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Baonza, J. Note: This simplification ignores the noble gases. Historically this is because they were believed not to form bonds - and if they do not form bonds, they cannot have an electronegativity value. Even now that we know that some of them do form bonds, data sources still do not quote electronegativity values for them. The positively charged protons in the nucleus attract the negatively charged electrons.
As the number of protons in the nucleus increases, the electronegativity or attraction will increase. Therefore electronegativity increases from left to right in a row in the periodic table. This effect only holds true for a row in the periodic table because the attraction between charges falls off rapidly with distance.
The chart shows electronegativities from sodium to chlorine ignoring argon since it does not does not form bonds. As you go down a group, electronegativity decreases. If it increases up to fluorine, it must decrease as you go down. The chart shows the patterns of electronegativity in Groups 1 and 7.
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. When atoms with an electronegativity difference of less than two units are joined together, the bond that is formed is a covalent bond , in which the electrons are shared by both atoms.
When two of the same atom share electrons in a covalent bond, there is no electronegativity difference between them, and the electrons in the covalent bond are shared equally — that is, there is a symmetrical distribution of electrons between the bonded atoms. These bonds are nonpolar covalent bonds. As an analogy, you can think of it as a game of tug-of-war between two equally strong teams, in which the rope doesn't move. For example, when two chlorine atoms are joined by a covalent bond, the electrons spend just as much time close to one chlorine atoms as they do to the other, and the resulting molecule is nonpolar:.
When the electronegativity difference is between 0 and 2, the more electronegative element attracts the shared more strongly, but not strongly enough to remove the electrons completely to form an ionic compound. The electrons are shared unequally — that is, there is an unsymmetrical distribution of electrons between the bonded atoms.
These bonds are called polar covalent bonds. For example, in the hydrogen chloride molecule, chlorine is more electronegative than hydrogen by 0. The shared electrons spend more time close to the chlorine atom, making the chlorine end of the molecule very slightly negative indicated in the figure below by the blue shaded region , while the hydrogen end of the molecule is very slightly positive indicated by the red shaded region , and the resulting molecule is polar:.
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