1.1 IMPORTANCE OF CHEMISTRY many big environmental problems continue to be matters of grave concern to the chemists.Chemistry plays a central role in science and One such problem is the management of theis often intertwined with other branches of Green House gases, like methane, carbonscience. dioxide, etc. Understanding of biochemical Principles of chemistry are applicable processes, use of enzymes for large-scale in diverse areas, such as weather patterns, production of chemicals and synthesis of new functioning of brain and operation of a exotic material are some of the intellectual computer, production in chemical industries, challenges for the future generation of manufacturing fertilisers, alkalis, acids, salts, chemists. A developing country, like India, dyes, polymers, drugs, soaps, detergents, needs talented and creative chemists for metals, alloys, etc., including new material. accepting such challenges. To be a good chemist and to accept such challanges, one Chemistry contributes in a big way to the needs to understand the basic concepts ofnational economy. It also plays an important chemistry, which begin with the concept ofrole in meeting human needs for food, matter. Let us start with the nature of matter.healthcare products and other material aimed at improving the quality of life. This 1.2 Nature of Matter is exemplified by the large-scale production You are already familiar with the term matter of a variety of fertilisers, improved variety from your earlier classes. Anything which has of pesticides and insecticides. Chemistry mass and occupies space is called matter. provides methods for the isolation of life- Everything around us, for example, book, pen, saving drugs from natural sources and pencil, water, air, all living beings, etc., are makes possible synthesis of such drugs. composed of matter. You know that they have Some of these drugs are cisplatin and mass and they occupy space. Let us recall the taxol, which are effective in cancer therapy. characteristics of the states of matter, which The drug AZT (Azidothymidine) is used for you learnt in your previous classes. helping AIDS patients. Chemistry contributes to a large extent in 1.2.1 States of Matter the development and growth of a nation. With You are aware that matter can exist in three a better understanding of chemical principles physical states viz. solid, liquid and gas. it has now become possible to design and The constituent particles of matter in these synthesise new material having specific three states can be represented as shown in magnetic, electric and optical properties. This Fig. 1.1. has lead to the production of superconducting Particles are held very close to each other ceramics, conducting polymers, optical fibres, in solids in an orderly fashion and there is not etc. Chemistry has helped in establishing much freedom of movement. In liquids, the industries which manufacture utility goods, particles are close to each other but they can like acids, alkalies, dyes, polymesr metals, move around. However, in gases, the particles etc. These industries contribute in a big way are far apart as compared to those present in to the economy of a nation and generate solid or liquid states and their movement is employment. easy and fast. Because of such arrangement of particles, different states of matter exhibit In recent years, chemistry has helped in the following characteristics:dealing with some of the pressing aspects (i) Solids have definite volume and definiteof environmental degradation with a fair shape.degree of success. Safer alternatives to environmentally hazardous refrigerants, (ii) Liquids have definite volume but do like CFCs (chlorofluorocarbons), responsible not have definite shape. They take the for ozone depletion in the stratosphere, have shape of the container in which they are been successfully synthesised. However, placed. Reprint 2025-26 Some Basic Concepts of Chemistry 5 Fig. 1.1 Arrangement of particles in solid, liquid Fig. 1.2 Classification of matter and gaseous state (iii) Gases have neither definite volume nor completely mix with each other. This means definite shape. They completely occupy particles of components of the mixture are the space in the container in which they uniformly distributed throughout the bulk of are placed. the mixture and its composition is uniform throughout. Sugar solution and air are the These three states of matter are examples of homogeneous mixtures. Ininterconvertible by changing the conditions contrast to this, in a heterogeneous mixture,of temperature and pressure. the composition is not uniform throughout Solid liquid Gas and sometimes different components are On heating, a solid usually changes to visible. For example, mixtures of salt and a liquid, and the liquid on further heating sugar, grains and pulses along with some changes to gas (or vapour). In the reverse dirt (often stone pieces), are heterogeneous process, a gas on cooling liquifies to the liquid mixtures. You can think of many more and the liquid on further cooling freezes to examples of mixtures which you come across the solid. in the daily life. It is worthwhile to mention here that the components of a mixture can 1.2.2. Classification of Matter be separated by using physical methods, In Class IX (Chapter 2), you have learnt that such as simple hand-picking, filtration, at the macroscopic or bulk level, matter can crystallisation, distillation, etc. be classified as mixture or pure substance. Pure substances have characteristics These can be further sub-divided as shown different from mixtures. Constituent particles in Fig. 1.2. of pure substances have fixed composition. When all constituent particles of a Copper, silver, gold, water and glucose are substance are same in chemical nature, it some examples of pure substances. Glucose is said to be a pure substance. A mixture contains carbon, hydrogen and oxygen in contains many types of particles. a fixed ratio and its particles are of same A mixture contains particles of two or composition. Hence, like all other pure substances, glucose has a fixed composition.more pure substances which may be present Also, its constituents—carbon, hydrogenin it in any ratio. Hence, their composition is and oxygen—cannot be separated by simplevariable. Pure substances forming mixture physical methods.are called its components. Many of the substances present around you are mixtures. Pure substances can further be classified For example, sugar solution in water, air, into elements and compounds. Particles tea, etc., are all mixtures. A mixture may of an element consist of only one type of be homogeneous or heterogeneous. In a atoms. These particles may exist as atoms or homogeneous mixture, the components molecules. You may be familiar with atoms Reprint 2025-26 6 chemistry and molecules from the previous classes; however, you will be studying about them in detail in Unit 2. Sodium, copper, silver, hydrogen, oxygen, etc., are some examples of elements. Their all atoms are of one type. Water molecule Carbon dioxideHowever, the atoms of different elements are different in nature. Some elements, (H2O) molecule (CO2) such as sodium or copper, contain atoms Fig. 1.4 A depiction of molecules of water and as their constituent particles, whereas, in carbon dioxide some others, the constituent particles are molecules which are formed by two or more elements are present in a compound in a fixed atoms. For example, hydrogen, nitrogen and and definite ratio and this ratio is characteristic oxygen gases consist of molecules, in which of a particular compound. Also, the properties two atoms combine to give their respective of a compound are different from those of its molecules. This is illustrated in Fig. 1.3. constituent elements. For example, hydrogen and oxygen are gases, whereas, the compound formed by their combination i.e., water is a liquid. It is interesting to note that hydrogen burns with a pop sound and oxygen is a supporter of combustion, but water is used as a fire extinguisher. 1.3 Properties of Matter and their Measurement 1.3.1 Physical and chemical properties Every substance has unique or characteristic properties. These properties can be classified into two categories — physical properties, such as colour, odour, melting point, boiling point, density, etc., and chemical properties, like composition, combustibility, ractivity with Fig. 1.3 A representation of atoms and molecules acids and bases, etc. When two or more atoms of different Physical properties can be measured elements combine together in a definite ratio, or observed without changing the identity the molecule of a compound is obtained. or the composition of the substance. The measurement or observation of chemicalMoreover, the constituents of a compound properties requires a chemical change tocannot be separated into simpler substances occur. Measurement of physical propertiesby physical methods. They can be separated does not require occurance of a chemicalby chemical methods. Examples of some change. The examples of chemical propertiescompounds are water, ammonia, carbon are characteristic reactions of differentdioxide, sugar, etc. The molecules of water substances; these include acidity or basicity,and carbon dioxide are represented in Fig. 1.4. combustibility, etc. Chemists describe, Note that a water molecule comprises interpret and predict the behaviour of two hydrogen atoms and one oxygen atom. substances on the basis of knowledge of their Similarly, a molecule of carbon dioxide physical and chemical properties, which are contains two oxygen atoms combined with determined by careful measurement and one carbon atom. Thus, the atoms of different experimentation. In the following section, we Reprint 2025-26 Some Basic Concepts of Chemistry 7 will learn about the measurement of physical properties. Maintaining the National Standards of Measurement 1.3.2 Measurement of physical properties The system of units, including unit Quantitative measurement of properties is definitions, keeps on changing with time. reaquired for scientific investigation. Many Whenever the accuracy of measurement of a properties of matter, such as length, area, particular unit was enhanced substantially volume, etc., are quantitative in nature. Any by adopting new principles, member nations quantitative observation or measurement is of metre treaty (signed in 1875), agreed represented by a number followed by units to change the formal definition of that in which it is measured. For example, length unit. Each modern industrialised country, including India, has a National Metrologyof a room can be represented as 6 m; here, Institute (NMI), which maintains standards of6 is the number and m denotes metre, the measurements. This responsibility has been unit in which the length is measured. given to the National Physical Laboratory Earlier, two different systems of (NPL), New Delhi. This laboratory establishes measurement, i.e., the English System experiments to realise the base units and derived units of measurement and maintainsand the Metric System were being used National Standards of Measurement. Thesein different parts of the world. The metric standards are periodically inter-compared system, which originated in France in late with standards maintained at other National eighteenth century, was more convenient as Metrology Institutes in the world, as well it was based on the decimal system. Late, as those, established at the International need of a common standard system was felt Bureau of Standards in Paris. by the scientific community. Such a system was established in 1960 and is discussed in governmental treaty organisation created by a detail below. diplomatic treaty known as Metre Convention, 1.3.3 The International System of Units (SI) which was signed in Paris in 1875. The International System of Units (in The SI system has seven base units French Le Systeme International d’Unités and they are listed in Table 1.1. These units — abbreviated as SI) was established by pertain to the seven fundamental scientific the 11th General Conference on Weights and quantities. The other physical quantities, Measures (CGPM from Conference Generale such as speed, volume, density, etc., can be des Poids et Measures). The CGPM is an inter- derived from these quantities. Table 1.1 Base Physical Quantities and their Units Base Physical Symbol for Name of Symbol for Quantity Quantity SI Unit SI Unit Length l metre m Mass m kilogram kg Time t second s Electric current I ampere A Thermodynamic T kelvin K temperature Amount of n mole mol substance Iv candela cd Luminous intensity Reprint 2025-26 8 chemistry The definitions of the SI base units are These prefixes are listed in Table 1.3. given in Table 1.2. Let us now quickly go through some of The SI system allows the use of prefixes to the quantities which you will be often using indicate the multiples or submultiples of a unit. in this book. Table 1.2 Definitions of SI Base Units The metre, symbol m is the SI unit of length. It is defined by taking the fixed numerical value of the speed of light in vacuum Unit of length metre c to be 299792458 when expressed in the unit ms–1, where the second is defined in terms of the caesium frequencyV Cs. The kilogram, symbol kg. is the SI unit of mass. It is defined by taking the fixed numerical value of the planck constant h to Unit of mass kilogram be 6.62607015×10–34 when expressed in the unit Js, which is equal to kgm2s–1, where the metre and the second are defined in terms of c and V Cs. The second symbol s, is the SI unit of time. It is defined by taking the fixed numerical value of the caesium frequency V Cs, Unit of time second the unperturbed ground-state hyperfine transition frequency of the caesium-133 atom, to be 9192631770 when expressed in the unit Hz, which is equal to s–1. The ampere, symbol A, is the SI unit of electric current. It is defined by taking the fixed numerical value of the elementary Unit of electric ampere charge e to be 1.602176634×10–19 when expressed in the unit current C, which is equal to As, where the second is defined in terms of V Cs. The kelvin, symbol k, is the SI unit of thermodynamic temperature. It is defined by taking the fixed numerical value Unit of of the Boltzmann constant k to be 1.380649×10–23 when thermodynamic kelvin expressed in the unit JK–1, which is equal to kgm2s–2k–1 where temperature the kilogram, metre and second are defined in terms of h, c and V Cs. The mole, symbol mol, is the SI unit of amount of substance. One mole contains exactly 6.02214076×1023 elementary entities. This number is the fixed numerical value of the Avogadro constant, NA, when expressed in the unit mol–1 and Unit of amount mole is called the Avogadro number. The amount of substance, of substance symbol n, of a system is a measure of the number of specified elementary entities. An elementary entity may be an atom, a molecule, an ion, an electron, any other particle or specified group of particles. The candela, symbol cd is the SI unit of luminous intensity in a given direction. It is defined by taking the fixed numerical value of the luminous efficacy of monochromatic radiation Unit of luminous Candela of frequency 540×1012 Hz, Kcd, to be 683 when expressed Intensity in the unit lm·W–1, which is equal to cd·sr·W–1, or cd sr kg–1 m–2s3, where the kilogram, metre and second are defined in terms of h, c and V Cs. Reprint 2025-26 Some Basic Concepts of Chemistry 9 Table 1.3 Prefixes used in the SI System Multiple Prefix Symbol 10–24 yocto y 10–21 zepto z 10–18 atto a 10–15 femto f 10–12 pico p 10–9 nano n 10–6 micro µ 10–3 milli m 10–2 centi c 10–1 deci d 10 deca da 102 hecto h 103 kilo k 106 mega M Fig. 1.5 Analytical balance 109 giga G 1012 tera T SI system, volume has units of m3. But again, 1015 peta P in chemistry laboratories, smaller volumes 1018 exa E are used. Hence, volume is often denoted in 1021 zeta Z cm3 or dm3 units. 1024 yotta Y A common unit, litre (L) which is not an SI unit, is used for measurement of volume 1.3.4 Mass and Weight of liquids. Mass of a substance is the amount of matter 1 L = 1000 mL, 1000 cm3 = 1 dm3present in it, while weight is the force Fig. 1.6 helps to visualise these relations. exerted by gravity on an object. The mass of a substance is constant, whereas, its weight may vary from one place to another due to change in gravity. You should be careful in using these terms. The mass of a substance can be determined accurately in the laboratory by using an analytical balance (Fig. 1.5). The SI unit of mass as given in Table 1.1 is kilogram. However, its fraction named as gram (1 kg = 1000 g), is used in laboratories due to the smaller amounts of chemicals used in chemical reactions. 1.3.5 Volume Volume is the amont of space occupied by a Fig. 1.6 Different units used to express volumesubstance. It has the units of (length)3. So in Reprint 2025-26 10 chemistry In the laboratory, the volume of liquids fahrenheit) and K (kelvin). Here, K is the or solutions can be measured by graduated SI unit. The thermometers based on these cylinder, burette, pipette, etc. A volumetric scales are shown in Fig. 1.8. Generally, flask is used to prepare a known volume of a the thermometer with celsius scale are solution. These measuring devices are shown calibrated from 0° to 100°, where these two in Fig. 1.7. temperatures are the freezing point and the boiling point of water, respectively. The fahrenheit scale is represented between 32° to 212°. The temperatures on two scales are related to each other by the following relationship: 9  F   C   32 5 The kelvin scale is related to celsius scale as follows: K = °C + 273.15 It is interesting to note that temperature below 0 °C (i.e., negative values) are possible in Celsius scale but in Kelvin scale, negative Fig. 1.7 Some volume measuring devices temperature is not possible. 1.3.6 Density The two properties — mass and volume discussed above are related as follows: Mass Density = Volume Density of a substance is its amount of mass per unit volume. So, SI units of density can be obtained as follows: SI unit of density = kg = or kg m–3 m3 This unit is quite large and a chemist often expresses density in g cm–3, where mass is expressed in gram and volume is expressed Fig. 1.8 Thermometers using differentin cm3. Density of a substance tells us about temperature scaleshow closely its particles are packed. If density is more, it means particles are more closely 1.4 U Many a time in the study of chemistry, one 1.3.7 Temperature has to deal with experimental data as well as There are three common scales to measure theoretical calculations. There are meaningful temperature — °C (degree celsius), °F (degree ways to handle the numbers conveniently and Reprint 2025-26 Some Basic Concepts of Chemistry 11 present the data realistically with certainty to the extent possible. These ideas are discussed Reference Standard below in detail. After defining a unit of measurement such as the kilogram or the metre, scientists agreed 1.4.1 Scientific Notation on reference standards that make it possible As chemistry is the study of atoms and to calibrate all measuring devices. For getting molecules, which have extremely low masses reliable measurements, all devices such as and are present in extremely large numbers, metre sticks and analytical balances have a chemist has to deal with numbers as large been calibrated by their manufacturers to give correct readings. However, each of as 602, 200,000,000,000,000,000,000 for the these devices is standardised or calibrated molecules of 2 g of hydrogen gas or as small as against some reference. The mass standard 0.00000000000000000000000166 g mass of is the kilogram since 1889. It has been a H atom. Similarly, other constants such as defined as the mass of platinum-iridium Planck’s constant, speed of light, charges on (Pt-Ir) cylinder that is stored in an airtight particles, etc., involve numbers of the above jar at International Bureau of Weights magnitude. and Measures in Sevres, France. Pt-Ir was chosen for this standard because it is highly It may look funny for a moment to resistant to chemical attack and its mass write or count numbers involving so many will not change for an extremely long time. zeros but it offers a real challenge to do Scientists are in search of a new simple mathematical operations of addition, standard for mass. This is being attempted subtraction, multiplication or division with through accurate determination of Avogadro such numbers. You can write any two constant. Work on this new standard focuses numbers of the above type and try any one on ways to measure accurately the number of the operations you like to accept as a of atoms in a well-defined mass of sample. challenge, and then, you will really appreciate One such method, which uses X-rays to the difficulty in handling such numbers. determine the atomic density of a crystal of ultrapure silicon, has an accuracy of about This problem is solved by using scientific 1 part in 106 but has not yet been adopted to notation for such numbers, i.e., exponential serve as a standard. There are other methods notation in which any number can be but none of them are presently adequate to represented in the form N × 10n, where n is an replace the Pt-Ir cylinder. No doubt, changes exponent having positive or negative values are expected within this decade. and N is a number (called digit term) which The metre was originally defined as the varies between 1.000... and 9.999.... length between two marks on a Pt-Ir bar kept at a temperature of 0°C (273.15 K). In Thus, we can write 232.508 as 1960 the length of the metre was defined as 2.32508 × 102 in scientific notation. Note that 1.65076373 × 106 times the wavelength of while writing it, the decimal had to be moved light emitted by a krypton laser. Although to the left by two places and same is the this was a cumbersome number, it preserved exponent (2) of 10 in the scientific notation. the length of the metre at its agreed value. Similarly, 0.00016 can be written as The metre was redefined in 1983 by CGPM 1.6 × 10–4. Here, the decimal has to be as the length of path travelled by light in vacuum during a time interval of 1/299 792 moved four places to the right and (–4) is the 458 of a second. Similar to the length and exponent in the scientific notation. the mass, there are reference standards for While performing mathematical operations other physical quantities. on numbers expressed in scientific notations, the following points are to be kept in mind. Reprint 2025-26 12 chemistry Multiplication and Division mass obtained by an analytical balance is These two operations follow the same rules slightly higher than the mass obtained by which are there for exponential numbers, i.e. using a platform balance. Therefore, digit 4 placed after decimal in the measurement by 10 5 .6  10 6 .9  10  5 8 5  8 platform balance is u 13 The u = 38.64  1013 mentioning the number of significant figures. Significant figures are meaningful = 3.864  1014 digits which are known with certainty plus 10 2  6   2 .5  10 6 9 .8  10 2     =  9 .8  2 .5    one which is estimated or u 2 6 u = 24.50  10 8 write a result as 11.2 mL, we say the 11 is = 2.450  10 7 certain and 2 is u 3 would be +1 in the last digit. Unless otherwise 2 .7  10 3 4 7 10 =0.4909  1 0 stated, an u =4.909  10 8 There are certain rules for determining the number of significant figures. These are Addition and Subtraction stated below: For these two operations, first the numbers are (1) All non-zero digits are significant. For written in such a way that they have the same example in 285 cm, there are three exponent. After that, the coefficients (digit significant figures and in 0.25 mL, there terms) are added or subtracted as the case are two significant figures. may be. (2) Zeros preceding to first non-zero digit Thus, for adding 6.65×104 and 8.95×103, are not significant. Such zero indicates exponent is made same for both the numbers. the position of decimal point. Thus, Thus, we get (6.65×104) + (0.895×104) 0.03 has one significant figure and 0.0052 has two significant figures.Then, these numbers can be added as follows (6.65 + 0.895)×104 = 7.545×104 (3) Zeros between two non-zero digits are significant. Thus, 2.005 has fourSimilarly, the subtraction of two numbers can significant figures.be done as shown below: (4) Zeros at the end or right of a number(2.5 × 10–2) – (4.8 ×10–3) are significant, provided they are on = (2.5 × 10–2) – (0.48 × 10–2) the right side of the decimal point. For = (2.5 – 0.48)×10–2 = 2.02 × 10–2 example, 0.200 g has three significant figures. But, if otherwise, the terminal 1.4.2 Significant Figures zeros are not significant if there is no Every experimental measurement has decimal point. For example, 100 has some amount of uwith it because of limitation of measuring three significant figures and 100.0 has instrument and the skill of the person making four significant figures. Such numbers the measurement. For example, mass of an are better represented in scientific object is obtained using a platform balance notation. We can express the number and it comes out to be 9.4g. On measuring 100 as 1×102 for one significant figure, the mass of this object on an analytical 1.0×102 for two significant figures and balance, the mass obtained is 9.4213g. The 1.00×102 for three significant figures. Reprint 2025-26 Some Basic Concepts of Chemistry 13 (5) Counting the numbers of object, for Here, 18.0 has only one digit after the decimal example, 2 balls or 20 eggs, have infinite point and the result should be reported only significant figures as these are exact up to one digit after the decimal point, which numbers and can be represented by is 31.1. writing infinite number of zeros after placing a decimal i.e., 2 = 2.000000 or Multiplication and Division of 20 = 20.000000. Significant Figures In numbers written in scientific notation, In these operations, the result must be all digits are significant e.g., 4.01×102 has reported with no more significant figures as three significant figures, and 8.256×10–3 has in the measurement with the few significant four significant figures. figures. However, one would always like the results to be precise and accurate. Precision and 2.5×1.25 = 3.125 accuracy are often referred to while we talk Since 2.5 has two significant figures, about the measurement. the result should not have more than two Precision refers to the closeness of significant figures, thus, it is 3.1. various measurements for the same quantity. While limiting the result to the required However, accuracy is the agreement of a number of significant figures as done in the particular value to the true value of the above mathematical operation, one has to result. For example, if the true value for a keep in mind the following points for rounding result is 2.00 g and student ‘A’ takes two off the numbers measurements and reports the results as 1.95 1. If the rightmost digit to be removed is g and 1.93 g. These values are precise as they more than 5, the preceding number is are close to each other but are not accurate. increased by one. For example, 1.386. If Another student ‘B’ repeats the experiment we have to remove 6, we have to round and obtains 1.94 g and 2.05 g as the results it to 1.39. for two measurements. These observations 2. If the rightmost digit to be removed is are neither precise nor accurate. When the less than 5, the preceding number is not third student ‘C’ repeats these measurements changed. For example, 4.334 if 4 is to and reports 2.01 g and 1.99 g as the result, be removed, then the result is rounded these values are both precise and accurate. upto 4.33. This can be more clearly understood from the 3. If the rightmost digit to be removed is 5, data given in Table 1.4. then the preceding number is not changed Table 1.4 Data to Illustrate Precision if it is an even number but it is increased and Accuracy by one if it is an odd number. For example, Measurements/g if 6.35 is to be rounded by removing 5, we 1 2 Average (g) have to increase 3 to 4 giving 6.4 as the result. However, if 6.25 is to be rounded Student A 1.95 1.93 1.940 off it is rounded off to 6.2. Student B 1.94 2.05 1.995 1.4.3 Dimensional Analysis Student C 2.01 1.99 2.000 Often while calculating, there is a need to Addition and Subtraction of convert units from one system to the other. Significant Figures The method used to accomplish this is called factor label method or unit factor methodThe result cannot have more digits to the right or dimensional analysis. This is illustratedof the decimal point than either of the original below.numbers. 12.11 18.0 Example 1.012 A piece of metal is 3 inch (represented by in) 31.122 long. What is its length in cm? Reprint 2025-26 14 chemistry Solution The above is multiplied by the unit factor We know that 1 in = 2.54 cm 3 1 m 3 2 m 3 3 3 2  1000 cm  6 3  3  2  10 m From this equivalence, we can write 10 cm 10 1 in 2 .54 cm Example = 1 = 2 .54 cm 1 in How many seconds are there in 2 days? Solution 1 in 2 .54 cm Here, we know 1 day = 24 hours (h) Thus, equals 1 and 2 .54 cm 1 in 1 day 24 h or = 1 = 24 h 1 day also equals 1. Both of these are called unit then, 1h = 60 min factors. If some number is multiplied by these 1 h 60 min unit factors (i.e., 1), it will not be affected or = 1 = 60 min 1 h otherwise. so, for converting 2 days to seconds, Say, the 3 in given above is multiplied by the unit factor. So, i.e., 2 days – – – – – – = – – – seconds 2 .54 cm The unit factors can be multiplied in3 in = 3 in × = 3 × 2.54 cm = 7.62 cm 1 in series in one step only as follows: 24 h 60 min 60 s Now, the unit factor by which multiplication 2 day × × × 1 day 1 h 1 min 2 .54 cm is to be done is that unit factor ( in 1 in = 2 × 24 × 60 × 60 s = 172800 sthe above case) which gives the desired units i.e., the numerator should have that part 1.5 Laws of Chemical which is required in the desired result. Combinations It should also be noted in the above The combination of elements example that units can be handled just like to form compounds is other numerical part. It can be cancelled, governed by the following five divided, multiplied, squared, etc. Let us study basic laws. Antoine Lavoisier one more example. (1743–1794) 1.5.1 Law of Conservation Example of Mass A jug contains 2L of milk. Calculate the This law was put forth by Antoine Lavoisier volume of the milk in m3. in 1789. He performed careful experimental Solution studies for combustion reactions and reached Since 1 L = 1000 cm3 to the conclusion that in all physical and and 1m = 100 cm, which gives chemical changes, there is no net change in mass duting the process. Hence, he reached 1 m 100 cm = 1 = to the conclusion that matter can neither be 100 cm 1 m created nor destroyed. This is called ‘Law To get m3 from the above unit factors, the of Conservation of Mass’. This law formed first unit factor is taken and it is cubed. the basis for several later developments in 3 3 chemistry. Infact, this was the result of exact  1 m  1 m 3  6 3 1 1 measurement of masses of reactants and  100 cm  10 cm products, and carefully planned experiments Now 2 L = 2 ×1000 cm3 performed by Lavoisier. Reprint 2025-26 Some Basic Concepts of Chemistry 15 1.5.2 Law of Definite Proportions are produced in a chemical reaction they do so in aThis law was given by, a simple ratio by volume,French chemist, Joseph provided all gases are atProust. He stated that a given the same temperature andcompound always contains pressure.exactly the same proportion of elements by weight. Thus, 100 mL of hydrogen Joseph Louis combine with 50 mL of Gay Lussac Proust worked with two Joseph Proust oxygen to give 100 mL ofsamples of cupric carbonate (1754–1826) water vapour.— one of which was of natural origin and the other was synthetic. He found Hydrogen + Oxygen → Water that the composition of elements present in it 100 mL 50 mL 100 mL was same for both the samples as shown below: Thus, the volumes of hydrogen and % of % of % of oxygen which combine (i.e., 100 mL and copper carbon oxygen 50 mL) bear a simple ratio of 2:1. Natural Sample 51.35 9.74 38.91 Gay Lussac’s discovery of integer ratio in volume relationship is actually the law of Synthetic Sample 51.35 9.74 38.91 definite proportions by volume. The law of Thus, he concluded that irrespective of the definite proportions, stated earlier, was with source, a given compound always contains respect to mass. The Gay Lussac’s law was same elements combined together in the same explained properly by the work of Avogadro proportion by mass. The validity of this law in 1811. has been confirmed by various experiments. It is sometimes also referred to as Law of 1.5.5 Avogadro’s Law Definite Composition. In 1811, Avogadro proposed that equal volumes of all gases at the same temperature1.5.3 Law of Multiple Proportions and pressure should contain equal number This law was proposed by Dalton in 1803. of molecules. Avogadro made a distinction According to this law, if two elements can between atoms and molecules which is combine to form more than one compound, quite understandable in present times. If the masses of one element that combine with we consider again the reaction of hydrogen a fixed mass of the other element, are in the and oxygen to produce water, we see that ratio of small whole numbers. two volumes of hydrogen combine with one For example, hydrogen combines with volume of oxygen to give two volumes of water oxygen to form two compounds, namely, water without leaving any unreacted oxygen. and hydrogen peroxide. Note that in the Fig. 1.9 (Page 16) each Hydrogen + Oxygen → Water box contains equal number of 2g 16g 18g molecules. In fact, Avogadro Hydrogen + Oxygen → Hydrogen Peroxide could explain the above result by considering the molecules 2g 32g 34g to be polyatomic. If hydrogenHere, the masses of oxygen (i.e., 16 g and 32 g), and oxygen were consideredwhich combine with a fixed mass of hydrogen as diatomic as recognised(2g) bear a simple ratio, i.e., 16:32 or 1: 2. now, then the above results Lorenzo Romano 1.5.4 Gay Lussac’s Law of Gaseous are easily understandable. Amedeo Carlo Volumes However, Dalton and others Avogadro di Quareqa edi This law was given by Gay Lussac in 1808. believed at that time that Carreto He observed that when gases combine or atoms of the same kind (1776–1856) Reprint 2025-26 16 chemistry Fig. 1.9 Two volumes of hydrogen react with one volume of oxygen to give two volumes of water vapour cannot combine and molecules of oxygen or Dalton’s theory could explain the laws hydrogen containing two atoms did not exist. of chemical combination. However, it could Avogadro’s proposal was published in the not explain the laws of gaseous volumes. It French Journal de Physique. In spite of being could not provide the reason for combining correct, it did not gain much support. of atoms, which was answered later by other After about 50 years, in 1860, the first scientists. international conference on chemistry was 1.7 Atomic and Molecular Massesheld in Karlsruhe, Germany, to resolve After having some idea about the termsvarious ideas. At the meeting, Stanislao atoms and molecules, it is appropriate hereCannizaro presented a sketch of a course of to understand what do we mean by atomicchemical philosophy, which emphasised on and molecular masses.the importance of Avogadro’s work.

4.16 Describe the preparation of potassium permanganate. How does the acidified permanganate solution react with (i) iron(II) ions (ii) SO2 and (iii) oxalic acid? Write the ionic equations for the reactions. 4.17 For M2+/M and M3+/M 2+ systems the E o values for some metals are as follows: Cr2+/Cr -0.9V Cr3/Cr2+ -0.4 V Mn 2+/Mn -1.2V Mn3+/Mn2+ +1.5 V Fe2+/Fe -0.4V Fe3+/Fe2+ +0.8 V Use this data to comment upon: (i) the stability of Fe3+ in acid solution as compared to that of Cr3+ or Mn3+ and (ii) the ease with which iron can be oxidised as compared to a similar process for either chromium or manganese metal. 4.18 Predict which of the following will be coloured in aqueous solution? Ti 3+, V3+, Cu+, Sc3+, Mn 2+, Fe3+ and Co 2+. Give reasons for each. 4.19 Compare the stability of +2 oxidation state for the elements of the first transition series. 4.20 Compare the chemistry of actinoids with that of the lanthanoids with special reference to: (i) electronic configuration (iii) oxidation state (ii) atomic and ionic sizes and (iv) chemical reactivity. 4.21 How would you account for the following: (i) Of the d4 species, Cr2+ is strongly reducing while manganese(III) is strongly oxidising. (ii) Cobalt(II) is stable in aqueous solution but in the presence of complexing reagents it is easily oxidised. (iii) The d1 configuration is very unstable in ions. 4.22 What is meant by ‘disproportionation’? Give two examples of disproportionation reaction in aqueous solution. 4.23 Which metal in the first series of transition metals exhibits +1 oxidation state most frequently and why? 4.24 Calculate the number of unpaired electrons in the following gaseous ions: Mn3+, Cr3+, V3+ and Ti3+. Which one of these is the most stable in aqueous solution? 4.25 Give examples and suggest reasons for the following features of the transition metal chemistry: (i) The lowest oxide of transition metal is basic, the highest is amphoteric/acidic. (ii) A transition metal exhibits highest oxidation state in oxides and fluorides. (iii) The highest oxidation state is exhibited in oxoanions of a metal. 4.26 Indicate the steps in the preparation of: (i) K2Cr2O7 from chromite ore. (ii) KMnO4 from pyrolusite ore. 4.27 What are alloys? Name an important alloy which contains some of the lanthanoid metals. Mention its uses. 4.28 What are inner transition elements? Decide which of the following atomic numbers are the atomic numbers of the inner transition elements : 29, 59, 74, 95, 102, 104. 4.29 The chemistry of the actinoid elements is not so smooth as that of the lanthanoids. Justify this statement by giving some examples from the oxidation state of these elements. 4.30 Which is the last element in the series of the actinoids? Write the electronic configuration of this element. Comment on the possible oxidation state of this element. Chemistry 116 Reprint 2025-26 4.31 Use Hund’s rule to derive the electronic configuration of Ce3+ ion, and calculate its magnetic moment on the basis of ‘spin-only’ formula. 4.32 Name the members of the lanthanoid series which exhibit +4 oxidation states and those which exhibit +2 oxidation states. Try to correlate this type of behaviour with the electronic configurations of these elements. 4.33 Compare the chemistry of the actinoids with that of lanthanoids with reference to: (i) electronic configuration (ii) oxidation states and (iii) chemical reactivity. 4.34 Write the electronic configurations of the elements with the atomic numbers 61, 91, 101, and 109. 4.35 Compare the general characteristics of the first series of the transition metals with those of the second and third series metals in the respective vertical columns. Give special emphasis on the following points: (i) electronic configurations (ii) oxidation states (iii) ionisation enthalpies and (iv) atomic sizes. 4.36 Write down the number of 3d electrons in each of the following ions: Ti 2+, V 2+, Cr3+, Mn 2+, Fe2+, Fe3+, Co2+, Ni2+ and Cu2+. Indicate how would you expect the five 3d orbitals to be occupied for these hydrated ions (octahedral). 4.37 Comment on the statement that elements of the first transition series possess many properties different from those of heavier transition elements. 4.38 What can be inferred from the magnetic moment values of the following complex species ? Example Magnetic Moment (BM) K4[Mn(CN)6) 2.2 [Fe(H2O)6]2+ 5.3 K2[MnCl4] 5.9 Answers to Some Intext Questions 4.1 Silver (Z = 47) can exhibit +2 oxidation state wherein it will have incompletely filled d-orbitals (4d), hence a transition element. 4.2 In the formation of metallic bonds, no eletrons from 3d-orbitals are involved in case of zinc, while in all other metals of the 3d series, electrons from the d-orbitals are always involved in the formation of metallic bonds. 4.3 Manganese (Z = 25), as its atom has the maximum number of unpaired electrons. 4.5 Irregular variation of ionisation enthalpies is mainly attributed to varying degree of stability of different 3d-configurations (e.g., d 0, d 5, d 10 are exceptionally stable). 4.6 Because of small size and high electronegativity oxygen or fluorine can oxidise the metal to its highest oxidation state. 4.7 Cr 2+ is stronger reducing agent than Fe 2+ Reason: d 4  d 3 occurs in case of Cr 2+ to Cr 3+ But d 6  d 5 occurs in case of Fe2+ to Fe 3+ In a medium (like water) d 3 is more stable as compared to d 5 (see CFSE) 4.9 Cu + in aqueous solution underoes disproportionation, i.e., 2Cu +(aq) ® Cu 2+(aq) + Cu(s) The E0 value for this is favourable.

6.10 ACIDS, BASES AND SALTS a very high dielectric constant of 80. Thus, Acids, bases and salts find widespread when sodium chloride is dissolved in water, occurrence in nature. Hydrochloric acid the electrostatic interactions are reduced by present in the gastric juice is secreted by the a factor of 80 and this facilitates the ions to lining of our stomach in a significant amount move freely in the solution. Also, they are of 1.2-1.5 L/day and is essential for digestive well-separated due to hydration with water processes. Acetic acid is known to be the main molecules. constituent of vinegar. Lemon and orange juices contain citric and ascorbic acids, and tartaric acid is found in tamarind paste. As most of the acids taste sour, the word “acid” has been derived from a latin word “acidus” meaning sour. Acids are known to turn blue litmus paper into red and liberate dihydrogen on reacting with some metals. Similarly, bases are known to turn red litmus paper blue, taste bitter and feel soapy. A common example of a base is washing soda used for washing purposes. When acids and bases are mixed in the right proportion they react with each Fig.6.10 Dissolution of sodium chloride in water.other to give salts. Some commonly known Na+ and Cl– ions are stablised by their examples of salts are sodium chloride, barium hydration with polar water molecules. sulphate, sodium nitrate. Sodium chloride (common salt) is an important component of Comparing, the ionization of hydrochloric our diet and is formed by reaction between acid with that of acetic acid in water we find hydrochloric acid and sodium hydroxide. It that though both of them are polar covalent Faraday was born near London into a family of very limited means. At the age of 14 he was an apprentice to a kind bookbinder who allowed Faraday to read the books he was binding. Through a fortunate chance he became laboratory assistant to Davy, and during 1813-4, Faraday accompanied him to the Continent. During this trip he gained much from the experience of coming into contact with many of the leading scientists of the time. In 1825, he succeeded Davy as Director of the Royal Institution laboratories, and in 1833 he also became the first Fullerian Professor of Chemistry. Faraday’s first important work was on analytical chemistry. After 1821 Michael Faraday much of his work was on electricity and magnetism and different electromagnetic (1791–1867) phenomena. His ideas have led to the establishment of modern field theory. He discovered his two laws of electrolysis in 1834. Faraday was a very modest and kind hearted person. He declined all honours and avoided scientific controversies. He preferred to work alone and never had any assistant. He disseminated science in a variety of ways including his Friday evening discourses, which he founded at the Royal Institution. He has been very famous for his Christmas lecture on the ‘Chemical History of a Candle’. He published nearly 450 scientific papers. Reprint 2025-26 190 chemistry molecules, former is completely ionized into its constituent ions, while the latter is only Hydronium and Hydroxyl Ions partially ionized (< 5%). The extent to which Hydrogen ion by itself is a bare proton with very ionization occurs depends upon the strength small size (~10–15 m radius) and intense electric of the bond and the extent of solvation field, binds itself with the water molecule at of ions produced. The terms dissociation one of the two available lone pairs on it giving and ionization have earlier been used with H3O+. This species has been detected in many different meaning. Dissociation refers to the compounds (e.g., H3O+Cl–) in the solid state. In process of separation of ions in water already aqueous solution the hydronium ion is further existing as such in the solid state of the solute, hydrated to give species like H5O2+, H7O3 + and as in sodium chloride. On the other hand, H9O4+. Similarly the hydroxyl ion is hydrated to give several ionic species like H3O2–, H5O3–ionization corresponds to a process in which – and H7O4 etc.a neutral molecule splits into charged ions in the solution. Here, we shall not distinguish between the two and use the two terms interchangeably. 6.10.1 Arrhenius Concept of Acids and Bases According to Arrhenius theory, acids are H9O4+ substances that dissociates in water to give hydrogen ions H+(aq) and bases are 6.10.2 The Brönsted-Lowry Acids and substances that produce hydroxyl ions Bases OH –(aq). The ionization of an acid HX (aq) can The Danish chemist, Johannes Brönsted and be represented by the following equations: the English chemist, Thomas M. Lowry gave HX (aq) → H+(aq) + X– (aq) a more general definition of acids and bases. or According to Brönsted-Lowry theory, acid HX(aq) + H2O(l) → H3O+(aq) + X –(aq) is a substance that is capable of donating a hydrogen ion H+ and bases are substances A bare proton, H+ is very reactive and capable of accepting a hydrogen ion, H+. Incannot exist freely in aqueous solutions. short, acids are proton donors and bases areThus, it bonds to the oxygen atom of a solvent proton acceptors.water molecule to give trigonal pyramidal Consider the example of dissolution of NH3hydronium ion, H3O+ {[H (H2O)]+} (see box). In in H2O represented by the following equation:this chapter we shall use H+(aq) and H3O+(aq) interchangeably to mean the same i.e., a hydrated proton. Similarly, a base molecule like MOH ionizes in aqueous solution according to the equation: MOH(aq) → M+(aq) + OH–(aq) The hydroxyl ion also exists in the hydrated form in the aqueous solution. Arrhenius concept of acid and base, however, suffers The basic solution is formed due to the from the limitation of being applicable only to presence of hydroxyl ions. In this reaction, aqueous solutions and also, does not account water molecule acts as proton donor and for the basicity of substances like, ammonia ammonia molecule acts as proton acceptor which do not possess a hydroxyl group. and are thus, called Lowry-Brönsted acid and Reprint 2025-26 EQUILIBRIUM 191 Arrhenius was born near Uppsala, Sweden. He presented his thesis, on the conductivities of electrolyte solutions, to the University of Uppsala in 1884. For the next five years he travelled extensively and visited a number of research centers in Europe. In 1895 he was appointed professor of physics at the newly formed University of Stockholm, serving its rector from 1897 to 1902. From 1905 until his death he was Director of physical chemistry at the Nobel Institute in Stockholm. He continued to work for many years on electrolytic solutions. In 1899 he discussed the temperature dependence of reaction rates on the basis of an equation, now usually known as Arrhenius equation. He worked in a variety of fields, and made important contributions to Svante Arrhenius immunochemistry, cosmology, the origin of life, and the causes of ice age. He was (1859-1927) the first to discuss the ‘green house effect’ calling by that name. He received Nobel Prize in Chemistry in 1903 for his theory of electrolytic dissociation and its use in the development of chemistry. base, respectively. In the reverse reaction, in case of ammonia it acts as an acid by H+ is transferred from NH4+ to OH–. In this donating a proton. case, NH4+ acts as a Bronsted acid while Problem 6.12OH– acted as a Brönsted base. The acid-base What will be the conjugate bases for thepair that differs only by one proton is called following Brönsted acids: HF, H2SO4 anda conjugate acid-base pair. Therefore, OH– – HCO3 ?is called the conjugate base of an acid H2O and NH4+ is called conjugate acid of the base Solution NH3. If Brönsted acid is a strong acid then The conjugate bases should have one its conjugate base is a weak base and vice- proton less in each case and therefore the versa. It may be noted that conjugate acid corresponding conjugate bases are: F –, has one extra proton and each conjugate base HSO4– and CO32– respectively. has one less proton. Problem 6.13 Consider the example of ionization of Write the conjugate acids for the following hydrochloric acid in water. HCl(aq) acts as Brönsted bases: NH2–, NH3 and HCOO–. an acid by donating a proton to H2O molecule which acts as a base. Solution The conjugate acid should have one extra proton in each case and therefore the corresponding conjugate acids are: NH3, NH4+ and HCOOH respectively. Problem 6.14 The species: H2O, HCO3–, HSO4– and NH3 can act both as Bronsted acids and bases. For each case give the corresponding conjugate It can be seen in the above equation, that acid and conjugate base. water acts as a base because it accepts the Solution proton. The species H3O+ is produced when The answer is given in the following Table:water accepts a proton from HCl. Therefore, Cl– is a conjugate base of HCl and HCl is the Species Conjugate Conjugate conjugate acid of base Cl –. Similarly, H2O is acid base a conjugate base of an acid H3O+ and H3O+ is H2O H3O+ OH–a conjugate acid of base H2O. – 2– HCO3 H2CO3 CO3 It is interesting to observe the dual role – 2– HSO4 H2SO4 SO4of water as an acid and a base. In case of + – reaction with HCl water acts as a base while NH3 NH4 NH2 Reprint 2025-26 192 chemistry 6.10.3 Lewis Acids and Bases (HClO4), hydrochloric acid (HCl), hydrobromic G.N. Lewis in 1923 defined an acid as a acid (HBr), hyrdoiodic acid (HI), nitric acid species which accepts electron pair and base (HNO3) and sulphuric acid (H2SO4) are termed which donates an electron pair. As far as bases strong because they are almost completely are concerned, there is not much difference dissociated into their constituent ions in an between Brönsted-Lowry and Lewis concepts, aqueous medium, thereby acting as proton as the base provides a lone pair in both the (H+) donors. Similarly, strong bases like cases. However, in Lewis concept many lithium hydroxide (LiOH), sodium hydroxide acids do not have proton. A typical example (NaOH), potassium hydroxide (KOH), caesium is reaction of electron deficient species BF3 hydroxide (CsOH) and barium hydroxide with NH3. Ba(OH)2 are almost completely dissociated into ions in an aqueous medium giving BF3 does not have a proton but still acts hydroxyl ions, OH–. According to Arrhenius as an acid and reacts with NH3 by accepting concept they are strong acids and bases asits lone pair of electrons. The reaction can be they are able to completely dissociate andrepresented by, produce H3O+ and OH– ions respectively in BF3 + :NH3 → BF3:NH3 the medium. Alternatively, the strength of an Electron deficient species like AlCl3, Co3+, acid or base may also be gauged in terms of Mg2+, etc. can act as Lewis acids while species Brönsted-Lowry concept of acids and bases, like H2O, NH3, OH– etc. which can donate a wherein a strong acid means a good proton pair of electrons, can act as Lewis bases. donor and a strong base implies a good proton acceptor. Consider, the acid-base dissociation Problem 6.15 equilibrium of a weak acid HA, Classify the following species into Lewis HA(aq) + H2O(l)   H3O+(aq) + A–(aq) acids and Lewis bases and show how these conjugate conjugate act as such: acid base acid base (a) HO– (b) F – (c) H+ (d) BCl3 In section 6.10.2 we saw that acid (or Solution base) dissociation equilibrium is dynamic involving a transfer of proton in forward and (a) Hydroxyl ion is a Lewis base as it can reverse directions. Now, the question arises donate an electron lone pair (:OH– ). that if the equilibrium is dynamic then with (b) Flouride ion acts as a Lewis base passage of time which direction is favoured? as it can donate any one of its four What is the driving force behind it? In order electron lone pairs. to answer these questions we shall deal (c) A proton is a Lewis acid as it can into the issue of comparing the strengths accept a lone pair of electrons from of the two acids (or bases) involved in the bases like hydroxyl ion and fluoride dissociation equilibrium. Consider the two ion. acids HA and H3O+ present in the above mentioned acid-dissociation equilibrium. (d) BCl3 acts as a Lewis acid as it can We have to see which amongst them is a accept a lone pair of electrons from stronger proton donor. Whichever exceeds species like ammonia or amine in its tendency of donating a proton over the molecules. other shall be termed as the stronger acid