Page 1 :
CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES OF ELEMENTS, By , ESHWARA K, M.Sc. , B.Ed., KSET, Lecturer in Chemistry
Page 2 :
Why do we need classification.? , Dobereniner’s triads , Newlands law of octave , Lother Meyer volume curve , Mendeleev’s periodic table , Modern periodic table , IUPAC nomenculature for elements Z >100, , , , 8. Periodic Trends in Physical Properties, Shielding effect & Effective Nuclear Charge, Atomic Radius, Ionic Radius, Ionization Enthalpy, Electron Gain Enthalpy, Electronegativity, 9. Periodic Trends in Chemical Properties, Periodicity of Valence or Oxidation States, Anomalous Properties of Second Period Elements, Chemical Reactivity, SESSION OBJECTIVES
Page 3 :
Lavoisier (1789) classified elements into metals, non-metals, gases and earths., 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.
Page 4 :
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., Model of triads, 2, Li,Na,K, 7, 23, 39, 7 + 39, 2, = 23
Page 5 :
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 (1866)]
Page 6 :
NEWLAND’S LAW OF OCTAVES [JOHN NEWLAND (1863)], Newland was first to publish the list of elements in increasing order of atomic masses.
Page 8 :
DMITRI MENDELEEV, 1834 - 1907, In 1869 he published a table of the elements organized by increasing atomic mass.
Page 9 :
The physical and chemical properties of elements are periodic function of their atomic masses., MENDELEEV’S , PERIODIC LAW
Page 10 :
MENDELEEV’S PERIODIC TABLE
Page 11 :
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., Only 63 elements were known., MENDELEEV’S PERIODIC TABLE
Page 12 :
MERITS OF MENDELEEV’S PERIODIC TABLE
Page 13 :
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.
Page 14 :
DEFECTS OF MENDELEEV’S PERIODIC TABLE, Anomalous pairs. (Ar and K, Co and Ni, Te and I), Position of isotopes e.g. 1H1, 1H2, 1H3
Page 15 :
Chemically dissimilar elements are grouped together. (Cu-IA and Na-IB), Chemically similar elements are placed in different groups. [Cu (I) and Hg (II)]., DEFECTS OF MENDELEEV’S PERIODIC TABLE
Page 16 :
Mendeleev’s periodic table was published in 1905 when no one had an idea of the structure of an atom., DO YOU KNOW?, 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
Page 17 :
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., MODERN PERIODIC LAW AND THE MODERN PERIODIC TABLE
Page 19 :
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.
Page 20 :
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.
Page 21 :
MODERN PERIODIC TABLE
Page 25 :
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.
Page 26 :
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., Periods, Horizontal column, Total 7 numbered from 1 to 7., Elements in a group have similar but not identical electronic configuration and properties, Contains 2,8,8,18,18,32 and 28 elements respectively.
Page 27 :
ELECTRONIC CONFIGURATIONS AND TYPES OF 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-,p-,d-,f- Block Elements
Page 28 :
Electronic configuration:, Groups:, All are metal, low ionisation energy and low, melting and boiling points, electropositive , elements., compounds are mostly ionic & colourless., IA (alkali metals )and , IIA(alkaline earth metals, ns1 or ns2, In these elements last electron enters the s-orbital, s-Block Elements
Page 29 :
Electronic configuration:, Groups:, Non-metals, electronegative., Form covalent compounds., ns2,np1 -6, III A to VII A and zero group (group 13-18)., In these elements last electron enters the p-orbital, p-Block Elements
Page 30 :
Electronic configuration:, Groups:, Variable valency high melting and boiling point., Coloured compounds and catalytic property., (n-1)d1-10 ns1or2, I B to VII B and VIII groups (Gr- 3-12)., In these elements last electron enters the d-orbital,, Also known as transition metals., d-Block Elements
Page 31 :
Electronic configuration:, Have high melting and boiling point., (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), In these elements last electron enters the f-orbital,, Also known as Inner-Transition Elements, f-Block Elements
Page 32 :
Representative elements, Transition elements, s and p block 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., FEATURES OF LONG FORM OF PERIODIC TABLE
Page 33 :
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., FEATURES OF LONG FORM OF PERIODIC TABLE
Page 34 :
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., MERITS OF LONG FORM OF PERIODIC TABLE
Page 35 :
DEFECTS OF LONG FORM OF PERIODIC TABLE, 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., It is unable to include lanthanides and , actinides in its main body.
Page 38 :
Periodic Properties, Periodic Trends in Physical Properties, Shielding effect & Effective Nuclear Charge, Atomic Radius, Ionic Radius, Ionization Enthalpy, Electron Gain Enthalpy, Electronegativity
Page 39 :
Periodic Trends in Chemical Properties, Periodicity of Valence or Oxidation States, Anomalous Properties of Second Period Elements, Chemical Reactivity, Periodic Properties
Page 40 :
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.
Page 41 :
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
Page 42 :
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
Page 43 :
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
Page 44 :
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
Page 45 :
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., Shielding effect & Effective Nuclear Charge
Page 46 :
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.
Page 49 :
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*
Page 50 :
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
Page 51 :
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
Page 53 :
Anionic Radii
Page 54 :
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------🡪
Page 55 :
Atomic and Ionic Radii
Page 57 :
Atomic Radius
Page 58 :
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.
Page 59 :
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
Page 61 :
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., Ionisation Energy (IE) or Ionisation Enthalpy (ΔiH )
Page 62 :
(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 )
Page 63 :
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 (ΔiH )
Page 64 :
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
Page 66 :
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 )
Page 67 :
(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., Periodic Trend of Ionisation Enthalpy (ΔiH )
Page 68 :
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)
Page 69 :
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)
Page 70 :
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)
Page 71 :
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., Factors Affecting E G E (ΔegH)
Page 72 :
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)
Page 73 :
First Electron Affinities
Page 74 :
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)
Page 75 :
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
Page 76 :
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
Page 77 :
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
Page 79 :
Periodicity of Valence or Oxidation States, Anomalous Properties of Second Period Elements, Chemical Reactivity, Periodic Trends in Chemical Properties
Page 80 :
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
Page 81 :
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, Periodicity of Valence or Oxidation States
