First Ionization Energies Definition The 1 st ionization

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First Ionization Energies ▪ Definition: The 1 st ionization energy of an element is

First Ionization Energies ▪ Definition: The 1 st ionization energy of an element is the energy required to remove the outermost electron from a gas-phase atom in its ground state. 1 st IE in k. J mol-1 ▪ Trends: 1 st ionization energy increases across a period and decreases down a group. ▪ Explanation: Ionization energy (I) is given by: RH = Rydberg Constant 1

First Ionization Energies ▪ Across a period, as Zeff increases and the valence-shell principal

First Ionization Energies ▪ Across a period, as Zeff increases and the valence-shell principal quantum number n remains constant, the ionization energy should increase. And down a group, as n increases and Zeff increases only slightly, the ionization energy should decrease. Exceptions from Expected Trends ▪ High ionization energies occur when electrons are removed from half-filled or full shells or subshells. ▪ Thus, the first ionization energy of nitrogen ([He]2 s 22 p 3) is 1402 k. J mol-1, which is higher than the value for oxygen ([He]2 s 22 p 4, 1314 k. J mol-1). ▪ Similarly, the ionization energy of phosphorus (1011 k. J mol-1) is higher than that of sulfur (1000 k. J mol-1). ▪ The 1 st ionization energy of Be is higher than B. This is due to to the relative stability of the filled shell 2 s 2 configuration compared with the 2 s 22 p 1 arrangement. 2

Second and Higher Ionization Energies ▪ Definition: Second Ionization Energy is the energy required

Second and Higher Ionization Energies ▪ Definition: Second Ionization Energy is the energy required to remove a second outermost electron from a ground state atom. ▪ Subsequent ionization energies increase greatly. More energy must be invested to remove an electron from a charged atom than from a neutral atom. This is a direct consequence of Coulomb’s law which states that the force of attraction between oppositely charged particles is directly proportional to the magnitudes of the charges. Ionization Energies (k. J mol-1) Element 1 st IE 2 nd IE 3 rd IE 4 th IE Na 495. 8 4562. 4 6912 9543 Mg 737. 7 1450. 6 7732. 6 10, 540 Al 577. 6 1816. 6 2744. 7 11, 577 3

Electron Affinities ▪ Definition: The energy released when a neutral atom in the gas

Electron Affinities ▪ Definition: The energy released when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion. EA in k. J mol-1 ▪ Trends: Electron affinity (EA) increases across a period and decreases down a group. 4

Electron Affinities ▪ Explanation: As Zeff increases across a series, the larger nuclear charge

Electron Affinities ▪ Explanation: As Zeff increases across a series, the larger nuclear charge exerts a stronger attraction on the electron, resulting in a large value for the E. A. . Progressing down a column, the electron is added to a quantum level that is, on average, farther from the nucleus (larger n). Thus, that electron will experience a smaller nuclear attraction, and the atom will have a low E. A. Exceptions from Expected Trends ▪ The repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. ▪ Among the nonmetals in Groups 16 and 17, this force of repulsion is largest for the very smallest atoms in these columns: oxygen and fluorine. ▪ As a result, these elements have a smaller electron affinity than the elements below them in these columns as shown in the figure below. From that point on, however, the electron affinities decrease as we continue down these columns. - 5

Electron Affinities ▪ Elements with electron configurations of ns 2, np 3, and np

Electron Affinities ▪ Elements with electron configurations of ns 2, np 3, and np 6 have electron affinities greater than zero because they are unusually stable. In other words, instead of energy being released, these elements actually require an input of energy in order to gain electrons. e. g. Be, Mg, N, Ne. - 6

Electronegativity ▪ Definition: The power of an atom of the element to attract electrons

Electronegativity ▪ Definition: The power of an atom of the element to attract electrons to itself when it is part of a compound. Pauling Electronegativity ▪ Because electronegativity is a qualitative property, there is no standardized method for calculating electronegativity. ▪ Trends: Electronegativity increases across a period and decreases down a group. 7

Electronegativity ▪ Explanation: According to Mulliken-Jaffe, electronegativity is the mean of IE and EA

Electronegativity ▪ Explanation: According to Mulliken-Jaffe, electronegativity is the mean of IE and EA of an element. If an atom has a high ionization energy (so it is unlikely to give up electrons) and a high electron affinity (so there are energetic advantages in its gaining electrons), then it is more likely to attract an electron to itself. So, the electronegativities of the elements follow the trends in ionization energies and electron affinities. Exceptions from Expected Trends ▪ There are some exceptions from the general trend of electronegativity values: ▪ This departure from a smooth decrease down the group is called the alternation effect. It is due to the intervention of the poorly shielding 3 d subshell earlier in Period 4. Thus, Ga has higher Zeff than Al and Ge has higher Zeff than Si. Hence, Ga is more electronegative than Al and Ge is more electronegative than Si. 8