Ionic and covalent bonding. In this lecture, we will discuss the basics of ionic and covalent bonding. We will pay special attention to the importance of electronegativity in determining the type of bond formed.
Finally, we will finish with a discussion of the octet rule. Previously, we learned the electronegativity of an atom is its tendency to attract electrons in a bond. They have been calculated for most atoms and are related to measured properties and are listed on this table.
using the Pauling scale. On this table, the electronegativity is the real number listed below the elemental symbol. Recall we've previously learned about electronegativity.
It tends to increase as we go up a group and to the right across a period. This will be important as we are going to use electronegativity in this lecture to determine the types of bonds formed. In this lecture, there are two broad categories of bonds we are going to discuss, ionic bonds and covalent bonds.
An ionic bond occurs when one atom or molecule gives up an electron or electrons to another atom or molecule. This results in the first species having a positive charge and the second species having a negative charge. The two then stick together because positive and negative attract. Covalent bonds occur when two atoms share a pair of electrons. The determining factor in the type of bond formed is the difference in electronegativity of the two atoms.
When two atoms have a difference in electronegativity that is greater than or equal to zero, but less than or equal to 0.4, they form a nonpolar covalent bond, and share the electrons more or less equally. When their difference is greater than 4, but less than or equal to 1.9, they form a polar covalent bond. In a polar covalent bond, the atom with a greater electronegativity has the lion's share of the electrons.
When the difference is greater than 1.9, They form an ionic bond. Let's try some examples. Why don't we start with one of the most abundant bonds in nature, the carbon-hydrogen single bond. First, find carbon and hydrogen on the periodic table.
Then record their respective electronegativities. Next, take the difference by subtracting the smaller electronegativity from the larger one. In this case, the difference is found to be 0.3, which, by our guidelines, indicates the bond is non-polar covalent and thus the electrons are shared equally. Now let's give the fluorine hydrogen bond a try.
Start as before by finding their electronegativities and recording them. Then take the difference and find it to be 1.9, which places this bond at the very far edge of polar covalent. Thus, the fluorine hydrogen bond is highly polar and a lot of the negative charge hangs out around fluorine. This is indicated by the funny little delta superscript negative symbol. Furthermore, the positive on hydrogen indicates the electrons aren't as strongly associated with hydrogen.
Why don't you try the oxygen-sodium bond? Did you find the difference to be 2.6, indicating an ionic bond? If so, excellent work.
At this point, we've addressed some of the types of bonds, but we haven't indicated why they form. Atoms form bonds in order to increase stability. Atoms with 8 valence electrons are more stable and hence unreactive.
Atoms without 8 valence electrons will form ions and bonds to get the 8 they want. The most obvious example of this behavior is the inert noble gases. Each one, except for the smallest helium, has 8 valence electrons, as indicated here by the noble gas configuration of neon.
This explains the bonding and ion formation patterns of various atoms. When chlorine gains an electron to become the chloride ion, the extra electron moves into the empty spot in the 3p subshell. Thus, it obtains a full valence shell of eight electrons. Similarly, when sodium loses an electron, it is removed from the 3s subshell. Thus, it now has a full outer 2s shell with eight electrons.
The tendency of atoms To have eight electrons in their outermost shell is known as the octet rule, and it will be much more important in lectures on drawing molecules.