Periodic Properties PDF

# PERIODIC PROPERTIES ## Syllabus (4 Lecturers) - Effective nuclear charge - Penetration of orbitals, variations of s, p, d, and f orbital energies of atoms in the periodic table - Electronic configurations - Atomic and ionic sizes - Ionization energies - Electron affinity and electronegativity -...

# PERIODIC PROPERTIES ## Syllabus (4 Lecturers) - Effective nuclear charge - Penetration of orbitals, variations of s, p, d, and f orbital energies of atoms in the periodic table - Electronic configurations - Atomic and ionic sizes - Ionization energies - Electron affinity and electronegativity - Polarizability - Oxidation states - Coordination numbers and geometries - Hard soft acids and bases - Molecular geometries ## 7.1 Introduction All the 112 known elements, either natural or synthetic, have been classified in the long form of a periodic table in accordance with their properties in such a way that elements with similar properties are grouped together in the same vertical column called group and dissimilar elements are separated from one another. According to modern periodic law the physical and chemical properties of the elements are a periodic function of their atomic numbers. It means that if the elements are arranged in increasing order of their atomic numbers, the elements with similar properties are repeated after certain regular intervals. The cause of periodicity in properties is the repetition of similar outer electronic configuration at certain regular intervals. Most of the physical and chemical properties of the elements change periodically with the atomic number. Some of these properties such as effective nuclear charge, atomic size, ionization energy and electron affinity are directly related to the electronic configurations of the atoms. On the other hand, properties such as melting point and density are only indirectly related to electronic configuration. The properties directly or indirectly related to their electronic configuration and show a regular gradation when we move from left to right in a period or from top to bottom in a group are called periodic properties. In this chapter, we discuss some of the periodic properties in accordance with the revised syllabus of AICTE adopted by I.K. Gujral, Punjab Technical University. ## 7.2 Effective Nuclear Charge In multi-electron atoms the electrons in the valence shell experience an attractive force from the nucleus and a repulsive force from the electrons in the inner shells. The overall effect of these two opposing forces is that the attractive force exerted by the nucleus on the valence shell electrons is somewhat reduced by the repulsive force exerted by the electrons present in the inner shells. In other words, the valence shell electrons do not feel the full charge of the nucleus. The actual charge felt by the valence electrons is called effective nuclear charge and the repulsive force felt by the valence shell electrons from the electrons present in the inner cells is called shielding effect or screening effect. Therefore, the effective nuclear charge $(Z_{eff})$ is given by the relation $Z_{eff} = Total nuclear charge (Z) - screening constant (S)$ where screening constant (S) takes into account the screening effect of the electrons present in the inner shells. Obviously, greater the number of electrons in inner shells, larger will be the screening effect. As a screening effect increases, the effective nuclear charge decreases. Consequently, the force of attraction by the nucleus for the valence electrons decreases and hence, the ionization energy decreases. ## Shielding and Effective Nuclear Charge An atom with its atomic number greater than 2 has core electrons that are extremely attracted to the nucleus in the center of the atom. However, the number of protons in the nucleus are never equal to the number of core electrons (relatively) adjacent to the nucleus. The number of protons increase by one across the periodic table, but the number of core electrons change by periods. The first period has no core electrons, the second has 2, the third has 10, and so on. This number is not equal to the number of protons. So that means that the core electrons feel a stronger pull towards the nucleus than any other electron within the system. The valence electrons are farther out from the nucleus, so they experience a smaller force of attraction. Shielding refers to the core electrons repelling the outer rings and thus lowering the 1:1 ratio. Hence, the nucleus has "less grip" on the outer electrons and are shielded from them. Electrons that have greater penetration can get closer to the nucleus and effectively block out the charge from electrons that have less proximity. For example, $Z_{eff}$ is calculated by subtracting the magnitude of shielding from the total nuclear charge. The value of $Z_{eff}$ will provide information on how much of a charge an electron actually experiences. Because the order of electron penetration from greatest to least is s, p, d, f; the order of the amount of shielding done is also in the order s, p, d, f. Since the 2s electron has more density near the nucleus of an atom than a 2p electron, it is said to shield the 2p electron from the full effective charge of the nucleus. Therefore, the 2p electron feels a lesser effect of the positively charged nucleus of the atom due to the shielding ability of the electrons closer to the nucleus than itself, (i.e. 2s electron). ## 7.3 Penetration of Orbitals The ability of an electron to get close to the nucleus is called penetration of the electron. The electrons in different orbitals have different penetration effect and even electrons present in the same shell differ in the penetration effect of their subshell i.e. electrons differ in the penetration effect of the orbital to which they belong. In simple words, the penetration effect of electrons of s, p, d and f orbitals of any one main shell is different. The electrons are negatively charged and are pulled towards the positively charged nucleus. The electrons are thus attracted to the nucleus, but at the same time electrons repel each other. These attractive and repulsive forces result in shielding. Penetration and shielding are two underlying principles in determining the physical and chemical properties of elements. We can predict basic properties of elements such as ionization energy and electron affinity by using shielding and penetration characteristics. Most of the physical and chemical properties of the elements depends upon these two periodic properties namely ionization energy and electron affinity. Electrons are negatively charged and are quite close to each other, which means that they can repel each other. The repulsion an electron feels is shielding and attraction it feels to the nucleus is penetration. The force that an electron feels is dependent on the distance from the nearest charge (i.e., an electron, usually with bigger atoms and on the outer shells) and the amount of charge. More distance between the charges will result in less force, and more charge will have more force of attraction or repulsion. This is in accordance with the Coulomb's law of classical mechanics. In an ideal system for atoms, every electron should feel the same amount of "attraction" from the nucleus. This means that the negative to positive charge ratio should be 1:1. However, that is not the case as observed in practice from the behaviour of atoms at least. While considering the core electrons i.e. the electrons closest to the nucleus, the ratio is 1:1, or at least close to it. When we proceed from the core electrons to the outer valence electrons, the negative to positive charge ratio falls below 1:1. This is because of shielding, or simply the electrons repelling each other. But the same core electrons penetrate and feel more of the nucleus than the other electrons, because they are attracted more by the nucleus as compared to repulsion from other electrons with the same atom. In a multi-electron system, the penetration of the nucleus by an electron is measured by the relative electron density near the nucleus of an atom for each shell and subshell of an electron. For example, we see that since a 2s electron has more electron density near the nucleus than a 2p electron, it is penetrating the nucleus of the atom more than the 2p electron. The penetration power of an electron, in a multi-electron atom, is dependent on the values of both the shell and subshell of an electron in an atom. Therefore, for the same shell value (n) the penetrating power of an electron in orbitals (subshells) is given as under: s > p > d > f And for different values of shell (n) and subshell (l), penetrating power of an electron follows this trend: 1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f and the energy of an electron for each shell and subshell follows the order ls < 2s < 2p <3s <3p < 4s <3d < 4p ## 7.3.1 Periodic Trends Due to Penetration and Shielding - **Effective Nuclear Charge ($Z_{eff}$):** The effective nuclear charge increases from left to right and increases from top to bottom in the periodic table. - **Atomic Radius:** The atomic radius decreases from left to right in a period, and increases from top to bottom in a group. - **Ionization Energies:** The ionization energies increase from left to right, and decrease from top to bottom. - **Electronegativity:** The electronegativity of the elements is highest near fluorine. In general, it increases from left to right in a period and decreases from top to bottom in a group. ## 7.4 Division of Elements into s, p, d and f Blocks The elements, as arranged in the long form of the periodic table, can also be divided into four main blocks known as s, p, d and f-blocks. The classification depends upon the type of the orbitals (s, p, d or f) into which the last electron of the atoms enters. - **s-Block elements:** The elements in which the last electron enters the s-orbital of their outermost energy level are called s-block elements. It consists of elements of groups 1 and 2 having the ground state electronic configuration of outermost shell as $ns^1$ and $ns^2$ respectively (where n stands for outermost energy shell). The elements corresponding to electronic configuration $ns^1$ are called alkali metals while those with $ns^2$ electronic configuration are called alkaline earth metals. Thus, the general electronic configuration of s-block elements may be expressed as: $ns^{1-2}$ The s-block contains 13 elements, except the first period, each other period of the block has 2 elements. - **p-block elements:** The elements whose atoms receive the last electron in their p-orbitals are called p-block elements. Therefore, elements of groups 13, 14, 15, 16, 17 and 18 constitute p-block. The general electronic configuration of p-block element is written as: $ns^2 np^{1-10}$ The p-block contains 31 elements with the exception of first period, each other period of the block has got 6 elements. - **d-block elements:** d-block elements are those elements in which the last electron enters the d-orbital of the last but one energy level (Penultimate shell). As a result, they have 1 or 2 electrons (zero in some cases) in their ns orbital while the last electron enters the (n-1)d-orbital. Thus, their outermost energy level remains incomplete while the (n-1)d-orbitals are progressively filled as we move from one element to the next. So, the elements lying between s and p-blocks are jointly known as d-block elements. Therefore, elements of group 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 constitute d-block. These elements are called transition elements. The general electronic configuration of d-block elements may be written as: $(n-1)d^{1-10} ns^{1 or 2}$ The d-block contains 33 elements. While the incomplete seventh period has got only 3 elements, each of the remaining period (i.e. 4th, 5th and 6th) of the block has got 10 elements. - **f-block elements:** The elements in which the last election enters the f-orbitals of their atoms are called f-block elements. In these elements, the last electron is added to the third to outermost (called antepenultimate) energy level; (n-2) f. These consist of two series of elements viz. Lanthanides [elements following La (Z = 57)] and Actinides [elements following Ac (Z = 89)]. The general electronic configuration of f block elements may be written as: $(n-2) f^{1-14} (n - 1) d^{0-1} ns^2$ The elements belonging to f-block are called Inner transition elements. The f-block contains 28 elements in the form of 2 rows of 14 elements each. ## 7.4.1 Difference between s-block and d-block elements It may be noted that both s-block and d-block have either 1 or 2 electrons in s-orbital of their outermost energy level. However, s-block elements do not have any electron in the d-orbital of last but one energy level while d-block elements have always got 1 to 10 electrons in the d-orbitals of last but one energy level. Thus $(n-1)d$ electrons serve as differentiating electrons between s-block and d-block elements. ## 7.5 VARIATION OF s, p, d AND f ORBITAL ENERGIES There are many physical and chemical properties of elements which show periodic variation with atomic number. Some of these properties such as valency, atomic size, ionization energy and electron affinity are directly related to the electronic configuration of the atoms. On the other hand, the other properties such as melting point, density, etc. are only indirectly related to electronic configurations. Periodic properties are repeated after a regular interval in a regular manner. The variation of some of the properties are illustrated below: - Valency - Atomic Size - Ionization Energy - Electron Affinity - Electronegativity ## 7.5.1 Valency The valency of an element is directly related to the electronic configuration of its atoms and is usually determined by the number of electrons in the outermost shell. This is because the outer electrons are the ones largely responsible for the chemical behavior, as the electrons in these orbitals usually participate in chemical bonding. The outer shell electrons are often referred to as valence shell electrons or simply valence electrons. **Representative elements.** In case of representative elements, the valency is generally equal to either the number of valence electrons or 18 minus the number of valence electrons. At the same time, the number of valence electrons for the representative elements of any group is equal to the number of the group. Therefore, the valence of representative elements is either equal to group number of the element or 18 minus group number. **Variation of valency in the periodic table** - **Variation in a period.** On moving from left to right across a period, the number of valence electrons increases from 1 to 8. Consequently, the usual valency of the elements increases from 1 to 4 and then decreases to 1 as shown in Table 7.2, for the elements of second period. It may be pointed out that there are some exceptions to this rule but we shall not consider them at this stage. It should also be remembered that elements of group 18 are chemically inert. - **Variation in a group.** On moving down a group the number of valence electrons remains the same Therefore, all the elements in a group have the same valency. For example, all the elements of group 1 have valency 1 and of group 2 have valency 2. **Transition elements.** The transition element have generally 1 to 2 electrons in their outermost energy level. They exhibit variable valency but their most common valency is 2. ## 7.5.2 Atomic and Ionic Sizes Atomic radius may be defined in different ways as follows: - If the atom is considered to be a sphere, the atomic radius is equal to the radius of the sphere. - Atomic radius is defined as the distance from the center of the, nucleus to the outermost shell containing the electrons. - Atomic radius may be defined as the distance from the center of the nucleus to the point upto which the density of the electron cloud (i.e. probability of finding the electron) is maximum. ## Difficulties in determining atomic radius - The exact size of the electron cloud cannot be determined as the probability of finding an electron even at large distances from the nucleus never becomes zero. - It is not possible to isolate a single atom for the purpose of determination of its radius. - The probability distribution of electrons of an atom is appreciably affected by the presence of other atoms within its molecule or in its neighbourhood. Thus, the atomic radius may change in going from one environment to the other. - The atomic radius also changes when the atom is present in different bonded states. ## Types of atomic radius Despite the above limitations, we need some operational definition for the radius of an atom or an ion. This is necessary in order to explain a number of chemical properties of different elements in terms of the size of their atoms or ions. There are three operational concepts which have been widely used. These are: - Covalent radii - Van der Waals' radii - Ionic radii. ## 1. Covalent Radii It is defined as one-half of the distance between the nuclei of two covalently bonded atoms of the same element in a molecule (see Fig. 7.3). Thus, for a homonuclear diatomic molecule, $r_{covalent} = \frac{1}{2}$ [Internuclear distance between two bonded atoms] Since the internuclear distance between two bonded atoms is called the bond length. Therefore, $r_{covalent} = \frac{1}{2}$ [bond length] For example, the internuclear distance between two hydrogen atoms in H₂ molecule is 74pm (0.74 A). Hence, the covalent radius of hydrogen is 74/2 = 37pm (0.37 Å). Similarly, the internuclear distance between chlorine atoms in Cl2 molecule is 198 pm (1.98 Å), the covalent radius of chlorine is 198/2 = 99pm (0.99Å). ## 2. Van der Weals' Radii It is defined as one-half of the distance between the nuclei of two non-bonded isolated atoms or two adjacent atoms belonging to two neighbouring molecules of an element in the solid state (Fig. 7.3). The name Van der Waals' radius simply means that the forces existing between the non-bonded isolated atoms or neighbouring molecules are the weak Van der Waals' forces of attraction. Consequently, the magnitude of the van der Waals' radii depends upon the packing of the atoms when the element is in the solid state. These are obtained from X-ray studies.of the various atoms in the solid state. For example, the internuclear distance between the adjacent chlorine atoms of the two neighbouring molecules in the solid state is 360 pm or 3.6 A (Fig. 7.3). Therefore, the Van der Waals' radius of chlorine atom = 360/2 = 180 pm or 1.8 A. Similarly, the internuclear distance between two adjacent hydrogen atoms of two neighbouring molecules in the solid state = 240 pm. Therefore, its Van der Waal's radius = 240/2 = 120 pm or 1.2 Å. ## Comparison of covalent radius and Van der Waals' radius. Van der Waals' radius of an element is always larger than its covalent radius because of the following two reasons: - Since the Van der Waals' forces of attraction are weak, therefore the internuclear distances in case of atoms held by Van der Waals' forces are much larger than those between covalently bonded atoms. Therefore, Van der Waals' radii are always larger than covalent radii. - Since a covalent bond is formed by overlap of two half-filled atomic orbitals, a part of the electron cloud becomes common (Fig. 7.4). Therefore, covalent radii are always smaller than the Van der Waals' radii. ## Periodic Trends in Atomic Radii - **Variation of atomic radii in a period:** In general, the atomic radii (covalent or Van der Waal's) decrease with increase in atomic number as we move from left to right in the periodic table. Along a period the electron is being added into the same principal shell and effective nuclear charge, $Z_{eff}$ increase, therefore, nuclear attraction on the valence electrons increases and the radii of atoms decrease. As we move from left to right in a period, nuclear charge increases by one unit in each succeeding element while the number of the shells remains the same. Due to this enhanced nuclear charge, the electrons of all the shells are pulled little closer to the nucleus thereby making each individual shell smaller and smaller. This results in a decrease of the atomic radius as we move from left to right in a period. Like covalent radii, Van der Waals' radii also decrease as we move from left to right in a period. For example, the Van der Waals' radii of N, O and F are 155 pm, 152 pm and 147 pm respectively. - **Variation of atomic radii in a group:** The atomic (covalent) radii of elements increase with increase in atomic number as we move from top to bottom in a group. As we move down in a group, a new enegy shell is added at each succedding element though the number of electrons in the valence shall remains to be the same. In other words, electrons in the valence shell of each succeeding element lie farther and farther away from the nucleus. As a result, the attraction of the nucleus for the electron decreases and hence the atomic number increases. However, with the increase in atomic number, the nuclear charge also increases and with its result the force of attraction of the nucleus for the electrons increases and hence the atomic radii should decrease. But the effect of the increased nuclear charge is reduced due to the screening or shielding effect on the valence eletrons by the electron present in the inner shells. Thus, the effect of adding a new energy shell is so large that it over weights the contractive effect of the increased nuclear charge. Hence, the increase in atomic radii as we move down the group, for example: from Li to Cs along alkali metals and from F to I among halogens, Is primilarily due to the addition of new energy shell. Like covalent radii, Van der Waals radii also increase as we move down the group. For example, the Van der Waals radii of Cl, Br and I are 175 pm, 185 pm and 198 pm respectively. ## 3. Ionic and Crystal Radii The ionic radii means the radii of ions in ionic crystals. The ions are formed as a result of addition or removal of electrons from the outermost shells of atoms. The ions formed by the loss of electrons acquire positive charge and are called cations while the ions formed by the gain of electrons acquire negative charge and are called anions. Ionic radius may be defined as the effective distance from the nucleus of the ion upto which it has an influence in the ionic bond. The internuclear distance between the nuclei of adjacent positive and negative ions in a crystal can be calculated by X-ray studies or spectroscopic studies. By knowing the radius of one ion, the radius of other ions can be calculated. **Size of the cation:** In general, the size of a cation is smaller than a neutral atom. This is due to the fact that with the removal of electrons from an atom, the magnitude of nuclear charge remains same while the number of electrons decreases. As a result, the nuclear charge now acts on lesser number of electrons. In other words, the effective nuclear charge per electron increases and the electrons are more strongly attracted and are pulled towards the nucleus. This causes a decrease in the size of the atom. For example, - Atomic radius of Na = 1.54 Å, - Ionic radius of $Na^+$ = 0.95 Å **Size of the anion:** In general; the negative ion or anion is always large than that of the corresponding atom. The negative ion is formed by the gain of one or more electrons in the neutral atom and the number of electrons increases while the magnitude of nuclear charge remains the same. As a result, the same nuclear charge acts on large number of electrons than are present in the neutral atom. In other words, effective nuclear charge per electron is reduced and the electron cloud is held less tightly by the nucleus. This causes an increase in size and thus, anions are larger in size than the corresponding atoms. For example, - Atomic radius of Cl = 0.99 Å, - Ionic radius of $Cl^-$ = 1.81 Å ## Periodic Trends in Ionic Radii With in a period, the ionic (cation) radii of elements of groups 1, 2, 13, 16 and 17 decrease with increase in atomic number and so do the ionic radii (anions) of groups 16 and 17. This is because the atomic size decreases along a period so do the ionic radii. - **Variation of ionic radii in a period.** For an isoelectronic series (ions having same number of electrons). The ionic radius decreases as the nuclear charge increase. In an isoelectronic series, there is an increase in nuclear charge and, therefore, the attraction for the same number of electrons increases. As a result, the electrons are pulled more and more strongly and thus ionic radius decreases. All ions have neon configuration containing 10 electrons. - **Variation of ionic radii in a group.** As we go down a group, the ionic radii increase with increase in atomic number like atomic radius. This is mainly due to increase in the principal quantum levels. The effect of increase in principal quantum level is more pronounced than the effect of increased nuclear charge. However, this decrease is not much amongest the last members. For example, the increase in ionic radius is quite rapid when we move from $Li^+$ to $Na^+$ and from $Na^+$ to $K^+$ but the increase is not so rapid as we move from $K^+$ to $Rb^+$ and from $Rb^+$ to $Cs^+$. This is explained on the basis that with in $K^+$ and $Rb^+$ comes the elements of the transition series and in these elements the last occupied shell is the same, their increasing nuclear charge with increase in atomic number tend to cause contraction of the atoms and ions. Thus, the ions which follow any of the transition series would be smaller than if only eight elements had separated them from the higher members of the family. ## 7.5.3 Ionization Energy or Ionization Potential The ionization energy gives a measure of the ease with which an atom can lose an electron and change into a cation. Ionization energy of an element is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom. $M (g) + Ionization Energy → M^+ (g) + e (g)$ Ionization energy is also called as ionization potential because it is measured as the amount of potential needed to remove the most loosely held electron from the gaseous atom. The experimental values of ionization energies are determined either by spectroscopic methods or by passing an electric current of gradually increasing intensity through the vapours of the element in a discharge tube. It is expressed in terms of either kcal/mol or kJ/mol or electron volt/atom. Thus, the ionization energy gives the ease with which electron can be removed from an atom. Evidently, the smaller the value of ionization energy, the easier it is to remove the electron from the atom. ## Factors on which ionization energy depends The magnitude of ionization energy for an atom depends upon the following factors: - **Size of the atom.** The ionization energy depends upon the distance between the electron and the nucleus, i.e. size of atom. As the size of the atom increases, the outermost electrons are less strongly attracted by the nucleus because the force of attraction is inversely proportional to the square of the distance between the charged particles. As a result, it becomes easier to remove the electron and, therefore, ionization energy would tend to decrease with the increase in the size of the atom. - **Charge on the nucleus.** The attractive force between the nucleus and the electrons increases with increase in nuclear charge. This is because, the force of attraction is directly proportional to the product of charges on the nucleus and that on the electrons. Therefore, with the increase in nuclear charge, it becomes more difficult to remove an electron and ionization energy increases. - **Screening effect of the inner electrons.** In multi-electron atoms, the outermost electrons are shielded or screened from the nucleus by the inner electrons. This is known as shielding or screening effect. As a result of this, the outermost electron does not feel the full charge of the nucleus. The actual charge felt by an electron is termed as effective nuclear charge. Effective nuclear charge ($Z_{eff}$) is $Z_{eff} = Total nuclear charge (Z) - screening constant (s)$ where screening constant takes into account the screening effect of the inner electrons. If the number of electrons-electron force will be less. Consequently, ionization energy will decrease. Thus, if other factors do not change, an increase in the number of electrons tends to decrease the ionization energy. - **Penetration effect of electrons.** It is well known that in case of multi-electron atoms, the electrons in the s-orbital have the maximum probability of being found near the nucleus and this probability goes on decreasing in case p, d and f orbitals. In other words, s-electrons are more penetrating towards the nucleus than p-electrons and the penetration power decreases in a given shell in the order : s > p > d > f Now, if the penetration of the electron is more, it will be closer to the nucleus and will be held firmly. Consequently, ionization energy will be high. This means that ionization energy increases with the increase in penetration power of the electrons. Thus, for the same shell, the required ionization energy would be more to remove an s-electron than to remove a p-electron, which in turn will be more than that for the removed of d-electron and so on. - **Electronic arrangement or configuration.** The ionization energy also depends upon the electronic configuration of the atom. It has been observed that ceratin electronic configurations are more stable than the others. For example, half filled and completely filled shells have extra stability associated with them. Consequently, it is difficult to remove electron from these stable configurations and ionization energy is high. This may be illustrated by the following examples: - The noble gases have most stable electronic configurations ($ns^2 np^6$) in each period and have highest ionization energies. - The elements like Be ($1s^2 2s^2$) and Mg ($1s^2 2s^2 2p^6 3s^2$) have completely filled orbitals and their ionization energies are large. - The elements like N ($1s^2 2s^2 2p^2 2p 2p^1$, 2p2 ) and P ($1s^2 2s^2 2p^6 3s^23p^1 3p 3p^1$, 3p2 ) have the configurations in which the orbitals of same sub-shell are exactly half filled and are also stable. Hence, they need large energy to remove the electron, i.e., their ionization energies are high. Thus, the more stable the electronic configuration, the greater is the ionization energy. ## Variation of lonization Energy in the Periodic Table Ionization energy provides another example for understanding periodicity among the elements. Fig. 7.5 shows the variation of first ionization energy of elements with atomic number upto atomic number 60. It can be seen that in each period, the maxima are found at the noble gases while minima are found at alkali metals. Thus, the metals of group I with one electron in outermost s-orbital are easy to ionize while the noble gases (group) with $ns^2 np^6$ configuration are the most difficult to ionize. The following periodic trends have been observed. - **Variation along a period.** In general, the ionization energy increases with increasing atomic number in a period. This is quite clear from the values of ionization energy of the second row elements. The general increase along a period can be explained on the basis of atomic size and nuclear charge. We know that - On moving across a period from left to right, the nuclear charge increases. - The atomic size decreases along a period though energy level remains the same. As a consequence of increased nuclear charge and simultaneous decrease in atomic size, the valence electrons are more and more tightly held by the nucleus. Therefore, more and more energy is needed to remove the electron and hence, ionization energy keeps on decreasing. However, some irregularities in the general trend have been noticed. These are due to half filled and completely filled configurations with extra stability. To illustrate this, let us consider the variation of ionization energy in second period (Fig. 7.6) going from one element to another. - **Variation down a group.** Within a group, there is a gradual decrease in ionization energy in moving from top to bottom. This is clear from the ionization energy values of the elements of the first group as given in Fig. 7.7. This decrease in ionization energy down a group can be explained in terms of net effect of the following factors: - In going from top to bottom in a group, the nuclear charge increases. - There is a gradual increase in atomic size due to an additional main shell (n). - There is an increase in shielding effect on the outermost electron due to an increase in the number of inner electrons. The effect of increase in atomic size and the shielding effect is much more than the effect of increase in nuclear charge. As a result, the electron becomes less and less firmly held to the nucleus as we move down the group. Hence, there is a gradual decrease in the ionization energies in a group. ## Successive Ionization Energies The atom may not only lose one electron but more than one electrons also. Therefore, it is essential to specify the definition of ionization energy for each electron which is to be removed. The energies required to remove subsequent electrons from the atom in the gaseous state, are known as successive ionization energies. The term first, second, third and more ionization energy refers to the removal of first, second, third or more electron respectively. These changes may be represented as follows: $M(g) \xrightarrow[]{IE_1}M^+(g)+e^-$ $M^+(g) \xrightarrow[]{IE_2}M^{2+}(g)+e^-$ $M^{2+}(g) \xrightarrow[]{IE_3}M^{3+}(g)+e^-$ Hence, $IE_1$, $IE_2$ and $IE_3$ represent first, second and third ionization energies respectively. For example, the first and second ionization energies for lithium amy be given as: $Li (g) → Li^+(g) + e^-; IE_1 = 520 kJ mol^{-1}$ $Li^+(g) → Li^{2+}(g) + e^-; IE_2 = 7289 kJ mol^{-1}$ As can be seen, the second ionization energies are very much higher than the first ionization energies. This is mainly due to the fact that after the removal of the first electron, the atom changes into monovalent positive ion. In the ion, the number of electrons decreases but the nuclear charge remains the same. As a result of this, the remaining electrons are held more tightly be the nucleus and it becomes difficult to remove the second electron. Hence, the value of second ionisation energy, $IE_2$ is greater than that of the first ($IE_1$). In the same way, the removal of the second electron will result in the formation of diapositive ion and attraction between the nucleus and remaining electrons increases further. This results into higher value of third ionisation energy ($IE_3$) than second ($IE_2$). ## 7.5.4 Electron Affinity The amount of energy released when an electron is added to an isolated gaseous atom is called electron affinity. When an electron is removed from an atom energy is required for the process of removal of electron but energy is released when electron is added in neutral atom. The process maybe expressed as: $X (g) + e^- \xrightarrow[]{} X^- (g) + Energy (Electron affinity)$ e.g. $Cl (g) + e^- (g) \xrightarrow[]{} Cl- (g) + EA$ As the definition implies, the magnitude of the electron affinity measures the ability of an atom to hold an additional electron. If an atom has more tendency to accept an electron, large energy will be released, consequently, electron affinity will be high. On the other hand if an atom has less tendency to hold the electron, small amount of energy will be released, leading to a small value of electron affinity. Electron affinities can be positive or negative. When energy is released during the addition of an electron to an atom, the electron affinity is taken as positive. Electron affinities are expressed in kJ mol⁻¹. ## Factors on which Electron Affinity depends There are many factors which govern the electron affinity but the following are some important factors on which it mostly depends: - **Nuclear charge.** The electron affinity increases as the nuclear charge increases. This is due to greater attraction for the incoming electron if nuclear charge is high. - **Size of the atom.** With the increase in size of the atom, the distance between the nucleus and the incoming electron increases and this results in lesser attraction. Consequently, the electron affinity value will decrease. - **Electronic configuration.** The element having stable electronic configurations of half and completely filled valence subshells show very small tendency to accept additional electron and thus, electron affinities are low or almost zero in certain. ## Variation of Electron Affinity Due to lack of sufficient data, the changing trends in electron affinities on moving along the periodic table are less well defined than those for ionisation energies. However, it has been observed that the electron affinity in general increases from left to right in a period and decreases from top to bottom in a group. These variations are discussed