the nature of charge on an atom when it loses an electron or when it gains an electron? 9. For what purpose U-235 is used? 10. A patient has goiter. How will it be detected? 11. Give three properties of positive rays. 12. What are the defects of Rutherford's atomic model? 13. As long as electron remains in an orbit, it does not emit or absorb energy. When does it emit or absorb energy? Long Answer Questions. 1. How are cathode rays produced? What are its five major characteristics? 2. How was it proved that electrons are fundamental particles of an atom? 3. Draw a labeled diagram to show the presence of protons in the discharge tube and explain how canal rays were produced. 4. How Rutherford discovered that atom has a nucleus located at the centre of the atom? 5. One of the postulates of Bohr's atomic model is that angular momentum of a moving electron is quantized. Explain its meaning and calculate the angular momentum of third orbit (i.e. n=3) 6. How did Bohr prove that an atom must exist? 7. What do you mean by electronic configuration? What are basic requirements while writing electronic configuration of an element (atom)? 8. Describe the electronic configuration of Na+, Mg2+ and Al3+ ions. Do they have the same number of electrons in the outermost shell? 9. Give the applications of isotopes in the field of radiotherapy and medicines. 10. What is an isotope? Describe the isotopes of hydrogen with diagrams.  Chapter3 Periodic Table and Periodicity of Properties Major Concepts Time allocation Teaching periods 12 3.1 Periodic Table Assessment periods 02 3.2 Periodic Properties Weightage 10% Students Learning Outcomes Students will be able to: Distinguish between period and group in the Periodic table. State the Periodic law. Classify elements (into two categories: groups and periods) according to the configuration of their outermost electrons. Determine the demarcation of the periodic table into s-block and /?-block. Explain the shape of the periodic table. Determine the location of families of the periodic table. Recognize the similarity in the physical and chemical properties of elements in the same family of the elements. Identify the relationship between electronic configuration and position of elements in the periodic table. Explain how shielding effect influences periodic trends. Describe how electronegativities change within a group and within a period in the periodic table. Introduction In nineteenth century, chemists devoted much of their efforts in attempts to arrange elements in a systematic manner. These efforts resulted in discovery of periodic law. On the basis of this law, the elements known at that time, were arranged in the form of a table which is known as periodic table. One of the significant features of the table was that it predicted the properties of those elements which were not eve