1. Ionic Nomenclature Recap ▪ Cation is ALWAYS written first. – If monoatomic, use the name of the element. – If polyatomic, use the name of the polyatomic ion. ▪ Anion is ALWAYS written second. – If monoatomic, use -ide as the suffix. – If polyatomic, use the name of the polyatomic ion. ▪ Examples: – NaCl = Sodium Chloride – NaOH = Sodium Hydroxide – NH4NO3 = Ammonium Nitrate

2. Recap from last semester… 3 types of bonding: Ionic Metallic Covalent

3. Covalent Bonding

4. Covalent Bonds ▪ What happens to electrons in covalent bond? – Electrons shared between atoms, not given or taken completely. – Atomic orbitals are combined to form molecular orbitals. ▪ Generally bond between nonmetal and nonmetal. – Why? 3 trends – Electron affinity, Ionization energy, and electronegativity ▪ Differences between atoms are small. ▪ One atom cannot completely take away electrons from the other ▪ Example: CH 4

5. ▪ Molecule—a neutral group of atoms that are held together by covalent bonds. ▪ Molecular formula—shows the types and numbers of atoms combined in a single molecule of a compound. – Example: CH 4

6. Formation of Covalent Bond ▪ Nature favors chemical bonding…why? – makes the atoms more stable. – How? Potential energy is lowered when the atoms are bonded. ▪ Lower energy = more stable. Negatively charged electrons interact with the positively charged nuclei. Lowest potential energy = H 2 molecule formed repulsion forces = attraction forces

7. Characteristics of a Covalent Bond ▪ Bond length—the average distance between two bonded atoms. – Depends on type of bond: single, double, or triple. Bond Energy—The energy required to make or break a chemical bond. It takes positive energy to break a bond, and negative energy to make a bond. Example: H—H +436 kJ is required to break H—H -436 kJ is required to make H—H Negative energy comes from the lowering of the potential energy when a bond is formed.

8. Octet Rule Atoms undergo bonding in order to satisfy the octet rule. Octet rule: Chemical compounds tend to form so that each atom has 8 electrons, either by gaining, losing, or sharing electrons. Atoms want to become “Noble-gas-like” with filled valence shells.

9. Exceptions to Octet Rule Hydrogen—forms only one bond to have two valence electrons. Group 13—Has three valence electrons. Tends to form three bonds. Some elements can form an expanded octet if bound to highly electronegative atoms. Example: SF 6 Expanded octet involves empty d orbitals to fit extra electrons.

10. Diatomic Molecules Diatomic molecule: A molecule in which there are only two atoms. All of the following elements can exist in nature as diatomic molecules. F 2 Cl 2 Br 2 I 2 H 2 O 2 N 2 Bonding can be shown through electron configurations.

11. Lewis Structures ▪ What are they?? – A formula where atomic symbols represent nuclei and inner shell electrons, and dot pairs represent valence and bonded electrons. ▪ What are they used for?? – Gives us a way to visualize bonding between atoms – Represents where electrons are located – Gives relative bond strengths to establish reactivity of molecules.

12. Lewis Structures ▪ Six Steps: 1.) Determine types of atoms in molecule. ▪ Example: CH3I 2.) Write electron dot notation for each atom. 3.) Determine the total number of valence electros available. 4.) Arrange atoms with LEAST electronegative atom in the center (exception: H), and place one shared pair of electrons between each of the atoms. 5.) Fill in valence shells of atoms with unshared electrons (lone pairs). 6.) Count electrons to make sure all available valence electrons are accounted for.

13. Practice with Lewis Structures ▪ Draw the Lewis structures for the following molecules: – NH 3 – H 2 S – SiH 4 – PF 3

14. Before we go any further… Try the Lewis structure for C2H4.

15. Single and Multiple Covalent Bonds – Single bond – two electrons shared ▪ Also called sigma bond, or  bond. ▪ Examples: H 2 – Double bond – four electrons shared ▪ Also called pi bond, or  bond. ▪ Example: O 2 – Triple bond – six electrons shared ▪ Consists of one  and one  bond. ▪ Example: N 2

16. Lewis structures with multiple bonds ▪ Multiple bonds become evident in lewis structures when there are not enough valence electrons after adding lone pairs. ▪ Examples: – CH 2 O – CO 2 – HCN

17. Lewis Structures of polyatomic ions ▪ Follow the 6 rules for writing lewis structures to make sure everything has a full octet. – Add extra electrons to most electronegative element. – If charge is positive, subtract the number of electrons specified by the charge. ▪ CO 3 2- ▪ NH4 -

18. Relative Bond Lengths and Strengths ▪ Bond length—the more shared pairs = the shorter the bond. – Single > Double > Triple ▪ Bond Strength—the more shared pairs = the stronger the bond. – Triple > Double > Single

19. Resonance Structures ▪ In molecules with multiple bonds, electrons are delocalized due to shared orbitals. ▪ Since electrons are delocalized, a single lewis structure cannot account for all of the possible locations of electrons. ▪ Most stable resonance forms have least amount of charges.

20. Non-polar and Polar Covalent Bonds ▪ Dipole—a molecule that contains an unequal sharing of electrons. – Dipole is represented by arrow pointing in the direction of the negatively charged region. – Created by a difference in electronegativity between bonded atoms. ▪ Nonpolar bonds do not contain a dipole.

21. Polar Covalent Bonds

22. Identify the bond as polar or nonpolar ▪ C—H ▪ O—Cl ▪ C—Cl ▪ H—N ▪ B—F ▪ F—F

23. VSEPR Theory ▪ What is it? – Valence shell electron pair repulsion theory – Helps us predict shapes of molecules – Shapes are determined by number of lone pairs and surrounding atoms on central atom.

24. Examples: CO 2 HCN C 2 H 2

25. Examples: BF 3 CH 2 O SO 2

26. Examples: CH 4 NH 3 H 2 O

27. Example: PCl 5

28. Example: SF 6

29. Identifying a polar or nonpolar molecule ▪ Draw Lewis structure based on VSEPR theory. ▪ Determine whether the molecule has polar bonds. ▪ Determine the dipole of each bond. – If dipoles all point to one general direction, the molecule is polar. – If dipoles point in opposite directions, molecule is nonpolar. – If no dipole, the overall molecule is nonpolar.

30. Determine if the following molecules are polar or nonpolar: ▪ NH 3 ▪ CH 3 F ▪ SF 6 ▪ AlCl 3 ▪ CCl 4

31. Covalent Compounds—Nomenclature ▪ Rules—Binary Compounds – Element with smaller group # is always given first (similar to cation in ionic bonding). – Second element combines prefix with suffix –ide. – If second element begins with vowel, the –o- or –a- in the prefix is dropped. ▪ Example—pentoxide Prefixes: Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10 Practice: N 2 O As 2 O 5 Carbon Tetrafluoride CO Sulfur Trioxide

32. Acids 2 types: Binary Acids—contain H and another element, usually a halogen. Put hydro- as prefix for H. Use –ic as suffix for second element. Example—HCl = hydrochloric acid Oxyacids—contain H, O, and a third element (mostly H paired with a polyatomic ion). Nomenclature based on third element or polyatomic ion. See list of common acids on page 230 in book.

33. Hydrocarbons ▪ Hydrocarbons = the simplest organic compounds – Contain only carbon and hydrogen – Can be straight-chain, branched chain, or cyclic molecules ▪ 3 types of straight-chain hydrocarbons – Alkanes—completely saturated hydrocarbons (no double or triple bonds) ▪ Go by the formula C n H 2n+2 ▪ End in suffix -ane – Alkenes—contain a double bond ▪ End in suffix -ene – Alkynes—contain a triple bond ▪ End in suffix -yne

34. Need to know first ten alkane hydrocarbons… ▪ CH 4 Methane ▪ C 2 H 6 Ethane ▪ C 3 H 8 Propane ▪ C 4 H 10 Butane ▪ C 5 H 12 Pentane ▪ C 6 H 14 Hexane ▪ C 7 H 16 Heptane ▪ C 8 H 18 Octane ▪ C 9 H 20 Nonane ▪ C 10 H 22 Decane