CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES OF ELEMENTS-2 BY- A.P.S. BHADOURIYA M.Sc. , B.Ed., NET PGT-CHEMISTRY K.V. BARABANKI 1

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

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

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 effect of inner shell or intervening electrons on outer shell electron.

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 бσ 5

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

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

Shielding effect & Effective Nuclear Charge Determination of ENC (Z*) e.g. Calculate the Z* for the 2p electron Fluorine (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

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. Li Be B C N O F Ne 1.3 1.95 2.6 3.3 3.9 4.6 5.2 5.9

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. Gr-1 H Li Na K Rb Cs Z* 1.0 1.3 2.2

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PERIODIC TREND OF ATOMIC RADIUS In A Period- atomic radius decreases with increase in 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*

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

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

Cationic Radii

Anionic Radii

o2->F- >Ne>Na+>Mg2+>Al3+ 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 of isoelectronic species decreases with increase in nuclear charge. e.g.- o2->F- >Ne>Na+>Mg2+>Al3+ ---------SIZE DECREASING------

Atomic and Ionic Radii

Atomic Radius

NOTE: Metallic radii in the third row d-block are similar to the second row d-block, but not larger as one would expect given their larger number of electrons. This is due to Lanthanide Contraction as f-orbitals have poor shielding properties.

Ionisation Energy (IE) or Ionisation 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 Δi H

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

Ionisation Energy (IE) or Ionisation 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.

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

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 which shield the valence electrons from the nucleus Thus IE decreases down the group

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.

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.

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.

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)

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.

Factors Affecting E G E (ΔegH) ENC- With increase in ENC, the force of attraction exerted by the nucleus on the electrons increases. Consequently, the atom has a greater tendency 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.-

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

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 the nucleus on the electrons increases. Consequently, the atom has a greater tendency 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 attraction decreases and thus electron affinity decreases

First Electron Affinities

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.

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 .

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

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

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

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 Group 1 2 13 14 15 16 17 Number of valence electron 3 5 6 7 Valence 4 3,5 2,6 1,7

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. In a group – it remains same (with few exceptions) In a period- the normal oxidation state first increases then decrease

Anomalous Properties of Second Period Elements The elements of 2nd period show some differences from rest members of their group members in some propertieses like valency, reacivity etc. Their anomalous behaviour is attributed to their small size, large charge/radius ratio, high electro negativity, non- availability of d- orbitals in their valence shell. the first member of each group of p-Block elements displays greater ability to form pp-pp multiple bonds to itself (e.g. C=C, C C O=O, N N) and to other second period elements (e.g. C=O, C N, N=O) compared to subsequent member of the group.

Trend in some chemical properties SOLUBILITY OF ALKALI METALS CARBONATES AND BICARBONATES: PERIODICITY IN GROUP: The solubility of alkali metal carbonates and bicarbonates in water increases down the group (From Lithium to Caesium). SOLUBILITY OF ALKALINE EARTH METAL HYDROXIDES AND SULPHATES: PERIODICITY IN GROUP: The solubility of alkaline earth metal hydroxide and sulphates in water increases down the group (From Beryllium to Barium).

BASIC STRENGTH OF ALKALINE EARTH METAL HYDROXIDES: PERIODICITY IN GROUP: The basic strength of alkaline earth metal hydroxide in water increases down the group (From Beryllium to Barium), i.e., Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2 -----------Basic strength increases

THERMAL STABILITY OF CARBONATES OF ALKALI AND ALKALINE EARTH METALS: Except lithium carbonate, (Li2CO3), the carbonates of all other alkali metals are stable towards heat, i.e., carbonates of alkali metals (except Li2CO3) do not decompose on heating. Li2CO3 decomposes on heating to give lithium oxide (Li2CO3). The carbonates of alkaline earth metals are relatively less stable. On heating, they decompose to give corresponding oxide and CO2 gas. The decomposition temperature for alkaline earth metal carbonates increases as we go down the group.

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