Notes On Periodic Trends: Physical Properties - II - CBSE Class 11 Chemistry
Ionization Enthalpy: The energy required to remove the most loosely bound electron from an isolated gaseous atom (X) in its ground state is called ionization enthalpy. Ionization enthalpy is expressed in the units of kilo joules per mole. The energy required to remove the most loosely bound electron is called the first ionization enthalpy. The energy required to remove the second and the third electron is termed as the second ionization enthalpy and the third ionization enthalpy respectively. The second ionization enthalpy (IE₂) will always be higher than the first ionization enthalpy (IE₁). IE₂>IE₁ This is because, on removing an electron, the unipositive ion formed will have more effective nuclear charge than the number of electrons. Due to this, the nuclear attraction on the electrons increases and hence, more energy is required to remove an electron from the unipositive ion. The graph below showed the variation in first ionization enthalpies (riH) of the first 60 elements in the periodic table. Noble gases with stable ns² np⁶ configuration have the maximum first ionization enthalpy and alkali metals, with one electron in their valence shell on the other hand, have the lowest first ionization enthalpy. The ionization enthalpy generally increases across a period because of the decrease in atomic size and increase in nuclear charge. Exceptions: (i) The first ionization enthalpy of boron is slightly less than that of beryllium even though the former has a greater nuclear charge due to completely filled shell in Beryllium. (ii) The other anomaly is the smaller first ionization enthalpy of oxygen as compared to nitrogen because of the half filled shell in the case of Nitrogen. As moving go down in a group, although the nuclear charge increases, the size of an atom and the shielding effect also increases with the increase in the number of shells. Thus, ionization energy decreases down a group. Electron gain enthalpy: The enthalpy change during the process of addition of an electron to a neutral gaseous atom to convert it into a negative ion is defined as the electron gain enthalpy. The process of adding an electron can be either endothermic or exothermic. It is exothermic when an element reaches a more stable state by the addition of an electron. Ex: Halogens attain a stable noble gas electronic configuration after gaining an electron. So they have very high negative electron gain enthalpy values. And noble gas elements have large positive electron gain enthalpies. In the case of noble gases, the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration. Electron gain enthalpy becomes more negative with each successive element across a period, thus increases across a period. This is due to increase in effective nuclear charge. The electron gain enthalpy decreases down a group because of increase in the size of the atom and also the added electron would be further away from the nucleus. Exceptions: The electron gain enthalpy of oxygen is less negative than that of the succeeding element sulphur and the electron gain enthalpy of fluorine is less negative than that of the succeeding element chlorine. This is because of smaller size of oxygen and fluorine. Electronegativity: It is the ability of an atom in a chemical compound to attract shared pair of electrons towards itself. Different scales have been developed to measure electronegativity, such as Pauling scale, Mullikan-Jaffe scale and Allred-Rocha scale. Of these scales, the Pauling scale, - named after its discoverer Linus Pauling - is the most widely used scale to measure electronegativity. Pauling assigned an arbitrary value of 4.0 to fluorine. The higher the difference in the electronegativity of two bonded atoms, the more is the probability of the bonding electron pair to localize around the more electronegative atom. Ex: In the case of water, oxygen, being more electronegative, pulls the shared pair of electrons closer to itself, thus causing an overall dipole moment in water. Electronegativity increases across a period, is due to the increase in effective nuclear charge and decreases down a group because of increase in atomic size and metallic nature. Non-metallic elements have stronger tendency to gain electrons thus, electronegativity increases with the increase in non-metallic nature of elements across a period.

#### Summary

Ionization Enthalpy: The energy required to remove the most loosely bound electron from an isolated gaseous atom (X) in its ground state is called ionization enthalpy. Ionization enthalpy is expressed in the units of kilo joules per mole. The energy required to remove the most loosely bound electron is called the first ionization enthalpy. The energy required to remove the second and the third electron is termed as the second ionization enthalpy and the third ionization enthalpy respectively. The second ionization enthalpy (IE₂) will always be higher than the first ionization enthalpy (IE₁). IE₂>IE₁ This is because, on removing an electron, the unipositive ion formed will have more effective nuclear charge than the number of electrons. Due to this, the nuclear attraction on the electrons increases and hence, more energy is required to remove an electron from the unipositive ion. The graph below showed the variation in first ionization enthalpies (riH) of the first 60 elements in the periodic table. Noble gases with stable ns² np⁶ configuration have the maximum first ionization enthalpy and alkali metals, with one electron in their valence shell on the other hand, have the lowest first ionization enthalpy. The ionization enthalpy generally increases across a period because of the decrease in atomic size and increase in nuclear charge. Exceptions: (i) The first ionization enthalpy of boron is slightly less than that of beryllium even though the former has a greater nuclear charge due to completely filled shell in Beryllium. (ii) The other anomaly is the smaller first ionization enthalpy of oxygen as compared to nitrogen because of the half filled shell in the case of Nitrogen. As moving go down in a group, although the nuclear charge increases, the size of an atom and the shielding effect also increases with the increase in the number of shells. Thus, ionization energy decreases down a group. Electron gain enthalpy: The enthalpy change during the process of addition of an electron to a neutral gaseous atom to convert it into a negative ion is defined as the electron gain enthalpy. The process of adding an electron can be either endothermic or exothermic. It is exothermic when an element reaches a more stable state by the addition of an electron. Ex: Halogens attain a stable noble gas electronic configuration after gaining an electron. So they have very high negative electron gain enthalpy values. And noble gas elements have large positive electron gain enthalpies. In the case of noble gases, the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration. Electron gain enthalpy becomes more negative with each successive element across a period, thus increases across a period. This is due to increase in effective nuclear charge. The electron gain enthalpy decreases down a group because of increase in the size of the atom and also the added electron would be further away from the nucleus. Exceptions: The electron gain enthalpy of oxygen is less negative than that of the succeeding element sulphur and the electron gain enthalpy of fluorine is less negative than that of the succeeding element chlorine. This is because of smaller size of oxygen and fluorine. Electronegativity: It is the ability of an atom in a chemical compound to attract shared pair of electrons towards itself. Different scales have been developed to measure electronegativity, such as Pauling scale, Mullikan-Jaffe scale and Allred-Rocha scale. Of these scales, the Pauling scale, - named after its discoverer Linus Pauling - is the most widely used scale to measure electronegativity. Pauling assigned an arbitrary value of 4.0 to fluorine. The higher the difference in the electronegativity of two bonded atoms, the more is the probability of the bonding electron pair to localize around the more electronegative atom. Ex: In the case of water, oxygen, being more electronegative, pulls the shared pair of electrons closer to itself, thus causing an overall dipole moment in water. Electronegativity increases across a period, is due to the increase in effective nuclear charge and decreases down a group because of increase in atomic size and metallic nature. Non-metallic elements have stronger tendency to gain electrons thus, electronegativity increases with the increase in non-metallic nature of elements across a period.

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