6. METODOLOGIA I FLUX DE TREBALL
6.2. P REPRODUCCIÓ
6.2.1. Procés d’escriptura del guió
It is based on the type of subshells which receives the differentiating electron (i.e. last electron).
(a ) s- Block Ele me nt s :
W hen last electron enters the s- orbital of the outermost (nth) shell, the elements of this class are called s- block elements.
Characteristics :
(i) Group 1 & 2 elements constitute the s - block.
(ii) General electronic configuration is ns1–2 . (iii) s - block elements lie on the extreme left of the periodic table.
(iv) This block includes metals only, except H.
Note : The total number of elements in s-block is 13 (including hydrogen).
(b) p-Block Elements :
When differentiating electron enters the p - orbital of the nth orbit, elements of this class are called p - block elements.
Characteristics :
(i) Elements of group 13 to 18 constitute the p - block.
(ii) General electronic configuration is ns2np1-6 . (iii) p - block elements lie on the extreme right of the periodic table.
(iv) This block includes some metals, all non-metals and metalloids.
Note : The total number of elements in p-block is 31.
2929 2929 (c) d - Bloc k Element s:
When differentiating electron enters the (n–1)d orbital, then elements of this class are called d - block elements.
Characteristics :
(i) Elements of group 3 to 12 constitute the d - block.
(ii) General electronic configuration is (n – 1)d1–10 ns0- 2. (iii) d - block elements lie in the centre of the periodic table.
(iv) All the d - block elements are metals and most of them form coloured complexes or ions.
Note : The total number of elements in d-block is 39.
(d) f - block eleme nts :
When last electron enters into f - orbital of (n – 2)th shell, elements of this class are called f - block elements.
Characteristics :
(i) All f - block elements belong to 3rd group.
(ii) General electronic configuration is (n – 2)f 1 – 14 (n– 1)d0-1 ns2.
(iii) f-block elements are present in two separate rows below the periodic table.
(iv) All f-block elements are metals only.
The elements of f- block have been classified into two series :
(A) Lanthanides : 14 elements present after element lanthanum (57La) are called lanthanides.1st inner transition series of metals or 4f - series, contains 14 elements i.e. 58Ce to 71Lu .
(B) Actinides : 14 elements present after element actinium (89Ac) are called actinides. 2nd inner transition series of metals or 5f- series, also contains 14 elements i.e. 90Th to 103Lr.
Note : The total number of elements in f-block is 28.
Division of the periodic table into various blocks :
Different Types of Elements :
(i) Noble gases : The elements belonging to group 18 are called noble gases or aerogenous. They are also known as inert gases because their outermost orbitals are completely filled. Helium (He) is an exception. It has only two electrons which are present in s-orbital.
(ii) Representative elements : Elements in which atoms have all shells complete except outermost shell which is incomplete.Except 18th group, all s - block and p - block elements are collectively called normal or representative elements.
(iii) Transition elements : Those elements which have partially filled d - orbitals in neutral state or in any stable oxidation state are called transition elements.
Note : Zn, Cd and Hg are d-block elements , but not transition elements, because they do not contain partially filled d-orbitals.
(iv) Inner transition elements :- They contain three incomplete outermost shell and were also referred to as rare earth elements, since their oxides were rare in earlier days.
(v) Diagonal relationship : Some elements of 2nd and 3rd period show diagonal relationship among them. They represent the same properties of two periods. This relation is known as diagonal relation.
(vi) Transuranium elements : The elements with atomic number greater than 92 (Z > 92) are known as transuranium elements.
3030 3030 PAGE # 30
1.00791 H Hydrogen 6.9413 Li Lithium 22.99011 Na Sodium 39.09819 K Potassium 85.46837 Rb Rubidium 132.9155 Cs Caesium (223)87 Fr Francium
9.01224 Be Beryllium 24.30512 Mg Magnesium 40.07820 Ca Calcium 87.6238 Sr Strontium 137.3356 Ba Barium (226)88 Ra Radium
44.95621 Sc Scandium 88.90639 Y Yttrium
47.86722 Ti Titanium 50.94223 V Vanadium 51.99624 Cr Chromium
54.93825 Mn Manganese 91.22440 Zr Zirconium 178.4972 Hf Hafnium (261)104 Rf Rutherfordium
92.90641 Nb Niobium 180.9573 Ta Tantalum (262)105 Db Dubnium
42 Mo Molybdenum 183.8474 W Tungsten (266)106 Sg Seaborgium
95.94(98)43 Tc Technetium 186.2175 Re Rhenium (264)107 Bh Bohrium
55.84526 Fe Iron 58.93327 Co Cobalt 58.96328 Ni Nickel
63.54629 Cu Copper 107.8747 Ag Silver 196.9779 Au Gold (272)111 Rg Roentgenium
(281)110 Ds Darmstadtium
(268)109 Mt Meitnerium
(277)108 Hs Hassium 190.2376 Os Osmium 101.0744 Ru Ruthenium
102.9145 Rh Rhodium 192.2277 Ir Iridium
195.0878 Pt Platinum 106.4246 Pd Palladium
65.3930 Zn Zinc 112.4148 Cd Cadmium 200.5980 Hg Mercury
69.72331 Ga Gallium 72.6432 Ge Germanium 74.92233 As Arsenic 78.9634 Se Selenium 79.90435 Br Bromine
83.8036 Kr Krypton 114.8249 In Indium
118.7150 Sn Tin
121.7651 Sb Antimony 127.6052 Te Tellurium 126.9053 I Iodine
131.2954 Xe Xenon 204.3881 Ti Thallium
207.282 Pb Lead
208.9883 Bi Bismuth (209)84 Po Polonium (210)85 At Astatine (222)86 Rn Radon
26.98213 Al Aluminium 10.8115 B
12.0116 C Carbon
14.0077 N Nitrogen 15.9998 O Oxygen 18.9989 F Fluorine
20.18010 Ne Neon 39.94818 Ar Argon
35.45317 Cl Chlorine
32.06516 S Sulphur
30.97415 P Phosphorus
28.08614 Si Silicon
Boron
4.00262 He Helium
B BoronSymbol
510.81113 Atomic number Element name
Relative atomic mass IIIA
Group IUPAC IIA
IA IIIBIVBVBVIBVIIB
VIIIB
1 2 3456789101112IBIIB
13IIIA14151617
18 IVAVAVIAVIIA
VIIIA 1 2 3 4 5 6 7
Period
Group LANTHANOIDES
(227)89 Ac Actinium 140.1258 Ce Cerium 232.0490 Th Thorium
140.9159 Pr Praseodymium 231.0491 Pa Protactinium
144.2460 Nd Neodymium 238.0392 U Uranium
(145)61 Pm Promethium
150.3662 Sm Samarium (244)94 Pu Plutonium
(237)93 Np Neptunium
151.9663 Eu Europium (243)95 Am Americium
157.2564 Gd Gadolinium (247)96 Cm Curium
158.9365 Tb Terbium (247)97 Bk Berkelium
162.5066 Dy Dysprosium (251)98 Cf Californium
164.9367 Ho Holmium (252)99 Es Einsteinium
197.2668 Er Erbium 168.9369 Tm Thulium 173.0470 Yb Ytterbium
174.9771 Lu Lutetium (257)100 Fm Fermium
(258)101 Md Mendelevium (259)102 No Nobelium (262)103 Lr Lawrencium
138.9157 La Lanthanum
s–BlockElements
T ra n s it io n M e ta ls ( d – B lo c k E le m e n ts )
p–Block Elements
In n e r - T ra n s it io n M e ta ls ( f- B lo c k E le m e n ts )
ACTINOIDESP E R IO D IC T A B L E O F T H E E L E M E N T S
3131 cyclically, i.e. it decreases at first for some elements, becomes minimum in the middle and then increases.
The following two factors explain this trend -(i) atomic radii decrease due to increase of nuclear charge.
(ii) the number of valence electrons increases in a period.
As to accommodate all the valence electrons, the volume increases. These two factors oppose each other. The effect of first factor is more on the left hand side and that of the second factor is more on the right hand side in the periodic table . The volumes are in cm3.
cm3) is observed in the case of francium.
( b) D e n s it y :
The density of the elements in solid state varies periodically with their atomic numbers . At first, the density increases gradually in a period and becomes maximum somewhere for the central members and then starts decreasing afterwards gradually. The value of densities in the table are in g/cc.
Period/ periodicity with rise of atomic number. It is observed that elements with low values of atomic volumes have high melting points, while elements with high values of atomic volumes have low melting points. In general, melting points of elements in any period at first melting point (3727ºC) amongst non-metal. Helium has the minimum melting point (–270ºC) amongst all elements . The metals Cs (m.p. = 28.5ºC), Ga (m.p. = 30ºC) and Hg (m.p. = –39ºC) are known in liquid state at 30ºC.
The boiling points of the elements also show similar trends, however, the regularities are not striking as noted in the case of melting points.
( d) At omic Ra dius :
(i) Covalent radius : It may be defined as one - half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond.
A B
1
2AB = rcovalent (of element X)
X X
e.g. The internuclear distance between two hydrogen atoms in H2 molecule is 74 pm. Therefore, the covalent radius of hydrogen atom is 37 pm.
Note : Covalent radius is generally used for non - metals.
(ii) Vander Waal’s radius : It may be defined as half of the internuclear distance between two adjacent atoms of the same element belonging to two nearest neighbouring molecules of the same substance.
E H
X X X X
1
2EH = rvander Waals
3232 overlapping of atomic orbitals, as a result of this, the internuclear distance between the covalently bonded atoms is less than the internuclear distance between the non bonded atoms.
e.g. Vander Waals radius of helium is 1.20 Å (iii) Metallic radius (Crystal radius) : Metallic radius may be defined as half of the internuclear distance between two adjacent atoms in a metallic lattice.
C
• The metallic radius of an atom is always larger than its covalent radius.
Note : The order of different radius is - r Vander Waals > rMetallic
> rCovalent
(iv) Variation of atomic radii in a period : As we move from left to right across a period, there is a regular decrease in atomic radii of the representative elements. This is due to the fact that number of energy shells remains the same in a period, but nuclear charge increases gradually as the atomic number increases. This increases the force of attraction towards nucleus which brings contraction in size.
This can also be explained on the basis of effective nuclear charge which increases gradually in a period i.e. electron cloud is attracted more strongly towards nucleus as the effective nuclear charge becomes more and more as we move in a period.
The increased force of attraction brings contraction in size.
(v) Variation of atomic radii in a group : Atomic radii in a group increase as the atomic number increases.
The increase in size is due to extra energy shells which outweigh the effect of increased nuclear charge.
The following table illustrates the periodicity in atomic radii (covalent radii) of representative elements. The radii given in the table are in angstrom (Å).
Periodicity in atomic radii (covalent radii)
Atoms of zero group elements do not form chemical bonds among themselves. Hence for them Vander Waals radii are considered.
Element He Ne Ar Kr Xe
Vander Waals
radii (in Aº ) 1.20 1.60 1.91 2.00 2.20 The sudden increase in atomic radii in comparison to the halogens (the elements of 7th group) in case of inert gases, is due to the fact that, Vander Waals radii are considered which always possess higher values than covalent radii.
The decrease in the size of transition elements is small since the differentiating electrons enter into inner ‘d’ levels. The additional electrons into (n–1)d levels effectively screen much of increased nuclear charge on the outer ns electrons and therefore, size remains almost constant.
However, in vertical columns of transition elements, there is an increase in size from first member to second member as expected, but from second member to third member, there is very small change in size and sometimes sizes are same. This is due to Lanthanide contraction (in the lanthanide elements differentiating electrons enter into 4f-levels).
Since these electrons do not effectively screen the valence electrons from the increased nuclear charge, the size graduallydecreases. This decrease is termed lanthanide contraction.
Conclusions
(i) The alkali metals which are present at the extreme left of the periodic table have the largest size in a period.
(ii) The halogens which are present at the extreme right of the periodic table have the smallest size.
3333 3333 (iii) The size of the atoms of inert gases are, however,
larger than those of preceding halogens because in inert gases van der W aals' radii are taken into consideration.
(iv) In a group of transition elements, there is an increase in size from first member to second member as expected but from second member to third member, there is very small change in size and sometimes sizes are same. This is due to Lanthanide contraction.
(iv) Ionic radius : It is the distance between the nucleus and outermost shell of an ion or it is the distance between the nucleus and the point where the nucleus exerts its influence on the electron cloud.
(A) The radius of the cation is always smaller than the atomic radius of its parent atom. This is due to the fact that nuclear charge in the case of a cation is acting on a lesser number of electrons and pulls them closer.
(B) The radius of the anion is always larger than the atomic radius of its parent atom. In an anion as electron or electrons are added to the neutral atom, the nuclear charge acts on more electrons so that each electron is held less tightly and thereby the electron cloud expands.
Comparative sizes of atoms and their cations
Atom
Atomic radii (crystal, Å)
Corresponding
cations Ionic radii (Å)
Li 1.52 Li+ 0.59
Na 1.86 Na+ 0.99
K 2.31 K+ 1.33
Mg 1.60 Mg2+ 0.65
Ba 2.22 Ba2+ 1.35
Al 1.43 Al3+ 0.50
Pb 1.75 Pb2+ 1.32
Conclusions
• The radius of cation (positive ion) is always smaller than that of the parent atom.
• The radius of anion (negative ion) is always larger than that of the parent atom.
• The ionic radii in a particular group increase in moving from top to bottom.
• In a set of species having the same number of electrons (isoelectronic), the size decreases as the charge on the nucleus increases.
• The size of the cations of the same element decreases with the increase of positive charge.
( e) Ionis a tion Ene rgy ( IE) :
Ionisation Energy (IE) of an element is defined as the amount of energy required to remove an electron from an isolated gaseous atom of that element resulting in the formation of a positive ion.
(i) Characteristics :
(A)The energy required to remove the outermost electron from an atom is called first ionisation energy (IE)1.
After removal of one electron, the atom changes into monovalent positive ion.
M(g) + IE1 M+(g) + e–
(B) The minimum amount of energy required to remove an electron from monovalent positive ion of the element is known as second ionisation energy (IE)2.
M+(g) + IE2 M2+(g) + e–
(C) The first, second etc. ionisation energies are collectively known as successive ionisation energies.
M2+(g) + IE3 M3+(g) + e–
In general (IE)1 < (IE)2 < (IE)3 so on, because, as the number of electrons decreases, the attraction between the nucleus and the remaining electrons increases considerably and hence subsequent ionisation energies increase.
(D) Units : Ionisation energy is expressed either in terms of electron volts per atom (eV/atom) or Kilojoules per mole of atoms (KJ mol – 1) or K cal mol – 1.
1 eV/atom = 96.49 KJ/mol = 23.06 Kcal/mol = 1.602 × 10–19 J/atom
(ii) Factors influencing ionisation energy :
(A) Size of the atom : Ionisation energy decreases with increase in atomic size. As the distance between the outermost electrons and the nucleus increases, the force of attraction between the valence shell electrons and the nucleus decreases. As a result, outermost electrons are held less firmly and lesser amount of energy is required to knock them out.
For example, ionisation energy decreases in a group from top to bottom with increase in atomic size.
(B) Nuclear charge : The ionisation energy increases with increase in the nuclear charge. This is due to the fact that with increase in the nuclear charge, the electrons of the outermost shell are more firmly held by the nucleus and thus greater amount of energy is required to pull out an electron from the atom.
For example, ionisation energy increases as we move from left to right along a period due to increase in nuclear charge.
(C) Shielding effect : The electrons in the inner shells act as a screen or shield between the nucleus and the electrons in the outermost shell. This is called shielding effect or screening effect. Larger the number of electrons in the inner shells, greater is the screening effect and smaller the force of attraction and thus ionisation energy decreases.
These electrons shield the outer electrons from the
nucleus
This electron does not feel the full inward pull
of the positive charge of the nucleus
3434 3434 PAGE # 34 (D) Penetration effect of the electrons : The
ionisation energy increases as the penetration effect of the electrons increases. It is a well known fact that the electrons of the s-orbital have the maximum probability of being found near the nucleus and this probability goes on decreasing in case of p, d and f orbitals of the same energy level.
Greater the penetration effect of electrons more firmly the electrons will be held by the nucleus and thus higher will be the ionisation energy of the atom.
For example, ionisation energy of aluminium is comparatively less than magnesium as outermost electron is to be removed from p-orbital (having less penetration effect) in aluminium, whereas in magnesium it will be removed from s-orbital (having larger penetration effect) of the same energy level.
Note : With in the same energy level,the penetration effect decreases in the order s > p > d > f
(E) Electronic Configuration : If an atom has exactly half-filled or completely filled orbitals, then such an arrangement has extra stability.The removal of an electron from such an atom requires more energy than expected. For example,
E1 of Be >
E1 of BBe (Z = 4) ) stable more (
orbitalCompletelyfilled s 2 , s
12 2
B (Z = 5) ) stable less ( orbitalPartiallyfilled
p 2 , s 2 , s
1 2 2 1
As noble gases have completely filled electronic configurations, they have highest ionisation energies in their respective periods.
(iii) Variation of ionisation energyinaperiod: In general, the value of ionisation energy increases with increase in atomic number across a period. This can be explained on the basis of the fact that on moving across the period from left to
right-(A) nuclear charge increases regularly.
(B) addition of electrons occurs in the same shell.
(C) atomic size decreases.
(iv) Variation of ionisation energy in a group : In general, the value of ionisation energy decreases while moving from top to bottom in a group.This is because
-(A) effective nuclear charge decreases regularly.
(B) addition of electrons occurs in a new shell.
(C) atomic size increases.
(v) Conclusions :
(i) In each period, alkali metals show lowest first ionisation enthalpy. Caesium has the minimum value.
(ii) In each period, noble gases show highest first ionisation enthalpy. Helium has the maximum value of first jonisation enthalpy.
(iii) The representative elements show a large range of values of first ionisation enthalpies, metals having low while non-metals have high values.
(iv) Generally. ionisation enthalpies of transition metals increase slowly as we move from left to right in a period. The f-block elements also show only a small variation in the values of first ionisation enthalpies.
(f) Ele ct ron Affinity (EA) :
Electron affinity is defined as the energy released in the process of adding an electron to a neutral atom in the gaseous state to form a negative ion.
X(g) + e– X–(g) + Energy (E.A.) Cl(g) + e– Cl– (g) + 349 KJ/mol
The electron affinity of chlorine is 349 KJ/mol.
The addition of second electron to an anion is opposed by electrostatic repulsion and hence the energy has to be supplied for the addition of second electron.
O(g) + e– O– (g) + Energy (EA -) O–(g) + e– O2– (g) – Energy (EA-)
(EA -) is exothermic whereas, (EA-) is endothermic.
(i) Units : Kilo joules per mole (KJ/mol) of atoms or electron volts per atom (eV/atom).
(ii) Factors affecting electron affinity:
(A) Nuclear charge : Greater the magnitude of nuclear charge greater will be the attraction for the incoming electron and as a result, larger will be the value of electron affinity.
Electron affinity Nuclear charge.
(B) Atomic size : Larger the size of an atom is, more will be the distance between the nucleus and the incoming electron and smaller will be the value of electron affinity.
size Atomic
1 E.A.
(C) Electronic configuration : Stable the electronic configuration of an atom lesser will be its tendency to accept the electron and lower will be the value of its electron affinity.
3535 3535 (iii) Variation of electron affinity in a period : On
moving across the period the atomic size decreases and nuclear charge increases. Both these factors result into greater attraction for the incoming electron, therefore electron affinity in general increases in a period from left to right.
(iv) Variation of electron affinity in a group : On moving down a group, the atomic size as well as nuclear charge increase, but the effect of increase in atomic size is much more pronounced than that of nuclear charge and thus, the incoming electron feels less attraction consequently, electron affinity decreases on going down the group.
(v) Some irregularities observed in general trend:
(A) Halogens have the highest electron affinities in their respective periods. This is due to the small size and high effective nuclear charge of halogens.
Halogens have seven electrons in their valence shell.
By accepting one more electron they can attain stable electronic configuration of the nearest noble gas. Thus they have maximum tendency to accept an additional electron.
(B) Due to stable electronic configuration of noble gases electron affinities are zero.
(C) Be, Mg, N and P also have exceptionally low values of electron affinities due to their stable electronic configurations.
Be = 1s2, 2s2 N = 1s2, 2s2, 2p3 Mg = 1s2, 2s2 , 2p6, 3s2 P = 1s2, 2s2, 2p6, 3s2, 3p3
Conclusion
(i) The electron gain enthalpies, in general, become
(i) The electron gain enthalpies, in general, become