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Chapter 8 PERIODIC PROPERTIES 8.1 Introduction The genesis of classificatio...

Chapter 8 PERIODIC PROPERTIES 8.1 Introduction The genesis of classification of elements dates back to early 1800s when the German chemist Johann Dobereiner made the first attempt to systematise the study of properties of elements and gave the Dobereiners law of triads. This was followed by the Newlands law of octaves. The first detailed classification of elements was proposed by Russian chemist Dmitri Mendeleev (1837–1907) and German chemist Lothar Meyer (1830–1895). Both these scientists worked independently and in 1869 proposed that when elements are arranged in the increasing order of their atomic weights, similarities appear in their physical and chemical properties at regular intervals. However, the Mendeleev’s periodic table had certain anomalies, which were addressed in due course of time. Gradual improvement saw the development of the modern periodic law (given by Henry Moseley) in which the elements are arranged in the order of increasing atomic numbers. The elements are arranged in the periodic table in the order of increasing atomic numbers. They are divided into 18 vertical columns called groups and seven horizontal rows termed as periods. According to the IUPAC recommendations, the groups are numbered from 1 to 18, replacing the older notations of groups IA…VIIA, IB…VIIB and zero. The seven periods have 2, 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and will theoretically consist of 32 elements. The properties of the elements vary periodically due to periodic variation in their electronic configuration. In this chapter, we will study the variations in electronic configurations, atomic and ionic sizes, ionisation energies and other related properties. 8.2 Basic Concepts Before studying the periodicity in properties of elements, it is important to have knowledge of some basic terms and concepts. t
In the atoms, the orbitals are filled in the order of increasing energy (Aufbau principle), that is, an electron will first enter the orbital of lowest energy. The sequence of orbitals in the order of increasing energy is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s. ")!"
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&&#%""$      Periodic Properties 457 t
The other rules governing the filling of orbitals are t
Pauli’s exclusion principle: It was given by Austrian scientist Wolfgang Pauli (1925). It states that in an atom, no two electrons can have the same set of all four quantum numbers. If the principal, azimuthal and magnetic quantum numbers are the same, the electrons will have different spins. t
Hund’s rule of maximum multiplicity: It states that electrons in the same subshell (p, d and f ) do not pair until all the orbitals of similar energy gain one electron each. t
Effective nuclear charge: The hydrogen atom has only one electron, the energy level of which is determined by its principal quantum number n. All the orbitals inside this shell have the same energy and the only force experienced by the electron is the attractive force towards the positively charged nucleus. The situation is, however, different in case of multielectron atoms. The electrons in such atoms experience not only attractive force of the nucleus but also repulsion among themselves and each of their neighbours. The repulsion of outer-shell electrons by the inner-shell electrons is particularly important, as the outer shell electrons are pushed away from the nucleus due to repulsion of the inner-shell electrons. As a result, the nuclear charge felt by the outer-shell electrons is less than what it would have been if there were no inner electrons. This net charge felt by the valence electrons is termed as the effective nuclear charge, represented by the symbol (Zeff), and the repulsion experienced due to the inner electrons is the shielding effect. The shielding effect is called so because the inner electrons shield/screen the outer electrons from experiencing the actual nuclear charge (Zeff = Zactual – electron shielding). t
Consider the example of fluorine (atomic number = 9); there are two electrons in shell 1 (n=1) and seven electrons in shell 2 (n = 2). The seven electrons in n = 2 will be repelled by two electrons in n = 1 hence they will experience less attraction than what they would have experienced in the absence of these two electrons. Hence, these two inner-shell electrons are shielding or screening the outer electrons from the attractive force of the nucleus, and because of this shielding, the effective attractive force felt by the valence electrons reduces. Shielding effect generally remains the same in the period, as in the period the number of shells is the same and it increases down the group. t
Penetration effect: It is the tendency of electrons in an orbital to be attracted towards the nucleus. In a multielectron system, penetration is the electron’s relative electron density (probability density) near the nucleus and can also be considered as the presence of electrons inside the shell of other electrons. This can be explained by considering the shapes of orbitals and their radial distribution curves (Fig 8.1). The 1s orbital is spherically symmetrical with no node, hence we can see from the curve (Fig 8.1) that it has maximum probability density near the nucleus. On comparing 2s and 2p orbitals, we see that 2s is spherically symmetrical having a sphere within a sphere. It has two regions of high probability separated by a surface of zero probability called the node, and 2p orbital is dumb-bell shaped with electron distribution in identical lobes on either side of the nucleus and separated by a node cutting through the nucleus. This can be seen in Fig 8.1, which shows that 2s penetrates inside 1s (see the small loop of 2s inside 1s orbital in Fig 8.1). The 2p does not penetrate effectively because its wave function goes to zero at the nucleus. Thus 2s is more tightly bound to the nucleus and has lower energy than the 2p orbital. Thus the penetration of 2s is greater than 2p and the electrons in 2s experience large value of Zeff than the electrons in the 2p orbital. For subshells ")!"
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&&#%""$      458 Engineering Chemistry: Fundamentals and Applications within the same shell (n), the penetrating power of electron follow the sequence s > p > d > f. When the value of shell (n) and subshell (l) are different the penetrating power of electrons follows the trend 1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f Figure 8.1 Radial distribution function of orbitals Greater the penetration power of an electron, closer it will be to the nucleus, it will experience greater effective nuclear charge and it will be held firmly. Hence the ionisation enthalpy would be high. Table 8.1 gives the effective nuclear charge of the orbitals of a few atoms. Table 8.1 Effective nuclear charge Zeff of the orbitals of few atoms H He Z 1 2 1s 1.00 1.69 Li Be B C N O F Ne Z 3 4 5 6 7 8 9 10 1s 2.69 3.68 4.68 5.67 6.66 7.66 8.65 9.64 2s 1.28 1.91 2.58 3.22 3.85 4.49 5.13 5.76 2p 2.42 3.14 3.83 4.45 5.10 5.76 Na Mg Al Si P S Cl Ar Z 11 12 13 14 15 16 17 18 1s 10.63 11.61 12.59 13.57 14.56 15.54 16.52 17.51 2s 6.57 7.39 8.21 9.02 9.82 10.63 11.43 12.23 2p 6.80 7.83 8.96 9.94 10.96 11.98 12.99 14.01 3s 2.51 3.31 4.12 4.90 5.64 6.37 7.07 7.76 3p 4.07 4.29 4.89 5.48 6.12 6.76 ")!"
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&&#%""$      Periodic Properties 459 8.3 Periodicity in Properties Electronic configuration in period: Each period corresponds to the value of n or principal energy level. First period (n = 1) There are only two elements in this period, hence the electronic configuration of the first element hydrogen is 1s1 and that of helium is 1s2. With this, the first shell or K shell is completely filled. Second period (n = 2) It starts with lithium 1s22s1, beryllium with four electrons has the configuration 1s22s2. There are eight elements in this period. The configuration of the last element neon is 1s22s22p6. The second shell is completely filled. Third period (n = 3) It starts with Na (1s22s22p63s1). This period also has eight elements. The 3s and 3p orbitals are completely filled; however, the 3d orbitals remain vacant as their energy is higher than the energy of 4s orbitals. The configuration of the last element [Ar] of this period is 1s22s22p63s23p6. Fourth period (n = 4) It starts with potassium, where the electron is added to the 4s orbital. The electronic configurations of K and Ca are [Ar]4s1 and [Ar]4s2, respectively. After filling of 4s orbital, the filling of 3d orbital becomes energetically favourable. It starts at Sc (Z = 21), which has the configuration [Ar]3d14s2 and ends at Zn(Z = 30) with electronic configuration [Ar]3d104s2. There are certain irregularities in filling of 3d orbitals, for example, Cr (Z = 24) and Cu (Z = 29) have the electronic configuration [Ar]3d54s1 and [Ar]3d104s1, respectively, against the expected configuration of [Ar]3d44s2 and [Ar]3d94s2, respectively. These ten elements from Sc to Zn constitute the first transition series. Filling of 4p orbital starts at Ga (Z = 31) and is completed at Kr (Z = 36). Fifth period (n = 5) It begins with rubidium (Z = 37) where the next electron enters the 5s orbital and this orbital is complete at Sr (Z = 38). After this, ten electrons fill the 4d shell Y (Z = 39) to Cd (Z = 48), where the valence shell configurations are 4d15s2 and 4d105s2, respectively. This is termed as the second transition series. However, molybdenum has the configuration 4d55s1 and Ag has the configuration 4d105s1. The filling of 5p orbitals begins at In (Z = 49) and is completed at Xe (Z = 54). Sixth period (n=6) It contains 32 elements. The orbitals are filled in the order 6s, 4f, 5d and 6p. The filling of 4f orbitals begins with cerium (Z = 58) and ends with lutetium (Z = 71) to give the 4f -inner transition series termed as the lanthanoid series. Seventh period (n = 7) The filling of orbitals is similar to the sixth period with successive filling of 7s, 5f, 6d and 7p orbitals. The filling of 5f orbitals starts with Th (Z = 90) and is completed at Lw (Z = 103). This series from Th to Lw is called the 5f-inner transition series or the actinoid series. The 4f and 5f inner transition series are placed separately below the periodic table to comply with the principle of placing elements with similar properties in a single column. Table 8.2 gives the IUPAC-recommended names for some groups and sections in the periodic table. ")!"
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&&#%""$      460 Engineering Chemistry: Fundamentals and Applications Table 8.2 IUPAC-recommended names of some groups and sections Group number IUPAC-recommended name 1 (except H) Alkali metals 2 Alkaline earth metals 15 Pnictogens 16 Chalcogens 17 Halogens 18 Noble gases 1(except H), 2, 13, 14, 15, 16, 17, 18 Main group elements 8.4 Types of Elements on the Basis of Their Electronic Configuration Depending on the orbital in which the last electron enters, elements are classified in to four blocks, viz s-block, p-block, d-block and f-block elements. 1. s-block elements: They are elements of groups 1 (alkali metals) and 2 (alkaline earth metals) in which the valence electrons enter the s-orbitals. Their outer electronic configuration is ns1 and ns2, respectively. They are soft metals with low melting and boiling points, reactive as they have low ionisation enthalpies and strong reducing agents as they lose electrons readily to acquire stable noble gas configuration. Their common oxidation states are +1 and +2, respectively. Their reactivity increases down the group and most of them form ionic compounds, with the exception of lithium and beryllium. 2. p-block elements: They are elements of groups 13 to 18, in which the valence electron enters the p-orbitals. Their outermost electronic configuration varies from ns2np1 to ns2np6. The p-block elements together with s-block elements are termed as the representative elements or the main group elements. The p-block elements include both metals and nonmetals. The nonmetallic character increases as we move from left to right across the period and the metallic character increases as we move from top to bottom in the group. Their ionisation enthalpies are higher than that of s block elements. The elements of group 18 (noble gases) with electronic configuration ns2np6 are very stable because of completely filled valence electrons. They have very low chemical reactivity. Elements of group 17 (halogens) with electronic configuration ns2np5 and group 16 (chalcogens) with electronic configuration ns2np4 have highly negative electron gain enthalpies (earlier termed as electron affinity) and they gain one and two electrons, respectively, to acquire the stable noble gas configuration. Some of these elements exhibit more than one oxidation state. 3. d-block elements: They are elements of group 3 to 12, in which the last electron enters the d-orbitals of their second last (penultimate) energy level. Their outer electronic configuration is (n-1) d1–10ns0–2 These elements are mostly metals, have high melting and boiling points, generally form coloured complexes, exhibit variable valency (oxidation states) and exhibit para magnetism, and many elements like Mn, Ni, Co, Cr, V and Pt and their compounds are used as catalysts. These are termed as transition elements as they act as a bridge between highly reactive elements of groups 1 and 2 and less reactive elements of groups 13 and 14. The d block consists of the following: ")!"
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&&#%""$      Periodic Properties 461 First transition series (4th period where 3d orbital is filled), Sc (Z=21) to Zn (Z=30) Second transition series (5th period where 4d orbitals is gradually filled) Yttrium (Z=39) to Cd (Z=48) Third transition series (6th period where 5d orbitals are gradually filled) Lanthanum (Z=57), Hafnium (Z=72) and mercury (Z=80) Fourth transition series (7th period where 6d orbitals are gradually filled) – It is incomplete series. 4. f-block elements: f-block elements comprises of two rows of elements placed at the bottom of the periodic table. In these elements, the last electron enters the f-orbital. However, the last electron is added to the third to the outermost (antepenultimate) energy level; (n-2)f, hence these elements are termed as inner transition element. These elements are characterised by the general electronic configuration (n-2) f 1–14(n – 1)d 0or1ns2 (a) The first series called lanthanoid series extends from from Ce(Z=58) to Lu(Z=71). They are also termed as rare earth elements and have the general electronic configuration is 4f 1–145d 0or1 6s2. (b) The second series of elements from Th(Z=90) – Lr(Z=103) is called the actinoid series; their general electronic configuration is 5 f 1–14(n – 1)d 0or1 7s2 All the f block elements are metals having high melting and boiling points. They exhibit variable oxidation state and form coloured ions. These elements also have the tendency to form complex compounds. Actinoid elements exhibit radioactivity. They have been prepared in nanogram quantities by nuclear reactions and their chemistry is not fully studied. The elements after uranium are called trans-uranium elements. Note: IUPAC recommends the use of the terms lanthanoids and actinoids over lanthanides and actinide, as ‘ide’ usually implies negatively charged ions. However, the terms lanthanides and actinides are still commonly used. 8.5 Periodic Trends in Properties of Elements The physical and chemical properties of elements vary periodically. This has been attributed to similarity in the outer electronic configuration after regular intervals. Let us discuss the periodic trends of some important properties: Atomic and Ionic Radii Atomic radius is the distance from the centre of the nucleus to the outermost shell of the electrons. Since it is difficult to obtain an isolated atom, and the electron cloud around an atom does not have a sharp boundary, it is not possible to obtain the atomic radii by direct measurement for an individual atom. For a nonmetal, atomic radius is considered half the distance between the nuclei of two similar atoms bonded together by a single covalent bond (covalent radius) and for a metal it is half the distance between centres of nearest neighbours in a metallic crystal (metallic radius). Both metallic and covalent radii are jointly referred to as the atomic radii.(Table 8.3). For example the inter-nuclear distance between Cl–Cl is 198 pm, hence its atomic radius is taken as 198/2 = ")!"
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&&#%""$      462 Engineering Chemistry: Fundamentals and Applications 99 pm and the distance between two adjacent copper atoms in solid copper is 256 pm, hence its metallic radius would be 256/2 = 128 pm Inter-nuclear distance between two similar bonded atoms rcovalent = 2 Noble gases do not form covalent bonds, hence their atomic radius is found out by calculating their van der Waals radius, which is defined as half the distance between the nuclei of two non- bonded neighbouring atoms of two adjacent molecules. This can be explained by considering the example of two chlorine molecules in close contact (Fig 8.2) Figure 8.2 Depiction of van der Waals radii Atomic radius can be found by X-ray diffraction, electron diffraction and other spectroscopic techniques. Trends in atomic and ionic radii: Table 8.3 and Figure 8.3 show the atomic radii of some elements in the period and group, respectively. It is observed that the atomic radii increases down the group and decreases from left to right across the period within s and p block. The atomic radius increases down the group because as one moves down the group the valence electrons occupy the orbitals of successively higher principal quantum number. The valence electrons in Li, Na, K, Rb, Cs, Fr occupy second (2s1), third (3s1), fourth (4s1), fifth (5s1), sixth (6s1), and seventh (7s1), shells, respectively. Since larger shells are occupied, hence the atomic size increases down the group. The atomic radius decreases across the period, because the valence electrons in a period occupy the same shell but the nuclear charge increases progressively thereby increasing the effective nuclear charge on the valence electrons drawing them closer to the nucleus and decreasing the atomic radii (Fig 8.3). However, it is observed that the atomic radius of second and third rows of d-block elements are almost similar. This is due to lanthanide contraction. The elements of the third row of d block elements (sixth period) are preceeded by 14 lanthanoids where 4f orbitals are occupied. Since these orbitals have poor shielding properties, the valence electrons experience greater attraction towards the nucleus. The repulsions between electrons fails to compensate for the increasing nuclear charge, so Zeff increases from left to right, resulting in the decrease in atomic radii, and the sizes of elements of second and third rows of d-block elements remain almost the same. ")!"
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&&#%""$      Periodic Properties 463 Figure 8.3 Trends in atomic radius and atomic number Table 8.3 Atomic radii of elements in pm Increase in nuclear charge: It results in greater attraction of the orbital electrons towards the nucleus thereby decreasing the atomic radius. As noticed in the Table 8.3, the radius of Ne and Ar is large. This is because it represents the van der Waals radius. The atomic radius increase down the group because while descending in a group two things happen: (i) increase in nuclear charge and (ii) addition of new shells. The second factor outweighs the first, leading to increase in atomic radius from top to bottom. Variation in ionic radii: It has been observed that the radius of cation (positive ion) is always smaller than that of the parent atom and the radius of anion (negative ion) is always larger than that of the parent atom. ")!"
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&&#%""$      464 Engineering Chemistry: Fundamentals and Applications The size of cation decreases not only due to the increase of effective nuclear charge Zeff due to loss of valence electrons thereby decreasing electron–electron repulsions but also due to the fact that generally formation of cation leads to the complete loss of the valence shell. For example the configuration of Li is 1s2 2s1 and that of Li ion is 1s2. Here the 2s shell is completely lost. Li → Li+ + e- (1s22s1) (1s2) 123 pm 60 pm Similarly, formation of Mg2+ ion (formed by the complete loss of 3s shell) is represented as follows: Mg → Mg2+ + 2e- (1s22s22p63s2) (1s22s22p6) 136 pm 65 pm The radius of cation increases down the group Size of anion increases because addition of electron increases inter-electronic repulsion in the orbital electrons, and hence the electron cloud expands leading to the increase in size. This is also associated with the decrease in Zeff. For example, Cl + e- → Cl- (1s 2s22p5) 2 (1s22s22p6) 99 pm 181 pm Figure 8.4 (a) and (b) Comparison of atomic and ionic radii of cations (group 1) and anions (group 17) ")!"
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&&#%""$      Periodic Properties 465 Trends in ionic radii: Like atomic radii, the ionic radii also increases down the group and decreases across the period. However, there are some exceptions to this trend. 8.6 Ionisation Enthalpy First ionisation energy is defined as the energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state at 0 K. A(g) → A+(g) + e–(g) IE1 = E(A+, g) – E(A, g) The second ionisation energy is the energy required to remove the most loosely bound electron from the resulting cation. A+(g) → A2+(g) + e– (g) IE2 = E(A2+, g) – E(A+, g) Energy required to convert A2+(g) in to A3+(g) is the third ionisation energy and so on. Ionisation energies are expressed in kJ mol–1 or electronvolts (eV), where 1 eV = 94.485 kJ mol–1. However, most chemical reactions occur at constant pressure and at a temperature of 298 K. In such a situation it is more appropriate to use the term ionisation enthalpy. If the gaseous atoms are 5 assumed to be ideal gases then the ionisation enthalpy is larger than ionisation energy by R DT , 2 where R is the molar gas constant (8.314 × 10–3 kJ K–1 mol–1 and DT is the difference between the two temperatures (0 K and 298 K). The ionisation energies for the removal electrons are of the order of 103 kJ mol–1 and the ionisation enthalpy is 5 DH (298K ) = DH (0 K ) + ¥ 8.314 ¥ 10-3 ¥ 298 2 DH (298K ) = DH (0 K ) + 6.2 kJ mol -1 The difference between ionisation energy and enthalpy is very small (≈ 6 kJ mol–1) and is often ignored. However, the term ionisation enthalpy is commonly used these days. The ionisation energy depends upon the size of the atom, charge on the nucleus, type of electrons involved (s, p, d, f) on the screening effect of the inner electrons and the penetration effect of electrons. The ionisation enthalpies of elements with Z=1 to 60 are shown in Fig 8.5. From the figure it is clear that alkali metals have low ionisation enthalpies and noble gases have very high ionisation enthalpies. This can be attributed to the fact the alkali metals have only one electron in their valence shell. All the inner shell electrons called core electrons shield this electron from the nucleus thereby reducing Zeff, hence they lose electrons easily and acquire a stable noble gas configuration. Noble gases because of stable electronic configuration do not lose electrons readily, which is reflected in their very high values of ionisation enthalpies. ")!"
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&&#%""$      466 Engineering Chemistry: Fundamentals and Applications Figure 8.5 Variation of first ionisation enthalpies (DiH) with atomic number for elements with Z = 1 to 60 Variation in a group The ionisation enthalpy decreases as we descend down the group. On moving down the group, the following factors become effective: t
Nuclear charge increases. t
New shells are added. t
The shielding of outermost electrons increases due to increase in the number of inner electrons. The first factor is outweighed by the second and third factors, as a result of which the ionisation enthalpy decreases on descending down the group. Fig 8.6(b) shows the variation of ionisation enthalpy of the elements of group 1. Table 8.4 gives the first and second ionisation enthalpies of elements. Variation in a period In a period the ionisation enthalpy generally increases from left to right. This is because nuclear charge increases in a period and the successive electrons are added in the orbitals with same principal quantum level resulting in the increase in Zeff. Consequently the electrons are more tightly held by the nucleus, resulting in an increase in ionisation enthalpy. Figure 8.6(a) shows the variation of ionisation enthalpy of the elements of the second period. However, there are some deviations in this general trend. Although the nuclear charge of B is greater than Be, the first ionisation enthalpy of B (Z=5) is slightly less than Be (Z=4). This is because of the fact that in beryllium the electron is removed from s orbital, whereas in boron the electron is removed from p orbital. Since the penetration of 2s electron to the nucleus is more than that of 2p electron, it is easier to remove electron from 2p orbital, and hence the ionisation enthalpy of boron is less than that of Be. Similarly O (Z=8) has lesser ionisation enthalpy than N (Z=7). This is because nitrogen has half-filled 2p orbitals which reside in different atomic orbitals whereas in oxygen, two of the four 2p orbitals reside in the same orbital thereby increasing the electron–electron repulsion. Hence it is easier to remove the 2p electron of oxygen. Moreover, the half-filled p orbitals of nitrogen is a stable configuration. ")!"
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&&#%""$      Periodic Properties 467 Table 8.4 First and second ionisation enthalpies of elements in kJ mol–1 H He 1312 2373 5259 Li Be B C N O F Ne 520 899 801 1086 1402 1314 1681 2080 7297 1757 2426 2352 2855 3386 3375 3952 Na Mg Al Si P S Cl Ar 495 737 577 786 1011 1000 1251 1520 4562 1476 1816 1577 1903 2251 2296 2665 K Ca Ga Ge As Se Br Kr 419 589 579 762 947 941 1139 1351 3051 1145 1979 1537 1798 2044 2103 3314 Rb Sr In Sn Sb Te I Xe 403 549 558 708 834 869 1008 1170 2632 1064 1821 1412 1794 1795 1846 2045 Cs Ba Tl Pb Bi Po At Rn 374 502 590 716 704 812 926 1036 2420 965 1971 1450 1610 1800 1600 Figure 8.6 (a) First ionisation enthalpies (DiH) of elements of the second period as a function of atomic number (Z) and (b) DiH of alkali metals as a function of Z. 8.7 Electron Gain Enthalpy Electron gain enthalpy is defined as the energy change when an electron is added to an isolated gaseous atom in ground state to give an anion represented by the equation ")!"
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&&#%""$      468 Engineering Chemistry: Fundamentals and Applications A(g) + e–(g) → A–(g) DEG = E(A–, g) – E(A, g) DEG is the electron gain enthalpy. If the energy is released during the addition of an electron, then the electron gain enthalpy is considered negative and if energy is supplied to add electron to the atom, then the electron gain enthalpy is considered as positive. The electron gain enthalpy is measured in kJ mol–1 or electronvolts (eV). Table 8.5 gives the electron gain enthalpies for few atoms. Electron affinity Ea on the other hand is the difference in energy between the gaseous atoms and the gaseous ions at T = 0 Ea = E(A, g) – E(A–, g) Hence the sign convention of electron gain enthalpy and electron affinity are opposite DEG = –Ea Note that the sign convention for electron affinity is the reverse of the convention used in thermodynamics. A positive value of electron affinity indicates that the ion A– has a lower energy than the atom A. Although the term electron affinity is widely used in inorganic chemistry, the terms electron gain enthalpy and electron affinity cannot be used interchangeably. In the present text, we shall use the term electron gain enthalpy. Table 8.5 Electron gain enthalpies (kJ mol–1) of some elements H He –72 +48 Li Be B C N O F Ne –60 ≈0 –27 –122 ≈0 –141 –328 +116 +780 Na Mg Al Si P S Cl Ar –53 ≈0 –43 –134 –72 –200 –349 +96 +492 K Ca Ga Ge As Se Br Kr –48 ≈0 –29 –116 –78 –195 –325 +96 Rb Sr In Sn Sb Te I Xe –47 ≈0 –29 –116 –103 –190 –295 +77 *The second value given for sulphur and oxygen is for the addition of second electron to the anion From Table 8.5, it is clear that the electron gain enthalpy increases in a period and decreases down the group. This is because as one moves from left to right in the period the effective nuclear charge increases hence an electron is easily added. Elements of group 17 add electron easily to acquire a stable noble gas configuration; hence they have high negative electron gain enthalpy. Electron gain enthalpy decreases down the group, because on moving down, the group the atomic radius increases and the added electron feels lesser attraction by the nucleus, that is, Zeff decreases. ")!"
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