Electronegativity varies in a predictable way across the periodic table. Electronegativity increases from bottom to top in groups , and increases from left to right across periods. Thus, fluorine is the most electronegative element, while francium is one of the least electronegative.
Helium, neon, and argon are not listed in the Pauling electronegativity scale, although in the Allred-Rochow scale, helium has the highest electronegativity.
The trends are not very smooth among the transition metals and the inner transition metals, but are fairly regular for the main group elements, and can be seen in the charts below. The difference in electronegativity between two bonded elements determines what type of bond they will form. When atoms with an electronegativity difference of greater than two units are joined together, the bond that is formed is an ionic bond , in which the more electronegative element has a negative charge, and the less electronegative element has a positive charge.
Electronegativity differs from electron affinity because electron affinity is the actual energy released when an atom gains an electron. Electronegativity is not measured in energy units, but is rather a relative scale. All elements are compared to one another, with the most electronegative element, fluorine, being assigned an electronegativity value of 3.
Fluorine attracts electrons better than any other element. The table below shows the electronegativity values for the elements. Figure 1. The largest electronegativity 3. Since metals have few valence electrons, they tend to increase their stability by losing electrons to become cations.
Consequently, the electronegativities of metals are generally low. Nonmetals have more valence electrons and increase their stability by gaining electrons to become anions.
The electronegativities of nonmetals are generally high. Electronegativities generally increase from left to right across a period. 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. Smith, Derek W. Russo, Steve, and Mike Silver. Introductory Chemistry. San Francisco: Pearson, Petrucci, Ralph H, et al.
General Chemistry: Principles and Modern Applications. New Jersey: Pearson, Atkins, Peter et. Consider sodium at the beginning of period 3 and chlorine at the end ignoring the noble gas, argon.
Think of sodium chloride as if it were covalently bonded. Both sodium and chlorine have their bonding electrons in the 3-level. 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.
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