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Classification of Elements and Periodicity in Properties
A large number of elements and compounds are known today. Therefore a systematic classification of these elements has made their study possible and easy. As we know it today Periodic Table is a well organized and tabulated classification of elements. It helps us to locate, identify and characterize the element and its properties and also points out the directions in which new investigations are made.
GENESIS OF PERIODIC CLASSIFICATION
The number of elements was limited in the 18th century. In the 19th century, Scientists began to find out ways of classifying the elements because of their rapidly increasing number.
It was the first attempt towards the classification of elements. He arranged similar elements in groups of three elements called triad and the atomic mass of the middle elements of the triad is approximately the arithmetic mean of the other two.
Merits: After Dobereiner, Chemists focused on chemicals in groups having similar physical and chemical properties.
Demerits: All the known elements did not follow this rule. Law of triads was rejected as some triads nearly had the same atomic masses, e.g., (Fe, Co, Ni), (Ru, Rh, Pd), (Os, Ir, Pt)
Newland’s Rule of Octaves
When the lighter elements are arranged in order of their increasing atomic weights, then every eighth element is similar in its properties to the first element, just like the eighth note of a musical scale is similar to the first one. e.g. 8th element (Na) resembles in their properties with Li. Similarly, the 8th element (K) with Na. This type of classification was limited up to only 20th elements (Ca).
(i) Law of octave worked quite well for lighter elements but failed with heavier elements.
(ii) Properties of elements were not taken into account and the elements were arranged in the order of their increasing atomic masses.
(iii) No places were left for unknown elements and so, many elements occupied wrong positions. Thus, resulted in the rejection of the attempt.
Lothar Meyer’s Volume Curve
The graphs of atomic volumes against weights are known as Lothar Meyer’s volume curves.
Demerits: It lacked practical utility as it is not easy to remember the position of diﬀerent elements on the curve.
Mendeleev’s Periodic Law
The physical and chemical properties of elements are a periodic function of their atomic weights.
Merits of Mendeleev’s periodic table
(a) Study of elements and their compounds becomes easy and systematic, as by knowing the property of one element in a group, then the properties of the other elements present in the same group can easily be predicted.
(b) It helps in the discovery of new elements. As Mendeleev left some blank spaces for some unknown elements and further, predicted the properties of these elements e.g. Eka-Aluminium, Eka-Silicon.
(c) Correction of doubtful atomic mass.
(d) Correction in the valency of some elements.
(e) Correction in the position of some elements.
(f) Classification of elements then known, was done for the first time and the elements having similar properties were kept in the same group.
(g) Mendeleev had even predicted the properties of many elements not discovered at that time. This helped in the discovery of these elements. For example, Mendeleev predicted the properties of the following elements.
(i) Eka-boron later called scandium (Sc)
(ii) Eka- aluminium later called gallium (Ga)
(iii) Eka-silicon later called germanium (Ge)
Limitations of Mendeleev’s periodic table
(a) The position of hydrogen was found to be anomalous due to its resemblance with the 1st group alkali metals and also with the 7th group halogens in their properties.
(b) The position of isotopes: Isotopes must have diﬀerent positions but they were placed in the same group.
(c) The position of isobars: They were placed in diﬀerent groups.
(d) Dissimilar elements were placed together in the same group as K and Cu in the 1st group.
(e) Similar elements were placed in diﬀerent groups.
(f) Some higher atomic weight elements were placed before the lower atomic weight elements.
(g) e.g. Ar40 precedes K39, Co58.9 precedes Ni58.7, Te127.6 precedes I127.
(h) The position of metals and non-metals: Both were placed together in the same group.
(i) The diagonal relationship could not be explained.
(j) The position of lanthanides and actinides was not properly specified.
(k) No proper position to VIII group elements.
(l) There was no indication whether lanthanides and actinides were associated with group IIIA or group IIIB.
(m) The position of Isobars- These elements had diﬀerent groups when mass remained the same.
(n) A lot of stress was given to the valence of elements.
Modern Periodic Law and Modern Periodic Table
Mosley: He proved that the square root of frequency (f) of the rays, which are obtained from a metal on showering high-velocity electrons is proportional to the nuclear charge of the atom.
This can be represented by the following expression.
f = a(Z-b)
where Z is the nuclear charge on the atom and a and b are constants. The nuclear charge on an atom is equal to the atomic number.
Modern Periodic Law
“The properties of elements are the periodic function of their atomic numbers”.
Period- The modern periodic table has 7 horizontal rows called periods.
Group: The modern periodic table has 18 vertical columns called groups and according to CAS system there are 16 groups having the following number of elements.
Advantages of the Long Form of the Periodic Table
The table is based on a more fundamental property i.e. atomic number. It correlates the position of elements with their electronic configuration more clearly. The completion of each period is more logical. The position of VIII group is also justified in this table. All the transition elements have been brought to the middle as the properties of transition elements are intermediate between s- and p-block elements. The table completely separates metals and non-metals. Non-metals are present in the upper right corner of the periodic table. The advantage of this periodic table is that it is divided into four blocks namely s-, p-, d- and f-block elements. This arrangement of elements is easy to remember and reproduce.
Defects of the Long Form of the Periodic Table
(a) The position of hydrogen is still disputable as it was there in MENDELEEV periodic table in group IA as well as IVA & VIIA.
(b) Helium is an inert gas but its configuration is diﬀerent from that of the other inert gas elements.
(c) Lanthanide and actinide series could not be adjusted in the main periodic table and therefore they had to be provided with a place separately below the table.
NAMING OF ELEMENTS HAVING ATOMIC NUMBER GREATER THAN 100
The names are derived directly from the atomic number of the elements using numerical roots for 0 and numbers from 1-9 and adding the suffix ‘ium’ to spell out the name. The names are shortened in certain cases. For example, bi ium and tri ium are shortened to bium and trium and enn nil is shortened to ennil. The symbol of the element is then obtained from the first letters of the roots of numbers which make up the atomic number and finally the name.
CLASSIFICATION OF PERIODIC TABLE BASED ON BLOCKS
Elements of groups 1 and 2 including He in which the last electron enters the s-orbitals of the valence shell are called s-block elements. There are only 14 s-block elements in the periodic table.
The electronic configuration of the outermost shell of s-block elements is ns1 (alkali metals; group1) or ns2 (alkaline earth metals; group 2). The valency of group I elements is +1 and those of group II elements is +2. These are soft metals having low melting points and boiling points. Most of these form ionic compounds on account of their lower ionization energy. Most of these metals (except Be & Mg) and their salts imparts characteristic colour to the ﬂame e.g., sodium imparts a golden yellow colour; potassium imparts violet colour to the ﬂame. These are highly reactive elements and are strong reducing agents. All are good conductors of heat and electricity.
Elements of groups 13-18 in which the last electron enters the p-orbitals of the valence shell are called p-block elements.
The electronic configuration of the outermost shell of p-block elements (group 13, 14, 15, 16, 17 and 18) is ns2 np1-6. These elements include metals and non-metals with a few metalloids. The metallic character, however, decreases along the period but increases down the group. These possess relatively higher ionization energy which tends to increase along the period but decreases down the group. Most of them form covalent compounds. Most of these elements show negative (except some metals) as well as positive oxidation states (except F). The oxidizing power of these elements increases along the period but decreases down the group.
There are three complete series and one incomplete series of d-block elements. These are 1st or 3d-transition series which contains ten elements with atomic number 21-30 (21Sc-30Zn). 2nd or 4d-transition series which contains ten elements with atomic numbers 39-48(39Y-48Cd). 3rd or 5d transition series which contains ten elements with atomic numbers 57 and 72-80 (57La, 72Hf-80Hg). 4th or 6d transition series which is incomplete at present and contains only nine elements. These are 89Ac, 104Rf, 105Ha, Unh (Seaborgium, Z = 106), 107Bh (Bohrium), 108Hs (Hassium), 109Mt (Meitnerium), Ds (Darmstadtium, Z= 110) and Cn (Copernicium, Z = 112) or Ekamercury. The element, Z = 111 has not been discovered yet. Thus, in all, there are 39 d-block elements.
The electronic configuration of the outermost shell of d-block elements is (n–1) d1–10 ns0-2. All (except Hg) are hard, ductile metals with high melting and boiling points. All of these are good conductors of heat and electricity. Their ionization energies are higher than s-block elements but lesser than p-block elements. Most of the transition metals form coloured ions (Zn2+, Hg2+, Cd2+ are colourless.) These elements show variable oxidation states. Most of these elements possess catalytic activity. Metals and their ions are generally paramagnetic due to the presence of unpaired electrons. Most of the transition metal ions possess the tendency to form complex ions. Most transition metals form alloys.
f -block Elements:
f-Block elements are also called inner-transition elements. In these elements, the f-subshell of the inner-penultimate is progressively filled up. There are two series of f-block elements each containing 14 elements. The fourteen elements from 58Ce – 71Lu in which, 4f-subshell is progressively filled up are called lanthanides or rare elements. Similarly, the fourteen elements from 90Th – 103Lr in which, 5f-subshell is progressively filled up are called actinides.
The electronic configuration of the outermost shell of f-block elements is ns2, followed with (n–2)f1–14, (n–1) d0–2ns2. All are metals. Lanthanoids are also known as rare earth elements whereas most of the members of actinoid series are known as transuranic elements (made artificially). These show variable valency. These form coloured ions. Actinoids are radioactive. These also form complexes.
Metals, Non-Metals and Metalloids in Periodic Table
(a) Trends in metallic character in a Periodic table
(i)The metallic character increases down the group and decreases along the period.
(ii) The non-metallic character decreases down the group and increases along the period.
Note: All the non-metals and metalloids belong to p-block (except H and He).
TRENDS IN PHYSICAL PROPERTIES
It refers to the distance between the centre of the nucleus of the atom to the outermost shell containing electrons. Since the absolute value of the atomic size cannot be determined, it is usually expressed in terms of the following operational definitions.
(a) Covalent Radius
Normally, this term is used for non-metals. It is defined as half of the distance between two successive nuclei of two covalently bonded atoms in a molecule.
Covalent radius = ½ × Internuclear distance between two covalently bonded like atom(d).
Covalent radius may be of following types – Single bond covalent radius and Double bond covalent radius.
(b) Crystal Radius or Metallic Radius
The term is usually used for metals. It is defined as half of the distance between two successive nuclei of two adjacent metal atoms in the metallic closed packed crystal lattice.
This term is used in case of ions. It is the distance of outermost shell of an anion or cation from its nucleus. In other words, it is defined as the eﬀective distance from the nucleus of the ion which is under inﬂuence in an ionic bond.
Trends of Atomic Radius
(a) Along the period: On moving across a period, atomic radii decreases because of eﬀective nuclear charge increases.
(b) Down the group: On moving down a group, atomic radii increases, because of the increase in the number of orbits.
Factors Aﬀecting Atomic Radii
(a) Eﬀective nuclear charge: As the eﬀective nuclear charge increases, the attractive force between the nucleus and valence electron increases. Thus, across a period, atomic size/atomic radii decreases.
So, Atomic radii ∝ 1/Zeff
(b) Size of valence shell: Atomic radii is the measure of the radius of valence shell. As the value of n (principal quantum no.) increases, for an orbit, its size increases, thus down a group, atomic radii increases.
(c) Multiplicity of bond: Covalent radii decreases, as the multiplicity of bond increases.
d) Percentage ionic character in bond: Covalent radii of an atom in a bond depends upon % of ionic character. Increase in ionic character % leads in shortening of the bond, decreasing the atomic radii.
(e) Cationic Radii: Size of the cation is always lesser than its parent atom and greater the charge on cation, smaller is its ionic radii. E.g. Fe > Fe+ > Fe2+ > Fe3+ (decreasing ionic radii).
Formation of cation involves loss of an electron. Thus, the eﬀective nuclear charge increases, pulling the remaining electrons more tightly towards the nucleus.
(f) Anionic Radii: Size of an anion is always larger than its parent atom. Formation of an anion involves gain of electrons by an atom and so, eﬀective nuclear charge decreases. Thus, the valence shell electrons are lless tightly held by the nucleus.
It is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom i.e.
M (g) + IE → M+ (g) + e–
The amount of energies required to remove the first, second, third etc. electrons from an isolated gaseous atom are called successive ionization energies and are designated as IE1, IE2, IE3 etc. It may be noted that IE2 is always greater than IE1.Thus, the order is- IE3 > IE2 > IE1.
Factors aﬀecting Ionisation Potential
(a) The number of shells: With the increase in the number of shells, the atomic radius increases i.e. the distance of outermost shell electron from the nucleus increases and hence the ionization potential decreases.
(b) Eﬀective Nuclear Charge: Atomic size decreases with the increase in eﬀective nuclear charge because, higher the eﬀective nuclear charge, stronger will be the attraction of the nucleus towards the electron of the outermost orbit and higher will be the ionization potential.
(c) Shielding Eﬀect: The electrons of the inner orbits repel the electrons of the outermost orbit due to which the attraction of the nucleus towards the electrons of the outermost orbit decreases and thus the atomic size increases and the value of ionization potential decreases.
(d) Stability of half-filled and fully filled orbitals: The atoms whose orbitals are half-filled (p3, d5, f7) or fully filled (s2, p6, d10, f14) have greater stability than the other. Therefore, they require greater energy to remove an electron. However, the stability of fully filled orbitals is greater than that of the half-filled orbitals.
(e) Penetration power: In any atom, the s-orbital is nearer to the nucleus in comparison to p, and f orbitals. Therefore, greater energy is required to remove an electron from s-orbital than from p, d and f orbitals. The order is as follows- s > p > d > f
Periodic Trends in Ionisation Potential
In a Period: The value of ionization potential normally increase across a period, because eﬀective nuclear charge increases and the atomic size decreases.
Exceptions: In the second period, the ionization potential of Be is greater than that of B, and in the third period, the ionization potential of Mg is greater than that of Al due to the high stability of fully filled orbitals. In the second period, the ionization potential of N is greater than O and in the third period, the ionization potential of P is greater than that of S, due to the stability of half-filled orbitals.
In a Group: The value of ionization potential normally decreases down the group because of both atomic size and shielding eﬀect increase.
Exception: The value of ionization potential remains almost constant from Al to Ga in the III A group. (B > Al, Ga > In). In IV B group i.e. Ti, Zr and Hf, the I.P. of Hf is higher than that of Zr due to Lanthanide contraction. Thus the I.P. of IV B group varies as Ti > Zr < Hf.
Electron Gain Enthalpy
It is the amount of energy released when a neutral isolated gaseous atom accepts an electron to form a gaseous anion.
X (g) + E → X– (g) + EA
Similarly, the second and third electron can be added to form gaseous di-negative and tri-negative ions. The energy changes accompanying the addition of first, second, third etc. electrons to neutral isolated gaseous atoms are called successive electron affinities and are designated as EA1, EA2, EA3, etc.
Since an atom has a natural tendency to accept an electron, therefore, the first electron affinity (EA1) is always taken as positive. However, the addition of the second electron to the negatively charged ion is opposed by coulombic repulsion. Hence, energy has to be supplied for the addition of the second electron. Thus, second electron affinity (EA2) of an element is taken as negative.
Factors Aﬀecting Electron Gain Enthalpy
Atomic size or atomic radius: When the atomic size/radius increases, the electrons entering the outermost orbit is more weakly attracted by the nucleus and the value of electron gain enthalpy is lower.
Eﬀective Nuclear charge: When the eﬀective nuclear charge is more, then, the atomic size is less. Then, the atom can easily gain an electron and possess a higher value of electron gain enthalpy.
Stability of Fully-Filled and Half-Filled orbitals: The stability of the configuration having fully-filled orbitals (p6, d10, f14) and the half-filled orbitals (p3, d5, f7) is relatively higher than that of other configurations. Hence, such type of atoms have a lesser tendency to gain an electron, therefore, their electron gain enthalpy values will be very low or zero.
Trends in Electron Gain Enthalpy:
In a period, atomic size decreases with the increase in eﬀective nuclear charge and hence, increases the electron gain enthalpy.
Ongoing from C6 to N7 in the second period, the values of electron gain enthalpy decrease instead of increasing. This is because there are half-filled (2p3) orbitals in the outermost orbit of N, which are more stable. On the other hand, the outermost orbit in C has a 2p2 configuration. In the third period, the value of electron gain enthalpy of Si is greater than that of P. This is because the electronic configuration of the outermost orbit in P atom is 3p3, which being half-filled, is relatively more stable. The values of electron gain enthalpy of inert gases are zero because their outermost orbit has fully-filled p orbitals.
In a period, the value of electron gain enthalpy goes on decreasing on moving from group IA to group IIA. The value of electron gain enthalpy of the elements of group IIA is zero because ns orbitals are fully-filled and such orbitals have no tendency to accept electrons.
In a Group: The values of electron gain enthalpy normally decrease down a group because the atomic size increases, decreasing the actual attractive force of the nucleus.
The value of the electron gain enthalpy of F is lower than that of Cl because the size of F is very small and compact and the charge density is high on the surface. Therefore, the incoming electron/s experience more repulsion in comparison to Cl accounting for the highest value of Cl in the periodic table.
The values of electron gain enthalpy of alkali metals and alkaline earth metals can be regarded as zero as they do not have the tendency to form anions by accepting electron/s.
It is the tendency of an atom to attract the shared pair of electrons of the covalent bond towards itself.
Factors Aﬀecting Electronegativity
Atomic size: Electronegativity of a bonded atom decreases with increase in size since the attractive force on the valence electrons decreases and hence electronegativity decreases.
Hybridisation state of atom: Electronegativity increases with increase in the s-character of the hybrid orbital. This is because, the s-orbital is nearer to the nucleus and thus, suﬀers greater attraction leading to increased electronegativity. The number of covalent bond present between two bonded atoms is known as its bond order. With the increase in the bond order, the bond distance decreases eﬀective nuclear charge increases and thus electronegativity increases. Increasing order of electronegativity is as follows: C – C < C = C < C ≡C. When the eﬀective nuclear charge is high, the nucleus will attract the shared electrons with greater strength to give high electronegativity.
Oxidation number: The electronegativity value increases with the increase in oxidation number since the radius decreases with the increase in oxidation number.
The increasing order of electronegativity is as follows: Fe < Fe+2 < Fe+3
Electronegativity does not depend on the stability of fully-filled or half-filled orbitals because it is simply the capacity of the nucleus to attract a bonded pair of electrons.
Trends in Electronegativity: Atomic size decreases across a period. Thus, electronegativity increases. Atomic size increases down a group decreasing the electronegativity. F has maximum electronegativity value in the periodic table, while Cs has the minimum. According to the Pauling scale, the electronegativity value of F is 4.0, O is 3.5 N is 3.0 and Cl is 3.1.
(a)The elements of group IIB i.e. Zn, Cd and Hg show increase in electronegativity value down the group.
(b) The elements of group IIIA, i.e Al to Ga show increase in electronegativity value down the group.
(c) The elements of group IVA, Si onwards, show no change in electronegativity value down the group.
TRENDS IN CHEMICAL PROPERTIES
Valency/valence, also known as valence number, is the number of valence bonds a given atom has formed or can form, with one or more atoms.
Variation of valence in a period: On moving along the period, the number of valence electrons increases from 1 to 8. Consequently, the valence of the elements with respect to hydrogen increases from 1 to 4 up to group IV and then decreases to 1. However, valence with respect to oxygen increases from one to seven along the period.
Variation of valence in group: On moving down the group, the number of valence electrons remains the same. Therefore, all the elements in a group have the same valence. For example, elements of group I have valency 1 and elements of group II have valency 2.
Variation of valence in transition elements: Transition metals show variable valence of 1, 2 or 3 as they can use electrons from their outermost as well as penultimate shell, during chemical reactions as energy diﬀerence between them is small.
Some anomalous properties of second-period elements
Consider the elements of II period.
II period: Li Be B C N O F Ne
The elements of group 1, 2 (Li & Be) and of group 13–17 (B to F) differ in many respects from other members of their group. (These points of difference will be later studied in detail.)Some anomalous properties of 2nd period elements are given below with their explanation.
(a) Covalence: The maximum covalency of 2nd period elements is four while, other members may also show higher covalency. e.g., BF4– exists but [BF6]3– is known. Similarly, OF2 is known but OF4, OF6 are not while, SF4 are SF6 are known, N is never pentavalent etc.
Explanation: These elements have only two shells in their atom and the valence shell contain 4 orbitals only (one 2s and three 2p) so, a maximum of four bonds can be formed. In the 3rd period and onwards, the valence shell contains empty d–orbitals also. So, covalency may be more than four.
(b) pπ ‒ pπ Multiple bonding: Bonds like C=C, C ≡ C, N ≡ N, C=O etc. exist due to p p π – π multiple bonding. These elements are smaller in size and mostly electronegative in their respective groups thus forming multiple bonds.
(c) Diagonal relationship: The diagonal relationship between elements of II & III periods.
The diagonal relationship between these elements can be explained on the basis of approximately similar charge/ size ratio of diagonally related elements.
This article has highlighted the Classification of Elements and Periodicity in Properties in the form of short notes for class 11 students in order to understand the basic concepts of the chapter. The notes on Classification of Elements and Periodicity in Properties have not only been prepared for class 11 but also for the different competitive exams such as iit jee, neet, etc.