Covalent+Bonding

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 **Covalent bonding** usually occurs between two elements which both have a high electronegativity. Thus both atoms in a covalent bond will be pulling the electrons towards itself and since both atoms have a high electronegativity the electrons are shared between the two atoms. Usually a covalent bond occurs between two non-metals. Therefore you can say that the non-metals in groups 4 to 7 of the periodic table form covalent bonds with each other. As said earlier, in a covalent bond the atoms share a pair of electrons. Thus the (positive) nucleus of each atom will be pulling the shared electron pair towards itself, As a result the two atoms are held together. In a covalent bond the //electrostatic force of attraction// between the //electron pair// and //nuclei// holds the electron pair together.



In order to represent a covalent bond we draw a diagram showing the outershell electrons of each atom and the electrons pair which is being shared between them. In order to draw such diagrams we use what is known as Lewis structures.

=Lewis Structure =

A Lewis structure is just a way of drawing a covalent molecule using dots, crosses, or a combination of the two to represent the electrons.

For example the covalent molecule Cl 2 can be represented as follows:
 * [[image:Cl2Lewis1.png width="181" height="112" caption="The dots and the crosses represent electrons. you can use dots and crosses to make it easy to distinguish which atoms the electron came from."]] || [[image:Cl2Lewis2.png width="196" height="127" caption="We can also use only dots to represent the electrons."]] || [[image:Cl2Lewis3.png width="196" height="125" caption="OR, likewise we could use crosses to represent all of the electrons."]] || [[image:Cl2Lewis4.png width="219" height="94" caption="Here, the line represents an electron pair. Since each Cl atom has four lines around it, it means that they both have the same number of outershell electrons that is 8 electrons each."]] ||

Remember that in a Lewis structure we can use a dot or a cross to represent one molecule. You can also use a line to represent an electron pair. But be careful that **Cl—Cl is not a Lewis structure** because it doesn’t show all of the outershell electrons on the Cl atom.

From now on we will use Lewis structures to draw diagrams of covalently bonded molecules.

Represent the methane molecule (CH 4 ) using Lewis structure.
 * [[image:ibchem4u/CH4_1.png width="164" height="174" caption="each line represents a pair of electrons."]] || [[image:CH4_2.png width="168" height="174" caption="Or you could use dot and crosses to represent electrons."]] ||

=Single, double or triple covalent bond =

Atoms bond covalently so that their outermost electron shell gets filled (so that they can become stable).

So far we have seen examples of single covalent bonds, where only one pair of electrons is shared between the atoms. Eg CH 4 in the previous example

However atoms can also share two or even three pairs of electrons with each other thus forming double or triple covalent bonds respectively.

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">For example Oxygen (O) has 6 outershell electrons, so it still has place for two more electrons in its outershell (8-6=2)

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Therefore in O 2 molecule, the two O atoms share two pairs of electrons so that their outershell is complete. They form a double covalent bond. <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Nitrogen atoms have 5 electrons in their outershell, therefore they have space for 3 more electrons. Thus two nitrogen atoms form a covalent bond where they share 3 electron pairs (so that their outershell becomes complete) form a triple covalent bond.

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Can you draw the Lewis structure for N 3 now???

=<span style="color: #362079; font-family: Tahoma,sans-serif; font-size: 14pt;">Dative covalent bond =

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">As we said earlier in a covalent bond, an electron pair is shared between the two molecules. Most of the times, each one of the atoms has contributed one of the electrons. For example when Cl atom reacts with another Cl atom to form a covalent bond, each Cl atom contributes one electron which forms a pair which is shared between the two atoms.

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">However this isn’t always the case. Sometimes in a covalent bond, both electrons in the shared pair come from the same atom. This is usually known as a dative (or coordinate) bond. It’s still a covalent bond but with the exception that both electrons in the shared electron pair have come from the same atom. This can be clarified through a few examples.

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Ammonia molecule (NH 3 ) can react with a hydrogen ion (H + ) to form ammonium (NH 4 ) <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">The NH 4 molecule can also be drawn using lines to represent one electron pair. But in that case be careful that the dative bond is represented as an arrow pointing in the direction of the atom that accepts the electron pair (or in other words the arrow points away from the atom where the electrons came from (N in this case)). <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">

//<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Did you know that carbon monoxide also contains a dative covalent bond? Can you draw the Lewis structure for CO (carbon monoxide)? //

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">
 * <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Carbon has 4 electrons in its outer shell so it still needs another 4 electrons in order to fill its outershell. Oxygen on the other hand has 6 valence electrons so it needs just two other electrons in order to fill its outershell. Thus carbon and oxygen share two electron pairs so that the outershell of O is filled. However C will still need two more electrons, so oxygen shares one of its electron pairs with carbon forming a dative covalent bond.

=<span style="color: #362079; font-family: Tahoma,sans-serif; font-size: 14pt;">Bond length and bond strength =

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Atoms can form a single/double or triple covalent bond. So do you think that a single, double and triple bond are all going to have the same strength and all be of the same length?

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">No. The more the number of bonds, it means that more electrons are being shared and therefore the electrostatic forces of attraction between the atoms and the shared electron pair(s) increases.
 * <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Therefore the ** //more the number of bonds//, ** the ** shorter the bond length ** will be, because the two atoms are more strongly attracted to each other, and as has been said the stronger the bond strength.
 * <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">Therefore, a triple covalent bond will have a much stronger bond and much shorter bond length than a double or even a single covalent bond.

By the end of this lesson you should be able to:
 * Describe a covalent bond as a sharing of electrons.
 * Know that the covalent molecule is held together due to the electrostatic forces of attraction between the positive nuclei and the shared electron pair.
 * Use Lewis structures to represent covalent molecules.
 * Explain the relationship between the number of bonds, bond strength and bond length.
 * Predict whether a compound of two elements is covalent from the position of the elements in the periodic table, or from their electronegativity values.
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