Periodic table notes part 1.pdf

F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta PERIODIC TABLE AND PERIODICITY OF PROPERTIES Q.1] Give Mendeleev’s and Modern Periodic law. Ans. Mendeleev’s Periodic Law:  Properties of elements are periodic functions of t...

F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta PERIODIC TABLE AND PERIODICITY OF PROPERTIES Q.1] Give Mendeleev’s and Modern Periodic law. Ans. Mendeleev’s Periodic Law:  Properties of elements are periodic functions of their atomic weights. Modern Periodic Law:  Properties of elements are periodic functions of their atomic numbers. Q.2] Discuss the main features of long form of periodic table. Ans. Main features of long form of periodic table are as follows: Periods:  Definition: Horizontal rows in a periodic table are called as Periods. There are seven such periods.  Indicates :Each period begins with an element whose outermost shell configuration is ns 1. A period ends with an element whose ns and np orbitals are fully occupied i.e. ns2np6. The period number indicates shell number (n) in each element.  First Period: It has only one principal shell (n=1) in this period. It has two elements H(Z=1) and He (Z=2)  Second Period: This period has eight elements. They are from Li (Z=3) to Ne (Z=10). There are 2 shells (n=2) in each element of this period.  Third Period: This period has eight elements. They are from Na (Z=11) to Ar (Z=18). There are 3 shells (n=3) in each element of this period. The first three periods are short periods.  Fourth Period : This period has eighteen elements. They are from K (Z=19) to Kr (Z=36). There are 4 shells (n=4) in each element of this period.  Fifth Period : This period has eighteen elements. They are from Rb (Z=37) to Xe (Z=54). There are 5 shells (n=5) in each element of this period.  Sixth Period : This period has thirty two elements. They are from Cs (Z=55) to Rn (Z=86). There are 6 shells (n=6) in each element of this period. It includes lanthanides, La (Z=57) to Lu (Z=71). The fourth, fifth and sixth periods are called long periods.  Seventh Period : There are 7 shells (n=7) in each element of this period. It includes all radioactive elements. This period has thirty two elements. They are from Fr (Z=87). It includes actinides, Ac (Z=89) to Lr (Z=103) and recently discovered elements upto atomic number 111. The elements beyond uranium (Z=92) which have been prepared artificially by nuclear reactors are called transuranic elements. It is now a complete period with 32 elements though the last few elements are not thoroughly studied. Groups: F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta  Definition: Vertical columns in a periodic table are called as Groups. There are in all sixteen such groups as given below:  Normal elements: The first two columns constitute IA and IIA groups and last six columns constitute IIIA, IVA, VA, VIA, VIIA and zero group. All ‘A’ groups (totally seven) are of normal elements.  Noble gases: The elements of zero group are noble gases.  Transition and Inner transition elements : The elements of seven groups of IB, IIB, IIIB, IVB, VB, VIB, VIIB (seven columns) and group VIII (three columns) are placed in the middle part of the periodic table and are called as transition elements. In group IIIB, some elements are transition elements and remaining are inner transition elements. Q.3] Write a short note on types of elements. Ans. Types of elements: The elements of long form of periodic table can be divided into the following four types, depending upon the number of incompletely filled shells in their atoms: 1. Inert elements (Noble gases) : Filling of electrons: These are the elements in which the s and p subshells of the outermost shell are completely filled with electrons. General electronic configuration: General electronic configuration of their outermost shell is ns2np6 (except for helium which has ns2 configuration). Elements: Elements of Zero group i.e. Group 18 of the periodic table comprise the inert elements. These elements are gases. These elements are chemically inert and are known as inert gases. 2. Normal (s and p block) elements: Filling of electrons: These are the elements in which only the s or p subshells of their outermost shell is incompletely filled with electrons. So they are also known as s or p block elements. General electronic configuration: General electronic configuration of their outermost shell is ns1-2for s-block elements and ns2np1-5 for p-block elements. Elements: The elements of group 1 and 2 in the periodic table comprise the s-block elements. They are all metals. The elements of group 13, 14, 15, 16 and 17 in the periodic table are p- block elements. Most of them are non-metals. 3. Transition (d-block) elements: Filling of electrons: These are the elements in which last two shells are incompletely filled with electrons. The d subshell of their penultimate shell (n-1) has vacancy for filling of electrons, hence they are known as d-block elements. General electronic configuration: General electronic configuration of their outermost shell is (n-1)d1-10ns-0-2. F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta Elements: The elements from group 3to 12 in the periodic table comprise the d-block elements. They are all metals. 4. Inner transition(f-block) elements: Filling of electrons: These are the elements in which last three shells are incompletely filled with electrons. The f subshell of their prepenultimate shell (n-2) has vacancy for filling of electrons, hence they are known as f-block elements. General electronic configuration: General electronic configuration of their outermost shell is (n-2)f1-14(n-1)d0-1 ns2. Elements: There are two series of fourteen elements. There are called as lanthanides and actinides. Elements of each series are extremely close to one another in their properties. They are shown in two rows below all the other elements in the periodic table. Q.4] Define periodic properties. Name some of the periodic properties. Ans. Periodicity in properties:  The physical and chemical properties of elements which vary steadily in a period and are repeated at certain regular intervals of atomic number in the periodic table are known as periodicity in properties.  Some of the periodic properties are: o Effective Nuclear charge o Atomic and ionic radii o Electron gain enthalpy o Ionization enthalpy o Electronegativity Q.5] Write a short note on screening effect and effective nuclear charge. Ans. Definition of screening or shielding effect: In a multi-electron atom the electrons in the innerlying shells act as a screen or shield between the nuclear attraction and the outermost electron. This effect of inner electrons on the outermost electron is known as screening or shielding effect. Extent of shielding:  The extent to which inner electrons shield the outermost electron depends on the following: o Shape of the orbital o Distance of the orbital from the nucleus.  Since s-electron is closer to the nucleus than p-electron, p-electron in turn is closer than d- electron and d-electron is closer than f-electron; for a given quantum shell, shielding ability of inner electrons decreases in the order of s > p > d > f.  In other words, ns electron shields the nucleus more than np electron and np electron shields more effectively than nd electron and so on. F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta Definition of effective nuclear charge: The actual charge or nuclear attraction which is felt by the outermost electron in an atom considering the screening or shielding effect of inner electrons is known as effective nuclear charge (Z* ). Relation between effective nuclear charge and shielding constant : The effective nuclear charge (Z* ) is less than the nuclear charge (Z) by an amount ‘S’ which is shielding constant. Therefore, Z* = Z – S. Calculation of shielding constant: Shielding constant can be calculated using Slater rules. Slater proposed a set of rules where electrons are divided into groups and each group is assigned a contribution per electron. Q.6] Discuss Salter rules for calculation of shielding constant. Ans. Slater’s Rules: Slater proposed a set of rules for calculating shielding constant (S). The electrons are divided into group as (1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4f) (5s, 5p) etc. The group to which the electron under consideration belongs is represented as group ‘n’. Each group is assigned a contribution per electron. The sum of all such contributions that affect the electron under consideration is the shielding constant. The assigned contributions are as given below: i. Groups to the right of n: Zero for each electron in the group that is right to the group having the electron under consideration ii. Same group (n): 0.35 for each electron in the same group(n) other than the electron under consideration. iii. (n-1) group: 0.85 for each electron in the (n-1) group. iv. Closer than (n-1) group: 1.00 for each electron closer than (n-1) group. v. If electron under consideration is ‘d’ or ‘f’ electron: 1.00 for each electron in all lower groups, if the electron under consideration is ‘d’ or ‘f’ electron. Q.7] Calculate the effective nuclear charge felt by 4s electron of potassium. Ans. Solution: Electronic configuration of Potassium: At No. (Z) = 19 = 1s2 2s2 2p6 3s2 3p6 4s1 Groups of electron: (1s2) (2s2, 2p6) (3s2, 3p6) (4s1) Effective nuclear charge: Z* = Z – S. where S is the screening constant. Applying Slater rules: S = 0.85 for each electron in (n-1) group, S = 1.00 for each electron in groups lower than n-1 Substitution: F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta Z* = Z – S. = 19 - {0.85 [no. of electron in group (n-1) i.e. 3s and 3p] + 1.00 [no. of electron in groups lower than (n-1) i.e. 1s, 2s and 2p]} = 19 – {(0.85 x 8) + (1.00 x 10)} = 19 – 16.8 = 2.20 Answer:Z* felt by 4s electron in K atom = 2.20 Q.8] Calculate the effective nuclear charge experienced by 3d electron of zinc. Ans. Solution: Electronic configuration of Potassium: At No. (Z) = 30 = 1s2 2s2 2p6 3s2 3p6 3d10 4s2 Groups of electron: (1s2) (2s2, 2p6) (3s2, 3p6) (3d10) Effective nuclear charge: Z* = Z – S. where S is the screening constant. Applying Slater rules: S = 0.35 for each electron in the same group other than the electron under consideration, S = 1.00 for each electron in all lower groups, if the electron under consideration is‘d’ or ‘f’ electron. Substitution: Z* = Z – S. = 30 - {0.35 (no. of electron in group ‘n’ i.e. 3d) + 1.00 (no. of electron in all groups lower than ‘n’ i.e. 1s, 2s, 2p, 3s and 3p)} = 30 – {(0.35 x 9) + (1.00 x 18)} = 30 – {3.15 + 18} = 30 - 21.15 = 8.85 Answer: Z* felt by 3d electron in Zn atom = 8.85 Q.9] Calculate the effective nuclear charge felt by 4s electron of copper. Ans. Solution: Electronic configuration of Potassium: At No. (Z) = 29 = 1s2 2s2 2p6 3s2 3p63d104s1 Groups of electron: (1s2) (2s2, 2p6) (3s2, 3p6) (3d10) (4s1) Effective nuclear charge: Z* = Z – S. where S is the screening constant. Applying Slater rules: S = 0.85 for each electron in (n-1) group, S = 1.00 for each electron in groups lower than n-1 Substitution: Z* = Z – S. F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta = 29 - {0.85 (no. of electron in group (n-1) i.e. 3d, 3s, 3p) + 1.00 (no. of electron in groups lower than (n-1) i.e. 1s, 2s, 2p)} = 29 – {(0.85 x 18) + (1.00 x 10)} = 29 – {15.3 + 10} = 29 – 25.3 = 3.70 Answer: Z* felt by 4s electron in Cu atom = 3.70. Q.10] Discuss atomic, covalent and vander Waals radii. Also, explain their effects on properties of elements. Atomic radii::  Magnitude of the radius of atom or ion represent the size of the atom or ion.  The distance between the nucleus of the atom and its electron in its highest energy level is considered as atomic radius. Problems associated: The absolute size of the atom or ion cannot be defined in an exact manner because of the following reasons:  The exact position occupied by an electron in an atom cannot be defined with certainty.  The probability of finding the electron round the nucleus in an atom is influenced by the presence of other atoms in its environment. As a result, the size of the atom changes with the change of the atoms surrounding it. Hence, it is not possible to measure the radius of an isolated atom or ion. Therefore internuclear distances are used to calculate the atomic and ionic radii. 3 operational concepts - : Covalent radii, van der Waals radii and ionic radii Covalent radii:  In those elements, where the atoms are covalently bonded in a diatomic molecule, half of the covalent bond length is the covalent radius.  It gives an approximate value as little overlapping of the outermost orbitals of the two atoms is likely to take place.  For example, bond length between two chlorine atoms in its molecule is 19.8 x 10—2 nm. Hence, the covalent radius of chlorine is 9.9 x 10—2 nm. Covalent radii F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta  For example, bond length between two chlorine atoms in its molecule is 19.8 x 10 —2 nm. Hence, the covalent radius of chlorine is 9.9 x 10—2 nm.  In a non-polar heteronuclear diatomic molecule, the single covalent bond length can be approximated as the sum of their non-polar covalent radii. Therefore if radius of one is known, the radius of other can be determined. For example, C-F bond length is 14.9 x 10-2 nm and the covalent radius of carbon atom is 7.7 x 10-2 nm. Hence, the atomic radius of fluorine is 7.2 x 10-2. van der Waals radii:  In the elements where the atoms are not chemically bound to each other, the only attractive forces are van der Waals forces.  In such elements, the shortest distance to which the atoms can approach before their electron clouds start repelling each other is called van der Waals radii. van der Waals radii  van der Waals radii are usually much larger than the covalent radii because in covalent radii the orbital overlap due to bond formation bring the atoms much closer. Ionic radii:  Ionic radii can be determined from the ionic compounds by measuring the ionic bond length  The atoms of metal lose electrons to form cation. M → M+ + e- The cations are considerably smaller than their respective atoms as the nuclear charge is greater than the extranuclear charge in them, which leads to contraction. Example: rNa = 15.4 x 10-2 nm while rNa+= 9.7 x 10-2 nm  The atoms of non- metal accept electrons to form anion. A+ e-→ A- The anions are considerably bigger than their respective atoms as the nuclear charge is lesser than the extranuclear charge in them, which leads to expansion. Example: rCl = 9.9 x 10-2 nm while rCl-= 18.7 x 10-2 nm  Therefore, cation is smaller than the atom from which it is formed while anion is larger than the atom from which it is formed. Effect of Atomic and Ionic radii on properties: Though atomic and ionic radii are physical characteristics, they have effect on the chemical behavior of an element F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta  Basicity: Smaller cations are less basic in nature. Example: Among the alkali metals, oxides of lithium are less basic.  Hydration: Smaller ions undergo hydration to a greater extent. Example: Among the alkali metals, Li+ undergoes maximum hydration.  Complexing power: Smaller ions have greater complexing power and their complexes are more stable.  Chemical reactivity: It depends on the ability of atoms to lose, gain or share electrons which depends on the atomic and ionic size.  Extent of ionicity: The extent of ionicity in a covalent bond depends on sizes of cation and anion. Q.11] Define atomic radii and discuss the trends in the atomic size across the period. Ans. Atomic radius: The distance between the nucleus of the atom and its electron in its highest energy level is considered as atomic radius. General trend across the period:  As we go across the period in a periodic table, atomic number increases, nuclear charge increases but the number of shells in the atoms of all elements in a period remains the same.  Therefore, the atomic size decreases as the nuclear attraction on the outermost electrons increases. Variation of atomic size:  In a period, atomic size is largest for alkali metals, decreases gradually and it becomes smallest for the noble gas elements  However, the extent of decrease of atomic size does differ for different types of elements depending upon the addition of electrons in different energy level. Variation in the elements of s and p-block:  Addition of electrons: In s and p block element, along a period with increase in atomic number, electrons are added to the outermost shell.  Electron density and screening effect: Thus, the electron density of inner shell remains unchanged and screening effect of the inner electrons remains the same.  Effective nuclear charge and atomic size: Effective nuclear charge on the outermost electron also increases. Hence, the atomic size decreases significantly.  Therefore, these elements show differences in their properties. Variation in d-block elements:  Addition of electrons: In d block element, along a period with increase in atomic number, electrons are added to the penultimate shell.  Electron density and screening effect: Thus, the electron density of inner shells increases steadily and screening effect of the inner electrons also increases. F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta  Effective nuclear charge and atomic size: Hence increase in the effective nuclear charge on the outermost electron decreases. Hence, the atomic size decreases to a lesser extent than what has been observed among the s and p-block element.  Therefore, d-block elements are more similar in their sizes and properties. Q.12] Define atomic radii and discuss the trends in the atomic size down the group. Ans. Atomic radius: The distance between the nucleus of the atom and its electron in its highest energy level is considered as atomic radius. General trend down the group:  As we go down the group in a periodic table, atomic number increases, nuclear charge increases and simultaneously number of shells in their atoms also increases.  Therefore, the atomic size increases down the group. Variation of atomic size:  The difference in the atomic size between the elements of 2nd and 3rd period is much greater than between subsequent periods in a group.  The extent of increase in the atomic size of the elements as we move down from one period to another in a group is different depending upon the number of electrons in the inner shells of the atoms. Variation in the elements of 2 and 3 period: nd rd  Electrons in the inner shells: There are only 2 electrons in inner shell.  Electron density and screening effect: Thus, the electron density of inner shells is comparatively low and screening effect of the inner electrons is less.  Effective nuclear charge and atomic size: Effective nuclear charge is comparatively more. Hence, the contraction in the atomic size is more. Variation in elements of subsequent periods:  Electrons in the inner shells: Electrons in inner shells are more.  Electron density and screening effect: Thus, the electron density of inner shells is comparatively more and screening effect of the inner electrons is more.  Effective nuclear charge and atomic size: Effective nuclear charge is comparatively less than what is observed in the 2nd period. Hence, the contraction in the atomic size is less. Therefore, the elements of 2nd period show less similarity with other elements of their group. Typical or representative elements : That is why elements of 3rd period are chosen as typical or representative elements of their groups as they show more resemblance in properties with other elements of their group. Q.13] Define atomic radii and state the difference between atomic and ionic radii. Ans. Atomic radius: The distance between the nucleus of the atom and its electron in its highest energy level is considered as atomic radius. F.Y.B.Sc., Chem-I, Sem-1, Unit 2.2 PERIODIC TABLE & PERIODICITY OF PROPERTIES By, Dr. Hetal Mehta Difference: Cation is smaller than the atom from which it is formed while anion is larger than the atom from which it is formed. Q.14] Explain isoelectronic species with suitable examples. Ans. Definition: Isoelectronic species (atoms or ions) are those which have the same number of electrons.  For such species as the atomic number increases, nuclear charge increases and the increasing nuclear charge acts on the same number of electrons in each species and therefore size decreases.  Examples : Ions Na+ Mg2+ Al3+ Si4+ No. of electrons 10 10 10 10 Nuclear charge (Z) 11 12 13 14  The nuclear charge increases from Na+ to Si4+ but the number of electrons remains the same in each ion. Hence, the size decreases from Na+ to Si4+.