- First and second electronic affinities
- First
- Second
- How electron affinity varies in the periodic table
- Variation by core and shielding effect
- Variation by electron configuration
- Examples
- Example 1
- Example 2
- References
The electronic affinity or electro-affinity is a measure of the energetic variation of an atom in gas phase when it incorporates an electron to its valence shell. Once the electron has been acquired by atom A, the resulting anion A - may or may not be more stable than its ground state. Therefore, this reaction can be endothermic or exothermic.
By convention, when the electron gain is endothermic, a positive sign "+" is assigned to the electron affinity value; On the other hand, if it is exothermic -that is, it releases energy- this value is given a negative sign "-". In what units are these values expressed? In kJ / mol, or in eV / atom.
If the element were in a liquid or solid phase, its atoms would interact with each other. This would cause the energy absorbed or released, due to the electronic gain, to be dispersed among all of these, yielding unreliable results.
In contrast, in the gas phase they are assumed to be isolated; in other words, they don't interact with anything. So, the atoms involved in this reaction are: A (g) and A - (g). Here (g) denotes that the atom is in the gas phase.
First and second electronic affinities
First
The electronic gain reaction can be represented as:
A (g) + e - => A - (g) + E, or as A (g) + e - + E => A - (g)
In the first equation, E (energy) is found as a product on the left side of the arrow; and in the second equation the energy is counted as reactive, being located on the right side. That is, the first corresponds to an exothermic electronic gain and the second to an endothermic electronic gain.
However, in both cases it is only one electron that is added to the valence shell of atom A.
Second
It is also possible that, once the negative ion A - is formed, it absorbs another electron:
A - (g) + e - => A 2– (g)
However, the values for the second electron affinity are positive, since the electrostatic repulsions between the negative ion A - and the incoming electron e - must be overcome.
What determines that a gaseous atom better "receives" an electron? The answer is essentially found in the nucleus, in the shielding effect of the inner electronic shells and in the valence shell.
How electron affinity varies in the periodic table
In the upper image, the red arrows indicate the directions in which the electronic affinity of the elements increases. From this, electron affinity can be understood as one more of the periodic properties, with the peculiarity that it has many exceptions.
Electron affinity increases ascending through the groups and also increases from left to right along the periodic table, especially around the fluorine atom. This property is closely related to the atomic radius and the energy levels of its orbitals.
Variation by core and shielding effect
The nucleus has protons, which are positively charged particles that exert an attractive force on the electrons in the atom. The closer the electrons are to the nucleus, the greater the attraction they feel. Thus, as the distance from the nucleus to the electrons increases, the lower the attractive forces.
In addition, the electrons in the inner shell help to "shield" the effect of the nucleus on the electrons in the outermost shells: the valence electrons.
This is due to the electronic repulsions themselves between their negative charges. However, this effect is counteracted by increasing the atomic number Z.
How does the above relate to electronic affinity? That a gaseous atom A will have a greater tendency to gain electrons and form stable negative ions when the shielding effect is greater than the repulsions between the incoming electron and those of the valence shell.
The opposite occurs when the electrons are very far from the nucleus and the repulsions between them do not disfavor the electronic gain.
For example, descending in a group "opens" new energy levels, which increase the distance between the nucleus and external electrons. It is for this reason that as you move up the groups, electronic affinities increase.
Variation by electron configuration
All orbitals have their energy levels, so if the new electron will occupy a higher-energy orbital, the atom will need to absorb energy for this to be possible.
Furthermore, the way in which electrons occupy the orbitals may or may not favor electronic gain, thus distinguishing differences between atoms.
For example, if all the electrons are unpaired in the p orbitals, the inclusion of a new electron will cause the formation of a paired pair, which exerts repulsive forces on the other electrons.
This is the case for the nitrogen atom, whose electron affinity (8kJ / mol) is lower than for the carbon atom (-122kJ / mol).
Examples
Example 1
The first and second electronic affinities for oxygen are:
O (g) + e - => O - (g) + (141kJ / mol)
O - (g) + e - + (780kJ / mol) => O 2– (g)
The electron configuration for O is 1s 2 2s 2 2p 4. There is already a paired pair of electrons, which cannot overcome the attractive force of the nucleus; therefore, the electronic gain releases energy after the stable O - ion is formed.
However, although O 2– has the same configuration as the noble gas neon, its electronic repulsions exceed the attractive force of the nucleus, and an energy supply is necessary to allow the electron to enter.
Example 2
If the electronic affinities of the elements of group 17 are compared, the following will be obtained:
F (g) + e - = F - (g) + (328 kJ / mol)
Cl (g) + e - = Cl - (g) + (349 kJ / mol)
Br (g) + e - = Br - (g) + (325 kJ / mol)
I (g) + e - = I - (g) + (295 kJ / mol)
From top to bottom - descending in the group - the atomic radii increase, as well as the distance between the nucleus and the external electrons. This causes an increase in electronic affinities; however, fluorine, which should have the highest value, is outnumbered by chlorine.
Why? This anomaly demonstrates the effect of electronic repulsions on the attractive force and low shielding.
Because it is a very small atom, fluorine "condenses" all its electrons in a small volume, causing a greater repulsion on the incoming electron than its more voluminous congeners (Cl, Br and I).
References
- Chemistry LibreTexts. Electron Affinity. Retrieved on June 4, 2018, from: chem.libretexts.org
- Jim Clark. (2012). Electron Affinity. Retrieved on June 4, 2018, from: chemguide.co.uk
- Carl R. Nave. Electron Affinities of the Main-Group Elements. Retrieved on June 4, 2018, from: hyperphysics.phy-astr.gsu.edu
- Prof. N. De Leon. Electron Affinity. Retrieved on June 4, 2018, from: iun.edu
- Helmenstine, Anne Marie, Ph.D. (May 27, 2016). Electron Affinity Definition. Retrieved on June 4, 2018, from: thoughtco.com
- Cdang. (October 3, 2011). Electron affinity periodic table.. Retrieved on June 04, 2018, from: commons.wikimedia.org
- Whitten, Davis, Peck & Stanley. Chemistry. (8th ed.). CENGAGE Learning, p 227-229.
- Shiver & Atkins. (2008). Inorganic chemistry. (Fourth edition., P. 29). Mc Graw Hill.