Periodic Properties

By-

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

Periodic Properties

Periodic PropertiesPeriodic Trends in Physical PropertiesShielding effect & Effective Nuclear ChargeAtomic RadiusIonic RadiusIonization EnthalpyElectron Gain EnthalpyElectronegativity

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

Elements Chemical Reactivity

Periodic Properties

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

Determination of ENC (Z*) If the electron resides in s or p orbital1. 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

Determination of ENC (Z*) If the electron resides in s or p orbital1. 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 ChargeShielding effect & Effective Nuclear Charge

Determination of ENC (Z*) If the electron resides in d or f orbital1. All e-s in higher principal shell contribute 0 to σ 2. Each e- in same shell contribute 0.35 to σ3. All inner shells in (n-1) and lower contribute 1.00 to σ

Determination of ENC (Z*) If the electron resides in d or f orbital1. All e-s in higher principal shell contribute 0 to σ 2. Each e- in same shell contribute 0.35 to σ3. All inner shells in (n-1) and lower contribute 1.00 to σ


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


Trend of ENC in Periodic TableIn 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


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 2.2 2.2 2.2

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

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

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 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

X(g) + energy → X+(g) + e–. 1st ionisation enthalpy X+(g) + energy → X++(g) + e–.

2nd ionisation enthalpy


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.


(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

Factors affecting Ionisation Enthalpy (ΔiH )

On moving down a group1. nuclear charge increases2. Z* due to screening is almost constant3. 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 (ΔiH )

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

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.


(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)

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.

Electron Gain Enthalpy (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)

A. 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

B. Halogens have high electron affinities. This is due to their strong tendency to gain an additional electron to change into the stable ns2np6 configuration.

Factors Affecting E G E (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

PERIODIC TREND OF EGE (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.

Electron Gain Enthalpy (ΔegH)

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 .

Electronegativity

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

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 trend of EN

Periodicity of Valence or Oxidation States Anomalous Properties of Second Period

Elements Chemical Reactivity

Periodic Trends in Chemical Properties

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

Periodicity of Valence or Oxidation States

Group 1 2 13 14 15 16 17

Number of valence electron

1 2 3 3 5 6 7

Valence 1 2 3 4 3,5 2,6 1,7

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.

Periodicity of Valence or Oxidation States

Periodic Properties

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