Presentation on theme: "Unit 8: Covalent Bonding. But first … Stretch your mind back to the last chapter… What is an ionic bond?"— Presentation transcript:
Unit 8: Covalent Bonding
But first … Stretch your mind back to the last chapter… What is an ionic bond?
An ionic bond is the force of attraction between two oppositely charged ions. + -
Which of these are not ionic compounds? MgCl 2 Al 2 O 3 SCl 2 K 2 S CO 2
SCl 2 can’t be an ionic compound because sulfur and chlorine are both non-metals: they both need to gain extra electrons to become stable. Remember, for an ionic bond to form, you need an electron giver and an electron taker.
Draw dot diagrams to show how one S atom and 2 Cl atoms can share electrons so they each get that octet that all atoms want.
Cl S Cl In a covalent compound, the shared electrons are counted by both atoms as part of their octets.
That’s what a covalent bond is: the bond that results when atoms share electrons. Sharing = Cooperating: When you see the “co” in covalent, think of cooperating and sharing.
How can you recognize a covalent compound when you see it? Hint: Look at these formulas again. MgCl 2 Al 2 O 3 SCl 2 K 2 S CO 2
Properties of Ionic Bonds 1.Most will dissolve in water 2.Will conduct electricity when liquid (molten) or when dissolved in water 3.Brittle 4.Solids at room temperature 5.High Melting and Boiling Points 6.Made of positive and negative ions 7.Electrons are localized on ions
Properties of Metallic Compounds 1.WILL NOT dissolve in water 2.Will conduct electricity when solid 3.Malleable and Ductile 4.Most Solids at well above room temperature 5.High Melting and Boiling Points 6.Made of Positive Ions and delocalized electrons
Properties of Covalent Compounds 1.Generally WILL NOT dissolve in water 1.Small molecules with OH will dissolve in water – called Hydrogen Bonding, we will discuss this later. 2.DO NOT conduct electricity 3.Most have low melting and Boiling Points 4.The larger the molecule the greater the chance it will be a solid at Room Temp 5.Made of Neutral Nonmetals and localized shared electrons
Drawing Lewis Structures
Procedure Step 1: Count the valence electrons for all the atoms in the compound, and add them up. Example: To draw the structure of H 2 O, you’d add: H: 1 valence electron each x 2 H’s O: 6 valence electrons Total: 8 valence electrons
Procedure Step 2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond. Our H 2 O example: H – O - H
Logical question at this point… How do you know which atoms are attached to which? Hints: a. Hydrogen can form only one bond, so it’s always on an end, never in the middle. b. Formulas are often written to show the order the atoms go in, so HCN is attached this way: H – C – N c. When a central atom has group of other atoms bonded to it, the central atom is usually written first, so CO 3 would be: O – C – O O
Procedure Step 3: Complete the octets of the atoms bonded to the central atom. (Remember, of course, that hydrogen can have only two electrons.) Each bond you already drew represents two electrons for each atom it’s connecting. So, for our H 2 O example, we’ve already completed this step. H – O - H
Procedure Step 4: Place any leftover electrons on the central atom, even if doing so results in more than an octet. If the central atom has 8 electrons, you’re done. With water, we are done: the oxygen atom has 4 unshared electrons, plus two covalent bonds, which count as 2 electrons each.
Procedure Step 5: If there are not enough leftover electrons to give the central atom an octet, try multiple bonds until it does have an octet. Use one or more pairs of unshared electrons on the side atoms to form an extra pair of shared electrons.
Let’s try another example Carbon dioxide, CO 2. Step 1: Count the valence electrons you have to work with. ( = 16) Step 2: Attach the atoms in a logical way: O – C – O You’ve now used 4 of your 14 available valence electrons, because each single bond represents 2 shared electrons. So there are 10 electrons left to use.
CO 2 example, continued Step 3: Complete the octets of the outer electrons. You’ll now have used all 16 valence electrons you have to work with. However, the carbon atom only has 4 electrons (two from each single bond).
CO 2 example, continued Step 4: Since the C atom doesn’t have an octet, move the unshared electrons on the oxygen atoms until it does: Now, all three atoms are surrounded by 8 valence electrons, so we’re done!
Draw a dot diagram showing how ammonia, NH 3, is a covalent compound.
H N H H
More practice! 1. PCl 3 2. H 2 3. HCN 4. phosphate ion 5. SF 2 6. CO 7. chlorite ion
New definition: A molecule consists of two or more atoms bonded covalently.
Identify these chemicals as molecules, compounds, neither, or both: a. CH 4 b. MgS c. I 2 d. CO 2 e. Hef. Fe 2 O 3
Hydrogen gas and chlorine gas both exist as diatomic molecules. “diatomic” means “2 atoms” Draw dot diagrams of H 2 and Cl 2.
H H Cl Cl These single bonds are also known as sigma bonds. (Both words start with “si”)
Oxygen also exists as a diatomic molecule: Draw a dot diagram for O 2.
The O 2 molecule is an example of a double bond: two pairs of electrons are shared. The second bond in a double bond is called a pi bond.
N 2 is another diatomic molecule. Use dot diagrams to see what type of bond is in N 2. N
Another definition: Bond length: the distance between the two bonding nuclei
Single bonds are the longest, double bonds are in the middle, triple bonds are the shortest. Single bond Double bond Triple bond
The shorter the bond is, the stronger it is. Triple bond is the strongest, so it’s the hardest to break.
The shorter the bond is, the stronger it is. Triple bond is the strongest, so it’s the hardest to break. In other words, it takes more energy to break a stronger bond.
One last definition: Bond dissociation energy: the energy required to break a specific covalent bond. Breaking a bond always requires energy.