Bonding

1. Bonding Ch 8

2. Objectives • SWBAT identify the bond type of a molecule by using electronegativity differences. • SWBAT identify the polarity in small molecules.

3. Types of Bonding • Ionic • Covalent • Metallic

4. Octet Rule • Atoms tend to gain, lose or share electrons until they are surrounded by 8 valence electrons.

5. Ionic bonding • Transfer of electrons from the metal to the non-metal. • Formation of cations and anions Na + Cl Na+ + Cl-

6. Ionic Bonding • When an extremely electronegative atom, like fluorine, bonds with an electropositive atom, like sodium, the resulting bond is ionic due to the huge difference in electronegativity (difference > ~1.7).

7. Ionic Bonding • The electronegative atom's pull on the bonding electrons is so strong that it pulls the bonding electron off the electropositive atom resulting in two oppositely charged ions which are held together by electrostatic attraction (an ionic bond).

8. Energetics of Ionic Bond Formation • If you look in Appendix C (∆Hf values) you will find that the heat of formation for ionic compounds is exothermic. • The removal (or loss) of electrons is always an endothermic process. • When a non-metal gains an electron the process is generally exothermic.

9. Lattice Energy • The principal reason that ionic compounds are stable is the attraction between ions of unlike charge. • Lattice energy – is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions.

10. Lattice Energy Problems • Look at the example on B&L page 267 • We will go through the example as a group. • Try problems # 19 and 20 at the end of the chapter.

11. Covalent Bonding • Bonding between non-metals consists of two electrons shared between two atoms. • In covalent bonding, the two electrons shared by the atoms are attracted to the nucleus of both atoms. Neither atom completely loses or gains electrons as in ionic bonding. • http://www.elmhurst.edu/~chm/vchembook/152Apolar.html

12. Types of Covalent Bonding • Polar Covalent Bonding results when two different non-metals unequally share electrons between them. • Non-Polar Covalent Bonding results when two identical non-metals equally share electrons between them.

13. Non-polar covalent bonds • Identical non-metallic atoms have identical electronegativity • For example in an H-H bond both H atoms have the same electronegativity so the bond is non-polar • Different atoms can have the same electronegativity, such as, N-Cl. • Both N and Cl have an electronegativity of 3.0 so the bonding electrons will be shared equally between the two atoms resulting in a non-polar covalent bond.

14. Polar Covalent Bonds • When atoms of similar, but different, electronegativities (a difference < ~1.7) bond, the more electronegative atom has a greater share of the bonding electrons than the less electronegative atom. • The more electronegative atom has a partial negative charge, and the less electronegative atom has a partial positive charge. The resulting covalent bond is called a polar covalent bond.

15. Show the Class • Show how to draw a partial positive and partial negative charge symbol.

16. Multiple Bonds • The length of the bond decreases with increasing numbers of bonds.

17. Comparison of Bonding Types

18. Compare Ionic, Polar and Non-Polar Bonding

19. Electronegativity • Electronegativity is the relative tendency of a bonded atom to attract electrons to itself. • An atom with extremely low electronegativity, like a Group I metal, is said to be electropositive since its tendency is to lose rather than to gain, or attract, electrons. • Non-metals are more electronegative than metals. • http://www.ausetute.com.au/bondpola.html

20. Electronegativity Values and Bond Types • Electronegativity values are useful in determining if a bond is to be classified as non-polar covalent, polar covalent or ionic. • What you should do is look only at the two atoms in a given bond. Calculate the difference between their electronegativity values. Only the absolute difference is important. • http://dbhs.wvusd.k12.ca.us/webdocs/Bonding/Electroneg-Bond-Polarity.html

21. Electronegativity and Associated Bond Type < 0.5 non-polar covalent 0.5 – 1.9 polar covalent ** 2.0 or greater ionic Note: To find the electronegativity values of many common elements, look at the chart printed on the bottom of your orbital diagram. ** there are some exceptions

22. Calculating Electronegativity Values • Calculate the electronegativity and determine the bond type of an O-H bond • O has an electronegativity of 3.5 • H has an electronegativity of 2.1 • The difference in electronegativity is: 3.5 - 2.1 = 1.4 1.4 is less than 1.7, so the resulting bond is polar covalent.

23. AP Students Use This Slide Here are the rules: • If the electronegativity difference (usually called ΔEN) is less than 0.5, then the bond is nonpolar covalent. • If the ΔEN is between 0.5 and 1.6, the bond is considered polar covalent • If the ΔEN is greater than 2.0, then the bond is ionic. That, of course, leaves us with a problem. What about the gap between 1.6 and 2.0? So, rule #4 is: 4. If the ΔEN is between 1.6 and 2.0 and if a metal is involved, then the bond is considered ionic. If only nonmetals are involved, the bond is considered polar covalent. http://dbhs.wvusd.k12.ca.us/webdocs/Bonding/Electroneg-Bond-Polarity.html

24. AP Example Sodium bromide (formula = NaBr; ENNa = 0.9, ENBr = 2.8) has a DEN = 1.9 Hydrogen fluoride (formula = HF; ENH = 2.1, ENF = 4.0) has the same DEN. We use rule #4 to decide that NaBr has ionic bonds and that HF has a polar covalent bond in each HF molecule.

25. Electronegativity Trend • Electronegativity decreases down a Group in the Periodic Table as the atomic radius and number of inner electron shells both increase. • Electronegativity increases across a Period of the Periodic Table, in general, due to increasing nuclear charge and decreasing atomic radius. • For the commonly encountered atoms in high school science, the order in decreasing electronegativity is:F > O > N ~ Cl > Br > C ~ S ~ I > P ~ H > Si

26. Metallic Bonding • Metallic bonding is characterized by a “sea of electrons.”

27. Dipole • Polarity results from the uneven partial charge distribution between various atoms in a compound. • Atoms, such as nitrogen, oxygen, and halogens, that are more electronegative have a tendency to have partial negative charges. • Atoms, such as carbon and hydrogen, have a tendency to be more neutral or have partial positive charges.

28. Dipole • Electrons in a polar covalent bond are unequally shared between the two bonded atoms, which results in partial positive and negative charges. The separation of the partial charges creates a dipole. • The word dipole means two poles: the separated partial positive and negative charges. A polar molecule results when a molecule contains polar bonds in an unsymmetrical arrangement.

29. Dipole • Non-polar molecules are of two types. • Molecules whose atoms have equal or nearly equal electronegativities have zero or very small dipole moments. • A second type of nonpolar molecule has polar bonds, but the molecular geometry is symmetrical allowing the bond dipoles to cancel each other.

30. Dipole Moment • The dipole moment is a measure of the unevenness, or lack of symmetry, of the charge distribution in a molecule. • The mathematical definition of the dipole moment involves adding up the size of each charge in the molecule multiplied by the average distance that charge is from an arbitrary origin. • www.chem.unsw.edu.au/.../dipolemoments.html

31. Dipole Moment • symbol for dipole moment →

32. Formal Charge • See page 280 in the “old” B & L text http://www.chemprofessor.com/bonding_files/image033.jpg