1. Chemical Bonds

2. Forming Chemical Bonds  The force that holds two atoms together is called a chemical bond.  The valence electrons are the electrons involved in forming chemical bonds.  Elements tend to react to acquire eight electrons. This is call a stable octet.  Noble gases (group VIIIA/18) have this structure (octet) and are inert (does not form bonds).  Atoms can gain, lose, or share electrons to reach an octet.

3. Forming Ions  Positive ions (cations) are formed when atoms lose one or more valence electrons. [Usually these atoms are metals.] Reactivity of metals (cations) are based on how easily they lose electrons. [ Ionization energy]  Negative ions (anions) are formed when atoms gain one or more valence electrons. [These atoms tend to be non-metals.] Electronegativity is the ability to attract or gain electrons.

5. Forming Ionic Bonds  Ionic bonds: Complete transfer of electrons between atoms (difference of electronegativity of 1.6 or greater).  Two neutral atoms will form ions. The resulting compound is called an ionic compound.

6. Properties of Ionic Compounds  Form crystal lattice bonding structure.  Example: NaCl (sodium chloride)  High melting point and boiling point due to strong electrostatic charge between the atoms (cation is attracted to anion after the transfer of electrons).  Hard, rigid and brittle solids.  Conducts electricity in liquid state or dissolved in water only.

7. Forming Covalent Bonds  Covalent bonds: Sharing of electrons between atoms (difference of electronegativity of less than 1.6).  Occurs usually between elements close to each other on the periodic table (mostly non- metals).  The resulting compound is called a molecule.

8. Properties of Covalent Compounds  Have definite and predictable shapes.  Low melting and boiling points.  Relatively soft solids.  Can exist as solids, liquids or gases.

9. ioniccovalent valence electrons Comparison of Bonding Types sharing of electrons transfer of electrons ions molecules  EN > 1.6  EN < 1.6 high mp low mp molten salts conductive non-conductive

10. Determining Number of Covalent Bonds  To determine how many bonds exist in a molecule, use the following formula N - A = # bonds 2 -- Where N is the # of needed electrons, which is 8 for all elements but H, which is 2. -- Where A is the # of available electrons, which is the number of valence electrons.

11. Lewis Dot Structures  A system of “drawing” bonds  Shows how valence electrons are arranged  Dots represent valence electrons  Pairs of dots represent bonding pairs of electrons.  Show the reaction between sodium and chlorine:  Show the reaction between calcium and bromine:

12. Drawing Lewis Structures 1.Determine the # of bonds need with the formula 2.Draw a skeleton structure. The least electronegative element is the central atom. 3.Connect the atoms with the # of bonds determined in step 1. 4.Finish by making sure all atoms in the structure have an octet of electrons. 5.Remember, if the structure has a positive charge, it has fewer available electrons, and if it has a negative charge it has more available electrons.

13. Lewis Structure Practice  Draw Lewis Dot Structures for the following molecules and polyatomics: CO 2 N 2 O 2 Cl 2 H 2 O SO 4 -2

14. Exceptions to the Octet Rule  There are two elements which don’t want an octet of electrons and those are Be (wants 4) and B (wants 6).  Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule using their empty valence d orbitals, especially when surrounded by highly electronegative atoms such as chlorine, bromine, and oxygen.

15. Exceptions to the Octet Rule  The exceptions are easy to recognize-- when you calculate the # of bonds, the formula will give you an answer that makes no sense. If this happens, do the following: 1. Connect all of the atoms to the central element with one bond. 2. Give all of the surrounding elements an octet of electrons. 3. Count the number of electrons you have put in the structure so far. - If it equals the available # of electrons you calculated, you’re done. - If it is less than than the available # of electrons you calculates, place the needed number of electrons around the central element.

16. Exceptions to the Octet Rule  Draw Lewis Dot Structures for the following molecules: BF 3 PCl 5 SF 6 XeF 2

17. Carbon Bonding  Carbon follows the rules previously discussed, but it has a few unique properties. Carbon is always the central atom in its compounds and it tends to undergo bonding to itself. This tendency of carbon to bond to itself is called catenation. Carbon always has four bonds. Anytime you have a choice for a skeleton structure, always go for the most symmetrical option.

18. Carbon Bonding  Draw Lewis Dot Structures for the following molecules: C 2 H 6 C 3 H 8 C 2 H 2

19. Let’s Practice! H2H2  SO 2  CCl 4  SiH 4  PCl 5  XeCl 6 H2OH2O  SO 4 -2  NO 3 - C2H3O2-C2H3O2-

20. Molecular Shape  Valence Share Electron-Pair Repulsion (VSEPR) model allows us to predict the molecular shape by assuming that the repulsive forces of electron pairs cause them to be as far apart as possible from each other. Only the valence electron pairs are considered in determining the geometry.

21. Effect of the number of electron pairs around the central atom 2 charge clouds, linear 3 charge clouds, trigonal planar 4 charge clouds, tetrahedral 5 charge clouds, trigonal bipyramidal 6 charge clouds, Octahedral

22. PREDICTING EXPECTED GEOMETRY ACCORDING TO VSEPR THEORY  Lewis dot structure determines the total # of electrons around the central atom. Multiple bonds (double and triple) count as one.  The number of bonding and nonbonding electron pairs around the central atom determines the geometry of electron pairs and the molecular geometry.  Lone e  pairs affect geometry more than bonding pairs.  Multiple bonds have larger affect on geometry than single bonds: H 2 C=O (116° instead of 120°); H 2 C=CH 2 (117° instead of 120°).

23. Molecular Shapes 2,3,4 Electron Pairs

24. Molecular Shapes 5, 6 Electron Pairs