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Periodic Properties

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Periodic Properties

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  1. Periodic Properties By- A.P.S. BHADOURIYA M.Sc. (L.U.), B.Ed. NET (CSIR-UGC) Qualified

  2. Periodic Properties • Periodic Trends in Physical Properties • Shielding effect & Effective Nuclear Charge • Atomic Radius • Ionic Radius • Ionization Enthalpy • Electron Gain Enthalpy • Electronegativity

  3. Periodic Properties • Periodic Trends in Chemical Properties • Periodicity of Valence or Oxidation States • Anomalous Properties of Second Period Elements • Chemical Reactivity

  4. Shielding effect & Effective Nuclear Charge The decrease in nuclear charge ( nuclear force of attraction) on outermost shell electrons due to repulsion caused by inner shell electron is known as shielding effectof inner shell or intervening electrons on outer shell electron.

  5. Shielding effect & Effective Nuclear Charge Due to shielding effect the nuclear charge is lowered on outermost shell electrons, the net nuclear charge acting on outermost shell electrons is known as Effective Nuclear Charge. It is denoted by Z* or Zeff. • Z* or Zeff. = Z - σ • where Z = nuclear charge( = atomic No.) & σ = shielding constant or screening constant , it is a measure of shielding effect

  6. Shielding effect & Effective Nuclear Charge Determination of ENC (Z*) If the electron resides in s or p orbital 1. Electrons in principal shell higher than the e- in question contribute 0 to σ . 2. Each electron in the same principal shell contribute 0.35 to σ (0.30 if it is 1S shell). 3. Electrons in (n-1) shell each contribute 0.85 to σ . 4. Eelectrons in deeper shell each contribute 1.00 to σ

  7. Shielding effect & Effective Nuclear Charge Determination of ENC (Z*) If the electron resides in d or f orbital 1. All e-s in higher principal shell contribute 0 to σ 2. Each e- in same shell contribute 0.35 to σ 3. All inner shells in (n-1) and lower contribute 1.00 to σ

  8. Shielding effect & Effective Nuclear Charge Determination of ENC (Z*) e.g.Calculate the Z* for the 2p electronFluorine (Z = 9) 1s2, 2s 2p5. Soln.Screening constant for one of the outer electron • 6 (six) (two 2s e- and four 2p e-) = 6 X 0.35 = 2.10 • 2 (two)1s e- = 2 X 0.85 = 1.70 • σ = 1.70+2.10 = 3.80 • Z* = 9 - 3.80 = 5.20

  9. Shielding effect & Effective Nuclear Charge Trend of ENC in Periodic Table • In a Period - Effective nuclear charge Z* increases increases rapidly along a period(0.65 per next group) e.g.

  10. Shielding effect & Effective Nuclear Charge Trend of ENC in Periodic Table • In a Group - Effective nuclear charge Z* increases slowly along a group. e.g.

  11. PERIODIC TREND OF ATOMIC RADIUS In A Period- • atomic radius decreases with increasein atomic number (in a period left to right) BECAUSE in a period left to right- • 1. n (number of shells) remain constant. • 2. Z increases (by one unit) • 3. Z* increases (by 0.65 unit) • 4. Electrons are pulled close to the nucleus by the increased Z*

  12. In a group- Atomic radius increases moving down the group • Because, along a group top to bottom 1. n increases 2. Z increases 3. No dramatic increase in Z* - almost remains constant

  13. IONIC RADII • All anions arelarger than theirparent atoms. because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in ENC. • The cations aresmaller than theirparent atoms because it has fewer electrons while its nuclear charge remains the same & hence ENC is greater in cation than its parent atom

  14. Cationic Radii

  15. Anionic Radii

  16. ISOELECTRONIC SPECIES • Atoms and ions which contain the same number of electrons, are called as isoelectronic species. For example, F–, Na+ and Mg2+ have the same number of electrons(=10). • The size ofisoelectronic species decreases with increase in nuclear charge. e.g.- o2->F- >Ne>Na+>Mg2+>Al3+ ---------SIZE DECREASING------

  17. Atomic and Ionic Radii

  18. Atomic Radius

  19. NOTE: Metallic radii in the third row d-block are similar tothe second row d-block, but not larger as one wouldexpect given their larger number of electrons. This is due toLanthanide Contractionasf-orbitals have poor shielding properties.

  20. Ionisation Energy (IE) orIonisation Enthalpy (ΔiH ) • Ionization: removing an electron from an atom or ion • Ionization energy: energy required to remove an electron from an isolated, gaseous atom or ion is called as Ionization energy or ionisation enthalpy. • If the atom is neutral the above defined ionisation energy is called as first ionisation enthalpy. • Energy required to remove an electron from an isolated, monovalent cation is called as second Ionization energy. • The ionization enthalpy is expressed in units of kJ /mol

  21. Ionisation Energy (IE) orIonisation Enthalpy (ΔiH ) X(g) + energy → X+(g) + e–. 1st ionisation enthalpy X+(g) + energy → X++(g) + e–. 2nd ionisation enthalpy

  22. Ionisation Energy (IE) orIonisation Enthalpy (ΔiH ) The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom because a cation has greater ENC than a neutral atom. In the same way the third ionization enthalpy will be higher than the second and so on.

  23. Factors affecting Ionisation Enthalpy (ΔiH ) (a) Size of the atom - IE decreases as the size of the atom increases (b) Nuclear Charge - IE increases with increase in nuclear charge (c) The type of electron - Shielding effect, Penetration effect (e)Electronic configuration: e.g. noble gases passes very high value of IE due to stable octet configuration

  24. Periodic Trend of Ionisation Enthalpy (ΔiH ) On moving down a group 1. nuclear charge increases 2. Z* due to screening is almost constant 3. number of shells increases, hence atomic size increases. 4. there is a increase in the number of inner electrons whichshield the valence electrons from the nucleus Thus IE decreases down the group

  25. Periodic Trend of Ionisation Enthalpy On moving across a period(L--->R) 1. the atomic size decreases 2. Effective nuclear charge increases Thus IE increases along a period However there are some exceptions also e.g. • IE of Be is higher than that of B. • IE of N is higher than that of O.

  26. Periodic Trend of Ionisation Enthalpy (ΔiH ) Explain why- (a). IE of Be is higher than that of B. Ans. - In beryllium(1s2,2s2 ), the electron removed during the ionization is an s-electron whereas the electron removed during ionization of boron(1s2,2s2,2p1) is a p-electron. The penetration of a 2s-electron to the nucleus is more than that of a 2p-electron; hence the 2p electron of boron is more shielded from the nucleus by the inner core of electrons than the 2s electrons of beryllium. Therefore, it is easier to remove the 2p-electron from boron compared to the removal of a 2s- electron from beryllium. Thus, boron has a smaller first ionization.

  27. Periodic Trend of Ionisation Enthalpy (ΔiH ) (b) Why IE of N is higher than that of O. Ans. The first ionization enthalpy of oxygen compared to nitrogen is smaller. This arises because in the nitrogen atom(1s2,2s2,2p3) three 2p-electrons reside in different atomic orbitals (Hund’s rule) whereas in the oxygen atom (1s2,2s2,2p4), two of the four 2p-electrons must occupy the same 2p-orbital resulting in an increased electron-electron repulsion. Consequently, it is easier to remove the fourth 2p-electron from oxygen than it is, to remove one of the three 2p-electrons from nitrogen.

  28. Electron Gain Enthalpy (ΔegH) • When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change accompanying the process is defined as the Electron GainEnthalpy (ΔegH) or Electron Affinity. • Electron gain enthalpy provides a measure of the ease with which an atom adds an electron to form anion as represented by equation – • X(g) + e --- X- (g)+energy (electrongainenthalpy)

  29. Electron Gain Enthalpy (ΔegH) • Depending on the element, the process of adding an electron to the atom can be either endothermic or exothermic. • For many elements energy is released when an electron is added to the atom and the electron gain enthalpy is negative.

  30. Factors Affecting E G E (ΔegH) • ENC- With increase in ENC, the force of attraction exerted by thenucleus on the electrons increases. Consequently, the atom has a greatertendency to attract additional electron i.e., its EGE increases i.e. become more negative. • ATOMIC SIZE- With decrease in size ENC increases & hence EGE increases. • ELECTRONIC CONFIGURATION- The value of EGE depends effectively upon electronic configuration of elements, elements with stable electronic configuration posses lower (less -ve) value of EGE, e.g.-

  31. Factors Affecting E G E (ΔegH) • Noble gases have practically zero or +ve EGEs. This is because they have no tendencyto gain an additional electron as they already have the stable ns2np6 configuration • Halogens have high electron affinities. This is due to their strong tendencyto gain an additional electron to change into the stable ns2np6 configuration.

  32. PERIODIC TREND OF EGE (ΔegH) IN A PERIOD- The EGE increases i.e. become more negative as we move across a period because the atomic size decreases and hence the force of attraction exerted by thenucleus on the electrons increases. Consequently, the atom has a greatertendency to attract additional electron i.e., its electron affinity increases IN A GROUP- The EGE decreases (-)vely because the atomic size increases and therefore, the effective nuclear attractiondecreases and thus electron affinity decreases

  33. First Electron Affinities

  34. Electron Gain Enthalpy (ΔegH) • Explain why – (a). electron gain enthalpy of O is less than that of the S. (b). electron gain enthalpy of F is less than that of the Cl. • Ans:- The electron gain enthalpy of O or F is less than that of the succeeding element. This is because when an electron is added to O or F, the added electron goes to the smaller n = 2 quantum level and suffers significant repulsion from the other electrons present in this level. For the n = 3 quantum level (S or Cl ), the added electron occupies a larger region of space and the electron-electron repulsion is much less.

  35. Electronegativity • The tendency of an element in a molecule to attract the shared pair of electrons towards itself is known as electronegativity. • It is measured on Pauling scale in which F (most EN element)is attributed to a value of 4 .

  36. Periodic trend of EN In a Group- on moving down the group, • Z increases but Z* almost remains constant • number of shells (n) increases • atomic radius increases • force of attraction between added electron and nucleus decreases • Therefore EN decreases moving down the group

  37. Periodic trend of EN In a Period- On moving across a period left to right • Z and Z* increases • number of shells remains constant • atomic radius decreases • force of attraction between shared electron and nucleus increases Hence EN increases along a period

  38. Periodic Trends in Chemical Properties • Periodicity of Valence or Oxidation States • Anomalous Properties of Second Period Elements • Chemical Reactivity

  39. Periodicity of Valence or Oxidation States The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outer most orbitals and / or equal to eight minus the number of outermost Electrons(w.r.t. H) Some periodic trends observed in the valence of elements (hydrides and oxides) are shown in Table

  40. Periodicity of Valence or Oxidation States The oxidation state of an element in a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule.