Electron Affinity
Electron affinity refers to the energy change that occurs when an atom gains an electron to form a negative ion. It is a measure of an atom's ability to attract and hold onto electrons. A higher electron affinity indicates a stronger attraction for electrons, while a lower electron affinity suggests a weaker attraction.
3 Key excerpts on "Electron Affinity"
eBook - ePub- Tony Cox(Author)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
...Thus period 2 p -block elements are in many ways different from those lower in the p block, and 3 d series elements distinct from those of the 4 d and 5 d series. States of ionization The successive energies required to create more highly charged ions, M 2+, M 3+ … are the second, third, … IEs. The values always increase with the degree of ionization. When electrons are removed from the same shell, the main effect is that with each successive ionization there is one less electron left to repel the others. The magnitude of the change therefore depends on the size of the orbital, as electrons in smaller orbitals are on average closer together and have more repulsion. Thus with Be (2 s) 2 the first two IEs are 9.3 and 18.2 eV, whereas with Ca (4 s) 2 the values are 6.1 and 11.9 eV, not only smaller to start with (see above) but with a smaller difference. The third IE of both elements is very much higher (154 and 51 eV, respectively) because now the outer shell is exhausted and more tightly bound inner shells (1 s and 3 p, respectively) are being ionized. The trends are important in understanding the stable valence states of elements. The Electron Affinity of an atom may be defined as the ionization energy of the negative ion, thus the energy input in the process: M – → 2σ + e – although some books use a definition with the opposite sign. Electron affinities are always less than ionization energies because of the extra electron repulsion involved (see Fig. 1). As with successive IEs, the difference depends on the orbital size. Some apparently anomalous trends can be understood in this way. For example, although the IE of F is greater than that of Cl (17.4 and 13.0 eV, respectively) the Electron Affinity of F is smaller (3.4 eV compared with 3.6 eV) partly because the smaller size of F provides more repulsion from the added electron. Some atoms have negative electron affinities, meaning that the negative ion is not stable in the gas phase...
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...The atomic radius also increases as you move down the periodic table. Cations have smaller radii than their corresponding neutral atoms. This is due to a greater effective nuclear charge (greater number of protons in the nucleus compared to electrons) and thus the positively charged nucleus pulls the valence electrons closer. Anions have larger radii than their corresponding neutral atoms. This can be attributed to the addition of valence electrons and their repulsion, thus increasing the atomic radius. Figure 5.5. TEST TIP For the AP Chemistry exam, be able to describe why the trends in the atomic radius exist. Ionization Energy The ionization energy is the energy required to remove an electron from an atom, resulting in the formation of a cation. The ionization energy decreases when moving from right to left across the periodic table. It also decreases when moving down a group within the periodic table. It should be noted that more than 1 electron may be removed to form ions of greater charge, but the energy required to remove successive electrons will increase exponentially. Cations that have electron configurations containing filled octets are extremely stable and require an enormous amount of energy to remove an electron. Figure 5.6. Electron Affinity Electron Affinity is the ability of an atom to gain electrons in order to form anions. Nonmetal atoms have a much higher Electron Affinity, because metals will not form anions. The Electron Affinity increases moving from left to right across the periodic table (see Figure 5.7). It also increases when moving up a group within the periodic table. Figure 5.7. TEST TIP An explanation of most of these trends on the periodic table can be linked to the concept of the atomic radius. Electronegativity Electronegativity is the ability of an atom to attract electron density to it when forming a covalent bond. The electronegativity of an atom exhibits the same trend as Electron Affinity in the periodic table...
eBook - ePub- Subrata Pal(Author)
- 2019(Publication Date)
- Academic Press(Publisher)
...The obvious question then would be why and how the bonds are formed. 4.1.1 Electron configuration of atoms We know that atoms consist of positively charged protons in the nucleus (leaving aside the uncharged neutrons) and negatively charged electrons in the orbitals. When two atoms approach each other, a complex set of electrostatic interactions develop among the charged particles. The electrons of one atom are attracted by the nucleus, but repelled by the electrons of the other atom and vice versa. At the same time, the two nuclei also repel each other. If these interactions result in the reduction in energy of each atom and, thus, the total system, a chemical bond is formed. Therefore, the interactions should be understood in terms of the electron configurations of the atoms. In Section 3.5.5, we have established four quantum numbers (n, l, m l, m s) for the hydrogen atom to identify the possible quantum states of its single electron. This quantum number principle can be extended to multielectron atoms as well. Further, in chemistry, it is customary to denote the angular momentum quantum number l = 0, 1, 2, 3, 4,… as s, p, d, f, g,…, respectively. As in the hydrogen atom, in a multielectron atom, too, increasing n is associated with increasing energy of the corresponding shell and, for a particular value of n, a subshell with a greater value of l is at a higher energy level. Thus, E 4 s < E 4 p < E 4 d < E 4 f where E is the energy. However, when we compare two subshells with different principal quantum numbers, as for example 3d with 4s, we find that E (4s) < E (3d). With an increase in the atomic number (denoted by Z), the number of electrons in the atom increases. It is expected that each electron would try to occupy the lowest energy state which, as stated above, is 1s (that is n = 1, l = 0). Does it mean that all would crowd into 1s? Not at all! Here, we have to take into account an extremely important quantum principle— Pauli exclusion principle...
