6.2 Lewis Structures

1. BIO 115-120 Week 2  -  Human Anatomy and Physiology  -  Franklin College Biology Announcements:  1. REMEMBER, YOUR FIRST HOMEWORK ASSIGNMENT IS DUE ON FRIDAY. SEE WEB PAGE FOR WEEK 1. 2. You have a quiz scheduled for Friday. It will cover all material from the first three lectures, your homework assignment, and many of the study questions on the web pages for weeks 1 and 2. TOPICS OF INTEREST This is Linus Pauling.  He is a two-time Nobel Prize winner, and one of the world's most famous chemists.  He made enormous contributions to our understanding of the processes of chemical bonding.  Later in his career he helped us to understand the function of complex protein chemistry.  As a point of interest, he also was one of the earliest and strongest advocates of using vitamin C as a preventative for the common cold. IONIC BONDSSpeaking of chemical bonding, the animation at right illustrates the concept of an ionic bond.  During ionic bonding, one atom gives up one (or more) electrons to another atom.  The donor atom becomes a positively charged "cation" and the recipient becomes a negatively charged "anion."  The two, oppositely charged atoms form a bond due to the force of attraction.  Salts such as Sodium Chloride result from the formation of ionic bonds.  The sodium atom gives up an electron (becoming sodium ion;  Na+ ) and the chlorine accepts an electron (becoming chloride ion;  Cl- ). COVALENT BONDSThe atoms at left are engaging in a covalent bond.  Notice that as the atoms apporach one another they begin sharing electrons.  The bond DOES NOT depend on the formation of ions.  The force of the bond is based on the result of SHARED electrons.  When two oxygen atoms combine to form an oxygen molecule, the atoms are held together with covalent bonds.  Most of the bonds in organic molecules are also covalent.  For example Carbon-Hydrogen (C-H) bonds are covalent and Carbon-Carbon (C-C) bonds are covalent. POLAR COVALENT BONDS 6.2 Lewis Structures

2. Lewis Structures-show arrangement of atoms and valence electrons Unshared pair or LONE pair of electrons Bonding pair of electrons The central idea is that stability of a compound relates to noble gas electron configuration.i.e. octet rule and duet rule

3. A strategy for drawing Lewis structures • Count all the valence electrons for the molecule. • Determine the central atom on the molecule. HOW? The central atom is often the first atom in the formula. Otherwise, choose the atom that forms the greatest number of bonds as a starting point. • Place other atoms around the central atom. • Draw a single bond to connect the central atom to the other atoms.

4. Strategy continued 5) Add the remaining electrons to the atoms to satisfy the octet rule (except for hydrogen which only needs 2 electrons). 6) Double check that you used the right number of electrons and the octet rule is followed. 7) If the central atom still does not achieve an octet, and it is expected to do so, move lone pairs from the outlying atoms to form a double or triple bond between atoms.

5. Helpful hints HONC rule: generally H bonds to 1 other atom (1 valence electron- needs 1 more) O bonds to 2 other atoms (has 6 valence electron-needs 2 more) N bonds to 3 other atoms (has 5 valence electrons- needs 3 more) C bonds to 4 other atoms (has 4 valence electrons-needs 4 more) Carbon, nitrogen and oxygencommonly form double bonds. Hydrogen, and the halogens usually share one pair of electrons (no double or triple bonds) Nitrogen and carbon can share three pairs of electrons to form a triple bond.

6. Drawing Lewis Structures • Draw Lewis Structures for the following compounds: • HCl NF3 • CH4 H2O • CO2 NO+ • Cl2 OH-

7. Lewis Structures • Draw Lewis Structures for the following compounds: • O2 N2 HCN • C2H4 C2H2 • Covalent bonds may be classified as single, double, or triple depending on the number of pairs of electrons shared.

8. Some molecules can be represented by more than one Lewis structure • They are called resonance structures. • Draw the structure for SO2 • The actual structure is an average of the resonance structures. • Experiments show that the two bonds are really of equal length and strength. The bond length is shorter than a single bond, but longer than a double bond.

9. Drawing Resonance Structures • Draw Lewis Structures for the following compounds • NO2-1 • SO3 • O3

10. 6.2 Writing Formulas & Naming Covalent Compounds • IONIC compounds are made from 2 ions (metal + nonmetal) • 2. COVALENT compounds are made from two or more non-metals. • ** Have different methods for writing formulas and naming.

11. 6.2 Naming Covalent Compounds • Nonmetal + nonmetals can form more than one compound with each other … • Ex) NO, NO2, NO5 • Cannot use Ionic rules … otherwise, all these would all have the same name • “Nitrogen Oxide”

12. 6.2 Writing Formulas & Naming Compounds Using Prefixes • Must use Prefixes to show how many atoms of each element are in the covalent compound. • Change second element to end in -ide • Examples: • NO2 – nitrogen dioxide • NO5 – nitrogen pentoxide

13. Using Prefixes 3. The last vowel is dropped when the element begins with a vowel (except “i”) Example – pentoxide not pentaoxide

14. Using Prefixes 4. “mono-” is NOT used on first element in formula (but is used on 2nd element) Example – NO – nitrogen monoxide not mononitrogen monoxide

15. Carbon dioxide CO2 N2O4 Dinitrogen tetroxide Sulfur trichloride SCl3

16. Naming Compounds Covalent Ionic (Non-metal + Non-metal) 1. Regular Metal UsePrefixes to show different # of atoms Ex) CO3 - Name 1st ion, Name 2nd ion with “ide” at end Ex) NaCl – Carbon trioxide Sodium Chloride 1 mono* 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca 2. Polyatomic Ion Name using Chart Ex) (NH4)(OH) – Ammonium Hydroxide 3. Transition Metal Find charge on metal, Place charge as Roman Numeral in name Ex) Fe2O3 – Iron (III) Oxide