Electron Affinity

up to fill up the orbitals that already have an electron in them. This leads to a total of 10 electrons in the five orbitals per period. As always, the spins of the electrons will be opposite when they do inhabit the same orbital.

You need to know that it isn’t exactly written in the way you’d think. Rather than write Scandium, for example as 1s22s22p63s23p64s23d1 , which would make the most sense, it is instead written with the 3 level orbitals clumped together and the four level orders clumped together, like this: 1s22s22p63s23p63d14s2 . This is made even more confusing by the fact that the sequence for chromium is written differently, according to the levels that get filled up first, which makes the chromium sequence go like this: 1s22s22p63s23p63d54s1 .

So, the rules for filling up electrons in orbitals are that first the Periodic Table is used to find the atomic number or the total electron number. Next, fill up the 1s, 2s, 2p, 3s, 3p, 4s, 4d, and 4p orbitals except that in copper and chromium, the pattern is broken up and has the d orbital and s orbital pattern switched as was described.

ELECTRON AFFINITY

The definition of electron affinity is the change in energy (as listed in kilojoules or kJ per mole) in the gaseous phase when an electron is added to the atom in order to form a negative ion. It is basically the likelihood that a neutral atom will gain an electron.

It is through chemical equations that atoms gain or lose electrons. A reaction that releases energy is called an exothermic reaction, while a reaction that absorbs energy is called an endothermic reaction. Because energy is released in an exothermic reaction, the energy is given a negative sign. The opposite is true of an endothermic reaction, which is given a positive sign. When an electron is added to a neutral atom, this results in the release of energy, giving the reaction a negative sign. This holds true only for the first electro affinity so first electron affinities are negative.

How it looks is this: Atom + an electron goes to a negatively charged ion plus energy (which makes it exothermic). When it comes to adding a second electron; however, there is more energy required to add an electron to what is already a negatively charged

ion. This overwhelms the energy that is given off when the electron is added so the net energy is positive or endothermic.

Ionization energy, which we’ve already talked about, is essentially the opposite of electron affinity. This concept of electron affinity is usually seen for those electrons in groups 16 and 17, as these are the most likely to gain an electron to fill their orbitals. The first electron affinity is described as the amount of energy released when one mole of a gaseous atom becomes a mole of a negatively charged gaseous atom, having taken on one electron. The number is different for every atom. An example is chlorine, which has an electron affinity of -349 kJ per mole. According to convention, this is a negative sign because energy is released.

While the halogens have a highly negative electron affinity, metals do not. They require a great deal of energy to take on an electron and are much more likely to form cations, which are positively charged ions. Metals, unlike halogens, have a weaker pull on their electrons and will more easily lose them.

When nonmetal elements gain electrons, the change in energy is usually negative because they give off energy in an exothermic reaction to form an anion (a negativelycharged ion). Nonmetals, as you know, have more valence electrons than metals do so it is easier to fill a stable octet in the outer shell. Nonmetals will have a greater electron affinity than is true of metals. The electron affinity is highest for chlorine than with any other atom—even higher than with fluorine, which technically should have the highest affinity when you look at the periodic table. We will talk about why this is true in a minute.

The general pattern of electron affinity is for the affinity to increase as one goes across the periodic table from left to right as well as when the atoms get smaller. The fewer valence electrons an atom has, the less likely it is to want to add another one and the more likely it is to want to give the ones it has away. Larger atoms do not have as strong of a pull from the protons because of the shielding effect of existing electrons. The stronger the attraction of the proton (the smaller the atom), the more energy is released when the electron is held onto.