Class 11 Chemistry-01 (1)

Class 11th Chemistry

The chemistry chapters in class 11 are an important part of clearing not just the board exams, but also competitive exams like JEE, NEET, etc. 

For students who want to go the extra mile and understand more and more concepts, NCERTis the one-stop solution. NCERT solutions for class 11 chemistry helps students cover all the topics in detail. Students have the liberty and can choose any topic of their own from Class 11 Chemistry. The best way to prepare for the class 11 chemistry syllabus is to refer to class 11 chemistry notes.

Here are the class 11 chemistry notes for chapter 1 and 2.

Chapter 1 Some Basics Concepts of Chemistry

This class 11 Chemistry chapter 1 explains the role that is played in everyday life by chemistry. Furthermore, this chapter will explain the laws of chemical combinations and the nature of the matter. In this chemistry class 11 chapter 1, students will go in details about Dalton’s atomic theory where the concepts of molecules and atoms are explained. Also, class 11 chemistry chapter 1 deals with the molecular masses and concepts of atomic masses. 

Chapter 2 Structure of Atoms

This Chemistry class 11 chapter 2 will be fundamental for atoms and thus students will get to know about the discovery of proton, electron, and neutrons. They will also study what isotopes, isobars, atomic numbers, etc is. This chapter also describes details of Thomson’s model along with its limitations. Besides this, it talks about Bohr’s model and Rutherford’s model and its limitations. There is some detail being thrown into the dual nature of light and matter, Heisenberg’s uncertainty principle, de Broglie’s principle, the shapes of s, d, and p orbitals, quantum numbers, etc. There are also topics like Hund’s rule and Pauli’s exclusion principle which are also discussed in this chapter. It is advisable that students create class 11 chemistry notes for this chapter.      

Chapter 3 Periodicity in Properties and Classification of Elements

In this chapter, you will learn about a brief history of the periodic table and its development, the significance of classifying the periodic table, how the present form of the periodic table was formed, etc. Furthermore, the information about the trends in the periodic table for atoms like ionic radii, radii, inert gas radii, electronegativity, electron gain, valency, etc. is more discussed in this chapter. A total of 40 questions are there in CBSE Class 11 Chemistry in the NCERT for students to practice.

Chapter 4 Chemical Bonding and Molecular Structure

This Chemistry class 11 chapter 4 will help you understand what a covalent bond and an ionic bond are. There are more details about the parameters of the bonds, covalent bond and its polar character, the bond theory of valence, covalent bond and it’s geometry, resonance, etc. Furthermore, this chapter discusses the VSEPR theory, concepts of hybridization that involve s, d, and p orbitals, various shapes of some molecules, and many more. There are a total of 40 questions in this chapter that can help students to practice. It is advisable that students create class 11 chemistry notes for this chapter.

Chapter 5 States of Matter – Liquid, and Gas

This chapter will make students understand three states of matter along with the types of bonding and intermolecular interactions. There are also some insights about the boiling point and melting points given in the chapter. Furthermore, the roles of gas laws are discussed and how Gay lussac’s law, Boyle’s law, Avogadro’s law, etc are helping students understand their ideal behavior. Along with this, the Avogadro’s number, empirical deviation in the gas equation, and the ideal equation required for the numerical are illustrated.

Chapter 6 Chemical Thermodynamics

This chemistry class 11 chapter 6 helps with the concepts of various systems and their different types. There is also discussion provided about the surroundings in the form of heat, work, energy, intensive and extensive properties, and state functions. There is a discussion about the first law of thermodynamics in this chapter. It involves enthalpy and internal energy, specific heat, heat capacity, measurement of heat, etc. It is advisable that students create or refer to class 11 chemistry notes for this chapter.

Chapter 7 Equilibrium

This chemistry class 11 chapter 7 talks about the concepts of equilibrium in chemical and physical processes and details related to the equilibrium’s dynamic nature. There are also some insights related to the law of mass action, the factors affecting equilibrium, and the equilibrium constant as per Le Chatelier’s principle. Furthermore, the information about the acid strength, ionization of polybasic acids, Henderson equation, the concept of pH, etc are also discussed.

Chapter 8 Redox Reaction

This chapter will provide in-depth knowledge to students about the reduction and oxidation and various insights about the redox reactions. Furthermore, information about balancing the redox reactions, oxidation number, etc will also be provided. There are a total of thirty questions in the chapter that also discusses the loss and gains of electrons.

Chapter 9 Hydrogen

Through this chapter, you will learn about the occurrence of hydrogen and its position in the periodic table. Along with this, there will be some information about the isotopes, their properties, and how they are prepared is also discussed in this chapter. Information related to interstitial and hydrogen ionic covalent bonds is also discussed in this chapter.

Chapter 10 S-block Elements

This chapter discusses the elements present in groups 1 and 2. It discusses the electronic configuration along with their occurrence. Every first element in the group shows some anomalous behavior which is also discussed in this chapter. There are diagonal relationships like atomic radii, variation in terms of properties in ionization enthalpy, ionic radii, etc is also discussed. How some of the important compounds like sodium chloride, sodium carbonate, sodium hydrogen carbonate, and sodium hydroxide are prepared is also discussed in this chapter.

Chapter 11 Some P-block Elements

This chapter provides more of a general view of the p-block elements to the students. There is also in-depth and detailed information about the elements in group 13 being discussed in this chapter. Also, the variation of oxidation states and their properties is also discussed. The chemical and physical properties of boron along with its important compounds like boric acid, borax, boron hydrides, etc are discussed in this chapter.

Chapter 12 Organic Chemistry – Some Basic Techniques and Principles

This chapter talks more about various purification methods along with quantitative and qualitative analysis being used for it. Furthermore, information related to IUPAC nomenclature and classification of various organic compounds is also discussed in this chapter. Along with this, the electronic displacements occurring in a covalent bond in the form of electromeric effect, inductive effect, hyperconjugation, resonance are also discussed in-depth. It is advisable that students create and refer to class 11 chemistry notes for this chapter.

Chapter 13 Hydrocarbons

In this class 11 chemistry chapter 13, students will get to know in detail about the classification of hydrocarbons and their uses, properties, and related reactions. Furthermore, this chapter talks about alkanes, alkynes, and alkenes. It also talks about related nomenclature, physical properties, IUPAC names, chemical reactions, combustion, isomerism, etc. This chemistry chapters class 11 is considered an important one. 

Chapter 14 Environmental Chemistry

This chapter 14 Environmental Chemistry will talk about the environmental part of chemistry like environmental pollution related to air, soil, and water. Furthermore, all the chemical reactions happening in the atmosphere due to smog, major atmospheric pollutants, etc are also discussed. This final chapter for class 11 chemistry further discusses ozone, acid rains, and their reactions. A total of 20 questions will help students understand various alternative tools required for reducing pollution.

Examples

Here are some questions with solutions to serve as chemistry class 11 notes:

Making chemistry class 11 notes is very important and advisable for students to ensure better understanding of the topics. 

Class 11 chemistry chapter 2 notes:

Practice Questions

Class 11 chemistry mcqs

Study Material

Making chemistry class 11 notes is very important for students to better grasp the concepts. Here are chemistry class 11 notes for chapter 1 and 2 to help you begin with.

class 11 chemistry chapter 1 notes

  • Importance of Chemistry

Chemistry has a direct impact on our life and has a wide range of applications in different fields. These are given below:

(A) In Agriculture and Food:

(i) It has provided chemical fertilizers such as urea, calcium phosphate, sodium nitrate, ammonium phosphate etc.

(ii) It has helped to protect the crops from insects and harmful bacteria, by the use of certain effective insecticides, fungicides and pesticides.

(iii) The use of preservatives has helped to preserve food products like jam, butter, squashes etc. for longer periods.

(B) In Health and Sanitation:

(i) It has provided mankind with a large number of life-saving drugs. Today, dysentery and pneumonia are curable due to discovery of sulpha drugs and penicillin life-saving drugs. Cisplatin and taxol have been found to be very effective for cancer therapy and AZT (Azidothymidine) is used for AIDS victims.

(ii) Disinfectants such as phenol are used to kill the microorganisms present in drains, toilet, floors etc.

(iii) A low concentration of chlorine i.e., 0.2 to 0.4 parts per million (ppm) is used for sterilization of water to make it fit for drinking purposes.

(C) Saving the Environment:

The rapid industrialisation all over the world has resulted in a lot of pollution.

Poisonous gases and chemicals are being constantly released in the atmosphere. They are polluting the environment at an alarming rate. Scientists are working day and night to develop substitutes which may cause lower pollution. For example, CNG (Compressed Natural Gas), a substitute for petrol, is very effective in checking pollution caused by automobiles.

(D) Application in Industry:

Chemistry has played an important role in developing many industrially ^ manufactured fertilizers, alkalis, acids, salts, dyes, polymers, drugs, soaps,

detergents, metal alloys and other inorganic and organic chemicals including new materials contribute in a big way to the national economy.

  • Matter

Anything which has mass and occupies space is called matter.

For example, book, pencil, water, air are composed of matter as we know that they have

mass and they occupy space.

  • Classification of Matter

There are two ways of classifying the matter:

(A) Physical classification

(B) Chemical classification

(A) Physical Classification:

Matter can exist in three physical states:

  1. Solids 2. Liquids 3. Gases
  1. Solids: The particles are held very close to each other in an orderly fashion and there is not much freedom of movement.

Characteristics of solids: Solids have definite volume and definite shape.

  1. Liquids: In liquids, the particles are close to each other but can move around. Characteristics of liquids: Liquids have definite volume but not definite shape.
  1. Gases: In gases, the particles are far apart as compared to those present in solid or liquid states. Their movement is easy and fast.

Characteristics of Gases: Gases have neither definite volume nor definite shape. They completely occupy the container in which they are placed.

(B) Chemical Classification:

Based upon the composition, matter can be divided into two main types:

  1. Pure Substances 2. Mixtures.
  2. Pure substances: A pure substance may be defined as a single substance (or matter) which cannot be separated by simple physical methods.

Pure substances can be further classified as (i) Elements (ii) Compounds

(i) Elements: An element consists of only one type of particle. These particles may be atoms or molecules.

For example, sodium, copper, silver, hydrogen, oxygen etc. are some examples of elements. They all contain atoms of one type. However, atoms of different elements are different in nature. Some elements such as sodium . or copper contain single atoms held together as their constituent particles whereas in some others two or more atoms combine to give molecules of the element. Thus, hydrogen, nitrogen and oxygen gases consist of molecules in which two atoms combine to give the respective molecules of the element.

(ii) Compounds: It may be defined as a pure substance containing two or more elements combined together in a fixed proportion by weight and can be decomposed into these elements by suitable chemical methods. Moreover, the properties of a compound are altogether different from the constituting elements.

The compounds have been classified into two types. These are:

(i) Inorganic Compounds: These are compounds which are obtained from non-living sources such as rocks and minerals. A few

examples are: Common salt, marble, gypsum, washing soda etc.

(ii) Organic Compounds are the compounds which are present in plants and animals. All the organic compounds have been found to contain carbon as their essential constituent. For example, carbohydrates, proteins, oils, fats etc.

  1. Mixtures: The combination of two or more elements or compounds which are not chemically combined together and may also be present in any proportion, is called mixture. A few examples of mixtures are: milk, sea water, petrol, lime water, paint glass, cement, wood etc.

Types of mixtures: Mixtures are of two types:

(i) Homogeneous mixtures: A mixture is said to be homogeneous if it has a uniform composition throughout and there are no visible boundaries of separation between the constituents.

For example: A mixture of sugar solution in water has the same sugar water composition throughout and all portions have the same sweetness.

(ii) Heterogeneous mixtures: A mixture is said to be heterogeneous if it does not have uniform composition throughout and has visible boundaries of separation between the various constituents. The different constituents of a heterogeneous mixture can be seen even with naked eye.

For example: When iron filings and sulphur powder are mixed together, the mixture formed is heterogeneous. It has a greyish-yellow appearance and the two constituents, iron and sulphur, can be easily identified with naked eye.

  • Differences between Compounds and Mixtures

Compounds

  1. In a compound, two or more elements are combined chemically.
  2. In a compound, the elements are present in the fixed ratio by mass. This ratio cannot change.
  3. CompoUnds are always homogeneous i.e., they have the same composition throughout.

4 In a compound, constituents cannot be separated by physical methods

  1. In a compound, the constituents lose their identities i.e., i compound does not show the characteristics of the constituting elements.

Mixtures

  1. In a mixture, or more elements or compounds are simply mixed and not combined chemically.
  2. In a mixture the constituents are not present in fixed ratio. It can vary
  3. Mixtures may be either homogeneous or heterogeneous in nature.
  4. Constituents of mixtures can be separated by physical methods.

5, In a mixture, the constituents do not lose their identities i.e., a mixture shows the characteristics of all the constituents.

We have discussed the physical and chemical classification of matter. A flowsheet representation of the same is given below.

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Properties of Matter and Their Measurements

Physical Properties: Those properties which can be measured or observed without changing the identity or the composition of the substance.

Some examples of physical properties are colour, odour, melting point, boiling point etc. Chemical Properties: It requires a chemical change to occur. The examples of chemical properties are characteristic reactions of different substances. These include acidity, basicity, combustibility etc.

  • Units of Measurement

Fundamental Units: The quantities mass, length and time are called fundamental quantities and their units are known as fundamental units.

There are seven basic units of measurement for the quantities: length, mass, time, temperature, amount of substance, electric current and luminous intensity.

Si-System: This system of measurement is the most common system employed throughout the world.

It has given units of all the seven basic quantities listed above.

 Definitions of Basic SI Units

  1. Metre: It is the length of the path travelled by light in vacuum during a time interval of 1/299792458 of a second.
  2. Kilogram: It is the unit of mass. It is equal to the mass of the international prototype

of the kilogram. ,

  1. Second: It is the duration of 9192631, 770 periods of radiation which correspond to the transition between the two hyperfine levels of the ground state of caesium- 133 atoms.
  2. Kelvin: It is the unit of thermodynamic temperature and is equal to 1/273.16 of the thermodynamic temperature of the triple point of water.
  3. Ampere: The ampere is that constant current which if maintained in two straight parallel conductors of infinite length, of negligible circular cross section and placed, 1 meter apart in vacuum, would produce between these conductors a force equal to 2 x 10-7 N per metre of length.
  4. Candela: It may be defined as the luminous intensity in a given direction, from a source which emits monochromatic radiation of frequency 540 x 1012 Hz and that has a radiant intensity in that direction of 1/ 683 watt per steradian.
  5. Mole: It is the amount of substance which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon -12. Its symbol is ‘mol’.
  • Mass and Weight

Mass: Mass of a substance is the amount of matter present in it.

The mass of a substance is constant.

The mass of a substance can be determined accurately in the laboratory by using an analytical

balance. SI unit of mass is a kilogram.

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Weight: It is the force exerted by gravity on an object. Weight of substance may vary from one place to another due to change in gravity.

Volume: Volume means the space occupied by matter. It has the units of (length)3. In SI units, volume is expressed in metre3 (m3). However, a popular unit of measuring volume, particularly in liquids is litre (L) but it is not in SI units or an S.I. unit.

Mathematically,

1L = 1000 mL = 1000 cm3 = 1dm3.

Volume of liquids can be measured by different devices like burette, pipette, cylinder, measuring flask etc. All of them have been calibrated.

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Temperature: There are three scales on which temperature can be measured. These are known as Celsius scale (°C), Fahrenheit scale (°F) and Kelvin scale (K).

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Thermometers with Celsius scale are calibrated from 0°C to 100°C.

-> Thermometers with Fahrenheit scale are calibrated from 32°F to 212°F.

-> Kelvin Scale of temperature is S.I. scale and is very common these days.Temperature on this scale is shown by the sign K.

The temperature on two scales are related to each other by the relationship

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Density: Density of a substance is its amount of mass per unit volume. So, SI unit of density can be obtained as follows:

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This unit is quite large and a chemist often expresses density in g cm3 where mass is expressed in gram and volume is expressed in cm3.

  • Uncertainty in Measurements

All scientific measurements involve certain degree of error or uncertainty. The errors which arise depend upon two factors.

(i) Skill and accuracy of the worker (ii) Limitations of measuring instruments.

  • Scientific Notation

It is an exponential notation in which any number can be represented in the form N x 10n where n is an exponent having positive or negative values and N can vary between 1 to 10. Thus, 232.508 can be written as 2.32508 x 102 in scientific notation.

Now let us see how calculations are carried out with numbers expressed in scientific notation.

(i) Calculation involving multiplication and division

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(ii) A calculation involving addition and subtraction: For these two operations, the first numbers are written in such a way that they have the same exponent. After that, the coefficients are added or subtracted as the case may be. For example,

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  • Significant Figures

Significant figures are meaningful digits which are known with certainty. There are certain rules for determining the number of significant figures. These are stated below:

  1. All non-zero digits are significant. For example, in 285 cm, there are three significant figures and in 0.25 mL, there are two significant figures.
  2. Zeros preceding to first non-zero digit are not significant. Such zeros indicates the position of decimal point.

For example, 0.03 has one significant figure and 0.0052 has two significant figures.

  1. Zeros between two non-zero digits are significant. Thus, 2.005 has four significant figures.
  2. Zeros at the end or right of a number are significant provided they are on the right side of the decimal point. For example, 0.200 g has three significant figures.
  3. Counting numbers of objects. For example, 2 balls or 20 eggs have infinite significant figures as these are exact numbers and can be represented by writing infinite number of zeros after placing a decimal.

i.e., 2 = 2.000000

or 20 = 20.000000

  • Addition and Subtraction of Significant Figures

In addition or subtraction of the numbers having different precisions, the final result should be reported to the same number of decimal places as in the term having the least number of decimal places.

For example, let us carry out the addition of three numbers 3.52, 2.3 and 6.24, having different precisions or the different numbers of decimal places.

The final result has two decimal places but the answer has to be reported only up to one decimal place, i.e., the answer would be 12.0.

Subtraction of numbers can be done in the same way as addition.

The final result has four decimal places. But it has to be reported only up to two decimal places, i.e., the answer would be 11.36.

  • Multiplication and Division of Significant Figures

In the multiplication or division, the final result should be reported upto the same number of significant figures as present in the least precise number.

Multiplication of Numbers: 2.2120 x 0.011 = 0.024332

According to the rule the final result = 0.024

Division of Numbers: 4.2211÷3.76 = 1.12263

The correct answer = 1.12

  • 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.

  • Laws of Chemical Combinations

The combination of elements to form compounds is governed by the following five basic laws.

(i) Law of Conservation of Mass

(ii) Law of Definite Proportions

(iii) Law of Multiple Proportions

(iv) Law of Gaseous Volume (Gay Lussac’s Law)

(v) Avogadro’s Law

(i) Law of Conservation of Mass

The law was established by a French chemist, A. Lavoisier. The law states:

In all physical and chemical changes, the total mass of the reactants is equal to that of the products.

In other words, matter can neither be created nor destroyed.

The following experiments illustrate the truth of this law.

(a) When matter undergoes a physical change.

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It is found that there is no change in weight though a physical change has taken place.

(b) When matter undergoes a chemical change.

For example, decomposition of mercuric oxide.

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During the above decomposition reaction, matter is neither gained nor lost.

(ii) Law of Definite Proportions

According to this law:

A pure chemical compound always consists of the same elements combined together in a fixed proportion by weight.

For example, Carbon dioxide may be formed in a number of ways i.e.,

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iii) Law of Multiple Proportions

If two elements combine to form two or more compounds, the weight of one of the elements which combines with a fixed weight of the other in these compounds, bears simple whole number ratio by weight.

For example,

(iv) Gay Lussac’s Law of Gaseous Volumes

The law states that, under similar conditions of temperature and pressure, whenever gases combine, they do so in volumes that bear simple whole-number ratio with each other and also with the gaseous products. The law may be illustrated by the following examples.

(a) Combination between hydrogen and chlorine:

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(b) Combination between nitrogen and hydrogen: The two gases lead to the formation of ammonia gas under suitable conditions. The chemical equation is

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Avogadro’s Law: Avogadro proposed that, equal volumes of gases at the same temperature and pressure should contain equal number of molecules.

For example,

If we consider the reaction of hydrogen and oxygen to produce water, we see that two volumes of hydrogen combine with one volume of oxygen to give two volumes of water without leaving any unreacted oxygen.

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Dalton’s Atomic Theory

In 1808, Dalton published ‘A New System of Chemical Philosophy’ in which he proposed the following:

  1. Matter consists of indivisible atoms.
  2. All the atoms of a given element have identical properties including identical mass. Atoms of different elements differ in mass.
  3. Compounds are formed when atoms of different elements combine in a fixed ratio.
  4. Chemical reactions involve reorganisation of atoms. These are neither created nor destroyed in a chemical reaction.
  • Atomic Mass

The atomic mass of an element is the number of times an atom of that element is heavier than an atom of carbon taken as 12. It may be noted that the atomic masses as obtained above are the relative atomic masses and not the actual masses of the atoms.

One atomic mass unit (amu) is equal to l/12th of the mass of an atom of carbon-12 isotope. It is also known as unified mass.

Average Atomic Mass

Most of the elements exist as isotopes which are different atoms of the same element with different mass numbers and the same atomic number. Therefore, the atomic mass of an element must be its average atomic mass and it may be defined as the average relative mass of an atom of an element as compared to the mass of carbon atoms (C-12) taken as 12w.

Molecular Mass

Molecular mass is the sum of atomic masses of the elements present in a molecule. It is obtained by multiplying the atomic mass of each element by number of its atoms and adding them together.

For example,

Molecular mass of methane (CH4)

= 12.011 u + 4 (1.008 u)

= 16.043 u

Formula Mass

Ionic compounds such as NaCl, KNO3, Na2C03 etc. do not consist of molecules i.e., single entities but exist “as ions closely packed together in a three-dimensional space.

In such cases, the formula is used to calculate the formula mass instead of molecular mass. Thus, formula mass of NaCl = Atomic mass of sodium + atomic mass of chlorine

= 23.0 u + 35.5 u = 58.5 u.

  • Mole Concept

It is found that one gram atom of any element contains the same number of atoms and one gram molecule of any substance contains the same number of molecules. This number has been experimentally determined and found to be equal to 6.022137 x 1023 The value is generally called Avogadro’s number or Avogadro’s constant.

It is usually represented by NA:

Avogadro’s Number, NA = 6.022 × 1023

  • Percentage Composition

One can check the purity of a given sample by analysing this data. Let us understand by taking the example of water (H20). Since water contains hydrogen and oxygen, the percentage composition of both these elements can be calculated as follows:

 Empirical Formula

The formula of the compound which gives the simplest whole number ratio of the atoms of various elements present in one molecule of the compound.

For example, the formula of hydrogen peroxide is H202. In order to express its empirical formula, we have to take out a common factor 2. The simplest whole number ratio of the atoms is 1:1 and the empirical formula is HO. Similarly, the formula of glucose is C6H1206. In order to get the simplest whole number of the atoms,

Common factor = 6

The ratio is = 1 : 2 : 1 The empirical formula of glucose = CH20

  • Molecular Formula

The formula of a compound which gives the actual ratio of the atoms of various elements present in one molecule of the compound.

For example, molecular formula of hydrogen peroxide = H202and Glucose = C6H1206

Molecular formula = n x Empirical formula

Where n is the common factor and also called the multiplying factor. The value of n may be 1, 2, 3, 4, 5, 6 etc.

In case n is 1, Molecular formula of a compound = Empirical formula of the compound.

  • Stoichiometry and Stoichiometric Calculations

The word ‘stoichiometry’ is derived from two Greek words—Stoicheion (meaning element) and metron (meaning measure). Stoichiometry, thus deals with the calculation of masses (sometimes volume also) of the reactants and the products involved in a chemical reaction. Let us consider the combustion of methane. A balanced equation for this reaction is as given below:

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Limiting Reactant/Reagent

Sometimes, in alchemical equations, the reactants present are not the amount as required according to the balanced equation. The amount of products formed then depends upon the reactant which has reacted completely. This reactant which reacts completely in the reaction is called the limiting reactant or limiting reagent. The reactant which is not consumed completely in the reaction is called excess reactant.

Reactions in Solutions

When the reactions are carried out in solutions, the amount of substance present in its given volume can be expressed in any of the following ways:

  1. Mass percent or weight percent (w/w%)
  2. Mole fraction
  3. Molarity
  4. Molality
  1. Mass percent: It is obtained by using the following relation:

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2.Mole fraction: It is the ratio of the number of moles of a particular component to the total number of moles of the solution. For a solution containing n2 moles of the solute dissolved in n1 moles of the solvent,

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  • All substances contain matter which can exist in three states — solid, liquid or gas.
  • Matter can also be classified into elements, compounds and mixtures.
  • Element: An element contains particles of only one type which may be atoms or molecules.
  • Compounds are formed when atoms of two or more elements combine in a fixed ratio to each other.
  • Mixtures: Many of the substances present around us are mixtures.
  • Scientific notation: The measurement of quantities in chemistry are spread over a wide range of 10-31to 1023. Hence, a convenient system of expressing the number in scientific notation is used.
  • Scientific figures: The uncertainty is taken care of by specifying the number of significant figures in which the observations are reported.
  • Dimensional analysis: It helps to express the measured quantities in different systems of units.

     

  • Laws of Chemical Combinations are:

(i) Law of Conservation of Mass

(ii) Law of Definite Proportions

(iii) Law of Multiple Proportions

(iv) Gay Lussac’s Law of Gaseous Volumes

(v) Avogadro’s Law.

  • Atomic mass: The atomic mass of an element is expressed relative to 12C isotope of carbon which has an exact value of 12u.
  • Average atomic mass: Obtained by taking into account the natural abundance of different isotopes of that element.
  • Molecular mass: The molecular mass of a molecule is obtained by taking the sum of atomic masses of different atoms present in a molecule.
  • Avogadro number: The number of atoms, molecules or any other particles present in a given system are expressed in terms of Avogadro constant.

= 6.022 x 1023

  • Balanced chemical equation: A balanced equation has the same number of atoms of each element on both sides of the equation.
  • Stoichiometry: The quantitative study of the reactants required or the products formed is called stoichiometry. Using stoichiometric calculations, the amounts of one or more reactants required to produce a particular amount of product can be determined and vice-versa.

class 11 chemistry chapter 2 notes

Structure of Atom Class 11 Chemistry Chapter 2 Notes. This section serves as a place for Class 11 Chemistry Chapter 2 Notes. 

  • Discovery of Electron—Discharge Tube Experiment

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In 1879, William Crooks studied the conduction of electricity through gases at low pressure. He performed the experiment in a discharge tube which is a cylindrical hard glass tube about 60 cm in length. It is sealed at both ends and fitted with two metal electrodes as shown in Fig. 2.1.

The electrical discharge through the gases could be observed only at very low pressures and at very high voltages.

The pressure of different gases could be adjusted by evacuation. When sufficiently high voltage is applied across the electrodes, current starts flowing through a stream of particles moving in the tube from the negative electrode (cathode) to the positive electrode (anode). These were called cathode rays or cathode ray particles.

  • Properties of Cathode Rays

(i) Cathode rays travel in a straight line.

(ii) Cathode rays start from cathode and move towards the anode.

(iii) These rays themselves are not visible but their behaviour can be observed with the help of certain kinds of materials (fluorescent or phosphorescent) which glow when hit by them.

(iv) Cathode rays consist of negatively charged particles. When an electric field is applied on the cathode rays with the help of a pair of metal plates, these are found to be deflected towards the positive plate indicating the presence of negative charge.

(v) The characteristics of cathode rays do not depend upon the material of electrodes and the nature of gas present in the cathode ray tube.

  • Determination of Charge/Mass (elm) Ratio for Electrons
  1. J. Thomson for the first time experimentally determined charge/mass ratio called elm ratio for the electrons. For this, he subjected the beam of electrons released in the discharge tube as cathode rays to influence the electric and magnetic fields. These were acting perpendicular to one another as well as to the path followed by electrons.

According to Thomson, the amount of deviation of the particles from their path in presence of electrical and magnetic field depends upon following factors:

(i) Greater the magnitude of the charge on the particle, greater is the interaction with the electric or magnetic field and thus greater is the deflection.

(ii) The mass of the particle — lighter the particle, greater the deflection.

(iii) The deflection of electrons from their original path increases with the increase in the voltage across the electrodes or strength of the magnetic field.

By carrying out accurate measurements on the amount of deflections observed by the electrons on the electric field strength or magnetic field strength, Thomson was able to determine the value of

e/me = 1.758820 x 1011 C kg-1 where me = Mass of the electron in kg

e = magnitude of the charge on the electron in coulomb (C).

  • Charge on the Electron

Discovery of Proton—Anode Rays

In 1886, Goldstein modified the discharge tube by using a perforated cathode. On reducing the pressure, he observed a new type of luminous rays passing through the holes or perforations of the cathode and moving in a direction opposite to the cathode rays. These rays were named as positive rays or anode rays or as canal rays. Anode rays are not emitted from the anode but from a space between anode and cathode.

  • Properties of Anode Rays

(i) The value of positive charge (e) on the particles constituting anode rays depends upon the nature of the gas in the discharge tube.

(ii) The charge to mass ratio of the particles is found to depend on the gas from which these originate.

(iii) Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge.

(iv) The behaviour of these particles in the magnetic or electric field is opposite to that observed for electron or cathode rays.

  • Proton

The smallest and lightest positive ion was obtained from hydrogen and was called proton. Mass of proton = 1.676 x 10-27 kg

Charge on a proton = (+) 1.602 x 10-19 C

  • Neutron

It is a neutral particle. It was discovered by Chadwick (1932).

By the bombardment of thin sheets of beryllium with fast moving a-particles he observed • that highly penetrating rays consist of neutral particles which were named neutrons.

  • Thomson Model of Atom

(i) J. J. Thomson proposed that an atom may be regarded as a sphere of approximate radius 1CT8 cm carrying positive charge due to protons and in which negatively charged electrons are embedded.

(ii) In this model, the atom is visualized as a pudding or cake of positive charge with electrons embedded into it.

(iii) The mass of atom is considered to be evenly spread over the atom according to this model.

Drawback of Thomson Model of Atom

This model was able to explain the overall neutrality of the atom, it could not satisfactorily explain the results of scattering experiments carried out by Rutherford in 1911.

  • Rutherford’s a-particle Scattering Experiment

Rutherford in 1911, performed some scattering experiments in which he bombarded thin foils of metals like gold, silver, platinum or copper with a beam of fast moving a-particles. The thin gold foil had a circular fluorescent zinc sulphide screen around it. Whenever a a-particles struck the screen, a tiny flash of light was produced at that point.

From these experiments, he made the following observations:

Class 11 Chemistry 23

(i) Most of the a-particles passed through the foil without undergoing any deflection,

(ii) A few a-particles underwent deflection through small angles.

(iii) Very few mere deflected back i.e., through an angle of nearly 180°.

From these observations, Rutherford drew the following conclusions:

(i) Since most of the a-particles passed through the foil without undergoing any deflection, there must be sufficient empty space within the atom.

(ii) A small fraction of a-particles was deflected by small angles. The positive charge has to be concentrated in a very small volume that repelled and deflected a few positively charged a-particles. This very small portion of the atom was called nucleus.

(iii) The volume of nucleus is very small as compared to total volume of atom.

  • Rutherford’s Nuclear Model of an Atom

(i) The positive charge and most of the mass of the atom was densely concentrated in an extremely small region. This very small portion of the atom was called nucleus by Rutherford.

(ii) The nucleus is surrounded by electrons that move around the nucleus with a very high speed in circular paths called orbits.

(iii) Electrons and nucleus are held together by electrostatic forces of attraction.

  • Atomic Number

The number of protons present in the nucleus is equal to the atomic number (z). For example, the number of protons in the hydrogen nucleus is 1, in sodium atom it is 11, therefore, their atomic numbers are 1 and 11. In order to keep the electrical neutrality, the number of electrons in an atom is equal to the number of protons (atomic number, z). For example, number of electrons in hydrogen atom and sodium atom are 1 and 11 respectively.

Atomic Number (z) = Number of protons in the nucleus of an atom.

= Number of electrons in a neutral atom.

  • Mass Number

Number of protons and neutrons present in the nucleus are collectively known as nucleons. The total number of nucleons is termed as mass number (A) of the atom.

Mass Number (A) = Number of protons (p) + Number of neutrons (n).

  • Isotopes

Atoms with identical atomic numbers but a different atomic mass numbers are known as Isotopes.

Class 11 Chemistry 24

These three isotopes are shown in the figure below:

Class 11 Chemistry 25

Class 11 Chemistry 26

Characteristics of Isotopes

(i) Since the isotopes of an element have the same atomic number, but different mass number, the nuclei of isotopes contain the same number of protons, but different number of neutrons.

(ii) Since, the isotopes differ in their atomic masses, all the properties of the isotopes depending upon the mass are different.

(iii) Since, the chemical properties are mainly determined by the number of protons in the nucleus, and the number of electrons in the atom, the different isotopes of an element exhibit similar chemical properties. For example, all the isotopes of carbon on burning give carbon dioxide.

  • Isobars

Class 11 Chemistry iso

Characteristics of a Wave

Drawbacks of Rutherford Model

(i) When a body is moving in an orbit, it achieves acceleration. Thus, an electron moving around nucleus in an orbit is under acceleration.

According to Maxwell’s electromagnetic theory, charged particles when accelerated must emit electromagnetic radiations. Therefore, an electron in an orbit will emit radiations, the energy carried by radiation comes from electronic motion. Its path will become closer to nucleus and ultimately should spiral into nucleus within . 10-8 s. But actually this does not happen.

Thus, Rutherford’s model cannot explain the stability of atom if the motion of electrons is described on the basis of classical mechanics and electromagnetic theory.

(ii) Rutherford’s model does not give any idea about distribution of electrons around the nucleus and about their energies.

  • Developments Leading to the Bohr’s Model of Atom

Two developments played a major role in the formulation of Bohr’s model of atom. These were:

(i) Dual character of the electromagnetic radiation which means that radiations possess both wave like and particle like properties.

(ii) Experimental results regarding atomic spectra which can be explained only by assuming quantized electronic energy levels in atoms.

  • Nature of Electromagnetic Radiation (Electromagnetic Wave Theory)

This theory was put forward by James Clark Maxwell in 1864. The main points of this theory are as follows:

(i) The energy is emitted from any source (like the heated rod or the filament of a bulb through which electric current is passed) continuously in the form of radiations and is called the radiant energy.

(ii) The radiations consist of electric and magnetic fields oscillating perpendicular to each other and both perpendicular to the direction of propagation of the radiation.

(iii) The radiations possess wave character and travel with the velocity of light 3 x 108 m/sec.

(iv) These waves do not require any material medium for propagation. For example, rays from the sun reach us through space which is a non-material medium.

  • Characteristics of a Wave

Wavelength: It is defined as the distance between any two consecutive crests or troughs. It is represented by X and its S.I. unit is metre.

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Frequency: Frequency of a wave is defined as the number of waves passing through a point in one second. It is represented by v (nu) and is expressed in Hertz (Hz).

1 Hz = 1 cycle/sec.

Velocity: Velocity of a wave is defined as the linear distance travelled by the wave in one second.

It is represented by c and is expressed in cm/sec or m/sec.

Amplitude: Amplitude of a wave is the height of the crest or the depth of the through. It is represented by V and is expressed in the units of length.

Wave Number: It is defined as the number of waves present in 1 meter length. Evidently, it will be equal to the reciprocal of the wavelength. It is represented by bar v (read as nu bar).

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Electromagnetic Spectrum: When electromagnetic radiations are arranged in order of their increasing wavelengths or decreasing frequencies, the complete spectrum obtained is called electromagnetic spectrum.

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Limitations of Electromagnetic Wave Theory

Electromagnetic wave theory was successful in explaining properties of light such as interference, diffraction etc; but it could not explain the following:

(i) The phenomenon of black body radiation.

(ii) The photoelectric effect.

(iii) The variation of heat capacity of solids as a function of temperature.

(iv) The line spectra of atoms with reference to hydrogen.

  • Black Body Radiation

The ideal body, which emits and absorbs all frequencies is called a black body and the radiation emitted by such a body is called black body radiation. The. exact frequency distribution of the emitted radiation from a black body depends only on its temperature

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At a given temperature, intensity of radiation emitted increases with decrease of wavelength, reaches a maximum value at a given wavelength and then starts decreasing with further decrease of wavelength as shown in Fig 2.6.

  • Planck’s Quantum Theory

To explain the phenomenon of ‘Black body radiation’ and the photoelectric effect, Max Planck in 1900, put forward a theory known as Planck’s Quantum Theory.

This theory was further extended by Einstein in 1905. The main points of this theory were as follows: 

(i) The radiant energy emitted or absorbed in the form of small packets of energy. Each such packet of energy is called a quantum.

(ii) The energy of each quantum is directly proportional to the frequency of the radiation

where h is a proportionality constant, called Planck’s constant. Its value is equal to 6.626 x 10-34 Jsec.

  • Photoelectric Effect

Hertz, in 1887, discovered that when a beam of light of certain frequency strikes the surface of some metals, electrons are emitted or ejected from the metal surface. The phenomenon is called photoelectric effect.

Observations in Photoelectric Effect

(i) Only photons of light of certain minimum frequency called threshold frequency (v0) can cause the photoelectric effect. The value of v0 is different for different metals.

(ii) The kinetic energy of the electrons which are emitted is directly proportional to the frequency of the striking photons and is quite independent of their intensity.

(iii) The number of electrons that are ejected per second from the metal surface depends upon the intensity of the striking photons or radiations and not upon their frequency.

Explanation of Photoelectric Effect

Einstein (1905) was able to give an explanation of the different points of the photoelectric effect using Planck’s quantum theory as under:

Photoelectrons are ejected only when the incident light has a certain minimum frequency (threshold frequency v0)

(ii) If the frequency of the incident light (v) is more than the threshold frequency (v0), the excess energy (hv – hv0) is imparted to the electron as kinetic energy.

K.E. of the ejected electron

energy of the emitted electron.

(iii) On increasing the intensity of light, more electrons are ejected but the energies of the electrons are not altered.

  • Dual Behaviour of Electromagnetic Radiation

From the study of behaviour of light, scientists came to the conclusion that light and other electromagnetic radiations have dual nature. These are wave nature as well as particle nature. Whenever radiation interacts with matter, it displays particle like properties in contrast to the wavelike properties (interference and diffraction) which it exhibits when it propagates. Some microscopic particles, like electrons, also exhibit this wave-particle duality.

  • Spectrum

When a ray of white light is passed through a prism the wave with shorter wavelength bends more than the one with a longer wavelength. Since ordinary white light consists of waves with all the wavelengths in the visible range, array of white light is spread out into a series of coloured bands called spectrum. The light of red colour which has longest wavelength is deviated the least while the violet light, which has shortest wavelength is deviated the most.

Continuous Spectrum

When a ray of white light is analysed by passing through a prism it is observed that it splits up into seven different wide bands of colours from violet to red (like rainbow). These colours are so continuous that each of them merges into the next. Hence, the spectrum is called continuous spectrum.

Emission Spectra

Emission Spectra is noticed when the radiations emitted from a source are passed through a prism and then received on the photographic plate. Radiations can be emitted in a number of ways such as:

(i) from the sun or glowing electric bulb.

(ii) by passing an electric discharge through a gas at low pressure.

(iii) by heating a substance to high temperature.

Line Spectra

When the vapours of some volatile substance are allowed to fall on the flame of a Bunsen burner and then analysed with the help of a spectroscope. Some specific coloured lines appear on the photographic plate which are different for different substances. For example, sodium or its salts emit yellow light while potassium or its salts give out violet light.

Absorption Spectra

When white light is passed through the vapours of a substance and the transmitted light is then allowed to strike a prism, dark lines appear in the otherwise continuous spectrum. The dark lines indicate that the radiations corresponding to them were absorbed by the substance from the white light. This spectrum is called the absorption spectrum.

Dark lines appear exactly at the same positions where the lines in the emission spectra appear.

  • Line Spectrum of Hydrogen

When an electric discharge is passed through hydrogen gas enclosed in a discharge tube under low pressure and the emitted light is analysed by a spectroscope, the spectrum consists of a large number of lines which are grouped into different series. The complete spectrum is known as the hydrogen spectrum.

On the basis of experimental observations, Johannes Rydberg noted that all series of lines in the hydrogen spectrum could be described by the following expression:

Rydberg in 1890, and has given a simple theoretical equation for the calculation of wavelengths and wavenumbers of the spectral lines in different series of the hydrogen spectrums. The equation is known as Rydberg formula (or equation).

This relation is valid for hydrogen atom only. For other species,

where Z is the atomic number of the species.

Here RH = constant, called Rydberg constant for hydrogen and n1 , n2 are integers (n2 > n1)

For any particular series, the value of n1 is constant while that of n2 changes. For example,

For Lyman series, n1= 1, n2= 2, 3, 4, 5………..

For Balmer series, n1 = 2, n2 = 3, 4, 5, 6………..

For Paschen series, n1= 3, n2 = 4, 5, 6, 7………..

For Brackett series,n1 = 4, n2 = 5, 6, 7, 8………..

For Pjund series, n1 =5, n2 = 6, 7, 8, 9………..

Thus, by substituting the values of n1 and n2 in the above equation, wavelengths and wavenumber of different spectral lines can be calculated. When n1 = 2, the expression gave above is called Balmer’s formula.

  • Bohr’s Model of Atom

Niels Bohr in 1913, proposed a new model of atom on the basis of Planck’s Quantum Theory. The main points of this model are as follows:

(i) In an atom, the electrons revolve around the nucleus in certain definite circular paths called orbits.

(ii) Each orbit is associated with definite energy and therefore these are known as energy

levels or energy shells. These are numbered as 1, 2, 3, 4……….. or K, L, M, N………..

(iii) Only those energy orbits are permitted for the electron in which the angular momentum of the electron is a whole number multiple of h/2π

Angular momentum of electron (mvr) = nh/2π (n = 1, 2, 3, 4 etc).

m = mass of the electron.

v = tangential velocity of the revolving electron.

r = radius of the orbit.

h = Planck’s constant.

n is an integer.

(iv) As long as the electron is present in a particular orbit, it neither absorbs nor loses energy, and its energy, therefore, remains constant.

(v) When energy is supplied to an electron, it absorbs energy only in fixed amounts as quanta and jumps to a higher energy state away from the nucleus known as the excited state. The excited state is unstable, the electron may jump back to the lower energy state and in doing so, it emits the same amount of energy. (∆E = E2 – E1).

  • Achievements of Bohr’s Theory
  1. Bohr’s theory has explained the stability of an atom.
  2. Bohr’s theory has helped in calculating the energy of electrons in the hydrogen atom and one electron species. The mathematical expression for the energy in the nth orbit is,
  3. Bohr’s theory has explained the atomic spectrum of the hydrogen atoms.
  • Limitations of Bohr’s Model

(i) The theory could not explain the atomic spectra of the atoms containing more than one electron or multielectron atoms.

(ii) Bohr7s theory failed to explain the fine structure of the spectral lines.

(iii) Bohr’s theory could not offer any satisfactory explanation of Zeeman effect and Stark effect.

(iv) Bohr’s theory failed to explain the ability of atoms to form molecules formed by chemical bonds.

(v) It was not in accordance with Heisenberg’s uncertainty principle.

  • Dual Behaviour of Matter (de Broglie Equation)

de Broglie in 1924, proposed that matter, like radiation, should also exhibit dual behavior i.e., both particle-like and wave-like properties. This means that like photons, electrons also have momentum as well as wavelength.

From this analogy, de Broglie gave the following relation between wavelength (λ) and momentum (p) of a material particle.

  • Heisenberg’s Uncertainty Principle

It states that, “It is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron”.

  • Significance of Uncertainty Principle

(i) It rules out the existence of definite paths or trajectories of electrons and other similar particles.

(ii) The effect of Heisenberg’s uncertainty principle is significant only for microscopic objects and is negligible for macroscopic objects.

  • Reasons for the Failure of Bohr Model

(i) The wave character of the electron is not considered in Bohr Model.

(ii) According to Bohr Model an orbit is a clearly defined path and this path can completely be defined only if both the position and the velocity of the electron are known exactly at the same time. This is not possible according to the Heisenberg’s uncertainty principle.

  • Quantum Mechanical Model of Atom

Quantum mechanics: Quantum mechanics is a theoretical science that deals with the study of the motions of microscopic objects that have both observable wave-like and particle-like properties.

Important Features of Quantum Mechanical Model of Atom

(i) The energy of electrons in an atom is quantized i.e., can only have certain values.

(ii) The existence of a quantized electronic energy level is a direct result of the wave-like properties of electrons.

(iii) Both the exact position and exact velocity of an electron in an atom cannot be determined simultaneously.

(iv) An atomic orbital has wave function φ. There are many orbitals in an atom. Electron occupy an atomic orbital which has definite energy. An orbital cannot have more than two electrons. The orbitals are filled in increasing order of energy. All the information about the electron in an atom is stored in orbital wave function φ.

(v) The probability of finding electron at a point within an atom is proportional to square of orbital wave function i.e., |φ2|at that point. It is known as probability density and is always positive.

From the value of φ2 at different points within atom, it is possible to predict the region around the nucleus where electron most probably will be found.

  • Quantum Numbers

Atomic orbitals can be specified by giving their corresponding energies and angular momentums which are quantized (i.e., they have specific values). The quantized values can be expressed in terms of quantum number. These are used to get complete information about electron i.e., its location, energy, spin etc.

Principal Quantum Number (n)

It is the most important quantum number since it tells the principal energy level or shell to which the electron belongs. It is denoted by the letter V and can have any integral value except zero, i.e., n = 1, 2, 3, 4……….. etc.

The various principal energy shells are also designated by the letters, K, L, M, N, O, P ….. etc. Starting from the nucleus.

The principal quantum number gives us the following information:

(i) It gives the average distance of the electron from the nucleus.

(ii) It completely determines the energy of the electron in hydrogen atom and hydrogen like particles.

(iii) The maximum number of electrons present in any principal shell is given by 2n2 where n is the number of the principal shell.

Azimuthal or Subsidiary or Orbital Angular Quantum Number (l)

It is found that the spectra of the elements contain not only the main lines but there are many fine lines also present. This number helps to explain the fine lines of the spectrum.

The azimuthal quantum number gives the following information:

(i) The number of subshells present in the main shell.

(ii) The angular momentum of the electron present in any subshell.

(in) The relative energies of various subshells.

(iv) The shapes of the various subshells present within the same principal shell.

This quantum number is denoted by the letter T. For a given value of n, it can have any value ranging from 0 to n – 1. For example,

For the 1st shell (k), n = 1, l can have only one value i.e., l = 0 For n = 2, the possible value of l can be 0 and 1.

Subshells corresponding to different values of l are represented by the following symbols:

value of l 0 1 2 3 4 5 ……………..

Notation for subshell s p d f g h ………………..

Magnetic Orbital Quantum Number (m or m1)

The magnetic orbital quantum number determines the number of preferred orientations of the electrons present in a subshell. Since each orientation corresponds to an orbital, therefore, the magnetic orbital quantum number determines the number of orbitals present in any subshell.

The magnetic quantum number is denoted by the letter m or ml and for a given value of l, it can have all the values ranging from – l to + l including zero.

Thus, for the energy value of l, m has 2l + 1 value.

For example,

For l = 0 (s-subshell), ml can have only one value i.e., m1 = 0.

This means that the s-subshell has only one orientation in space. In other words, s-subshell has only one orbital called s-orbital.

Spin Quantum Number (S or ms)

This quantum number helps to explain the magnetic properties of the substances. A spinning electron behaves like a micro magnet with a definite magnetic moment. If an orbital contains two electrons, the two magnetic moments oppose and cancel each other.

  • Shapes of s-orbitals

s-orbital is present in the s-subshell. For this subshell, l = 0 and ml = 0. Thus, s-orbital with only one orientation has a spherical shape with uniform electron density along all three axes.

The probability of Is electron is found to be maximum near the nucleus and decreases with the increase in the distance from the nucleus. In 2s electron, the probability is also maximum near the nucleus and decreases to zero probability. The spherical empty shell for 2s electron is called nodal surface or simply node.

  • Shapes of p-orbitals

p-orbitals are present in the p-subshell for which l = 1 and m1 can have three possible orientations – 1, 0, + 1.

Thus, there are three orbitals in the p-subshell which are designated as px, py, and Pz orbitals depending upon the axis along which they are directed. The general shape of a p-orbital is dumb-bell consisting of two portions known as lobes. Moreover, there is a plane passing through the nucleus along which the finding of the electron density is almost nil. This is known as nodal plane as shown in the fig.

From the dumb-bell pictures, it is quite obvious that unlike s-orbital, a p-orbital is directional in nature and hence it influences the shapes of the molecules in the formation of which it participates.

  • Shapes of d-orbitals

d-orbitals are present in d-subshell for which l = 2 and m[ = -2, -1, 0, +1 and +2. This means that there are five orientations leading to five different orbitals.

  • Aufbau Principle

The principle states: In the ground state of the atoms, the orbitals are filled in order of their increasing energies.

In other words, electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled.

The order in which the energies of the orbitals increase and hence the order in which the orbitals are filled is as follows:

Is, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, id, 5p, 6s, if, 3d, 6p, 7s, 5f 6d, 7p

The order may be remembered by using the method given in fig. 2.11.

  • Pauli Exclusion Principle

According to this principle, no two electrons in an atom can have the same set of four quantum numbers.

Pauli exclusion principle can also be stated as: Only two electrons may exist in the same orbital and these electrons must have opposite spins.

  • Hund’s Rule of Maximum Multiplicity

It states that: pairing of electrons in the orbitals belonging to the same subshell (p, d or f) does not take place until each orbital belonging to that subshell has got one electron each i.e., it is singly occupied.

  • Electronic Configuration of Atoms

The distribution of electrons into orbitals of an atom is called its electronic configuration. The electronic configuration of different atoms can be represented in two ways.

For example:

  • Causes of Stability of Completely Filled and Half Filled Subshells

The completely filled and half filled subshells are stable due to the following reasons:

  1. Symmetrical distribution of electrons: The completely filled or half filled subshells have symmetrical distribution of electrons in them and are therefore more stable.
  2. The stabilizing effect arises whenever two or more electrons with same spin are present in the degenrate orbitals of a subshell. These electrons tend to exchange their positions

and the energy released due to their exchange is called exchange energy. The number of exchanges that can takes place is maximum when the subshell is either half filled or completely filled.

-As a result the exchange energy is maximum and so is the stability.

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Examples

Q1 .Why should a magnesium ribbon be cleaned before burning in air ?

Solution:

Magnesium gets covered with a layer of magnesium oxide when kept in air for a long time. This layer hinders the burning of magnesium. Hence, it is to be cleaned before burning.

Q2. Write the balanced equation for the following chemical reactions.

Solution:

(i) Hydrogen + Chlorine → Hydrogen chloride

(ii) Barium chloride + Aluminium sulphate → Barium sulphate + Aluminium chloride

(iii) Sodium + Water → Sodium hydroxide + Hydrogen

Answer:

(i) H2 + Cl2 → 2HCl

(ii) 3 BaCl2 + Al2(SO4)3 → BaSO4 + 2 AlCl3

(iii) 2Na + 2H2O → 2NaOH + H2↑

Q3.: Write a balanced chemical equation with state symbols for the following reactions :

Solution:

(i) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride.

(ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water.

Answer:

(i) BaCl2 (aq) + Na2SO4 (aq) → BaSO4(s) + 2NaCl (aq)

(ii) NaOH (aq) + HCl(aq) → NaCl(aq) + H2O(l)

Practice Questions

Chemistry Class 11 Practice Questions

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Physical state of water at is respectively

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Which of these elements is a nonmetal?

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The symbol Ag stands for which element

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What do you call an atom that has more protons than electrons

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Which of this is not a Types of Chemical Reactions

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What is its chemical formula of water?

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FAQ

What are chemical reactions and equations class 10?

  1. Chemical reactions: The transformation of chemical substance into another chemical substance is known as Chemical Reaction. For example: Rusting of iron, the setting of milk into curd, digestion of food, respiration, etc.
  2. Chemical equations: Representation of chemical reaction using symbols and formulae of the substances is called Chemical Equation.

Name and state the law which is kept in mind while we balance a chemical equation.

 Law of conservation of mass. Mass can neither be created nor be destroyed during a chemical reaction.

State one basic difference between a physical change and a chemical change.

In a physical change, no new substance is formed. In a chemical change, a new substance is formed.

What happens when quicklime is added to water?

Quicklime reacts with water vigorously to produce slaked lime and a large amount of heat.



   
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