Title 1 some basic concepts of chemistry.pdf ✅

Page 1 of 8 www.bankofchemistryfacts.blogspot.com IMPORTANCE OF CHEMISTRY Chemistry is the study of the composition, properties and interaction of matter, i.e., the science of atoms & molecules. Chemistry plays an important role in daily life. E.g. o Manufacture of fertilizers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metals, alloys etc. o Production of food, health care products etc. o Production of fertilizers, pesticides and insecticides. o Isolation of drugs from plants and animals or prepared by synthetic methods. E.g. Cisplatin & taxol (for cancer therapy) and AZT (Azidothymidine- for AIDS victims). o Synthesis of materials having specific magnetic, electric and optical properties. It helps to make superconducting ceramics, conducting polymers, optical fibres and miniaturization of solid state devices. o Tackling with environmental pollution. E.g. alternatives to environmentally hazardous refrigerants like CFCs. NATURE OF MATTER - Matter is anything which has mass and occupies space. - Matter can exist in 3 physical states: solid, liquid & gas. Solid Liquid Gas Particles are very close to each other in an orderly fashion. No much freedom of movement. Particles are close to each other but they can move around. Particles are far apart as compared to solid or liquid. Their movement is easy and fast. Definite volume and definite shape. Definite volume but no definite shape. They take shape of the container. Neither definite volume nor definite shape. They completely occupy the container. - States of matter are inter-convertible by changing the temperature and pressure. - At macroscopic or bulk level, matter is classified as follows: Mixture: Contains two or more substances (components) in any ratio and their composition is variable. o Homogeneous mixture: The components completely mix with each other and its composition is uniform throughout. E.g. Sugar solution and air. o Heterogeneous mixture: The composition is not uniform throughout and sometimes different components are seen. E.g. mixtures of salt & sugar, grains & pulses along with some dirt pieces. Components of a mixture can be separated by physical methods (hand picking, filtration, crystallisation, distillation etc.). Pure substances: - Have fixed composition. E.g. Cu, Ag, gold, water, glucose. - Glucose contains C, H & O in a fixed ratio. - The constituents of pure substances cannot be separated by simple physical methods. - Pure substances are 2 types: Elements and Compounds. Elements: - An element consists of only one type of particles. These particles may be atoms or molecules. E.g. Sodium, copper, silver, hydrogen, oxygen etc. - Atoms of different elements are different in nature. - Some elements (e.g. Na, Cu) contain single atoms held together as their constituent particles. In some others (e.g. H2, N2, O2), two or more atoms combine to give molecules of the element. Compounds: - When two or more atoms of different elements combine, the molecule of a compound is obtained. E.g. water, ammonia, carbon dioxide, sugar etc. Water molecule (H2O) Carbon dioxide molecule (CO2) - In a compound, atoms of different elements are present in a fixed and definite ratio. E.g. A water molecule has 2 hydrogen atoms and one oxygen atom. A carbon dioxide molecule contains two oxygen atoms and one carbon atom. - Properties of a compound are different from those of its constituent elements. E.g. H and O are gases but water is a liquid. Hydrogen burns with a pop sound and oxygen is a supporter of combustion, but water is a fire extinguisher. - Constituents of a compound cannot be separated into simpler substances by physical methods. They can be separated by chemical methods. Matter Mixtures Homogeneous mixtures Heterogeneous mixtures Pure substances Elements Compounds
Page 2 of 8 www.bankofchemistryfacts.blogspot.com PROPERTIES OF MATTER AND THEIR MEASUREMENT - Physical properties: They can be measured or observed without changing the identity or composition of substance. E.g. colour, odour, melting point, boiling point, density etc. - Chemical properties: Here, measurement or observation requires a chemical change. E.g. acidity or basicity, combustibility etc. - English System & Metric System are 2 systems of measurement of matter. - Metric system (originated in France in late 18th century) was more convenient as it was based on decimal system. - A common standard system was established in 1960. The International System of Units (SI) - The International System of Units (in French Le Systeme International d’Unités) was established by the 11th General Conference on Weights and Measures (CGPM from Conference Generale des Poids et Measures). 7 base units (fundamental quantities) of SI system Base physical quantity Symbol for quantity Name of SI unit Symbol for SI unit Length l metre m Mass m kilogram kg Time t second s Electric current I ampere A Thermodynamic temperature T kelvin K Amount of substance n mole mol Luminous intensity I candela cd • Metre: It is the length of the path travelled by light in vacuum in 1/299792458 second. • Kilogram: It is equal to the mass of the international prototype of the kilogram. • Second: It is the duration of 9,192,631,770 periods of the radiation corresponding to the transition between the two levels of the ground state of caesium-133 atom. • Ampere: It is the constant current which, if maintained in two straight parallel conductors of infinite length, of negligible circular cross-section, and placed 1 metre apart in vacuum, would produce between these conductors a force equal to 2 × 10–7 newton per metre of length. • Kelvin: It is the fraction 1/273.16 of the thermodynamic temperature of the triple point of water. • Mole: It is the amount of substance which contains as many elementary entities (atoms, molecules, ions, electrons etc.) as there are atoms in 12g of carbon-12. Its symbol is “mol.” • Candela: It is the luminous intensity, in a given direction, of a source that emits monochromatic radiation of frequency 540×1012 hertz and that has a radiant intensity in that direction of 1/683 watt per steradian. - The other physical quantities such as speed, volume, density etc. can be derived from these quantities. Mass and Weight - Mass is the amount of matter present in a substance. - Weight is the force exerted by gravity on an object. - Mass of a substance is constant whereas its weight may vary from one place to another due to change in gravity. - Mass of a substance can be determined very accurately by using an analytical balance. - The SI unit of mass is kilogram. However, its fraction gram (1 kg = 1000 g) is used in laboratories due to the smaller amount of chemicals used in chemical reactions. Volume - Volume has the units of (length) 3 . - So in SI system, volume has units of m3 . - In chemistry laboratories, smaller volumes are used. Hence, volume is often denoted in cm3 or dm3 units. - Litre (L) is measure of volume of liquids. It is not an SI unit. - 1 L = 1000 mL, 1000 cm3 = 1 dm3 Density - It is the mass of a substance per unit volume. SI unit of density = SI unit of mass SI unit of volume = kg m3 or kg m−3 - A chemist often expresses density in g cm–3 . Temperature There are three scales to measure temperature: • °C (degree Celsius) scale: It is calibrated from 0° to 100°. (0° = freezing point of water, 100° = boiling point of water). Celsius scale has temperature below 0 °C (negative values). • °F (degree Fahrenheit) scale: It is represented between 32° to 212°. It is related to Celsius scale as follows: °F = 9 5 (°C) + 32 • K (kelvin) scale: It has no negative temperature. Kelvin scale (SI unit) is related to Celsius scale as follows: K = °C + 273.15 Prefixes to indicate the multiples or submultiples of a unit Multiple Prefix Symbol Multiple Prefix Symbol 10-24 yocto y 10 deca da 10-21 zepto z 102 hecto h 10-18 atto a 103 kilo k 10-15 femto f 106 mega M 10-12 pico p 109 giga G 10-9 nano n 1012 tera T 10-6 micro  1015 peta P 10-3 milli m 1018 exa E 10-2 centi c 1021 zeta Z 10-1 deci d 1024 yotta Y
Page 3 of 8 www.bankofchemistryfacts.blogspot.com UNCERTAINTY IN MEASUREMENT Scientific Notation - It is the exponential notation in which a number is represented as N × 10n . Here, n is an exponent and N is a number (digit term). - E.g. 232.508 is written as 2.32508×102 . 0.00016 is written as 1.6 × 10–4 . Multiplication, Division, Addition & Subtraction - Multiplication of two numbers: (5.6 ×105 ) × (6.9 x108 ) = (5.6 × 6.9) (105+8) = (5.6 × 6.9) x1013 = 38.64 × 1013 = 3.864 × 1014 (9.8 ×10-2 ) × (2.5 × 10-6 ) = (9.8 × 2.5) (10-2+ (-6)) = (9.8 × 2.5) (10-2-6 ) = 24.50 × 10-8 = 2.450 × 10-7 - Division of two numbers: 2.7 x 10−3 5.5 x 104 = (2.7 ÷ 5.5)(10−3−4 ) = 0.4909 × 10−7 = 4.909 × 10-8 - Addition of two numbers: 6.65 × 104+ 8.95 × 103 = 6.65 × 104+ 0.895 × 104 = (6.65 + 0.895) × 104= 7.545 × 104 - Subtraction of two numbers: 2.5 × 10–2 – 4.8 × 10–3 = (2.5 × 10–2 ) – (0.48 × 10–2 ) = (2.5 – 0.48) × 10–2= 2.02 × 10–2 Significant Figures - Every experimental measurement has some uncertainty. - Precision: It refers to closeness of various measurements for the same quantity. - Accuracy: It is the agreement of a particular value to the true value of the result. - E.g., the true value for a result is 2.00 g. Measurements of three students A, B and C are given in the table: Measurements / g 1 2 Average Precise / accurate Student A 1.95 1.93 1.940 Precise but not accurate Student B 1.94 2.05 1.995 Neither precise nor accurate Student C 2.01 1.99 2.000 Precise & accurate - Significant figures are meaningful digits which are known with certainty. The uncertainty is indicated by writing the certain digits and the last uncertain digit. E.g. in 11.2 mL, 11 is certain and 2 is uncertain and the uncertainty would be ±1 in the last digit. Rules to determine the number of significant figures: 1. All non-zero digits are significant. E.g. 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. E.g. 0.03 has one significant figure. 0.0052 has two significant figures. 3. Zeros between two non-zero digits are significant. E.g. 2.005 has four significant figures. 4. Zeros at the end or right of a number are significant provided they are on the right side of the decimal point. E.g. 0.200 g has three significant figures. Terminal zeros without decimal point are not significant. E.g. 100 has only one significant figure, but 100. has three significant figures and 100.0 has four significant figures. In scientific notation: 100 = 1×102 100. = 1.0×102 100.0 = 1.00×102 5. Counting numbers of objects(e.g. 2 balls or 20 eggs), have infinite significant figures. It is represented by infinite number of zeros after a decimal i.e., 2 = 2.000000 or 20 = 20.000000 6. In numbers written in scientific notation, all digits are significant. E.g. 4.01×102 has three significant figures, and 8.256 × 10–3 has four significant figures. Addition and Subtraction of Significant Figures - The result cannot have more digits to the right of the decimal point than either of the original numbers. E.g. Multiplication and Division of Significant Figures - In these operations, the result must be reported with no more significant figures as are there in the measurement with the few significant figures. 2.5×1.25 = 3.125 - Since 2.5 has two significant figures, the result should not have more than two significant figures, thus, it is 3.1. - While limiting the result, the following points for rounding off the numbers should be considered: 1. If the rightmost digit to be removed is more than 5, the preceding number is increased by one. E.g. 1.386 is rounded to 1.39. 2. If the rightmost digit is lessthan 5, the preceding number is not changed. E.g. 4.334 is rounded up to 4.33. 3. If the rightmost digit is 5, then the preceding number is not changed if it is an even number but it is increased by one if it is an odd number. E.g. 6.35 is rounded to 6.4. 6.25 is rounded to 6.2. Dimensional Analysis - Dimensional analysis (factor label or unit factor method) is a method to convert units from one system to other. Example 1 A piece of metal is 3 inch long. What is its length in cm? We know that 1 in = 2.54 cm.  1 in 2.54 cm = 1 = 2.54 cm 1 in Both of these are called unit factors. 3 in = 3 in × 2.54 cm 1 in = 3 × 2.54 cm = 7.62 cm Example 2 A jug contains 2L milk. Calculate the volume of milk in m3 . Here, 18.0 has only one digit after the decimal point and the result should be reported only up to one digit after the decimal point which is 31.1.