Which groups readily accept electrons

That is, the Group 7A nonmetals form 1- charges, the Group 6A nonmetals form 2- charges, and the Group 5A metals form 3- charges. The Group 8A elements already have eight electrons in their valence shells, and have little tendency to either gain or lose electrons, and do not readily form ionic or molecular compounds.

Ionic compounds are held together in a regular array called a crystal lattice by the attractive forces between the oppositely charged cations and anions. These attractive forces are very strong, and most ionic compounds therefore have very high melting points. Ionic compounds are typically hard, rigid, and brittle.

Ionic compounds do not conduct electricity, because the ions are not free to move in the solid phase, but ionic compounds can conduct electricity when they are dissolved in water. When nonmetals combine with other nonmetals, they tend to share electrons in covalent bonds instead of forming ions, resulting in the formation of neutral molecules. Keep in mind that since hydrogen is also a nonmetal, the combination of hydrogen with another nonmetal will also produce a covalent bond.

Molecular compounds can be gases, liquids, or low melting point solids, and comprise a wide variety of substances. See the Molecule Gallery for examples. When metals combine with each other, the bonding is usually described as metallic bonding you could've guessed that. In this model, each metal atom donates one or more of its valence electrons to make an electron sea that surrounds all of the atoms, holding the substance together by the attraction between the metal cations and the negatively charged electrons.

The number of protons gives the element its identity. Look at the different groups for clues to how the elements will react. Noble Gases usually do not react because they do not tend to gain or lose electrons. Alkali and Alkaline Earth Metals are soft and melt at low temperatures. They react well with nonmetals because they can easily give up electrons to form ions. Transition Metals can have a slight charge which lets them bond easily to nonmetals.

Thus, nonmetals have a higher electron affinity than metals, meaning they are more likely to gain electrons than atoms with a lower electron affinity. For example, nonmetals like the elements in the halogens series in Group 17 have a higher electron affinity than the metals. This trend is described as below. Notice the negative sign for the electron affinity which shows that energy is released.

As the name suggests, electron affinity is the ability of an atom to accept an electron. Unlike electronegativity, electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom's affinity for electrons. Electron affinity increases upward for the groups and from left to right across periods of a periodic table because the electrons added to energy levels become closer to the nucleus, thus a stronger attraction between the nucleus and its electrons.

Remember that greater the distance, the less of an attraction; thus, less energy is released when an electron is added to the outside orbital. In addition, the more valence electrons an element has, the more likely it is to gain electrons to form a stable octet. The less valence electrons an atom has, the least likely it will gain electrons.

Electron affinity decreases down the groups and from right to left across the periods on the periodic table because the electrons are placed in a higher energy level far from the nucleus, thus a decrease from its pull.

However, one might think that since the number of valence electrons increase going down the group, the element should be more stable and have higher electron affinity. One fails to account for the shielding affect. As one goes down the period, the shielding effect increases, thus repulsion occurs between the electrons. This is why the attraction between the electron and the nucleus decreases as one goes down the group in the periodic table.

As you go down the group, first electron affinities become less in the sense that less energy is evolved when the negative ions are formed. Fluorine breaks that pattern, and will have to be accounted for separately. The electron affinity is a measure of the attraction between the incoming electron and the nucleus - the stronger the attraction, the more energy is released. The factors which affect this attraction are exactly the same as those relating to ionization energies - nuclear charge, distance and screening.

The increased nuclear charge as you go down the group is offset by extra screening electrons. A fluorine atom has an electronic structure of 1s 2 2s 2 2px 2 2py 2 2pz 1. It has 9 protons in the nucleus. The incoming electron enters the 2-level, and is screened from the nucleus by the two 1s 2 electrons. In contrast, chlorine has the electronic structure 1s 2 2s 2 2p 6 3s 2 3p x 2 3p y 2 3p z 1 with 17 protons in the nucleus.

There is also a small amount of screening by the 2s electrons in fluorine and by the 3s electrons in chlorine. This will be approximately the same in both these cases and so does not affect the argument in any way apart from complicating it!

The over-riding factor is therefore the increased distance that the incoming electron finds itself from the nucleus as you go down the group. The greater the distance, the less the attraction and so the less energy is released as electron affinity. Comparing fluorine and chlorine is not ideal, because fluorine breaks the trend in the group. However, comparing chlorine and bromine, say, makes things seem more difficult because of the more complicated electronic structures involved.

What we have said so far is perfectly true and applies to the fluorine-chlorine case as much as to anything else in the group, but there's another factor which operates as well which we haven't considered yet - and that over-rides the effect of distance in the case of fluorine. The incoming electron is going to be closer to the nucleus in fluorine than in any other of these elements, so you would expect a high value of electron affinity.

However, because fluorine is such a small atom, you are putting the new electron into a region of space already crowded with electrons and there is a significant amount of repulsion. This repulsion lessens the attraction the incoming electron feels and so lessens the electron affinity.

A similar reversal of the expected trend happens between oxygen and sulfur in Group The first electron affinity of oxygen kJ mol -1 is smaller than that of sulfur kJ mol-1 for exactly the same reason that fluorine's is smaller than chlorine's.