Why does electronegativity decrease from top to bottom

From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one. From top to bottom down a group, electronegativity decreases.

This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius. Important exceptions of the above rules include the noble gases, lanthanides , and actinides.

The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values. As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.

Ionization Energy Trends Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous phase. Trends The ionization energy of the elements within a period generally increases from left to right. This is due to valence shell stability. The ionization energy of the elements within a group generally decreases from top to bottom. This is due to electron shielding. The noble gases possess very high ionization energies because of their full valence shells as indicated in the graph.

Note that helium has the highest ionization energy of all the elements. Electron Affinity Trends As the name suggests, electron affinity is the ability of an atom to accept an electron. Electron affinity increases from left to right within a period.

This is caused by the decrease in atomic radius. Electron affinity decreases from top to bottom within a group. This is caused by the increase in atomic radius. Atomic Radius Trends The atomic radius is one-half the distance between the nuclei of two atoms just like a radius is half the diameter of a circle. Atomic radius decreases from left to right within a period.

This is caused by the increase in the number of protons and electrons across a period. One proton has a greater effect than one electron; thus, electrons are pulled towards the nucleus, resulting in a smaller radius. Atomic radius increases from top to bottom within a group.

This is caused by electron shielding. Melting Point Trends The melting points is the amount of energy required to break a bond s to change the solid phase of a substance to a liquid.

Metals generally possess a high melting point. Most non-metals possess low melting points. The non-metal carbon possesses the highest melting point of all the elements.

The semi-metal boron also possesses a high melting point. Metallic Character Trends The metallic character of an element can be defined as how readily an atom can lose an electron.

Metallic characteristics decrease from left to right across a period. This is caused by the decrease in radius caused by Z eff , as stated above of the atom that allows the outer electrons to ionize more readily. Metallic characteristics increase down a group. Electron shielding causes the atomic radius to increase thus the outer electrons ionizes more readily than electrons in smaller atoms. Metallic character relates to the ability to lose electrons, and nonmetallic character relates to the ability to gain electrons.

Problems The following series of problems reviews general understanding of the aforementioned material. Nitrogen has a larger atomic radius than oxygen. True B. False 3. Which has more metallic character, Lead Pb or Tin Sn? Which element has a higher melting point: chlorine Cl or bromine Br? Which element is more electronegative, sulfur S or selenium Se?

Oxygen O B. Chlorine Cl C. Calcium Ca D. Lithium Li E. None of the above 10 A nonmetal has a smaller ionic radius compared with a metal of the same period. Solutions 1. References Pinto, Gabriel. Iqbal M. Why is electronegativity a factor that influences NMR spectra? What is shielding and deshielding in NMR? Can you give me an example? How is it called the effect of electronegative atoms on their neighbours?

What is the pi -bond effect? What happens if the electron density around a nucleus is decreased? The electron pair is screened from both nuclei by the 1s, 2s and 2p electrons, but the chlorine nucleus has 6 more protons in it.

It is no wonder the electron pair gets dragged so far towards the chlorine that ions are formed. Electronegativity increases across a period because the number of charges on the nucleus increases. That attracts the bonding pair of electrons more strongly. As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus. Consider the hydrogen fluoride and hydrogen chloride molecules:. The bonding pair is shielded from the fluorine's nucleus only by the 1s 2 electrons.

In the chlorine case it is shielded by all the 1s 2 2s 2 2p 6 electrons. But fluorine has the bonding pair in the 2-level rather than the 3-level as it is in chlorine. If it is closer to the nucleus, the attraction is greater. At the beginning of periods 2 and 3 of the Periodic Table, there are several cases where an element at the top of one group has some similarities with an element in the next group.

Three examples are shown in the diagram below. Notice that the similarities occur in elements which are diagonal to each other - not side-by-side. For example, boron is a non-metal with some properties rather like silicon. Unlike the rest of Group 2, beryllium has some properties resembling aluminum. And lithium has some properties which differ from the other elements in Group 1, and in some ways resembles magnesium.

There is said to be a diagonal relationship between these elements. There are several reasons for this, but each depends on the way atomic properties like electronegativity vary around the Periodic Table. So we will have a quick look at this with regard to electronegativity - which is probably the simplest to explain. Electronegativity increases across the Periodic Table. So, for example, the electronegativities of beryllium and boron are:.

Electronegativity falls as you go down the Periodic Table. So, for example, the electronegativities of boron and aluminum are:. So, comparing Be and Al, you find the values are by chance exactly the same. The increase from Group 2 to Group 3 is offset by the fall as you go down Group 3 from boron to aluminum. Something similar happens from lithium 1. In these cases, the electronegativities are not exactly the same, but are very close.

Similar electronegativities between the members of these diagonal pairs means that they are likely to form similar types of bonds, and that will affect their chemistry. You may well come across examples of this later on in your course. Jim Clark Chemguide. What if two atoms of equal electronegativity bond together? If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms: To get a bond like this, A and B would usually have to be the same atom.

What if B is slightly more electronegative than A? B will attract the electron pair rather more than A does. A "spectrum" of bonds The implication of all this is that there is no clear-cut division between covalent and ionic bonds. Summary No electronegativity difference between two atoms leads to a pure non-polar covalent bond.

A small electronegativity difference leads to a polar covalent bond. A large electronegativity difference leads to an ionic bond.