CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES of elements

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

Session Objectives Why do we need classification.?
8. Periodic Trends in Physical Properties Shielding effect & Effective Nuclear Charge Dobereniner’s triads Atomic Radius Newlands law of octave Ionic Radius Ionization Enthalpy Lother Meyer volume curve Electron Gain Enthalpy Electronegativity Mendeleev’s periodic table 9. Periodic Trends in Chemical Properties Modern periodic table Periodicity of Valence or Oxidation States IUPAC nomenculature for elements Z >100 Anomalous Properties of Second Period Elements Chemical Reactivity

During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table. Lavoisier (1789) classified elements into metals, non-metals, gases and earths.

Dobereiner’s triads [ John Dobereiner (1817)]
In 1829, he classified some elements into groups of three, which he called triads. The elements in a triad had similar chemical properties and orderly physical properties. S.No Triad Atomic masses of elements of triad Arithmetic mean of atomic masses of first and third element 1 Cl,Br,I 35.5, 80, 127 2 = 81.25 2 Li,Na,K 7, 23, 39 7 + 39 = 23 2 3 Ca,Sr,Ba 40,87.5,137 40+137 2 = 88.5 Model of triads

Newland’s law of octaves [John Newland (1866)]
In 1866, he suggested that elements be arranged in “octaves” because he noticed (after arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th element.

Newland’s law of octaves [John Newland (1863)]
Element Atomic mass I II III IV V VI VII Li Be B C N O F Na Mg Al Si p S Cl K Ca Newland was first to publish the list of elements in increasing order of atomic masses.

Lother-Meyer’s atomic volume curve

Dmitri Mendeleev In 1869 he published a table of the elements organized by increasing atomic mass.

Mendeleev’s periodic law
The physical and chemical properties of elements are periodic function of their atomic masses.

Mendeleev’s periodic table

Mendeleev’s periodic table
Only 63 elements were known. Groups 8 vertical rows. 7 groups were subdivided in A and B. 8th group has 9 elements in the group of 3 each. Periods 7 horizontal rows.

Merits of Mendeleev’s periodic table
Prediction of new elements (Ge, Ga, Sc) 1 Systematic study of elements 2 Correction of atomic mass (Be, Au, Pt) 3

Mendeleev stated that if the atomic weight of an element caused it to be placed in the wrong group, then the weight must be wrong. (He corrected the atomic masses of Be, In, and U) was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown. After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions for Sc, Ga, and Ge were amazingly close to the actual values, his table was generally accepted.

Defects of Mendeleev’s periodic table
Position of hydrogen. Position of isotopes e.g. 1H1, 1H2, 1H3 Anomalous pairs. (Ar and K, Co and Ni, Te and I)

Defects of Mendeleev’s periodic table
Chemically dissimilar elements are grouped together. (Cu-IA and Na-IB) Chemically similar elements are placed in different groups. [Cu (I) and Hg (II)].

Do you know? Mendeleev’s periodic table was published in 1905 when no one had an idea of the structure of an atom. Mendeleev’s name has been immortalized by naming the element with atomic number 101, as Mendelevium. This name was proposed by American scientist Glenn T. Seaborg, the discoverer of this element, “in recognition of the pioneering role of the great Russian Chemist who was the first to use the periodic system of elements to predict the chemical properties of undiscovered elements, a principle which has been the key to the discovery of nearly all the transuranium elements

MODERN PERIODIC LAW AND THE MODERN PERIODIC TABLE
English physicist, Henry Moseley observed regularities in the characteristic X-ray spectra. A plot of f against atomic number (Z ) of the elements gave a straight line and not the plot of f vs atomic mass He thereby showed that the atomic number is a more fundamental property of an element than its atomic mass.

Mendeleev’s Periodic Law was, therefore, accordingly modified.
This is known as the Modern Periodic Law and can be stated as : The physical and chemical properties of the elements are periodic functions of their atomic numbers.

Henry Moseley In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number. *“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.” His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.

Modern periodic table

Contains elements arranged in increasing order of atomic numbers.
Features of long form of periodic table Contains elements arranged in increasing order of atomic numbers. Explains the position of an element in relation to other elements. Consists of groups and periods.

Features of long form of periodic table
Groups Vertical column Total 18. Numbered 1-18 or IA to VII A, IB to VII B, VIII and zero. Elements in a group have similar but not identical electronic configuration and properties Periods Horizontal column Total 7 numbered from 1 to 7. Contains 2,8,8,18,18,32 and 28 elements respectively.

ELECTRONIC CONFIGURATIONS AND TYPES OF ELEMENTS:
s-,p-,d-,f- Block Elements On the basis of the nature of sub-shell in which last electron of atom enters, elements are divided into 4 blocks s-Block Element p-Block Element d-Block Element f- Block Element

s-Block Elements ns1 or ns2
In these elements last electron enters the s-orbital Electronic configuration: ns1 or ns2 IA (alkali metals )and IIA(alkaline earth metals Groups: All are metal, low ionisation energy and low melting and boiling points, electropositive elements. compounds are mostly ionic & colourless.

Electronic configuration:
p-Block Elements In these elements last electron enters the p-orbital Electronic configuration: ns2,np1 -6 Groups: III A to VII A and zero group (group 13-18). Non-metals, electronegative. Form covalent compounds.

d-Block Elements In these elements last electron enters the d-orbital,
Also known as transition metals. Electronic configuration: (n-1)d1-10 ns1or2 Groups: I B to VII B and VIII groups (Gr- 3-12). Variable valency high melting and boiling point. Coloured compounds and catalytic property.

f-Block Elements In these elements last electron enters the f-orbital,
Also known as Inner-Transition Elements Electronic configuration: (n-2)f1-14(n-1)d0-1ns2 Present below the periodic table in two rows Lanthanides-elements after lanthanum(Gr.-3, Pd.-6) Actinides-elements after actinium. (Gr.-3, Pd.-7) Have high melting and boiling point.

Representative elements s and p block elements . Transition elements d-block elements. Valence shell and penultimate Shell both are incomplete. Inner Transition elements f-block elements. Valence shell, penultimate shell antipenultimate shell are incomplete.

Metals Present on left hand side of periodic table. Solid,malleable,ductile and conductors . Non-metals Present on right hand side of periodic table. Solid or liquid or gas. Metalloids Present on zig-zag between metals and non-metals. e.g. B,Si,Ge,As,Sb and Te.

Merits of long form of periodic table
Based on a more fundamental basis - the atomic number Position of an element is related to the electronic configuration of its atom. Due to separation of elements into groups, dissimilar elements (e.g. alkali metals I A and coinage metals I B) do not fall together.

Defects of long form of periodic table
It is unable to include lanthanides and actinides in its main body. The problem of the position of hydrogen in the table has not been solved completely Configuration of Helium(1s2 ) is different from inert gases (ns2,np6) but are placed in the same group.

Nomenclature of the elements with atomic number >100
Digit Name Abbreviation nil n 1 un u 2 bi b 3 tri t 4 quad q 5 pent p 6 hex h 7 sept s 8 oct o 9 enn e Name =digits name + ium e.g. atomic number 115 Will be named as un+un+pent+ium =ununpentium and symbol is Uup

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.

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 41

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 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 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

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 = = 3.80 Z* = = 5.20

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

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

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

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.

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.

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.

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

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. Each group has a common (+)ve or (-)ve oxidation state And it show gradual change in oxidation state in a period

CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES of elements

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