By Ron B. Davis Jr., Georgetown University
In the real world, where atoms are far more commonly found chemically bonded to other atoms, electronegativity gives us a more usable measure of how a tug-of-war for electrons between atoms plays out. Read on to know more.
Electronegativity
Electron-gaining measurement that describes how atoms attract electrons in compounds is called electronegativity. It measures how intensely a given element will fight with other elements for electrons in a chemical bond.
For decades chemists devised ways to try to describe and measure the force with which atoms of various elements pull on electron pairs in the bonds that join them.
New Short-form Periodic Table
In 1908, such measurements led to a truly remarkable version of the relatively new short-form periodic table of the elements.
British chemist, Sidney Young, produced this illustration of the table, in which each element is placed within its box, based on what Young called the elements’ ‘electrochemical character’—more to the left for less electronegative elements, and farther to the right for more electronegative elements.
Sorting the Transition Metals
The illustration encapsulated a clear and remarkable pattern emerging, in which one can see the beginnings of the transition metals being sorted from the others within their groups.
Copper, silver and gold, for example, began to stand out from the group 1 alkali metals lithium, sodium, potassium, sodium, rubidium and cesium. It was clear that there was something quite different about each group of metals.
That is, even though their valence appeared to be the same, the extent to which they pulled, bonding electrons toward themselves, was different. This was 15 years before the first long-form periodic table, but these trends involving electrons were beginning to hint that certain metals needed their own home in the table.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
Linus Pauling
It wasn’t until 1932, when the structure of the atom and the nature of chemical bonding had been well characterized and the periodic table commonly extended to its long form, that another famous chemist by the name of Linus Pauling devised the scale of electronegativity we use today.
Pauling’s system runs on a scale from 0.0 to 4.0, with 4.0 being the most electronegative and 0.0 the least. And when we take a look at the main block of the periodic table, the trend is undeniable, with fluorine the undisputed king of electron acceptors.
Trends in Electronegativity
Trends in electronegativity are reminiscent of our trend in atomic radius, and for a good reason. Atoms whose nuclei pull the hardest on their own outer electrons should by reason pull on shared electrons in their chemical bonds hardest as well.
So it comes as no surprise that fluorine is the undisputed king of electronegativity, with its proton-packed nucleus attracting electrons in the second valence shell very close by.
Descending the row we see once again that chlorine, bromine and iodine nuclei must attract electrons in more distant energy levels, progressively more and more screened from the pull of their nucleus.
And yet, as we move from left to right across the table, we are creating ever more positively charged nuclei, but adding electrons to the same valence shell which doesn’t create significant screening of the nucleus.
Electronegativity Scale
To sum it up, noble gases excepted, elements from the top right of the table, like fluorine and oxygen boast the highest electronegativities of all, and those from the bottom left like cesium and rubidium have the lowest electronegativity.
The electronegativity scale gives us a powerful tool for the prediction of how elements will choose to bond to one another. Pairs of atoms nearby one another on the table are usually of similar or equal electronegativity and are most unwilling to trade electrons.
This leads most metals to form metallic bonds with one another and many nonmetals to form covalent bonds with one another.
Growing Difference in Electronegativity
But, as elements get more and more distant on the periodic table, the difference in their electronegativities grows. At first, one atom simply pulls the covalent bonding pair a little bit closer to itself, creating what we call a polar covalent bond. This is similar to what we might find in hydrochloric acid.
However, when the difference is very large, sharing stops and an electron trade takes place instead, leading to an ionic bond like those we find in sodium chloride.
In conclusion, one can clearly observe many periodic patterns that help us understand some of the most fundamental characteristics of atoms and how they interact, attract and bond chemically.
Common Questions about Electronegativity and the Periodic Table
Electron-gaining measurement that describes how atoms attract electrons in compounds is called electronegativity. It measures how intensely a given element will fight with other elements for electrons in a chemical bond.
The electronegativity scale gives us a powerful tool for the prediction of how elements will choose to bond to one another. Pairs of atoms nearby one another on the table are usually of similar or equal electronegativity and most unwilling to trade electrons.
It wasn’t until 1932, when the structure of the atom and the nature of chemical bonding had been well characterized and the periodic table commonly extended to its long form, that a famous chemist by the name of Linus Pauling devised the scale of electronegativity we use today.
