1. Using Lewis Dot Structures to show Bonding

2. Remember Lewis Dot Structures?  Lewis dot structures are a way to represent the valence electrons for a particular atom.  Since the valence shell is where all the action happens during bonding, Lewis structures are a great way to help us predict how different elements will bond.  The key thing to remember is that reactions need to occur to reach the same valence arrangement as the nearest noble gas.

3. Ionic bonding  Remember that ionic bonding requires the transfer of electrons from a metal atom to a non-metal atom.  We can use Lewis dot structures to help identify how many electrons each component needs to gain or lose to reach its stable octet.  Let’s look at the reaction between sodium and chlorine as an example.

4.  Sodium has one valence electron, which it needs to lose for the atom to become stable.  Chlorine, on the other hand, has seven valence electrons, so it needs to gain one.  The dot structures give us a nice visual to help us understand what needs to happen.

5.  To show how the electrons will transfer, we simply circle the lone electron on the sodium and draw an arrow to show how it will transfer to the chlorine.

6.  Now that the sodium has lost an electron, it becomes positively charged.  The chlorine has gained an electron and so becomes negatively charged.  Notice that the charges on the two ions are equal and opposite, making a neutral combination when they come together.

7.  When ionic compounds form, there are two rules that will be observed:  1. The metals will lose all of their valence electrons and non-metals will gain enough to fill their outer shells.  2. When the ions come together to make their salt, they will always combine to create a neutral compound.

8. Covalent Bonding  Covalent compounds are a bit trickier to predict. Since the electrons are shared, not transferred, there are many more ways for non-metals to combine to meet the octet rule.  For our current lessons, we will focus on very simple combinations of two non-metals.  We can use the simple example of two hydrogen atoms bonding to each other as a starting point.

9.  Each hydrogen atom has a single valence electron.  Since it is one of the small elements, it only needs to get to two electrons to become stable.  So, by sharing their lone electron with each other, they both reach the same electron configuration as helium, the smallest noble gas.  We show the sharing of electrons by drawing a single large oval around both of the valence electrons.

10. Lewis Structures for covalent compounds  Once covalent compounds have formed, we have two common ways to represent them.  One is to place the shared pair or pairs of electrons between the elemental symbols.  The second is to draw a single line between the elemental symbols for each pair of shared electrons.

11.  Since each atom normally needs to complete its octet, it is possible that they will need to share more than a single pair of electrons.  Good examples of this are two more diatomic elements, oxygen (O 2 ) and nitrogen (N 2 ).  Since oxygen atoms contain six valence electrons, they need to gain or share two to reach their octet. When two oxygen atoms come together, they share both of their lone electrons, creating two shared pairs.  In the same way, nitrogen atoms need three more valence electrons, so they share all three, creating three shared pairs.