![]() ![]() The densities of the Group 1 elements increase down the group (except for a downward fluctuation at potassium). As before, the trend is determined by the distance between the nucleus and the bonding electrons. The increased charge on the nucleus down the group is offset by additional levels of screening electrons. As previously discussed, each atom exhibits a net pull from the nuclei of +1. The atoms are more easily pulled apart to form a liquid, and then a gas. As the atoms increase in size, the distance between the nuclei and these delocalized electrons increases therefore, attractions fall. The atoms in a metal are held together by the attraction of the nuclei to electrons which are delocalized over the whole metal mass. The decrease in melting and boiling points reflects the decrease in the strength of each metallic bond. When any of the Group 1 metals is melted, the metallic bond is weakened enough for the atoms to move more freely, and is broken completely when the boiling point is reached. Both the melting and boiling points decrease down the group. The figure above shows melting and boiling points of the Group 1 elements. The iodine atom is so large that the pull from the iodine nucleus on the pair of electrons is relatively weak, and a fully-ionic bond is not formed. Lithium iodide, for example, will dissolve in organic solvents this is a typical property of covalent compounds. In some lithium compounds there is often a degree of covalent bonding that is not present in the rest of the group. That means that the electron pair is going to be more strongly attracted to the net +1 charge on the lithium end, and thus closer to it. The net pull from each end of the bond is the same as before, but the lithium atom is smaller than the sodium atom. Now compare this with a lithium-chlorine bond. This strong attraction from the chlorine nucleus explains why chlorine is much more electronegative than sodium. The electron pair is so close to the chlorine that an effective electron transfer from the sodium atom to the chlorine atom occurs-the atoms are ionized. The electron pair will be pulled toward the chlorine atom because the chlorine nucleus contains many more protons than the sodium nucleus. The bond can be considered covalent, composed of a pair of shared electrons. Picture a bond between a sodium atom and a chlorine atom. Each of these elements has a very low electronegativity when compared with fluorine, and the electronegativities decrease from lithium to cesium. ![]() It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0 ( Table A2).Ī graph showing the electronegativities of the Group 1 elements is shown above. However, the distance between the nucleus and the outer electrons increases down the group electrons become easier to remove, and the ionization energy falls.Įlectronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. As mentioned before, in each of the elements Group 1, the outermost electrons experience a net charge of +1 from the center.
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