COVALENT BONDING NOTES (Ch 7 Section 2)

Transcription

1 COVALENT BONDING NOTES (Ch 7 Section 2) I. Introduction: Remember from the last section we talked about how an ionic bond forms when one atom loses and electron and one atom gains an electron. Bonds can also form when atoms electrons. This type of bond is called a. A. Terms used when describing covalent molecules: Molecule: a group of atoms that are by covalent bonds. Molecular Substance: substances which are made of. Molecular Formula: formulas used to describe the make-up of a molecular compound. These formulas tell are in a single molecule of a compound. Examples: C 12 H 22 O 11 (sucrose or table sugar) You can also write empirical formulas for molecules. Remember the empirical formula represents the of atoms in a compound. Examples: Glucose C 6 H 12 O 6 (molecular formula) The empirical formula would be (divide all subscripts by 6) Lactic Acid C 3 H 6 O 3 ( molecular formula) The empirical formula would be (divide all subscripts by 3) From these two examples you can see that empirical formulas can be the for different compound. Molecular formulas can also sometimes be the for different compounds; therefore there is another type of formula, a. Structural Formula: shows the of atoms in a molecule. This formula shows which atoms are bonded and in what arrangement. The ones we will draw are called and are based on Lewis Dot Diagrams. The main difference is that multiple atoms are combined with dots showing electrons as well as electrons.

2 Drawing Lewis Structures: When drawing Lewis Structures we must remember the which tells us that atoms will share electrons so that each atoms in the compound has valence electrons. You must check each atom to make sure it has its eight electrons. Example: Draw the Lewis Structure for F draw the lewis dot diagram for each atom in the compound. 2. Put the atoms together in a way which allows both atom to have 8 electrons, including the shared electrons. When putting the pieces together, you can move the electrons around, as long as you keep at most 2 dots on a side (exception is multiple bonds). The electrons found between the atoms are called of electrons. The other pairs are called pairs. Example: Draw the Lewis Structure for NH draw the lewis dot diagrams for each atom. 2. Put atoms together. I can look at N and see that it has three spots which need another electron. I also have three H atoms which each need one more electron ( remember, H needs only 2 valence electrons to have a full set). Both of these examples showed only molecules with bonds which are bonds where share of electrons. ( in F 2, these is one single covalent bond between the two F atoms. In NH 3, there are three single covalent bonds found between the N and H atoms.

3 Double Covalent Bonds: when share of electrons. Example: CH 2 O (formaldehyde) 1. Draw lewis dot diagrams for atoms. 2. Put pieces together. Usually when C is in a comp[ound, it will be found as the central atom. So start by putting C in the center and put the other three pieces around C. The problem now is that C and O each still have an empty spot, where there is neither a shared nor an unshared pair. What happens is that O and C will shift these unshared electrons between the C and O (where there is already a pair). Now we will have two shared pairs and C and O will each have 8 valence electrons ( sharing 4, plus the other shared or unshared pairs), and each H has 2 valence electrons. Triple Covalent Bonds: when share of electrons. Example: C 2 H 2 (ethyne) 1. Draw lewis dot diagrams for atoms. 2. Put pieces together. Since this molecule has two C atoms, put them next together in the center. Put one H on each C.

4 The problem is that each C has two extra unshared electrons. These 4 electrons will shift between the two C atoms so now there will be three pairs between the C atoms which is a triple covalent bond. Each C has 8 valence electrons and each H has 2 valence electrons. Sometimes you will see dashes used instead of dots for the shared electron pairs. For now you will need to use the dots. There are also some elements which will be stable with less than 8 valnce electrons. B needs only 6 valence electrons ( as well as other is group 3A). II. Properties of Covalent Bonds: As we have said, a covalent bond results from the of electrons, but the atoms do not always share electrons fairly. Just like two children will rarely share a bag of M & M s equally, atoms also have trouble sharing electrons equally. The unequal sharing results from the difference between atoms. The atom with the electronegativity will pull harder on the shared electrons. Because of this, the atom pulling harder on the electrons will have a slight charge (since it has the negative electrons closer to its side of the molecule). The atom with the lower electronegativity will have a slight charge. This slight charge is indicated on the molecule using either an arrow or a lower case delta with a + or -. We will talk more about this later. When one atom is more elctronegative than another, this results in what we call a bond. When two atoms have similar electronegativities, each atom is pulling equally on the electrons and there is no + or region. This is called a bond. We can actually calculate the type bond which will result between two atoms by their electronegativities (see chart for electronegativities).

5 After finding the difference between the electronegativities, use this guide to determine the type of bond which will result. >= 2.0 bond covalent covalent covalent III. Naming Covalent Compounds: We will be looking at (or two part) compounds. Naming these compounds is similar to ionic compounds, but it uses prefixes to indicate the number of atoms in the molecule. Prefixes: Steps: 1. Name the first element in the compound, using a to indicate how many atoms. If there is only a single atom of the first element, use the prefix mono, just write the element name. 2. Name the second element using the and modify name ending with, like with ionic compounds. Example: CO 2 CCl 4

6 Exceptions: Sometimes we will shorten the prefix to make the name easier. For example, CO would be carbon monoxide (not carbon monooxide). Also sometimes use common names, like H 2 O is water not dihydrogen monoxide. Practice: IV. Writing Molecular Formulas for Covalent Compounds: To write the formula for a covalent compound, you just need to know the prefixes and rules we used to name covalent compounds. For example, carbon tetrafluoride would be Other Practice: