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MODERN PERIODIC LAW, According to modern periodic law: “The properties of elements are the periodic functions of their atomic numbers”, Cause of periodicity: It is due to the repetition of similar outer shell electronic configuration at a certain regular intervals., Structural features of the long form of the periodic table., On the basic of the modern periodic law, a scientist named Bohr proposed a long form of periodic table that was prepared by Rang and Warner., It consists of 18 vertical columns called groups and 7 horizontal rows called periods., Elements of groups 1, 2, 13 – 17 are called normal or representative elements., Elements of groups 3 – 12 are called transition elements., The 14 elements with atomic numbers (Z) = 58 – 71 (occurring after lanthanum 57La in the periodic table) are called lanthanides or rare earth elements and are placed at the bottom of the periodic table. The 14 elements with atomic numbers (Z) = 90 – 103 (Occurring after actinium 89Ac in the periodic table) are called actinides and are placed at the bottom of the periodic table., The Eleven elements with Z = 93 – 103 (93Np – 103Lr) which occur in the periodic table after uranium and have been prepared from it by artificial means are called transuranic. These are all radioactive elements., The elements belonging to a particular group are said to constitute a chemical family which is usually named after the name of the first element. For example, Boron family (group 13), carbon family (group 14), nitrogen family (group 15), and oxygen family (group 16). In addition to this, some groups have typical names. For example,, Elements of group 1 are called alkali metals., Elements of group 2 are called alkaline earth metals., Elements of group 16 are called chalcogens., Elements of group 17 are called halogens., Elements of group 18 are called zero group or noble gases., The long form of the periodic table contains seven periods. These are :, 1st period (1H – 2He) contains only two elements. This is the shortest period., 2nd period (3 Li – 10Ne) contain 8 elements each and are called short periods., Third period (11Na – 18Ar) contain 8 elements each and are called short periods., 4th period (19K – 36Kr) and 5th period (37Rb – 54Xe) contain 18 elements each and are called long periods., 6th period (55Cs – 86Rn) contains 32 elements and is the longest period., 7th period (87Fr –) is, however, incomplete and contains at present only 24 elements., The long form of the periodic table has been divided into four blocks (i.e., s, p, d and f) depending upon the subshell to which the last electron enters., CLASSIFICATION IN BLOCKS, s-block elements. Elements of groups 1 and 2 in which the last electron enters the s-orbital of the valence shell are called s-block elements. There are only 12 s-block elements in the periodic table., p-block elements. Elements of groups 13–18 in which the last electron enters the p-orbitals of the valence shell are called p-block elements., d-block elements (Transition elements): There are three complete series and one incomplete series of d-block elements. These are:, 1st or 3d transition series which contains ten elements which atomic numbers 21–30 (21Sc – 30Zn)., 2nd or 4d-transition series which contains ten elements with atomic numbers 39 – 48 (39Y – 48Cd)., 3rd or 5d transition series which also contains ten elements which atomic numbers 57 and 72 – 80 (57La, 72Hf – 80Hg)., 4th or 6d transition series which is incomplete at present and contains only nine elements. These are 89Ac, 104Rf, 105Ha, Unh (Unnihexium, Z = 106), 107Ns (Neilsobohrium), 108Hs (Hassium), 109Mt (Meitherium), Uun (Ununnilium, Z = 110) and Uud (Unundium, Z = 112) or Ekamercury. The element, Z = 111 has not been discovered so far. Thus, in all there are 39 d-block elements., f-Block elements are called inner-transition elements. In these elements, the f-subshell of the anti-penultimate is being progressively filled up. There are two series of f-block elements each containing 14 elements. The fourteen elements from 58Ce – 71Lu in which 4 f-subshell is being progressively filled up are called lanthanides or rare elements. Similarly, the fourteen elements from 90Th – 103Lr in which 5 f-subshell is being progressively filled up are called actinides., Periodic Properties:, Properties which are directly or indirectly related to their electronic configuration and show a regular gradation when we move from left to right in a period or from top to bottom in a group are called periodic properties., Some important periodic properties are atomic size, ionization energy, electron affinity, electronegativity, valency, density, atomic volume, melting and boiling points etc., Effective Nuclear Charge, • In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons., • The nuclear charge that an electron “feels” depends on both factors., • It’s called Effective nuclear charge., • Electrons in lower energy levels “shield” outer electrons from positive charge of nucleus., The effective nuclear charge, Zeff is: Zeff = Z − S, Where: Z = atomic number, S = screening constant, Atomic radius or Atomic Radii is the total distance from the nucleus of an atom to the outermost orbital of its electron., Types of Atomic Radii, Atomic radii are divided into three types:, Metallic radius, Van der Waals radius, METALLIC RADIUS, Half of the distance between the nuclei of two nearest packed atom is called metallic radius, metallic radius > covalent radius, For example:, K(metallic) = 2.34 Å, K(covalent) = 2.03 Å, ATOMIC OR COVALENT RADIUS:, It may be defined as one half of the distance between the centers of the nuclei of two similar atoms bonded by a single covalent bond., VANDER WALL’S RADIUS:, Vander wall’s radius may be defined as one half of the distance between the nuclei of two adjacent atoms belonging to two neighboring molecules of an element in solid state., Thus rcovalent < rmetallic < rvan der walls, VARIATION OF ATOMIC RADIUS IN GROUP AND PERIOD, IN PERIODS, On moving from left to right along the period:-, Magnitude of nuclear charge increases with increase in atomic number, Addition of electrons takes place in same principal shell (Subshell), Electrons in the same shell do not screen each other from the nucleus, Increase in nuclear charge is not neutralized by the extra valence electron, Consequently, the effective nuclear charge increases steadily form left to right., As the effective nuclear charge increases, the outermost electrons are held more strongly by the nucleus, Thus the atomic radius decreases steadily form left to right., IN GROUPS, On moving from top to bottom in a group:-, Nuclear charge increases with increase in atomic number, Also there is increase in the value of Principle Quantum Level, although the number of electrons in outermost shell remains the same, The effect of increase in the size of electron cloud is more pronounced than the effect of increased nuclear charge, Distance between nucleus and outermost electron gradually increases down the group, Consequently, the effective nuclear charge decreases steadily down the group., Thus, atomic radius increases on moving from top to bottom in a group., IONIC RADIUS, Ions are formed when an atom loses or gains electrons. When an atom loses an electron it forms a cation and when it gains an electron it becomes an anion., Ionic radius is the radius of a cation or an anion., Ionic radius is the distance from the nucleus of an ion up to which it has an influence on its electron cloud., When a neutral atom is converted to an ion, there would be change in size, as the Zeff will change but the number of protons in nucleus remains same., If the atom forms an anion, its size (or radius) increases, since the nuclear charge remains the same but the repulsion resulting from the additional electron(s) enlarges the domain of the electron cloud., On the other hand, if one or more electrons are removed from an atom, it reduced the electron-electron repulsion but the nuclear charge remains the same, so the electron cloud shrinks and the cation is smaller than the atom., When a lithium atom reacts with a fluorine atom to form a LiF unit, the changes in size are very peculiar. Out of Li and F, Li is bigger in size. When lithium changes to Li+, its size decreases and when F changes to F-, its size increases. In LiF, Li+ is smaller than F- (Note that Li+ is smaller than F atom and F- is smaller than Li atom)., Cation is smaller than its corresponding atom, reason        1. Increase in effective nuclear charge., 2. Vanishing of outer most orbit.                        i.e. Li+ < Li, Anion is bigger than corresponding atom, reason       1. Decrease in effective nuclear charge., 2. Increase in interelectronic repulsion.              i.e. F- > F, VARIATION OF IONIC RADIUS IN GROUP AND PERIOD, Variation of Ionic Radius in a Group, In a periodic table while moving down in a group, atoms add extra shell (number of electrons) due to which ionic radius of elements increases down a group., Ions Configuration Ionic radii (nm) Ions Configuration Ionic radii (nm), Li+ 2 0.076 F– 2, 8 0.133, Na+ 2, 8 0.102 Cl– 2, 8, 8 0.181, K+ 2, 8, 8 0.138 Br– 2, 8,18,8 0.196, Variation of Ionic Radius in a period, For an isoelectronic series, the ionic radius decreases as the nuclear charge increases, Size in an isoelectronic series 1/Nuclear charge (No. of protons), Isoelectronic ions:- Ions having same number of electrons are known as isoelectronic ions., For example, N3-, O2- , F-, Na+, Mg2+, Al3+ are isoelectronic ions. All ions have 10 electrons each with same electronic configuration 1s2, 2s2, 2p6., As we move from left to right:-, nuclear charge increases, number of electrons remains same, effective nuclear charge per electron increases, electrons are attracted and pulled more and more strongly towards the nucleus, ionic radii decreases, In isoelectronic ions when nuclear charge increases, nuclear attraction for same number of electrons also increases. In other words we can say for isoelectronic ions, magnitude of effective nuclear charge per electron (nuclear charge actually experienced by electrons) increases with increase in nuclear charge. As a result electrons are attracted and pulled more and more strongly towards the nucleus. This causes a decrease in the ionic radii., If we examine isoelectronic ions, cations are smaller than anions., Let us compare the radius of Na+ ion and F- ion., Both ions have same number of electrons (10)., But Na (Z = 11) has more protons than F (Z = 9)., Thus, Zeff of Na+ is more than that of F-., So Na+ ion is smaller in size than F- ion., Similarly, for the three isoelectronic ions of third period, Al3+, Mg2+ and Na+ they all have the same number of electrons (10) but their number of protons are 13, 12 and 11 respectively. Thus, the electron cloud in Al3+ is pulled inward more than that in Mg2+ and Mg2+ would be smaller than Na+., Thus in general in an isoelectronic cation series, the radii of tripositive ions are smaller than those of dipositive ions, which in turn are smaller than unipositive ions., Similarly, in isoelectronic anions series, the radius, increases as we go from uninegative ion to dinegatie ion and so on., IONIZATION ENTHALPY, Ionization energy can be described as a measure of the difficulty in removing an electron from an atom or ion or the tendency of an atom or ion to surrender an electron. The loss of electron usually happens in the ground state of the chemical species., Alternatively, we can also state that ionization or ionisation energy is the measure of strength (attractive forces) by which an electron is held in a place., Ionization energy is the minimum energy that an electron in a gaseous atom or ion has to absorb to come out of the influence of the nucleus. It is also sometimes referred to as ionization potential and is usually an ., Ionization energy is the amount of  required to remove the most loosely bound  from an isolated gaseous, . H(g)    H+(g) + e- Ho = 1312.0 kJ/mol, It is measured either in units of electron volts or kJ/mol., Factors on which Ionization Energy depends, Size of the atom, Charge on the nucleus, Screening effect of the inner electrons, Penetration effect of electrons, Electronic arrangement, 1. Size of atom (distance of outermost electron from the nucleus), As atomic size increases, the electrostatic force of attraction between the positive nucleus and the outermost electron decreases and less energy is required to remove an electron., This means that the ionisation energy decreases., 2. Effective Nuclear charge, As the nuclear charge increases, electrostatic force of attraction between the nucleus and the outermost electron increases. So the greater energy will be required to remove the electron. This means that the ionisation energy increases., Ionisation potential ∝ Effective nuclear charge (Z eff), 3. Screening (shielding) effect of inner shell electrons, The shielding or screening effect increases if the number of electrons in the inner shells between the nucleus and the outermost electrons increases. This results in decrease of force of attraction between the nucleus and the outermost electron and lesser energy is required to separate the electron. Thus the value of I.P. decreases., Ionisation potential ∝ (1/Shielding or screening effect), 4. Penetration effect of orbitals:, The order of energy required to remove electron from s, p, d-and f-orbitals of a shell is s > p > d > f because the distance of the electron from the nucleus increases., For example:, The value of ionisation potential of Be (Z = 4, Is2 2s2) is more than the ionisation potential of B (Z = 5, 1s2 2s2 2p1x)., Value of ionisation potential of Mg (Z = 12, 1s2 2s2 2p6 3s2) is more than the ionisation potential of Al (Z=13,1s2 2s2 2p6 3s23p1), This is because the penetration power of 2s and 3s electrons is more than 2p and 3p orbitals respectively. More energy will be required to separate the electrons from 2s and 3s orbitals., 5. Electronic arrangement (Stability of half−filled and fully-filled orbitals):, According to Hund’s rule the stability of half-filled or completely filled degenerate orbitals is comparatively high. So comparatively more energy is required to separate the electron from such atoms., For example:, (a) Removal of electron is comparatively difficult from the half-filled configuration of N (Z = 7, Is2 2s22p3)., (b) The ionisation potential of inert gases is very high due to most stable ns2p6 electronic configurations., Periodicity in ionisation potential, On moving from left to right in a period,, Effective nuclear charge increases., Atomic size decreases, Value of ionisation potential of elements increases., Exceptions, In a period, the ionisation energy of group 2 elements is more than the elements of group 3 because penetration power of s-orbitals electrons., The value of ionisation energy of Be (1s2 2s2) is more than B (1s2 2s2 2p1) because the penetration power of 2s-electrons of Be is more than the 2p electrons of B., In a period, the ionisation energy of group 15 elements is more than the elements of group16 because the half-filled (N: [He] 2s2 2p3) configuration of group 15 elements is comparatively of higher stability., Group 16 elements (O: [He] 2s2 2p4) have the tendency to acquire comparatively more stable (O+: [He] 2s2 2p3) configuration by the loss of one electron., Ionisation energy of N(1s2 2s2 2p3) > O (1s2 2s2 2p4), Similarly Ionisation energy of P>S & As>Se., Variation of Ionization Energy in a group, On moving from top to bottom in a group:-, Nuclear charge increases with increase in atomic number, There is gradual increase in atomic size due to increase in the value of Principle Quantum number (shell), There is an increase in shielding effect on the outermost electron due to an increase in the number of inner electrons., The effect of increase in atomic size and the shielding effect are much more than the effect of increase in nuclear charge. As a result, the electron becomes less and less firmly held to the nucleus., Effective Nuclear charge decrease, Hence, there is gradual decrease in the ionization energy in a group., What are successive ionisation enthalpies? Explain why the second ionisation enthalpy is higher than the first ionisation enthalpy?, Successive ionisation enthalpies: The enthalpy required to remove the first loosely bound electron from gaseous isolated atom is called the first ionisation enthalpy (IE1). The enthalpies required to remove second, third and fourth electrons are called second (IE2), third (IE3) or fourth (IE4) ionisation enthalpies respectively. The amounts of enthalpies required to remove first, second and subsequent electrons from the gaseous atom one after the other are collectively called successive ionisation enthalpies., The first four ionisation energies of aluminium, for example, are given by, Theoretically speaking, there are as many ionisation enthalpies for an atom as there are electrons in it. In the above example IE1 < IE2 < IE3< IE4., Second ionisation enthalpy is always greater than the first ionisation enthalpy. After the removal of first electron, the atom changes into monovalent positive ion (M+). In the ion (M+), the number of electrons decreases but the nuclear charge remains the same as the parent atom. As a result the attraction of the nuclear charge (protons) increases over the remaining electrons. Hence, more enthalpy is required to remove the second electron. In other words, the second ionisation enthalpy is greater than the first ionisation enthalpy (IE2 > IE1). Similarly, the removal of second electron results in the formation of divalent positive ion (M2+ ) and the attraction between the nucleus and remaining electrons increases further. This accounts for the progressive increase in the value of ionisation enthalpies., Electron Affinity, What is Electron Affinity?, is the energy required to remove an electron from a gaseous atom. Energy is supplied for removing an atom implies that energy will be released if an extra electron is added to the atom., Electron affinity is defined as, The amount of energy released when an electron is added to a neutral gaseous atom to form an anion., The electron affinity is the potential energy change of the atom when an electron is added to a neutral gaseous atom to form a negative ion., So the more negative the electron affinity the more favourable the electron addition process is., Not all elements form stable negative ions in which case the electron affinity is zero or even positive., X (g)  +  e−  →  X− (g)  +  E.A., For example,, Cl (g)  +  e− → Cl− (g)  +  349 KJ/mol, The electron affinity of chlorine is – 349 KJ/mol., The unit of electron affinity is electron volts per atom or kilojoule per mole. It is represented by a negative sign [-]., Energy is released when the first electron is added to an atom and leads to the formation of a monovalent anion and this is known as the first electron affinity. Now if we add another electron to this anion, a force of repulsion is experienced by the electron. The second electron has to be forced to enter the mono negative ion, and energy is absorbed. Therefore, the second electron affinity and further affinities are positive in nature., The first electron affinity is always exothermic that is negative. The second electron affinity of the same element will be positive or endothermic., The electron affinity cannot be determined directly but is obtained indirectly from the Born-Haber cycle., Factors Affecting Electron Affinity, The general factors that affect the electron affinity are listed below., Atomic size: If the atomic size is small, then there will be greater electron gain enthalpy because the effective nuclear forces will be greater in the smaller atoms and the electrons will be held firmly., Nuclear charge: The greater the nuclear charge more will be the value for electron gain enthalpy because an increase in nuclear charge will increase the effective nuclear force on valence electrons., ∝ Effective nuclear charge (Z eff), Stability of half−filled and fully-filled orbitals: The stability of the configuration having fully-filled orbitals (p6, d10, f14) and half-filled orbital (p3, d5, f7 ) is relatively higher than that of other configurations., Screening Effect (Shielding Effect): Shielding effect is directly proportional to atomic size and atomic size is inversely proportional to electron affinity., Electron affinity Reactivity of non−metals, Electron affinity Oxidising power of element, VARIATION OF ELECTRON AFFINITY IN GROUP AND PERIOD, IN PERIOD, On moving from left to right along the period:-, Effective nuclear charge increases., As the effective nuclear charge increases, the outermost electrons are held more strongly by the nucleus, Thus the atomic radius decreases., Incoming electron experience greater attraction, Electron affinity increases as we move from left to right across a period., IN GROUP, On moving from top to bottom in a group:-, Nuclear charge increases with increase in atomic number., Atomic size increases due to increase in the value of Principle Quantum number (shell)., The effect of increase in atomic size is much more than the effect of increase in nuclear charge., Incoming electron experience less attraction, Hence, there is gradual decrease in the value of  down the group., In general electron affinity follows the following trends:, Halogens > Oxygen family > Carbon family > Nitrogen family > Metals of group 1 and 13 > Metals of group 2, IMPORTANT TRENDS IN ELECTRON GAIN ENTHALPIES, Why do halogens have very higher value of electron gain enthalpy?, This is due to the fact that halogens have the general electronic configuration of ns2 np5 and have only one electron less than the stable noble gas (ns2 np6) configuration. In order to acquire the noble gas configuration, halogens have a very strong tendency to accept an additional electron and their electron gain enthalpies become more negative., Why is electron affinity of Fluorine less than that of Chlorine?, Electron affinity of fluorine is less than chlorine because the atomic size of fluorine is very small and compact than chlorine as a result there is a large electronic repulsions between the electrons of fluorine. Hence, in fluorine atom the incoming electron experiences lesser attractions towards the nucleus than in chlorine atom. That is why the value of electron affinity of Cl is highest in the periodic table., Why do noble gases have positive electron gain enthalpy?, Noble gases have general electronic configuration of:, The sub-shells of noble gases are fully filled and their octet is complete. Due to the stable electronic configuration, they do not accept an extra electron to form negative ion., Also electron has to enter the next higher principal quantum level which leads to a very unstable electronic configuration., Therefore they have large positive electron gain enthalpy., Electron affinity of Be, Mg, N and P are almost zero., The value of electron affinity of the elements of group 2 is zero because ns orbitals are fully-filled and such orbitals have no tendency to accept electrons., Electronic configuration of Be: 1s2 2s2, Electronic configuration of Mg: 1s2 2s22p63s2, The value of electron affinity of the elements of group 15 is zero because of half-filled (np3) orbitals in the outermost orbit of group 15 elements, which are more stable., Electronic configuration of N: 1s2 2s22p3, Electronic configuration of P: 1s2 2s2 2p63s2 3p3, ELECTRONEGATIVITY, What is Electronegativity?, The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is known as electronegativity., Electronegativity is a relative value that indicates the tendency of an atom to attract shared electrons more than the other atom bonded to it. Therefore it does not have any unit or it is a dimensionless property., Pauling was the first scientist to put forward the concept of electronegativity., We measure electronegativity on several scales. The most commonly used scale was designed by Linus Pauling. According to this scale, fluorine is the most electronegative element with a value of 4.0 and Cesium is the least electronegative element with a value of 0.7., The numerical value of electronegativity of an atom depends on its ionisation potential and electron affinity values., Factors Affecting Electronegativity, 1. Size of an Atom:, A greater atomic size will result in less value of electronegativity, this happens because electrons being far away from the nucleus will experience a lesser force of attraction., 2. Nuclear Charge:, A greater value of nuclear charge will result in a greater value of electronegativity. This happens because an increase in nuclear charge causes electron attraction with greater force., 3. Effect of Substituent:, The electronegativity of an atom depends upon the nature of the substituent attached to that atom. For example, the carbon atom in CF3I acquires a greater positive charge than CH3I. Therefore, C-atom in CF3I is more electronegative than in CH3I. The difference in electronegativity of an atom caused by substituents results in different chemical behaviour of that atom., Periodic Trends in the Electronegativity of Elements, IN PERIOD, As we move across a period from left to right, Nuclear charge increases, Atomic size decreases, Therefore the value of electronegativity increases across a period., For example, the electronegativity trend across period 3 in the periodic table is depicted below., IN GROUP, On moving from top to bottom in a group:-, Nuclear charge increases with increase in atomic number., Atomic size increases due to increase in the value of Principle Quantum number (shell)., The effect of increase in atomic size is much more than the effect of increase in nuc lear charge., Hence, the value of electronegativity decreases as we move down the group. For example, in the halogen group as we move down the group from fluorine to astatine the electronegativity value decreases and it is shown in the diagram below., It is a general observation that, Metals show a lower value of electronegativity as compared to the non-metals., Metals are electropositive and non-metals are electronegative in nature., The elements in period two differ in properties from their respective group elements due to the small size and higher value of electronegativity., The elements in the second period show resemblance to the elements of the next group in period three. This happens due to a small difference in their electronegativity. This leads to the formation of a diagonal relationship., Most and Least Electronegative Elements, Fluorine is the most electronegative element on the periodic table. Its electronegativity value is 3.98., Cesium is the least electronegative element. Its electronegativity value is 0.79., Electro positivity is the exact opposite of electronegativity; therefore, we can say that Cesium is the most electropositive element., Impact of Electronegativity on Covalent Bonding, All covalent bonds between dissimilar species have some ionic character. Similarly, all ionic bonds have some covalent character as well. The ionic character of the covalent bond is determined by the difference in electronegativity. If the difference in electronegativity is greater than 1.7, the character of the bond will be ionic. If the difference in electronegativity is between 0.4 and 1.7, the character of the bond is polar covalent., Explain the Pauling scale for the determination of electronegativity., Pauling scale is based on an empirical relation between the energy of a bond and the electronegativity of bonded atoms., Consider a bond A−B between two dissimilar atoms A and B of a molecule AB., Let the bond energies of A−A, B−B and A−B bonds be represented as EA−A​, EB−B​ and EA−B​ respectively., It may be seen that the bond dissociation energy of A−B is almost higher than the geometric mean of the bond dissociation energies of A−A and B−B bonds i.e.,, EA−B​ > (EA−A​×EB−B)​​1/2, Their difference (ΔE) is related to the difference in the electronegativity of A and B according to the following equation:, ΔE = EA−B​ − (EA−A​×EB−B)​​1/2 = (XA​−XB​)2, or XA​−XB ​= 0.208 (ΔE)1/2, Here XA​and XB​are the electronegativity of A and B respectively., The factor 0.208 arises from the conversion of Kcals to electron volt., Considering arbitrarily the electronegativity of hydrogen to be 2.1, Pauling calculated electronegativity’s of other elements with the help of this equation.