CP Covalent Bonds Ch. 8 &

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1 CP Covalent Bonds Ch. 8 & Why do atoms bond? Atoms want stability- to achieve a noble gas configuration ( ) For bonds there is a transfer of electrons to get an octet of electrons For covalent bonds there is a of electrons to get an octet What is a covalent bond? Covalent bond - is the chemical bond that results from sharing of electrons Occurs with elements to each other on the periodic table Between a nonmetal and a nonmetal Molecule is two or more atoms are bonded Examples of Molecules F 2 H 2O 3 Pencil Demo NH 3 ( ) CH 4 (methane) Observations: Notice there are no, only non-metals Diatomic molecules Some atoms do not exist as a atom Atoms that exist as two: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2 HONClBrIF Magnificent -don t forget H 3 Types of Covalent Bonds: Single, Double and Triple Strength of Covalent Bond Several factors control bond strength Number of shared electrons-the electrons shared, greater the bond strength of the atom Single Covalent Bonds Each atom shares one (2 total) of electrons bond Weakest bond of the three Double Covalent Bonds In a double bond, each atom shares pairs (4 total) of electrons Medium length bond Medium bond Triple Covalent Bond In a triple covalent bond, each atom shares three pairs (6 total) of electrons The shortest bond The strongest bond Carbon and nitrogen can form covalent bonds Covalent Molecule Properties Covalent molecular solids tend to be solids, liquids, or gases at room temperature melting and boiling points conductors of heat and electricity Non-electrolytes do not conduct electricity in water 1

2 Naming Binary Molecular Compounds 3 Rules to Name 1. Name the first element using the name 2. Second element in the formula- use the root word and in ide ex: Oxygen Oxide Sulfur Sulfide Hydrogen 3. Add a prefix to both words to indicate the number of atoms 1. mono- 6. hexa- 2. di- 7. hepta- 3. tri- 8. octa- 4. tetra- 9. nona- 5. penta- 10. deca- 2 Exceptions to the rule 1. When the formula contains one atom of the element, omit mono ex: CO 2 Carbon dioxide, not Dioxide 2. Drop the final letter in the prefix ending in a or o if the element begins with a ex: CO Carbon monoxide, not Carbon Example 1 Example 2 Example 3 Example 4 S 4N 2 SO 3 P 4S 5 CO Example 5 NH 3 common name: Example 6 CH 4 common name: Example 7 Example 8 Example 9 As 2O 3 N 2O 5 P 4O 10 2

3 Lewis Structures: Electron dot structures for molecules 5 rules for Lewis Structures 1. Find the total # of electrons for the molecule 2. Find the center atom (the element with the # of atoms) 3. Draw bonds. Connect the other atoms to the center atom. Then subtract electrons from the total # of valence electrons for each bond drawn. 4. Distribute electrons around each atom to give a total of electrons except H, Al, B, & Be 5. If there are not enough to give 8 around each atom, create & triple bonds. Neutral Molecules Ex 1: CF 4 Ex 2: NH 3 Ex 3: H 2S Ions: With Ions we add or take away. Example 2: O 2 2- Double Bonds: Put ions in and the charge on the outside Ex 3: CO 2 Triple Bond: Try doubling or tripling the bonds if you do not have an. Example 8: CO Resonance: Try drawing NO 3-1. Does your diagram look like the model A, B or C? Incomplete Octet: Boron does not have 8e- around it. It is but it is okay; it is one of our exceptions. What elements are the exceptions to the octet rule? Example 6: BCl 3 2. Are the other diagrams wrong? Discuss this with your neighbor why or why not? 3. Do you think there are more molecules that have more than one valid structure like NO 3-? 4. Do you know what molecule this is? 3

4 NT: VSEPR- V S E P R Lewis Structures are, whereas VSEPR Molecules are VSEPR predicts the or of the molecule pairs of electrons influence the shape by pushing other atoms as far apart from each other as possible VSEPR NT Molecule Lewis Structure Number of (shared) bonding pairs of electrons H 2 Number of lone pairs of electrons on central atom Total Number of electron pairs Molecular Shapes (look at lone & bond pairs) Linear VSEPR Model (3-D Model) Bond Angle Polar or nonpolar molecule BeH 2 Linear BH 3 Trigonal planar CH 4 Tetrahedral NH 3 Trigonal Pyrimidal H 2O Bent 4

5 Electronegativity and Polarity Types of covalent bonds Non-polar covalent- sharing of electrons Polar covalent- unequal of electrons Non-polar covalent Polar covalent Sharing of electrons Usually occurs when two atoms are bonded together. Examples: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2 Polar covalent bond occurs when there is an sharing of electrons Unequal sharing caused by 2 elements with different (different abilities to attract electrons). The bond is called a dipole (two poles) Creates molecule with charges Polar covalent continued... Polar molecule or not? Partial charges symbolized by (delta) + and - The electronegative atom is located at the partially negative end Example: + - H-Cl or H-Cl The shape of a molecule tells if a molecule is polar or not If the VSEPR shape is symmetric it is usually If the molecule is asymmetric it is polar Name Polar Example Non-polar example Linear HCl Trigonal Planar AlH 3 Tetrahedral CH 3OH Trigonal pyramidal Always Bent always 5

6 Intermolecular Forces The force that exists between molecules. This force attracts molecules to each other. Intermolecular Forces (cont.) IMF s 3 types of forces Dispersion force or induced dipole moment between molecules; only force in non-polar molecules; force (caused by the motion of electrons) (ex: CH 4) Dipole-dipole the force between two molecules; stronger force (ex: HCl) Hydrogen bond the intermolecular force forms between hydrogen end of one dipole and fluorine, oxygen, or nitrogen (that have at least one lone pair) end of another dipole; force (ex: H 2O) molecules have dispersion forces All polar molecules have dipole-dipole forces and dispersion forces Molecules that H-bond have all 3 Solubility of polar molecules Properties are due to intermolecular forces Like dissolves like Polar substances will dissolve molecules (and ionic compounds) Non-polar substances will dissolve non-polar molecules 6