1. Chemical Bonding And Molecular Geometry

2. What is a chemical bond? - mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.  Individual atoms have high potential energy.  Atoms are less stable by self than when combined.  Chemical bonding minimizes the potential energy.

3. Chemical Bonding

4. Ionic Bond versus Covalent Bond Ionic Bond – chemical bonding that results from the electrical attraction between a large number of cations and anions. Covalent Bond – bonding resulting from the sharing of electron pairs between two atoms.

5. Electronegativity and its usefulness  Defined as the measure of the ability of an atom to attract electrons  Can be used to determine ionic or covalent bond. If the difference in electronegativity between two atoms is:  Greater than 1.7 (>50% ionic)Ionic bond  1.7 or Less (50% ionic or less)Covalent bond  1.7 to 0.4Polar Covalent bond  0.0 to 0.3 (0-5% ionic)Nonpolar Covalent bond

6. Electronegativity and Bond Type

7. Electronegativity of the Elements

8. Types of Covalent Bonds Nonpolar Covalent bond – a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge.  Bonding between two atoms of the same element is purely covalent; for example H 2  Polar – an uneven distribution of charge.  Polar Covalent bond –a covalent bond which contains bonded atoms that have an unequal attraction for the shared electrons.  element with high electronegativity contains a partial negative charge.  element with the lower electronegativity contains a partial positive charge  Dipole Moment – a center of positive charge and negative charge on a molecule

9. 3 Possible types of bonds

10. A polar molecule - Water

11. The effects of Polarity Negative side of water attracted to Positive ion. Positive side of water attracted to Negative ion.

12. How does Covalent Bonding take place?  Electrons and Nuclei are attracted which results in a decrease in potential energy.  At the same time, two nuclei repel and electron clouds repel leading to an increase in potential energy.  when attractions are equal to repulsions and optimal distance is reached the potential energy is at a minimum resulting in a stable molecule.  as bond forms potential energy is released to obtain stable molecule and minimum potential energy.  as bond breaks energy is needed to drive this to happen.  through covalent bonds noble gas configurations are obtained.

13. How Covalent Bonding takes place

14. More Information on Covalent Bonds  Bond length – distance between two bonded atoms at their minimum Potential energy is the average distance between two bonded atoms.  Bond Energy – energy required to break a chemical bond and form neutral isolated atoms. Reported in kilojoules/mole.  Bond Order- determined by the type of bonds present in a molecule.

15. The Relationship between Order, Length, and Energy Type of Bond Bond Bond Bond Order Length Energy Single 1 long low Double 2 medium medium Triple 3 short high

16. Ionic Bonding and Stable e - Configurations

17. Common Charges on Ions

18. Ionic compounds at the Molecular Level (Formula Unit)

19. Ionic Size Ionic Radii –The radius exhibited by an ion in an ionic crystal where the ions are packed together to a point where their outermost electronic orbitals are in contact with each other.  Cation = positive ion --> decrease in size  Anion = negative ion --> increase in size. Period Trends  Cations – decrease across period  Anions – larger in size than cations decrease across period. Group Trends – increase down a group.  Anions – ionic radii greater than atomic radii  Cations – ionic radii lesser than atomic radii

20. Cations and Anions