Chemistry Class 11 Chapter 1 Notes PDF

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Please let us know if the content in this file infringes any of your copyright by writing to us at: [email protected] and we will take appropriate action.  1 Chapter notes Class: XI Chapter 1: Some Basic Concepts of Chemistry Top concepts 1. The SI system has seven base units which pertain to the 7 fundamental scientific quantities Base Physical Symbol for Name of SI Unit Symbol for SI Quantity quantity Unit Length l metre m Mass m kilogram kg Time t second ms Electric current I ampere c o A. Thermodynamic T a Kelvin K temperature m Amount of n a mole mol substance yn Luminous Iv d candela cd intensity tu S 2. The unit is written on the right of the quantity with a space between them. 3. The SI system allows the use of prefixes to indicate the multiples or submultiples of a unit. Multiple Prefix Symbol 10-1 deci d  2 10 deca da 102 hector h 3 10 kilo k 106 mega M 4. To indicate very small numbers, we use negative exponents. 5. To indicate large numbers, we use positive exponents. 6. Scientific notation is the proper representation of a number in exponential form. 7. Precision indicates how closely repeated measurements match each other. m c o. a 8. Accuracy indicates how closely a measurement matches the correct or expected value. m a yn 9. A result is valid only if it is both accurate and precise. d tu 10. Significant figures are meaningful digits which are known with S certainty. 11. There are certain rules for determining the number of significant figures: i) All non-zero digits are significant ii) Zeros preceding the first non-zero digit are not significant iii) Zeros between two non-zero digits are significant.  3 iv) Zeros at the end or right of the number are significant provided they are on the right side of the decimal point. But, if otherwise, the zeros are not significant. 14. During addition and subtraction, the result cannot have more digits to the right of the decimal point than either of the original numbers. 15. In multiplication and division with significant figures, the answer cannot have more significant figures than either of the original numbers. 16. There are 5 basic laws of chemical combinations that govern every m reaction: Law of conservation of mass, law of definite proportions, c o law of multiple proportions, Gay Lussac’s law of gaseous volumes,. a and lastly, Avogadro law. m a 17. Law of Conservation of Mass: Antoine Lavoisier established the yn Law of Conservation of Mass. It states that matter can neither be d created nor destroyed. In other words, we can say that during any tu physical or chemical change, the total mass of reactants is equal to S the total mass of products. 18. Law of definite proportions: Joseph Proust showed that a given compound always contains exactly the same proportion of elements by weight. 19. Law of multiple proportions: Dalton proposed the law of multiple proportions. According to this law if two elements can combine to form more than one compound, the mass of one element that combines with the fixed mass of the other element is in the ratio of small whole numbers  4 20. Gay Lussac’s Law of gaseous volumes: When gases combine or are produced in a chemical reaction they do so in a simple ratio by volume, provided all the gases are at same temperature and pressure. 21. Avogadro law: At the same temperature and pressure, equal volumes of gases contain equal number of molecules. 22. Dalton’s atomic theory: In 1808, Dalton published ‘A New System of Chemical Philosophy’ in which he proposed the following :  Matter consists of indivisible atoms.  All the atoms of a given element have identical properties m including identical mass. Atoms of different elements differ in mass. c o. a  Compounds are formed when atoms of different elements combine in a fixed ratio. m  a Chemical reactions involve reorganisation of atoms. These are yn neither created nor destroyed in a chemical reaction. d tu 23. Dalton’s theory could explain the laws of chemical combination. S 24. The number 6.022  1023 is called Avogadro’s constant or Avogadro’s number. 25. A mole is a collection of 6.022  1023 particles. 26. One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12 g (or 0.012 kg) 12 of the C 27. The mass of one mole of a substance in grams is called its molar mass. 28. The molar mass in grams is numerically equal to the atomic/molecular/formula mass in u.(u is the unified mass)  5 29. Molarity is the number of moles of solute in per liter of solution. Unit is moles per liter. 30. Molality is the number of solute present in 1kg of solvent. 31. Atomic Mass: Average relative mass of an atom of an element as compared with the mass of a carbon atom taken as 12 amu 32. Atomic mass expressed in grams is called gram atomic mass 33. Molecular Mass: Sum of the atomic masses of elements present in a molecule 34. Molecular mass expressed in grams is called gram molecular mass 35. Formula Mass: Sum of atomic masses of all atoms in a formula unit of the compound 36. Following relations given below can be summarized  One mole of atoms = 6.022  1023 atoms=Gram atomic mass of element m c o One mole of molecules= 6.022  1023 molecules= Gram . molecular mass of substance a m a yn d 37. An empirical formula represents the simplest whole number ratio of tu various atoms present in a compound. S 38. Molecular formula shows the exact number of different types of atoms present in a molecule of a compound. 39. If the mass per cent of various elements present in a compound is known, its empirical formula can be determined. 40. Molecular formula = n (Empirical formula) , where n is a simple number and may have values 1, 2, 3…. 41. Following steps should be followed to determine empirical formula of the compound  Step 1: Conversion of mass per cent of various elements into grams.  Step 2: Convert mass obtained in step1 into number of moles  6  Step 3: Divide the mole value obtained in step 2 by the smallest mole value (out of the mole value of various elements calculated)  Step 4: In case the ratios are not whole numbers, then they may be converted into whole number by multiplying by the suitable coefficient.  Step 5: Write empirical formula by mentioning the numbers after writing the symbols of respective elements. 42. Anything that has mass and occupies space is called matter 43. Matter can either be a mixture or be a pure substance 44. Homogenous mixtures are those whose components completely mix with each other to make a uniform composition m c o 45. Heterogeneous mixtures are not uniform, and their components are. a separable through physical methods 46. Pure substances can be elements or compounds m 47. a An element consists of only one type of particles yn 48. Two or more atoms of different elements combine to form a molecule of a compound d tu 49. The constituents of a compound can be separated only by chemical S methods. 50. A compound has properties different from its constituent elements 51. Isotopes are elements with same atomic number but different mass number. 52. Atomic mass is donated by “u” – unified mass. 53. One mole is the amount of a substance that contains as many 12 particles as there are atoms in exactly12 g of the C isotope 54. The mass of one mole of a substance in grams is called its molar mass 55. Out of various reactants in a reaction ,a reactant that is completely consumed in a chemical reaction is called limiting reagent  7 56. Stoichiometry gives a quantitative relation between reactant and product in a reaction. It also helps us in identifying limiting reagents Top Formulae Mass % of an element 1. Mass of that element in the compound   100 Molar mass of compound Mass of solute 2. Massper cent   100 Mass of solution No. of mole of a particular component 3. Mole Fraction= Total No. of moles of solution m 4. Molarity= No. of moles of solute Volume of solution in litres c o. 5. Molality  No. of moles of solute Mass of solvent inkg a m 6. Moles of an element = a Mass of element yn Atomic mass Atomic mass 7. Mass of one atom= d 6.022  1023 tu Mass of compound S 8. Moles of a compound = Molecular mass Molecular mass 9. Mass of one molecule= 6.022  1023  Points To Remember Class: XI Ch 2: Structure O Atom Top Concepts 1. Atomic theory of matter was proposed by John Dalton 2. Electrons were discovered by Michael Faraday. 3. Electrons were discovered using cathode ray discharge tube experiment. 4. Cathode ray discharge tube experiment: A cathode ray discharge tube made of glass is taken with two electrodes. At very low pressure and high voltage, current starts flowing through a stream of particles moving in the tube from cathode to anode. These rays were called cathode rays. When a perforated anode was taken, the cathode rays struck the other end of the glass tube at the fluorescent coating and a bright spot on the coating was developed m Results: c o. a. Cathode rays consist of negatively charged electrons. a b. Cathode rays themselves are not visible but their behavior can be observed with help of fluorescent or phosphorescent materials. m c. In absence of electrical or magnetic field cathode rays travel in straight lines a yn d. In presence of electrical or magnetic field, behaviour of cathode rays is similar to that shown by electrons d e. The characteristics of the cathode rays do not depend upon the tu material of the electrodes and the nature of the gas present in the cathode ray tube. S 5. Charge to mass ratio of an electron was determined by Thomson. The charge to mass ratio of an electron as 1.758820 x 10 11 C kg-1 6. Charge on an electron was determined by R A Millikan by using an oil drop experiment. The value of the charge on an electron is -1.6 x 10-19 C. 7. The mass on an electron was determined by combining the results of Thomson’s experiment and Millikan’s oil drop experiment. The mass of an electron was determined to be 9.1094 x 10-31 kg. 8. Discovery of protons and canal rays: Modified cathode ray tube experiment was carried out which led to the discovery of protons. 9. Canal rays are positively charged particles called protons 10.Characteristics of positively charged particles: a. Charge to mass ratio of particles depends on gas from which these originate b. The positively charged particles depend upon the nature of gas present in the cathode ray discharge tube c. Some of the positively charged particles carry a multiple of fundamental of electrical charge. d. Behaviour of positively charged particles in electrical or magnetic field is opposite to that observed for cathode rays 11.Neutrons were discovered by James Chadwick by bombarding a thin sheet of beryllium by  - particles. They are electrically neutral particles having a mass slightly greater than that of the protons. m 12.Thomson model of an atom: This model proposed that atom is considered o as a uniform positively charged sphere and electrons are embedded in it. c. 13. An important feature of Thomson model of an atom was that mass of atom a is considered to be evenly spread over the atom. m 14.Thomson model of atom is also called as Plum pudding, raisin pudding or watermelon model a yn 15.Thomson model of atom was discarded because it could not explain certain d experimental results like the scattering of  - particles by thin metal foils tu 16.Observations from  - particles scattering experiment by Rutherford: S a. Most of the  - particles (nearly 99 %) passed through gold foil undeflected b. A small fraction of  - particles got deflected through small angles c. Very few  - particles did not pass through foil but suffered large deflection nearly 180 o 17.Observations Rutherford drew from  - particles scattering experiment: a. Since most of the  -particles passed through foil undeflected, it means most of the space in atom is empty b. Since some of the  -particles are deflected to certain angles, it means that there is positively mass present in atom c. Since only some of the  -particles suffered large deflections, the positively charged mass must be occupying very small space d. Strong deflections or even bouncing back of  -particles from metal foil were due to direct collision with positively charged mass in atom 18.Rutherford’s model of atom: This model explained that atom consists of nucleus which is concentrated in a very small volume. The nucleus comprises of protons and neutrons. The electrons revolve around the nucleus in fixed orbits. Electrons and nucleus are held together by electrostatic forces of attraction. 19.Drawbacks of Rutherford’s model of atom: a. According to Rutherford’s model of atom, electrons which are negatively charged particles revolve around the nucleus in fixed orbits. Thus, the electrons undergo acceleration. According to electromagnetic theory of Maxwell, a charged particle undergoing acceleration should emit electromagnetic radiation. Thus, an electron in an orbit should emit radiation. Thus, the orbit should shrink. But this does not happen. b. The model does not give any information about how electrons are m distributed around nucleus and what are energies of these electrons c o. a 20. Atomic number (Z): It is equal to the number of protons in an atom. It is also equal to the number of electrons in a neutral atom. m a 21.Mass number (A): It is equal to the sum of protons and neutrons. yn d 22.Isotopes: These are the atoms of the same element having the same tu atomic number but different mass number. S 23. Isobars: Isobars are the atoms of different elements having the same mass number but different atomic number. 24. Isoelectronic species: These are those species which have the same number of electrons. 25.Electromagnetic radiations: The radiations which are associated with electrical and magnetic fields are called electromagnetic radiations. When an electrically charged particle moves under acceleration, alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the form of waves. These waves are called electromagnetic waves or electromagnetic radiations. 26.Properties of electromagnetic radiations:  a. Oscillating electric and magnetic field are produced by oscillating charged particles. These fields are perpendicular to each other and both are perpendicular to the direction of propagation of the wave. b. They do not need a medium to travel. That means they can even travel in vacuum. 27.Characteristics of electromagnetic radiations : a. Wavelength: It may be defined as the distance between two neighbouring crests or troughs of wave as shown. It is denoted by. b. Frequency (): It may be defined as the number of waves which pass through a particular point in one second. c. Velocity (v): It is defined as the distance travelled by a wave in one second. In vacuum all types of electromagnetic radiations travel with the same velocity. Its value is 3 X108 m sec-1. It is denoted by v d. Wave number: Wave number ( ) is defined as the number of wavelengths per unit length. 28.Relationship between velocity, frequency and wavelength Velocity = frequency x wavelength m c =  c o. a 29.Black body: An ideal body, which emits and absorbs all frequencies, is called a black body. The radiation emitted by such a body is called black body radiation. m a yn 30.Planck’s quantum theory: Max Planck suggested that atoms and molecules d could emit or absorb energy only in discrete quantities and not in a tu continuous manner. Planck gave the name quantum, meaning ‘fixed amount’ to the smallest quantity of energy that can be emitted or absorbed S in the form of electromagnetic radiation. E v hc E = hv  Where: E is the energy of a single quantum is the frequency of the radiation h is Planck’s constant h= 6.626 X 10–34 Js 31.Quantisation of energy: Energy is always emitted or absorbed as integral multiple of this quantum.  E = nhv Where n 1, 2,3, 4,..... 32. Photoelectric effect: The phenomenon of ejection of electrons from the surface of metal when light of suitable frequency strikes it is called photoelectric effect. The ejected electrons are called photoelectrons. 33. Experimental results observed for the experiment of Photoelectric effect observed Hertz: a. When beam of light falls on a metal surface electrons are ejected immediately i.e. there is not time lag between light striking metal surface and ejection of electrons b. Number of electrons ejected is proportional to intensity or brightness of light c. Threshold frequency ( vo ): For each metal there is a characteristic minimum frequency below which photoelectric effect is not observed. This is called threshold frequency. m d. If frequency of light is less than the threshold frequency there is no ejection intensity. c o of electrons no matter how long it falls on surface or how high is its. a m 34.Photoelectric work function (Wo): The minimum energy required to eject a electrons is called photoelectric work function. yn Wo  hvo d 35.Energy of the ejected electrons : tu 1 h(v  v0 )  me v 2 2 S 36.When a white light is passed through a prism, it splits into a series of coloured bands known as spectrum. 37.Spectrum is of two types: continuous and line spectrum a. The spectrum which consists of all the wavelengths is called continuous spectrum. b. A spectrum in which only specific wavelengths are present is known as a line spectrum. It has bright lines with dark spaces between them. 38.Electromagnetic spectrum is a continuous spectrum. It consists of a range of electromagnetic radiations arranged in the order of increasing wavelengths or decreasing frequencies. It extends from radio waves to gamma rays.  39. Spectrum is also classified as emission and line spectrum. c. Emission spectrum: A substance absorbs energy and moves to a higher energy state. The atoms, molecules or ions that have absorbed radiation are said to be excited. Since the higher energy state is unstable they return to the more stable energy state by emitting the absorbed radiation in various regions of electromagnetic spectrum. The spectrum of radiation emitted by a substance that has absorbed energy is called an emission spectrum. d. Absorption spectrum is the spectrum obtained when radiation is passed through a sample of material. The sample absorbs radiation of certain wavelengths. The wavelengths which are absorbed are missing and come as dark lines. 40.The study of emission or absorption spectra is referred as spectroscopy. 41.Spectral Lines for atomic hydrogen: Series n1 n2 Spectral Region m Lyman 1 2, 3, 4, 5 … Ultraviolet c o. Balmer 2 3, 4, 5 … a Visible m a Paschen 3 4, 5 … Infrared yn Brackett 4 5, 6 … Infrared Pfund 5 d 6, 7… Infrared tu S 42.Rydberg equation: It allows the calculation of the wavelengths of all the spectral lines of hydrogen. 43. Bohr’s model for hydrogen atom: a. An electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy. These paths are called orbits or energy levels. These orbits are arranged concentrically around the nucleus. b. As long as an electron remains in a particular orbit, it does not lose or gain energy and its energy remains constant. c. When transition occurs between two stationary states that differ in energy, the frequency of the radiation absorbed or emitted can be calculated.  E E 2 -E1 v  h h v = Frequency of radiation h = Planck's constant E1  Energy of lower energy state E 2  Energy of higher energy state d. An electron can move only in those orbits for which its angular momentum is an integral multiple of h/2 43.Bohr’s theory for hydrogen atom: a. Stationary states for electron are numbered in terms of Principal Quantum numbered as n=1, 2, 3… m b. For hydrogen atom: The radii of the stationary states is expressed as rn = n2a0 where a0= 52.9 pm c o. c. Energy of stationary state  1  En   R H  2  n  a m where R H  2.18  1018 J(Rydberg cons tan t) a yn n  1, 2,3,.... d 1  tu E n   2.18 x1018  2  J S n  d. For ions containing only one electron:  Z2  E n   2.18 x1018  2  J n  where n  1, 2,3,.... rn = n2a0 pm Z Where Z is the atomic number 44.Limitations of Bohr’s model of atom:  a. Bohr’s model failed to account for the finer details of the hydrogen spectrum. For instance splitting of a line in the spectrum into two closely spaced lines. b. Bohr’s model was also unable to explain spectrum of atoms containing more than one electron. c. Bohr’s model was unable to explain Zeeman effect i.e. splitting of spectral line in presence of magnetic effect. d. Bohr’s model also failed to explain Stark effect i.e. splitting of spectral line in presence of electric field. e. Bohr’s model could not explain the ability of atoms to form molecules by chemical bonds 45.Dual behavior of matter: de Broglie proposed that matter exhibits dual behavior i.e. matter shows both particle and wave nature. 1. de Broglie’s relation: h h   mv p Where: m  - Wavelength c o p - Momentum. a v - Velocity h – Planck’s constant m a 2. According to de Broglie, every object in motion has a wave character. yn Wavelengths of macroscopic objects cannot be detected but for microscopic particles it can be detected. This is because for microscopic d objects, the mass is less. Since mass and wavelength are inversely tu proportional to each other, the wavelength will be more. But for macroscopic objects, the mass is large. Therefore, wavelength will be S too short to be detected. 3. Heisenberg’s uncertainty principle: It states that it is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron. h  x. p x  4 h  x.  (m v x )  4 h  x.  vx  4m Where ∆x – Uncertainty in position ∆vx - Uncertainty in velocity  ∆px - Uncertainty in momentum This means that if the position of electron is known, the velocity of electron will be uncertain. On the other hand, if the velocity of electron is known precisely, the position of electron will be uncertain. 4. Heisenberg’s uncertainty principle rules our the existence of definite paths or trajectories of electrons and other similar particles 5. Failure of Bohr’s model: a. It ignores the dual behavior of matter. b. It contradicts Heisenberg’s uncertainty principle. 46.Classical mechanics is based on Newton’s laws of motion. It successfully describes the motion of macroscopic particles but fails in the case of microscopic particles. Reason: Classical mechanics ignores the concept of dual behaviour of matter especially for sub-atomic particles and the Heisenberg’s uncertainty principle. m 47.Quantum mechanics is a theoretical science that deals with the study of the particle like properties. c o motions of the microscopic objects that have both observable wave like and. a 48.When quantum mechanics is applied to macroscopic objects (for which m wave like properties are insignificant) the results are the same as those from the classical mechanics. a yn 49.Quantum mechanics is based on a fundamental equation which is called Schrodinger equation. d tu 50.Schrodinger’s equation: For a system (such as an atom or a molecule whose energy does not change with time) the Schrödinger equation is S written as: Ĥ  E Where: Ĥ is the Hamiltonian operator E is the total energy of the system  represents the wave function which is the amplitude of the electron Wave 51.When Schrödinger equation is solved for hydrogen atom, the solution gives the possible energy levels the electron can occupy and the corresponding wave function(s) of the electron associated with each energy level.  Out of the possible values, only certain solutions are permitted. Each permitted solution is highly significant as it corresponds to a definite energy state. Thus, we can say that energy is quantized. That is, it can have only certain specific values. 52.  gives us the amplitude of wave. The value of  has no physical significance. 53.  2 gives us the region in which the probability of finding an electron is maximum. It is called probability density. 54.Orbital: The region of space around the nucleus where the probability of finding an electron is maximum is called an orbital. 55.Quantum numbers: There are a set of four quantum numbers which specify the energy, size, shape and orientation of an orbital. These are: a. Principal quantum number (n) b. Azimuthal quantum number (l) m c. Magnetic quantum number (ml) d. Electron spin quantum number (ms) c o. 56.Principal quantum number (n): It determines the size and to a large extent the energy of the orbital. a m a n 1 2 3 4 yn Shell no.: K L M N d Total number of orbitals in a 1 4 9 16 tu shell = n2 S Maximum number of electrons = 2n2 2 8 18 32  It can have positive integer values of 1, 2, 3 and so on.  It also identifies the shell.  As the value of n increases, the energy also increases. Hence, the electron will be located far away from the nucleus. 57.Azimuthal quantum number (l): Azimuthal quantum number. ‘l’ is also known as orbital angular momentum or subsidiary quantum number. It identified the sushell and the three dimensional shape of the orbital.  It also determines the number of subshells or sub levels in a shell. Total number of subshells in a particular shell is equal to the value of n. l = 0, 1, 2… (n-1)   Each subshell corresponding to different values of l are represented by different symbols: Value of l 0 1 2 3 Notation of s p d f symbol 58.Magnetic quantum number or Magnetic orbital quantum number (ml): It gives information about the spatial orientation of the orbital with respect to standard set of co-ordinate axis. For any sub-shell (defined by ‘l’ value) 2l+1 values of ml are possible. For each value of l, ml = – l, – (l –1), – (l–2)... 0,1... (l – 2), (l–1), l 59.Electron spin quantum number (ms): It refers to orientation of the spin of the electron. It can have two values +1/2 and -1/2. +1/2 identifies the clockwise spin and -1/2 identifies the anti- clockwise spin. m 60.An orbital is identified by the set of 3 quantum numbers: Principal quantum number, Azimuthal quantum number and magnetic quantum number. c o. a 61.An electron is identified by a set of four quantum numbers: Principal quantum number, azimuthal quantum number, magnetic quantum number and spin quantum number. m a yn 62.Sub-shell notation: Notation of a sub-shell is written as the Principal quantum number followed by the symbol of the respective sub-shell. d tu 63.Plots of the orbital wave function (r ) and probability density 2(r) Vs distance r of the electron from the nucleus for 1s orbital: S  For 1s orbital the probability density is maximum at the nucleus and it decreases sharply as we move away from it(which is not possible).Hence plot of probability density 2(r) Vs distance r of the electron from the nucleus was drawn as shown below.   The orbital wave function  for an electron in an atom has no physical meaning. It is simply a mathematical function of the coordinates of the electron. 64.Plots of the orbital wave function (r ) and probability density 2(r) Vs distance r of the electron from the nucleus for 2s orbital: m c o. a m a yn d tu  S For 2s orbital the probability density is maximum at the nucleus and it decreases sharply as we move away from it(which is not possible).Hence plot of probability density 2(r) Vs distance r of the electron from the nucleus was drawn as shown below.  For 2s orbital, the probability density first decreases sharply to zero and again starts increasing. After reaching small maxima it decreases again and approaches zero as the value of r increases further. 65.The region where this probability density function reduces to zero is called nodal surfaces or simply nodes. 66.Charge cloud diagrams: In these diagrams, dots represent the electron probability density. The density of the dots in a region represents electron probability density in that region. 67.Boundary surface diagram: In this representation, a boundary surface or contour surface is drawn in space for an orbital on which the value of probability density 2(r) is constant. However, for a given orbital, only that boundary surface diagram of constant probability density is taken to be good representation of the shape of the orbital which encloses a region or volume in which the probability of finding the electron is very high, say, 90%. m c o 68.Radial nodes: Radial nodes occur when the probability density wave. function for the electron is zero on a spherical surface of a particular radius. Number of radial nodes = n – l – 1 a m a 69.Angular nodes: Angular nodes occur when the probability density wave yn function for the electron is zero along the directions specified by a particular angle. Number of angular nodes = l d tu 70.Total number of nodes = n – 1 S 71.Degenerate orbitals: Orbitals having the same energy are called degenerate orbitals. 72.The stability of an electron in a multi electron system is because of: a. The repulsive interaction of the electrons in the outer shell with the electrons in the inner shell. b. The attractive interactions of electron with the nucleus. These attractive interactions increase with increase of positive charge (Ze) on the nucleus. a. The stability of an electron in multi-electron atom is because total attractive interactions are more than the repulsive interactions. 73. Shielding effect or screening effect: Due to the presence of electrons in the inner shells, the electron in the outer shell will not experience the full positive charge on the nucleus.  So due to the screening effect, the net positive charge experienced by the electron from the nucleus is lowered and is known as effective nuclear charge. Effective nuclear charge experienced by the orbital decreases with increase of azimuthal quantum number (l). 74. Orbitals have different energies because of mutual repulsion between electrons in a multi- electron atom. 75.Orbitals with lower value of (n+l) are filled first as they have lower energy. 76.If two orbitals have the same value of (n+l) then orbital with lower value of n will have lower energy. 77.Energies of the orbitals in the same subshell decrease with increase in atomic number. 78.Filling of electrons: The filling of electrons into the orbitals of different atoms takes place according to Aufbau principle ,Pauli’s exclusion principle, the Hund’s rule of maximum multiplicity m c o. 79. Aufbau Principle: In the ground state of the atoms, the orbitals are filled in a order of their increasing energies. The order in which the orbitals are filled is as follows: m 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f, 5d, 6p, 7s... a It is based on (n+ l) rule. It states that the orbital with lower value of (n + yn l) has lower energy. d 80.Pauli Exclusion Principle: No two electrons in an atom can have the same tu set of four quantum numbers. Only two electrons may exist in the same orbital and these electrons must have opposite spin. S 81. Hund’s rule of maximum multiplicity: 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. 82. Electronic configuration of atoms: The electronic configuration of different atoms can be represented in two ways. a. sapbdc...... notation: In the first notation, the subshell is represented by the respective letter symbol and the number of electrons present in the subshell is depicted, as the super script, like a, b, c,... etc. The similar subshell represented for different shells is differentiated by writing the principal quantum number before the respective subshell. b. Orbital diagram: In the second notation, each orbital of the subshell is represented by a box and the electron is represented by an arrow () a positive spin or an arrow () a negative spin. 83. Stability of completely filled and half filled subshells: a. Symmetrical distribution of electrons b. Exchange energy m c o. a m a yn d tu S  Class: XI Chapter 3: Classification of Elements Chapter Notes Top Concepts 1. Johann Dobereiner classified elements in group of three elements called triads. 2. In Dobereiner’s triad the atomic weight of the middle element is very close to the arithmetic mean of the other two elements. 3. Dobereiner’s relationship is referred as Law of triads. 4. Since Dobereiner’s Law of triads worked only for few elements, it was dismissed. m c o 5. Chancourtois arranged elements in order of increasing atomic weights and made a cylindrical table of elements.. a 6. John Newland arranged the elements in the increasing order of atomic m weight and noted that the properties of the every eighth element are similar a to the first one. This relationship is called as “Law of octaves” yn 7. Lothar Meyer proposed that on arranging the elements in order of d increasing atomic weights similarities appear at a regular interval in physical tu and chemical properties. S 8. According to Mendeleev’s periodic law the physical and chemical properties of elements are periodic functions of their atomic weights. 9. Merits of Mendeleev’s periodic table:  Mendeleev’s periodic table was very helpful in remembering and studying the properties of large number of elements  Mendeleev’s periodic table helped in correcting the atomic masses of some of the elements like gold, beryllium and platinum based on their positions in the periodic table  Mendeleev could predict the properties of some undiscovered elements like scandium, gallium and germanium. By this intuition, he had left gaps for the undiscovered elements while arranging elements in his periodic table. 10. Demerits of Merits of Mendeleev’s periodic table:  Position of hydrogen is not correctly defined in periodic table. It is placed in group I though it resembles both group 1 and 17.  In certain pairs of elements increasing order of atomic masses was not obeyed. For example argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1)  Isotopes were not given separate places in the periodic table although Mendeleev's classification is based on the atomic masses.  Some similar elements are separated and dissimilar elements are grouped together. For example copper and mercury resembled in their properties but had been placed in different groups. On the other hand lithium and copper were placed together although their properties are quite different.  m Mendeleev did not explain the cause of periodicity among the elements. c o . Lanthanoids and actinoids were not given a separated position in the table a m a yn 11. Moseley performed experiments and studied the frequencies of the X- d rays emitted from the elements. With these experiments he concluded that tu atomic number is more fundamental property of an element than its atomic mass. S 12. After Moseley’s experimental results Mendeleev’s periodic law was modified to modern periodic law. 13. According to Modern periodic law the physical and chemical properties of the elements are periodic functions of their atomic numbers. 14. Modern periodic table is also referred to as long form of periodic table 15. Horizontal rows in the periodic table are called periods. 16. Vertical columns in the periodic table are called groups. 17. In the modern periodic table there are 7 periods and 18 groups. 18. The period number corresponds to highest principal quantum number of elements. 19. First period contains 2 elements 20. Second and third period contains 8 elements 21. Fourth and fifth period contains 18 elements 22. Sixth period contains 32 elements 23. In the modern periodic table, 14 elements of both sixth and seventh periods i.e. lanthanoids and actinoids respectively are placed separately at the bottom of the periodic table. 24. Elements with atomic number greater than 92 are called transuranic elements. 25. According to IUPAC, until a new element’s discovery is proved and its m name is officially recognized it is given a temporary name. This nomenclature is based Latin words for their numbers. c o. 26. The interim names of the newly discovered elements are derived by a combining together the roots in order of digits which make up the atomic number and ium is added at the end. m a 27. Notation for the IUPAC nomenclature of elements Digit yn Name Abbreviation 0 d nil n tu 1 un u S 2 bi b 3 tri t 4 quad q 5 pent p 6 hex h 7 sept s 8 oct o 9 enn e 28. The distribution of electrons into orbitals of an atom is called its electronic configuration. 29. The electrons in an orbital are filled according to n+l rule. 30. The number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled. 31. On moving down a group in a periodic table the number of shell increases from 1 to 7 32. Elements in the same group have same number of valence electrons 33. Value of the principal quantum number for the valence or outermost shell gives the period. 34. The first period has principal quantum number n=1, contains two elements and corresponds to K-shell. 35. Since K-shell contains only one orbital (1s) it can accommodate two electrons. Thus there are two elements in K-shell. 36. The second period has principal quantum number n=2, contains eight elements and corresponds to L-shell. m three 2p (6 electrons). c o 37. The 4 orbitals filled in second period are one 2s (with 2 electrons) and. a 38. The third period has principal quantum number n=3, contains eight elements and corresponds to M-shell. m a 39. The four orbitals filled in third period are one 3s ( 2 electrons) and three yn 3p ( 6 electrons). d 40. The fourth period has principal quantum number n=4, contains eighteen tu elements. S 41. The 9 orbitals filled in fourth period are one 4s (2 electrons), five 3d (with 10 electrons) and three 4p (with 6 electrons). 42. Elements from Scandium (Z=21) to Zinc (Z=30) are called 3d transition series of elements or first transition series. 43. The fifth period has principal quantum number n=5, contains eighteen elements. 44. The nine orbitals filled in fifth period are one 5s (2 electrons), five 4d (10 electrons) and three 5p (6 electrons). 45. Elements from Yttrium (Z=39) to Cadmium (Z=48) are called 4d transition series of elements or second transition series. 46. The sixth period has principal quantum number n=6, contains 32 elements. 47. The 16 orbitals filled in sixth period are one 6s (2 electrons), seven 4f (14 electrons), five 5d (10 electrons) and three 6p (6 electrons). 48. The orbitals filled in seventh period are 7s,5f,6d and 7p 49. Elements from lanthanum (Z=57), Hafnium (Z=72) to mercury (Z=80) are called 5d transition series of elements or third transition series. 50. Fourteen elements from Cerium (Z=58) to Lutetium (Z=71) are called elements of inner transition series or lanthanoid series. 51. Fourteen elements from Thorium (Z=90) to Lawrencium (Z=103) are called elements of 5f inner transition series or actinoid series. 52. The 4f and 5f series of elements are placed separately in periodic table m to provide a theoretical justification for periodicity occurring at regular intervals. c o. 53. The modern periodic table is divided into four main blocks – s -block, p- a block, d-block and f-block depending on the type of orbital that are being filled with exception of hydrogen and helium. m a 54. The elements in which last electron enter the s-orbital of their outermost yn energy level are called s-block elements. d 55. The s-block consists of two groups, Group-1 and Group-2. tu 56. The elements of Group-1 are called alkali metals and have ns1 as the S general outer electronic configuration. 57. The elements of Group-2 are called alkaline earth metals and have ns2 as the general outer electronic configuration. 58. The elements in which last electron enter the p-orbital of their outermost energy level are called p-block elements. 59. The p-block elements constitute elements belonging to group 13 to 18. 60. Elements of s-block and p-block are collectively called representative element 61. The outermost electronic configuration of p-block elements varies from ns2np1 to ns2np6 62. Elements of group 18 having ns2np6 configuration are called noble gases. 63. Elements of group 17 are called halogens 64. Elements of group 16 are called chalcogens 65. Number of valence electrons in group =Group number -10 for elements belonging to group 13 to 18 66. Elements in which the last electron enters d-orbitals of penultimate energy level constitute d-block elements. 67. Elements of group 3 to 12 in the centre of periodic table constitute the d-block elements 68. General outer electronic configuration of d-block elements is (n-1)d1-10 ns1-2 69. d-block elements constitute transition series elements. The name m “transition series” is derived from the fact the d-block elements c o represent transition in character from reactive metals (belonging to group1 and 2 constituting s-block) on one side of the periodic table to. non-metals (belonging to group 13 to 18 constituting p-block) on other side of the periodic table. a m 70. Elements in which last electron enters f-orbitals are called f-block elements a yn 71. Elements of Lanthanoid series have general outer electronic d configuration of 4f1-14 5d0-1 6s2 tu 72. Elements of Actinoid series have general outer electronic configuration S of 5f1-14 6d0-1 7s2 73. Elements in lanthanoid and actinoid series are called inner transition series. 74. Metals comprise more than 78 % of all known elements and appear on left hand side of periodic table 75. Non-metals are placed on right hand side of periodic table 76. Metal are characterized by having a tendency to loose electron 77. Non-metals are characterised by having tendency to gain electron 78. In general metallic character increases down the group and decreases along period 79. In general non-metallic increases along a period and increases along group 80. Elements showing properties of both metals and non-metals are called metalloids or semi-metals 81. The recurrence of similar properties of elements after certain regular intervals when they are arranged in order of increasing atomic number is called periodicity. 82. The cause of periodicity of properties of elements is due to the repetition of similar electronic configuration of their atoms in the outermost energy shell after certain regular interval. 83. Covalent radius for a homonuclear molecule is defined as one half of the distance between the centres of nuclei of two similar atoms bonded by single covalent bond. m 84. For heteronuclear molecule covalent radius may be defined as the c o distance between the centre of nucleus of atom and mean position of the shared pair of electrons between the bonded atoms.. a 85. Metallic radius is defined as the one half of the internuclear distance two neighbouring atoms of a metal in a metallic lattice. m a 86. For simplicity term atomic radius is used for both covalent and metallic yn radius depending on whether element is non-metal or a metal. d 87. Atomic radius decrease with increase in atomic number on going from tu left to right in a period. S 88. Atomic radius of elements increase from top to bottom in a group. 89. van der Waals radius is half of the distance between two similar atoms in separate molecules in a solid. 90. Ionic radius may be defined as the effective distance from the nucleus of the ion upto which it has an influence in the ionic bond. 91. A cation is smaller than the parent atom. 92. An anion is larger than the parent atom. 93. On moving from top to bottom in a group in a periodic table ionic radius increases. 94. On moving from left to right in a period in a periodic table ionic radius decrease. 95. Atoms or ions which contain same number of electrons are called isoelectronic species. 96. In case of isoelectronic cations, the cation with a greater positive charge will have a smaller radius because of greater attraction of electrons to nucleus. 97. In case of isoelectronic anions, the anions with a greater negative charge will have a larger radius because repulsion of electrons will outweigh the nuclear charge. 98. A quantitative measure of the tendency of an element to lose electron is given by ionization enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom in ground state. m 99. o The energy required to remove second most loosely bound electron c is called second ionization energy. The value of second ionization. enthalpy is higher than first ionization enthalpy because it is more a difficult to remove an electron from a positively charged ion than a neutral atom. m a yn 100. The effective nuclear charge experienced by the valence electron in an atom will be less than actual nuclear charge on nucleus because d of shielding or screening of valence electron from the nucleus by tu inner core electrons. 101. SOn moving from left to right along a period in periodic table ionisation enthalpy increases. On moving along a period successive electrons are added to orbitals in same quantum level and shielding of nuclear charge by inner core of electrons does not increase to an extent to compensate for increased attraction of electron to nucleus. Thus increasing nuclear charge outweighs shielding across a period. Eventually more energy is required to remove outermost electron. 102. On moving from top to bottom in a group in periodic table ionisation enthalpy decreases. On moving down a group successive shells and are added and outermost electron move further away from nucleus. Due to electrons present in inner shells shielding of nuclear charge increases. Thus along a group shielding outweighs increasing nuclear charge. Eventually less energy is required to remove an outermost electron. 103. Among various groups in a periodic table, Group 18 elements have highest ionization enthalpy and because of stable electronic configuration 104. When an electron is added to neutral gaseous atom to convert it into a negative ion, the enthalpy change accompanying the process is called electron gain enthalpy. 105. Group 17 elements have high negative electron gain enthalpy because they can attain a stable electronic configuration as of noble gases by accepting an electron. 106. In general electron gain enthalpy becomes more negative from left to right in a period. The effective nuclear charge increases from left to right across a period. Thus it is easier to add to add an electron to a smaller atom because the added electron would be on an average closer to positively charged nucleus. 107. m In general electron gain enthalpy becomes less negative as we go be farther away nucleus. c o from top to bottom in a group. This is because added electron would. 108. a Elements like O and F have less negative electron gain enthalpy than the succeeding elements like S and Cl respectively of the same m group. This is because in case of O and F electron is added to smaller a quantum number (n=2) and suffers greater repulsion from electrons yn present in that level. In case of succeeding elements like S and Cl electron is added to n=3.The added electron occupies large region of d space and experiences less repulsion from electrons in that level. 109. tu Ability of an atom in a chemical compound to attract shared electrons S to itself is called electronegativity. 110. The electrons present in the outermost shell are called valence electrons and these electrons determine the valence of atom. 111. Valence of representative element is usually equal to  The number of electrons in the valence shell  8-(the number of electrons in the valence shell) 112. On moving down a group since the number of valence electrons remains the same, all elements exhibit same valence. 113. Oxidation state of an element in a particular compound gives the charge acquired by its atoms on basis of electronegativity consideration from other atoms in the molecule. 114. It is observed that some elements of second period show similarities with elements of third period present diagonally to each other though belonging to different group. This similarity in properties of elements present diagonally is called diagonal relationship. 115. Lithium is diagonally related to Magnesium, beryllium is diagonally related to aluminium and boron is diagonally related to silicon. 116. The anomalous behavior of first element of s and p block elements of each group as compared to other group members is due to following reasons:  Small size of atom  Large charge/radius ratio  High electronegativity  m Non availability of d-orbitals in their valence shell 117. c o In second period, the first element of each group has 4 valence. orbital (2s and 2p) which are available for bonding. Therefore a covalence of first member of each group is only 4. 118. m Elements of p-block in the second period displays greater ability to a yn form p  p multiple bond to itself (e.g., C=C, , N=N, ) and to other second period elements (e.g. C=O,C=N, ,N=O) d compared to subsequent members of the same group. tu S 119. In a periodic table there is high chemical reactivity at two extreme ends and lowest in centre. Maximum chemical reactivity at extreme left (alkali metals) is exhibited by the easy loss of electrons forming a cation and at extreme right (among halogens) shown by gain of electrons forming anion. 120. The elements which readily loose electrons act as strong reducing agent. 121. The elements which readily accept electrons acts as strong oxidizing agent. 122. Tendency of an element to lose or gain electrons is also related to metallic or non-metallic character. 123. Elements in extreme left of the periodic table have a tendency to lose electron and become positively charged. Hence they show metallic character. 124. Elements in extreme right of the periodic table have a tendency to gain electron. Hence they show non-metallic character. 125. Metallic character decreases along a period on moving from left to right in a periodic table. 126. Non-metallic character increases along a period on moving from left to right in a periodic table. 127. Since elements at extreme left of periodic table show metallic character, oxides formed by them are basic. 128. Since elements at extreme right of periodic table show non-metallic character, oxides formed by them are acidic. 129. m Oxides of elements in centre of periodic table are amphoteric or neutral. c o 130.. Transition metals of 3d series are less electropositive than group 1 a and 2 metals. This is because their size is small as compared to group 1 and 2 elements accompanied by higher ionization enthalpy m as compared to group 1 and 2 elements. a yn d tu S  Points to Remember Class: XI Chapter Name: Chemical Bonding and Molecular Structure Top Concepts 1. The attractive force which holds together the constituent particles (atoms, ions or molecules) in chemical species is known as chemical bond. 2. Tendency or urge atoms of various elements to attain stable configuration of eight electrons in their valence shell is cause of chemical combination. 3. The principle of attaining a maximum of eight electrons in the valence shell or outermost shell of atoms is known as octet rule. 4. Electronic Theory: Kossel-Lewis approach to chemical Bonding: Atoms achieve stable octet when they are linked by chemical bonds. The atoms m do so either by transfer or sharing of valence electrons. Inner shell c o electrons are not involved in combination process.. a 5. Lewis Symbols or electron dot symbols: The symbol of the element represents the whole of the atom except the valence electrons (i.e. m a nucleus and the electrons in the linear energy shells). The valence yn electrons are represented by placing dots (.) or crosses (x) around the symbol. d tu 6. Significance of Lewis Symbols: The Lewis symbols indicate the number of S electrons in the outermost or valence shell which helps to calculate common or group valence. 7. The common valence of an element is either equal to number of dots or valence electrons in the Lewis symbol or it is equal to 8 minus the number of dots or valence electrons. 8. The bond formed by mutual sharing of electrons between the combining atoms of the same or different elements is called a covalent bond. 9. If two atoms share one electron pair, bond is known as single covalent bond and is represented by one dash (–). 10.If two atoms share two electron pairs, bond is known as double covalent bond and is represented by two dashes (=) 11. If two atoms share three electron pairs, bond is known as triple covalent bond and is represented by three dashes ( ). 12.The formal charge of an atom in a polyatomic ion or molecule is defined as the difference between the number of valence electrons in an isolated (or free) atom and the number of electrons assigned to that atom in a Lewis structure. It may be expressed as: Formal charge on an atom = in free atom  Number of   Number of   Number of      1   valence electrons    nonbonding m    bonding    (lone pair) electrons  2  (shared) electrons  o  in free atom      . c a 13.Significance of Formal charge: The formal charges help in selection of m lowest energy structure from a number of possible Lewis structures for a a given molecule or ion. Lowest energy structure is the one which has yn lowest formal charges on the atoms. d tu 14.Expanded octet: Compounds in which central atom has more than eight electrons around it, atom is said to possess an expanded octet. S 15.Exceptions to the Octet Rule: Hydrogen molecule: Hydrogen has one electron in its first energy shell (n = 1). It needs only one more electron to fill this shell, because the first shell cannot have more than two electrons. This configuration (1s2) is similar to that of noble gas helium and is stable. In this case, therefore, octet is not needed to achieve a stable configuration. Incomplete octet of the central atom: The octet rule cannot explain the formation of certain molecules of lithium, beryllium, boron, aluminum, etc.  (LiCl, BeH2, BeCl2, BH3, BF3) in which the central atom has less than eight electrons in the valence shell as shown below: Expanded octet of the central atom: There are many stable molecules which have more than eight electrons in their valence shells. For example, PF5, has ten; SF6 has twelve and IF7 ha fourteen electrons around the central atoms, P, S, and I respectively. Odd electron molecules: There are certain molecules which have odd number of electrons, like nitric oxide, NO and Nitrogen dioxide, NO2. In these cases, octet rule is not satisfied for all the atoms. It may be noted that the octet rule is based upon the chemical inertness of noble gases. However, it has been found that some noble gases (especially xenon and krypton) also combine with oxygen and fluorine to form a large m o number of compounds such a XeF2, KrF2, XeOF2, XeOF4, XeF6, etc. c. a This theory does not account for the shape of the molecules. m It cannot explain the relative stability of the molecule in terms of the a yn energy. d 16.General Properties of Covalent Compounds tu 1. The covalent compounds do not exist as ions but they exist as S molecules. 2. The melting and boiling points of covalent compounds are generally low. 3. Covalent compounds are generally insoluble or less soluble in water and other polar solvents. However, these are soluble in non- polar solvents. 4. Since covalent compounds do not give ions in solution, these are poor conductors of electricity in the fused or dissolved state.  5. Molecular reactions are quite slow because energy is required to break covalent bonds. 6. Since the covalent bond is localized in between the nuclei of atoms, it is directional in nature. 17.Co- Ordinate Covalent Bond: Covalent type bond in which both the electrons in the shared pair come from one atom is called a coordinate covalent bond. Co- Ordinate Covalent Bond is usually represented by an arrow (→) pointing from donor to the acceptor atom. Co- Ordinate Covalent bond is also called as dative bond, donor – acceptor m bond, semi- polar bond or co-ionic bond. c o. a 18. The electrostatic force of attraction which holds the oppositely charged m ions together is known as ionic bond or electrovalent bond. a yn 19.Ionic compounds will be formed more easily between the elements with comparatively low ionization enthalpy and elements with comparatively d tu high negative value of electron gain enthalpy. S 20.A quantitative measure of the stability of an ionic compound is provided by its lattice enthalpy and not simply by achieving octet of electrons around the ionic species in the gaseous state 21.Lattice enthalpy may also be defined as the energy required to completely separating one mole of a solid ionic compound into gaseous ionic constituents. 22. Factor affecting lattice enthalpy: Size of the ions: Smaller the size of the ions, lesser is the internuclear distance and higher will be lattice enthalpy.  Larger the magnitude of charge on the ions, greater will be the attractive forces between the ions. Consequently, the lattice enthalpy will be high. 23. General Properties of Ionic Compounds Ionic compounds usually exist in the form of crystalline solids. Ionic compounds have high melting and boiling points. Ionic compounds are generally soluble in water and other polar solvents having high dielectric constants. Ionic compounds are good conductors of electricity in the solutions or in their molten states. The chemical reactions of ionic compounds are characteristic of the constituent ions and are known as ionic reactions. m c o. In ionic – compounds, each ion is surrounded by oppositely charged ions a uniformly distributed all around the ion and therefore, electrical field is non- directional. m a yn 24.Bond length: It is defined as the average distance between the nuclei of d the nuclei of two bonded atoms in a molecule. tu 25. Covalent radius is half of the distance between two similar atoms joined S by single covalent bond in same molecule. 26. Van der Waals radius is one half of the distance between two similar adjacent atoms belonging to two nearest neighbouring molecules of the same substance in the solid state. It is always larger than covalent radii. 27. Bond angle: It is defined as the average angle between orbitals containing bonding electron pairs around the central atom in a molecule. 28. Bond enthalpy: It is defined as amount of energy required to break one mole of bonds of a particular type between atoms in gaseous state. 29.Bond order: The bond order is defined as the number of bonds between two atoms in a molecule. 30. When a single Lewis structure cannot determine a molecule accurately, concept of resonance is used wherein a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as canonical structures of hydrid which describes molecule accurately. 31. Resonance: When a molecule cannot be represented by a single structure but its characteristic properties can be described by two or more than two structures, then the actual structure is said to be a resonance hybrid of these structure. m c o 32. Polarity of Bonds: In reality no bond is completely covalent or completely. a ionic. m 33.Non-polar covalent bond: When a covalent bond is formed between two a yn similar atoms, the shared pair of electrons is equally attracted by the two atoms and is placed exactly in between identical nuclei. Such a bond is d called non-polar covalent bond tu S 34.Molecules having two oppositely charged poles are called polar molecules and the bond is said to be polar covalent bond. Greater the difference in the electro-negativity of the atoms forming the bond, greater will be the charge separation and hence greater will be the polarity of the molecule. 35.Dipole moment is defined as the product of the magnitude of the charge and the distance of separation between the charges. Dipole moment (µ) = charge (q) x distance of separation (d) 36. Partial Covalent Character in Ionic Bonds: When two oppositely charge ions A+ and B- are brought together; the positive ion attracts the outermost electrons of the negative ion. This results in distortion of  electron clouds around the anion towards the cation. This distortion of electron cloud of the negative ion by the positive ion is called polarization. 37. Tendency of cation to polarize and polarisability of anion are summarized as Fajan’s rules: a. Smaller the size of the cation, greater is its polarizing power. b. Polarisation increases with increase in size of anion. This is because the electron cloud on the bigger anion will be held less firmly by its nucleus and, therefore, would be more easily deformed towards the cation. c. Larger the charge on cation greater is polarizing power and larger the charge on anion greater is its tendency to get polarized. m c o. 38.Valence Shell Electron Pair Repulsion (VSEPR) Theory: a Since Lewis symbols were unable to explain shapes of certain molecules, m VSEPR theory was introduced.The basic idea of this theory is that a yn bonded atoms in a molecule adopt that particular arrangement in space around the central atom which keeps them on the average as far apart as possible. d tu S 39.Geometry and shapes of molecules in which central atom has no lone pair of electrons Number of Arrangement of Molecular Examples electron pairs electron pairs geometry BeCl2,HgCl2 2 Linear 3 Trigonal planar BF3 4 Tetrahedral CH4,NH4+ m c o. a m 5 a Trigonal PCl5 yn bipyramidal d tu S 6 Octahedral SF6 Shapes of simple molecules/Ions with central ions having one or more lone pairs of electrons Molecule No. of No. of Arrangement of Shape Example type bonding lone pairs electron pairs pairs AB2E 2 1 Bent SO2,O3 Trigonal pyramidal NH3 AB3E 3 1 m c o. a m a yn AB2E2 2 d 2 Bent H2O tu S AB4E 4 1 See saw SF4  AB3E2 3 2 T-shape ClF3 Square AB5E 5 1 pyramidal BrF5 Square AB4E2 4 2 planar XeF4 m c o. a A: Central atom, B is surrounding atoms, E is the lone pair m a yn 40. Valence Bond Approach of Covalent Bond d The VSEPR theory gives the geometry of simple molecules but theoretically, tu it does not explain them and also has limited applications. To overcome S these limitations, two important theories based on quantum mechanical principles are commonly used. These are Valence bond (VB) theory and Molecular orbital (MO) theory. 41. Valence Bond Theory  A discussion of valence bond theory is based on the knowledge of atomic orbitals, electronic configuration of elements, overlap criteria of atomic orbitals and principles of variation and superposition.  Orbital Overlap Concept of Covalent Bond: When two atoms approach each other, partial merger of two bonding orbitals, known as overlapping of the orbitals occurs.   Depending upon the type of overlapping, the covalent bonds may be divided as sigma (σ) bond and Pi (  ) bond.  Sigma (σ) bond: This type of covalent bond is formed by the end to end (hand on) overlapping of bonding orbitals along the inter-nuclear axis. The overlap is known as head on overlap or axial overlap. The sigma bond is formed by any one of the following types of combinations of atomic orbitals. Sigma (σ) bond can be formed by s – s overlapping, s – p overlapping, p – p Overlapping etc.  Pi (  ) Bond: This type of covalent bond is formed by the sidewise overlap of the half- filled atomic orbitals of bonding atoms. Such an overlap is known as sidewise or lateral overlap. 42.Hybridization:  m In order to explain characteristic geometrical shapes of polyatomic molecules concept of hybridization is used. c o.  a The process of intermixing of the orbitals of slightly different energies so as m to redistribute their energies resulting in the formation of new set of a orbitals of equivalent energies and shape. yn d tu 43. Atomic orbtials used in different types of hybridization. Shapes of S molecules/io Hybridisa tion type Atomic orbitals Examples ns Linear sp one s + one p BeCl2 Trigonal sp2 one s + two p BCl3 planar Tetrahedral sp3 one s + three p CH4,NH3 Square [Ni(CN)4]2- planar dsp2 one d +one s +two p [Pt(Cl)4] 2- Trigonal sp3d one s+ three p + one d PF5,PCl5 bipyramidal Square sp3d2 one s +three p +two d BrF5 pyramidal Octahedral sp3d2 one s +three p +two d SF6,[CrF6]3- , d2sp3 two d + one s +three p [Co(NH3)6]3+ 44.Molecular Orbital Theory (MOT):  Basic idea of MOT is that atomic orbitals of individual atoms combine to form molecular orbitals. Electrons in molecule are present in the molecular orbitals which are associated with several nuclei.  The molecular orbital formed by the addition of atomic orbitals is called the bonding molecular orbital(  ). m o  The molecular orbital formed by the subtraction of atomic orbital is called c antibonding molecular orbital (  *).. a  The sigma (  ) molecular orbitals are symmetrical around the bond-axis m while pi (  ) molecular orbitals are not symmetrical. a yn  Sequence of energy levels of molecular orbitals changes for diatomic d molecules like Li2, Be2, B2, C2, N2 is 1s < *1s < 2s < *2s < (2px = 2py) tu 1.This implies m reaction is spontaneous or the reaction proceeds in the forward direction to form products predominantly. a yn G G - -   If, G >0 then RT is negative, and e RT 7.00 tu Neutral [H+] = 1.0 × 10–7M pH = 7.00 S 35. pKw  pH  pOH  14 36.Dissociation constant or ionization constant ( K a ) of a weak acid (HX) in water is given by equation shown below. [H+ ][X- ]  2c Ka=  [HX] (1- ) At a given temperature T, K a is a measure of the strength of the acid HX i.e., larger the value of K a ,the stronger is the acid. 37. Relation between K a and pK a :  pKa   logKa Larger the value of pK a , weaker is the acid 38.Dissociation constant or ionization constant ( Kb ) of a weak base (MOH) is water is given by equation shown below. [M+ ][OH- ]  2c Kb =  [MOH] (1- ) At a given temperature T, Kb is a measure of the strength of the base MOH i.e., larger the value of Kb , the stronger is the base. 39. The degree of ionization is defined as the extent of ionization 40.Relation between Kb and pKb pKb   logKb m Larger the value of pKb , weaker is the base. 41. Relation between K a and Kb c o. Ka  Kb  Kw a m 42. pKa  pKb  pKw  14 (at 298K) a yn 43.Bronsted acids which can donate more than one protons are called polyprotic d acids or polybasic acids. tu 44.Higher order higher order ionization constants ( K a2 K a3 ) are smaller than S , the lower order ionization constant ( Ka1 ) of a polyprotic acid. The reason for this is that it is more difficult to remove a positively charged proton from a negative ion due to electrostatic forces. 45.Factors affecting acid strength:  Extent of dissociation of acid (HA) depends on strength and polarity of H-A bond in acid.  As strength of H-A bond decreases acid strength increases  As the electronegativity difference between the atoms H and A increases ,charge separation between atoms increases, cleavage of the bond becomes easy. Eventually acid dissociates easily increasing acidity. 46. Common ion Effect: If to an ionic equilibrium, a salt containing a common ion is added, the equilibrium shifts in the backward direction. The shift in equilibrium position caused by the addition or presence of an ion involved in the equilibrium reaction is known as common ion effect 47.Interaction of anion or cation of the salt with water to produce an acidic or basic solution is called hydrolysis of salt. 48.Hydrolysis of Salts and the pH of their Solutions Type of acid and base Example Type of solution formed used to form salt by hydrolysis of salt Salt of strong acid and NH4Cl Acidic solutions weak base Salt of weak acid and CH3COONa Basic solutions strong base Salt of weak acid and CH3COONH4 Neutral solutions weak base m Salt of strong acid and strong base NaCl c o Neutral solutions. a m a yn 49. A buffer solution resists a change in pH caused by dilution or the addition of limited amounts of acid or base. d tu 50. Acidic buffers contain equimolar quantities of a weak acid and one of its salt S with a strong base 51. Basic buffers contain equimolar quantities of a weak base and one of its salt with a strong acid 52.Solubility is a measure of amount of solute that can be dissolved in a given amount of solvent at a specific temperature 53. In a saturated solution equilibrium exists between the undissolved solute and the solute in the solution. 54. Classification of salts on the basis of solubility 55. Solubility product of a salt at a given temperature is equal to product of concentration of the ions in a saturated solution with each concentration term raised to power equal to number of ions produced on dissociation of one mole of substance. In general if sparingly soluble salt , AxBy is in equilibrium with saturated solution of its ions, then AxBy  xAy   yBx  (solid) Solubility expression is: Ksp  [Ay  ]x [Bx  ]y m c o. 56.Ionic product is defined as the product of the molar concentrations of the a constituent ions, each raised to the power of its stoichiometric coefficient in a any solution. m a yn 57. Precipitation will occur only when ionic product exceeds the solubility product d tu S  Subject: Chemistry Class: XI Chapter: Redox Reactions Top concepts 1. Redox reactions are those reactions in which oxidation and reduction takes place simultaneously 2. Classical view of redox reactions  Oxidation is addition of oxygen / electronegative element to a substance or removal of hydrogen / electropositive element from a substance  Reduction is removal of oxygen / electronegative element from a substance or addition of hydrogen / electropositive element to a substance m c o 3. Redox reactions in terms of Electron transfer. a  Oxidation is defined as loss of electrons by any species  Reduction is defined as gain of electrons by any species m a yn 4. In oxidation reactions there is loss of electrons or increase in positive charge or decrease in negative charge d tu S 5. In reduction reactions there is gain of electrons or decrease in positive charge or increase in negative charge 6. Oxidising agents are species which gain one or more electrons and get reduced themselves 7. Reducing agents are the species which lose one or more electrons and gets oxidized themselves 8. Oxidation number denotes the oxidation state of an element in a compound ascertained according to a set of rules. These rules are formulated on the basis that electron in a covalent bond belongs entirely to the more electronegative element.  9. Rules for assigning oxidation number to an atom  Oxidation number of Hydrogen is always +1 (except in hydrides, it is -1).  Oxidation number of oxygen in most of compounds is -2. In peroxides it is (- 1). In superoxides, it is (-1/2). In OF2 oxidation number of oxygen is +2.In O2F2 oxidation number of oxygen is +1  Oxidation number of Fluorine is -1 in all its compounds  For neutral molecules sum of oxidation number of all atoms is equal to zero  In the free or elementary state, the oxidation number of an atom is always m zero. This is irrespective of its allotropic form c o . For ions composed of only one atom, the oxidation number is equal t the charge on the ion a m a yn  The algebraic sum of the oxidation number of all the atoms in a compound must be zero d tu  S For ions the sum of oxidation number is equal to the charge on the ion  In a polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must be equal to the charge on the ion 10. Oxidation state and oxidation number are often used interchangeably 11. According to Stock notation the oxidation number is expressed by putting a Roman numeral representing the oxidation number in parenthesis after the symbol of the metal in the molecular formula 12.Types of Redox Reactions   Combination Reactions: Chemical reactions in which two or more substances (elements or compounds) combine to form a single substance  Decomposition Reactions: Chemical reactions in which a compound break up into two or more simple substances  Displacement Reactions: Reaction in which one ion(or atom)in a compound is replaced by an ion(or atom) of other element a. Metal Displacement Reactions: Reactions in which a metal in a compound is displaced by another metal in the uncombined state m c o b. Non-metal Displacement Reactions: Such reactions are mainly. hydrogen displacement or oxygen displacement reactions a m  a Disproportionation Reactions: Reactions in which an element in one yn oxidation state is simultaneously oxidized and reduced d tu 13.Steps involved in balancing a Redox reactio