Electronegativity Values of Elements:
Simple To Understand
Electronegativity values of elements are a result of the relationship between the number of protons in the nucleus, the total number of electrons and the distance of the the outermost electron shell from the nucleus.
If we have a clear understanding of the electron details of subshell filling for elements we understand two things.
First, all elements have between one and 8 electrons in their outer shell. This results in the electron dot diagram and electron dot structure for the element, which is a useful tool for predicting the bonds it will form with other elements.
Secondly, elements have progressively larger nuclei as we proceed across the Periodic Table from left to right, and also as we move down the columns of the table. This is seen most easily in the atomic numbers of the elements in the table as shown below.
With the exception of the elements in the right most column, the Noble Gases, electronegativity generally increases from left to right across the Periodic Table and also increases up the columns. This means that Fluorine is the most electronegative element, the element that has the strongest pull on its outer shell electrons and other electrons it encounters. The values for electronegativities of elements is presented below.
Note that the values given in this table have been rounded to the nearest 0.1.
Factors Affecting Electronegativity Values
Electronegativity values are tied up with two other properties of elements, being their atomic radius (size) and the core charge felt by the outer shell electrons, which is also the same as their old 1-8 group number. These three factors are interdependent as shown to the right. In order to understand all three components, we need to break open this cycle and start somewhere. Core Charge is a good place to start.
Briefly: This needs to be understood as the charge felt by the outer shell electrons. Since for our purposes all outer shell electrons are at roughly the same distance from the nucleus, they all experience the same core charge and are pulled towards the nucleus because of it. For example, if an element has a core charge of +1 then ALL the outer shell electrons are affected equally by this +1 charge.
In More Detail: We can calculate the core charge felt by the outer shell of an atom by considering the number of protons in the nucleus and how the electrons are arranged in their subshells. This is best demonstrated through examples, and elements in the second row of the Periodic Table are the best for this purpose.
Example 1: Lithium
Lithium has 3 protons and three electrons. This gives the nucleus an overall pull of +3. If all of Lithium’s electrons were in the same shell, they would all feel this +3 pull. However, they are not all in the same shell. Lithium has the electron configuration of 1s2 2s1, meaning it has a full first shell and one electron in its second, outermost shell as shown to the right.
The core charge of +3 radiates outward from the nucleus. As it encounters the full first shell, two of the 3 positives from the nucleus are negated by the oppositely charged protons. The remaining +1 charge is free to move on to the outer shell. So in Lithium all the outer shell electrons only feel an overall charge of +1 from the nucleus, as demonstrated below:
That’s quite a low charge and the outer shell electron does not feel that great a pull from it. For this reason Lithium readily loses its outer shell electron to other more electronegative elements. It is this low pull on electrons that gives Lithium one of the lowest electronegativity values of all elements, a value of 1.0.
Example 2: Beryllium
The element Beryllium is next to Lithium in the Periodic Table. It has one more proton and one more electron. It has two electrons in its first shell and two in its outer shell for a total of four, as shown to the left. Similar to Lithium, the core charge of +4 (1 from each proton) radiates outward and affects the electrons.
The full first shell negates two of the 4 positive charges from the nucleus, leaving a core charge of +2 to be felt by the outer shell electrons. Both these outer shell electrons feel the +2 charge, which is double the core charge felt by the outer shell electron of Lithium.
For this reason the outer shell electrons are pulled closer to the nucleus. Any electron in the vicinity will also feel a greater pull from a Beryllium atom than from a Lithium atom. This results in a higher electronegativity value for Beryllium, which has a value of 1.6 compared to 1.0 for Lithium.
The extra pull on the outer shell electrons also affects the size of the atom. As it experiences twice the pull compared with the Lithium atom, the outer shell of Beryllium is pulled closer to the nucleus and is therefore smaller than the Lithium atom.
This seems counter intuitive at first; we would expect that more electrons means a bigger atom but this is not the case.
So by comparing Lithium with its neighbor Beryllium, we can see that core charge increasesfrom left to right across the table. At the same time, atomic radius decreases, meaning the atoms become smaller from left to right across the table.
Here is a summary of the core charges felt by the outer shell electrons and relative atomic sizes for the elements in the second row of the Periodic Table being Lithium through to Fluorine. Note also the corresponding increase in electronegativity values in these elements.
Moving DOWN the Periodic Table
So what effect does moving down the Periodic Table, to the third row, have on electronegativity values, core charge and atomic radius? Since we are already familiar with Lithium we can consider its vertical neighbor, Sodium.
The element Sodium has 11 protons and electrons. It has an electron configuration of 1s2 2s2 2p6 3s1 which tells us that, like Lithium, it has ONE electron in its outer shell. It differs from Lithium in that is has two full shells which contain 2 and 8 electrons respectively. But with all those extra protons, you’d expect it to have a greater pull on its electrons. Let’s have a look at how the core charge of Sodium is affected by those full outer shells:
We see that Sodium’s outer shell electron feels only a +1 core charge, as with Lithium. However that outer shell is considerably further away from the nucleus than in Lithium. This results in a much weaker pull on that one electron. Think of two magnets; when close they pull toward each other but at a slightly greater distance the pull is virtually zero. This is also true of atoms.
The result is that Sodium has a weaker grip on its outer shell electron than Lithium. This means that its electronegativity is lower, and if you look at the electronegativity values for Lithium (Li) and Sodium (Na) on the table you can see that Sodium’s value is lower.
The extra electron shell also means that Sodium is considerably bigger than Lithium, in addition to the relatively weaker core charge felt by the outer shell electron.
So now we can understand the patterns that are generally seen in electronegativity values and the other factors tied in with them They can be summarized as follows:
* Electronegativity values increase ACROSS the table from left to right
* Atomic size decreases ACROSS the table from left to right
* Electronegativity values decrease DOWN the table from top to bottom
* Atomic size increases DOWN the table from top to bottom.
What about those ones in the middle?
You will have noticed that some elements in the center, sunken part of the table seem to disobey the general rule. This section of the table is called the D-Block because they have an incompletely filled d-subshell underneath their outer shell. This gives rise to a myriad of properties, including their electronegativity values, that are beyond the charter of this site as they would need an entire site of their own to do them justice.