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Matter and Energy Mat ter is ev ery thing that oc cu pies space and has mass (or mo ment of in er tia). In the uni verse, mat ter ex ists in dif fer ent forms. On the ba sis of phys i cal state, mat ter may ex ist as solid, liq uid or gas. These states are com monly present on earth in the pro cesses tak ing place on earth. Three additional forms of matter are also known to exist in thermonuclear processes. They are plasma, Bose-Einstein condensate and Fermionic condensate. Measurements in Chemistry Measurements provide the macroscopic informations that are the basis of several laws, hypotheses or theories, given to explain the behaviour of matter and energy on both macroscopic and microscopic levels. (a) Unit (i) A unit is the standard of comparison for different measurements. (ii) Units are essential for presenting a measured quantity correctly. (b) Metric System (i) Many properties of matter are quantitative, that is, they are associated with numbers. When a number represents a measured quantity, the units of that quantity must always be specified. (ii) The metric system ( the more logical decimal system, first developed in France during the late eighteenth century (1791),) is used as the system of measurements in most of the countries throughout the world. (c) SI system In this system, there are seven fundamental scientific quantities. The SI units has seven basic units. Length (m) Mass (kg) Time (s) Electric current (A) Temperture (K) Luminous intensity (cd) Some prefix used in SI system 10 −12 (pico) 10 −9 (nano) 10 −6 (micro) 10 −3 (milli) 10 −2 (centi) 10 3 (kilo) 10 6 (mega) 10 9 (giga) 1. Mass and weight : (i) Mass of a system represents the amount of matter present in a system, while weight represents the force that gravity exert on that system (or object). (ii) These terms are often used interchangeably, although strictly speaking, they are different quantities. (iii) The SI base unit of mass is kilogram (kg) but in chemistry unit gram (g) is more convenient and frequently used. 1 kg 1000g 1 10 g 3 = = × Mass of a system is constant while weight varies from place to place. 2. Volume: Volume of a body = (length) 3 Because in SI system, unit of length in metre, so SI unit of volume is m 3 . Other smaller non SI units are dm 3 and cm 3 . 1 cm (1 10 m) 1 10 m 3 2 3 6 3 = × = × − − 1 dm (1 10 m) 1 10 m 3 1 3 3 3 = × = × − − Mixtures Pure substances Heterogeneous mixtures Homogeneous mixtures Elements Compounds Matter Some Basic Concepts of Chemistry 1
Significant Figures : Uncertainty in Measurements (i) All the measurements have some degree of uncertainty. (ii) A counted quantity is exact provided that the objects counted do not change while they are being counted. (iii) Defined quantities (like-1km =1000m, 1 ft=12 inches etc.) and ex po nen tial part in ex po nen tial no ta tion (e. g., 10 5 in 4.21 10 ) 5 × are also ex act. (iv) On the other hand, mea sured quan ti ties are not ex act and have some de gree of un cer tainty as so ci ated with them. Any mea sure ment has an un cer tainty of at least one unit in the last digit of the re ported value. Thus in a mea sure ment, last digit is un cer tain and oth ers are cer tain, e.g., 5.2g has un cer tainty of 1 g, 5.22g has un cer tainty of 0.01 g, 162 m has un cer tainty of 1 m etc. Rules for Determination of Significant Figures 1. All non-zero digits are significant. For example: 3.14 has 3 significant figures; 2 certain+1uncertain. 05.153 has 4 significant figures ; 3 certain+1 uncertain. 2. The zeros to the right of the decimal point are significant. For example : 3.0 has 2 significant figures ; 1 certain+1 uncertain digit. 3.10 has 3 significant figures; 2 certain+1 uncertain digit. 3. The zero to the left of the first non-zero digit in a number are not significant. They merely represent the position of a decimal point. For example: 0.02 g has 1 significant figure 0.002 g has 1 significant figure 4. The zeros between two non-zero digits are also significant. For example: 6.01 has 3 significant figures; 2 certain and last digit uncertain. 6.001 has 4 significant figures; 3 certain and last digit uncertain. 5. Leading zero before the decimal point are never significant. For example : 0.618 has 3 significant figures; 2 certain and last digit uncertain. 6. When a number ends in zeros that are not to the right of decimal point, the zeros are not necessarily significant. For example: 180 cm has two or three significant figures, 18600 g has three or four or five significant figures. This ambiguity has been removed by using exponential notation. 7. Exponential notation. The use of exponential notation avoids the potential ambiguity of whether the zeros at the end of a number are significant (rule 6) or not. For example, a mass of 19400 g can be written in exponential notations showing three, four or five significant figures as: 1 94 10 4 . × ( 3 significant figures) 1 940 10 4 . × ( 4 significant figures) 1 9400 10 4 . × ( 5 significant figures) Thus, all significant figures lying before the exponent and the exponential term does not add to number of significant figures. Similarly significant numbers in numerical value of Avogadro number (6.023 10 ) 23 × are four and in Planck's constant (6.626 10 ) 34 × − Js are four. 8. Exact numbers can be treated as if they have an infinite numbers of significant figures. This rule ap plies to many def i ni tions be tween units. Thus when we say 1 foot has 12 inches, the num ber 12 is ex act and we need not worry about the num ber of sig nif i cant fig ures in it. 9. Rounding off the uncertain digit: The rounding off of uncertain digit of significant figures is made as follows: (a) If the digit to be rounded off is more than 5, the preceding number is increased by one, e.g., 2 16 2 2. . , ⇒ 3 58 3 6. . ⇒ (b) If the digit to be rounded off is less than 5, the preceding number is retained, e. g., 2.14 ⇒ 2.1, 4.13 ⇒ 4.1 (c) If the digit to be rounded off is equal to 5, the preceding number is not changed if it is even and increased by one, if it is odd, e.g., 3.25 ⇒ 3 2. (Preceding digit 2 is even) 2.35 ⇒ 2.4 (Preceding digit 3 is odd) CHEMICAL STOICHIOMETRY ‘‘Chemical stoichiometry describes the quantitative relationships that exist between substances undergoing chemical changes.’’ To make the study of this wide chapter easy and systematic, it has been divided into four sub-parts : A. Laws of chemical combinations B. The mole concept C. Concentration units and their calculations D. Stoichiometric calculations Laws of Chemical Combinations 1. Law of Conservation of Mass (Lomonosov, 1744) (a) According to this law, ‘‘there is no detectable change in the total quantity of matter present in a system when it undergoes a chemical or physical change.’’ or ‘‘matter can neither be created nor destroyed during the course of a chemical reaction, however its physical or chemical nature may change.” 2 Crash Course Chemistry for NEET
(b) On the basis of this law, we may conclude that for a reaction, Total mass of reactants before the reaction = Total mass of products after the reaction. (c) After the origin of energy-mass relationship E = mc 2 by Einstein, this law needed a slight modification. 2. Law of Constant (or Definite) Proportions (a) According to this law, ‘‘a chemical compound always contains the same elements combined together in the same proportion by mass. or A compound always contains exactly the same proportions of elements by mass.’’ (b) For example, compound CO2 can be formed by either of the processes : By heating CaCO3 : CaCO3 CaO CO2 ∆→ + 3. Law of Multiple Proportions (Dalton 1803) (a) According to this law, ‘‘when two elements combine with each other to form two or more chemical compounds then different amounts of one element which react with definite amount of second element, bear a simple whole numbers ratio’’. (b) Let two elements A and B combine to form two compounds. In these compounds different amounts of B reacts with same amount of A, then these different amounts of B bear a simple whole number ratio with each other. For example Combination of C and O may form CO and CO2 . In CO, ratio C : O is 12 : 16, while in CO2 it is 12 : 32. Thus ratio of O in CO and CO2 is 16 : 32 or 1 : 2, i. e., a whole number ratio. (c) However, after the discovery of isotopes, some modifications were made in the explanations. 4. Law of Equiv a lent Pro por tions or Law of Re cip ro cal Pro por tions (Rich ter 1792-94) (a) According to this law, ‘‘the ratio of masses of two substance A and B which separately reacts with a fixed mass of other substance C is either simple multiple or fraction of the ratio of masses of A and B when they reacts mutually’’. (b) It can be explained with the following example : (i) Nitrogen (A) and oxygen (B) reacts with hydrogen (C ) to form ammonia and water respectively. (ii) 1g hydrogen reacts with 4. g 66 nitrogen to form ammonia and 8 g oxygen to form water. Thus, the ratio of masses reacting with a fixed mass of H is 4 66 8 0 583 . = . 5. Law of Gaseous Volumes (Gay Lussac 1808) (a) According to this law, ‘‘when gases combine, they do so in volumes which bear a simple ratio to each other and also to the product formed provided all gases are measured under similar conditions.’’ (b) (i) Consider the reaction between H2 and O2 to form water. H2 O2 H2O 1 2 ( g ) + ( g ) → ( v) The equation states that one volume of H2 reacts with half volume of O2 to give one volume of H2O vapours. 6. Avogadro’s Hypothesis Under similar conditions of pressure and temperature, equal volumes of gases contain equal number of molecules. Salient Features of Avogadro’s Hypothesis The significance of Avogadro’s hypothesis is summarised below : (a) Distinction between atom and molecule : It removed the anomaly between Dalton’s atomic theory and Gay Lussac’s law of volume by making a clear distinction in between atoms and molecules. (b) It provided a method to determine the atomic masses of gaseous elements. (c) It reveals that common elementary gases ( like- hydrogen, nitrogen, oxygen, chlorine etc.) are diatomic. (d) Relation between vapour density (V.D.) and molecular mass of gaseous substances : The vapour density of a gas is defined as the ratio of the mass of a certain volume of the gas to the mass of the same volume of hydrogen at same temperature and pressure. V.D. with respect to hydrogen, Vapour density Molecular mass 2 = (e) Molar volume of gases : One g molecular mass (or mass one mole molecules) of a gas occupies 22.4 L at NTP i. e., 2 g of H2 , 32 g O2 or 16 g CH4 at NTP occupies 22.4 litre volume. It may be calculated as given below : Vapour density = Mass of one molecule of gas Mass of one molecule of H2 = Mass of one mL of gas at STP Mass of one mL of H2 = Mass of one mL of gas at STP 0.0000897 (∵1 mL H2 has mass = 0.0000897 g at STP) ∴ Mass of 1 mL gas = × V. D. 0.0000897 g But V. D. × 0.0000897 g gas has volume = 1 mL ∴ 2 × V. D. (i. e., mol. mass) g gas has volume = × × × = 1 2 0 0000897 22400 V D V D mL . . . . . i e. ., g-molar mass of a gas or its 1 mole occupies 22.4 litre at STP or NTP. (f) It may be used in determination of M.F. of gases and in gas analysis. Some Basic Concepts of Chemistry 3
Atoms and Molecules (i) An atom is the smallest particle of an element that takes part in a chemical reaction. (ii) A molecule is the smallest particle of matter (element or compound) which is capable of independent existence. A molecule is generally (except inert gases) an assembly of two or more tightly bonded atoms. Atomic Mass Now a days, all the atomic masses have been measured with respect to 12C whose atomic mass is arbitrarily taken as 12.000 u. Relative atomic mass of an element is defined as the ratio of the masses of that atom to (1 12) th of the mass of one C 12 -atom, thus Relative atomic mass Mass of an atom of element 1 12 = × Mass of an atom of C 12 On the basis of this scale atomic masses of H and O-atoms are 1.008 u and 15.995 u respectively. Atomic mass unit (amu) is used to express the atomic masses. 1 amu 1.66056 10 g 24 = × − Average atomic mass (A) (% abundance isotopic mass) 100 = Σ × or A = ×Σ(Isotopic mass Fractional abundance) Molecular Mass and Formula Mass (i) A molecule is composed of several atoms and sum of atomic masses of atoms present in a molecule, refers to its molecular mass, e. g., consider a molecule AxByCz . Molecular mass = ( x × at. mass of at. mass of A) + ( y × B) + × ( ) z C at. mass of or Molecular mass = × + Σ atomic mass n atomic mass × n Element Elem 1 1 2 644474448 6 7444 8444 ent 2 + ... n1 and n2 are number of atoms respectively. e. g., for C6H12O6 M = [(6 × 12.0011) + (12 × 1.008) + (6 × 16.00)]u = 180.162 u (ii) For ionic solids term molecular mass is replaced by more common term ‘‘for mula mass’’. e. g., for NaCl, formula mass = (23.0 + 35.5) = 58.5 u The Mole Concept (i) Mole (symbol mol) is the seventh base S.I. unit. It expresses the amount of sub stance, in terms of both, its mass and the num ber of en ti ties (at oms/ions/mol e cules etc.) pres ent in it. (ii) By definition, one mole is the amount of substance that contains as many as particles or entities as there are atoms in exactly 12.00 g (or 0.012 kg) of the C 12 -isotope. (iii) One mole of all the substances (atoms/molecules/ions or anything other) contain same number of entities, and now this number (i. e., number of atoms in 12 0 12 . g C -isotope) has been determined exactly by using mass spectrometer. With this technique, mass of one carbon atom has been found to be 1 992648 10 23 . × − g. Thus, number of atoms in 12.0 g of C 12 = N A [Avogadro’s constant (or number)] = × − − 12 0 1 992648 10 12 23 . . g mol C g atom = 6.0221367 10 atoms mol 23 × Thus, 1 mole atoms = N A atoms (i. e., 6 023 10 23 . × atoms) 1 mole ions = N A ions (i. e., 6 023 10 23 . × ions) 1 mole electrons = N A electrons (i. e., 6 023 10 23 . × electrons) 1 mole compound = N A molecules of that compound (i. e., 6 023 10 23 . × molecules) (iv) Mass of one mole atoms of an element is known as g-atomic mass, while those of one mole mol e cules of a com pound is known as g-mo lar mass (or g-for mula mass). How ever, terms g-atomic mass or g-mo lar mass have been given for el e ments and com pounds, but in com mon prac tice, com mon term mole (sym bol mol) is used for all pur pose and sym bol ‘ n’ is used for ex press ing the num ber of moles. (v) Moles (n) of compound or number of g-molecules (simply mole) in wg of substance are given by n = mass Molar mass (w) (m) (vi) 1 mole molecules of any ideal gas at STP = 6.023 × 10 molecules of gas = 22.4 L 23 4 Crash Course Chemistry for NEET