Chapter 5 Chemical Bonding

1. Chapter 5Chemical Bonding

2. Poison or Seasoning? • How can two poisons (elemental sodium and elemental chlorine) combine to form a flavor enhancer (sodium chloride) that tastes great on steak? • Answer: By an exchange of electrons that stabilizes both atoms—the formation of a chemical bond.

3. G. N. Lewis and Bonding • Lewis theory fundamental ideas: • Valence electrons are the most important. • Two bond types • Ionic: valance electrons are transferred • Covalent: valence electrons are shared • Bond formation results in formation of full outer Bohr orbits. Because this stable configuration typically involves eight electrons, this is commonly known as the OCTET rule.

4. Lewis Structure • Element symbol surrounded by a number of dots equal to the number of valence electrons • Ignore the inner or CORE electrons. • The order in which electrons (dots) are drawn and their exact locations are not critical. • Chemical bonding brings together elements in the correct ratios so that all of the atoms involved form an octet.

5. Ionic Lewis Structures • Since ionic bonding involves the transfer of electrons from a metal to a nonmetal, the Lewis structure for an ionic compound involves moving dots. The metal becomes a cation and the nonmetal becomes an anion.

6. Ionic Lewis Structures: Charges The metal and the nonmetal each acquire a charge in the formation of an ionic bond. We indicate the magnitude of the charge in the upper right corner of the symbol. We enclose the anion in brackets. Charges on anions and cations within an ionic formula sum to zero.

7. Concept Check 5.1 Draw the Lewis structure for CaO.

8. Concept Check 5.1 Solution Calcium is a metal and when it forms ionic bonds with nonmetals, it loses both of its valence electrons to form a Ca2+ ion with a stable noble gas configuration. Oxygen is a nonmetal and when it forms ionic bonds with metals, it gains electrons to form an O2- ion with a stable noble gas configuration. The Lewis structure for CaO

9. Covalent Lewis Structures • Covalent bonds involve the sharing of electrons. • Covalent Lewis structures contain dots that count for the octet of more than one atom.

10. Types of Electron Pairs • The electrons between two atoms are called bonding pairs. • Electrons on a single atom are called lone pairs. • Only bonding electrons count toward the octet of both atoms.

11. Water • The Lewis structure shows a 2:1 ratio of hydrogen to oxygen. • H2O • Observed in nature to have two hydrogens for every oxygen

12. Multiple Bonds • Sometimes multiple bonding pairs are necessary to complete the octets for each atom in the Lewis structure. • Two bonding pairs are called a double bond. Three bonding pairs are a triple bond.

13. Writing Lewis Structures • Write the skeletal structure of the molecule. • Determine the total number of electrons for the molecule. • Place the electrons as dots to give octets to as many atoms as possible. • If the central atom has not obtained an octet, form multiple bonds as necessary to complete its octet.

14. Concept Check 5.2 • Draw the Lewis structure for the covalent molecule methane, CH4.

15. Concept Check 5.2 Solution • Draw the skeletal structure of the molecule. • Count the number of valence electrons in CH4. C 4 valence electrons × 1C = 4 electrons H 1 valence electrons × 4H = 4 electrons 8 electrons total in structure

16. Concept Check 5.2 Solution 3.Put a single bond between each C and H atom, using all eight of the valence electrons. Place the electrons as dots to give octets to as many atoms as possible. In this case, there are no unpaired electrons.

17. Chemical Bonding in Ozone • Ozone is an atmospheric gas that protects life on Earth from excessive exposure to UV light.

18. Ozone: Lewis Structure • There appear to be two equally valid Lewis structures. • The ozone molecule is actually best represented with two identical bonds, each one shorter than a single bond, but longer than a double bond.

19. Ozone: Resonance Structures • Resonance in ozone is the averaging of two identical Lewis structures. • Resonance structures are usually represented by a double-centered arrow between them.

20. Oxygen vs. Ozone • Molecular oxygen’s double bond is stronger than ozone’s “bond and a half.” • UV light is not energetic enough to break oxygen’s strong bond. • Oxygen does not absorb UV light. • UV light will break one of ozone’s two bonds. • Ozone absorbs UV light, effectively preventing much of the UV light from reaching the Earth’s surface.

21. The Shapes of Molecules • Molecular shape is an important factor in determining the properties of substances. • Valence shell electron pair repulsion theory (VSEPR theory) allows us to predict molecule shapes from their Lewis structures.

22. VSEPR Theory • A simple model • Based on the idea that centers of negative charge created by both bonding and lone pair electrons repel each other and determine the shape of the resulting molecule • Minimization of repulsions (maximization of distance between like charges) determines three-dimensional molecular geometry.

23. The Ideal Shape • Three-dimensional geometry • Greatest angle (minimum repulsions, maximum distance) between four bonding electron pairs on a central atom is tetrahedral. • The combination of lone pairs and multiple bonds with bonded pairs makes possible other electron pair and molecular geometries. • It is important to distinguish between electron geometry and molecular geometry. In CH4, there are no lone pairs of electrons, therefore, electron pair and molecular geometries are the same.

24. Water: Electron vs. Molecular Geometry • Electron geometry is tetrahedral. • Molecular geometry is bent.

25. Molecular Geometry

26. Concept Check 5.3 • Use Lewis theory and VSEPR to predict the electron and molecular structure of BF3. Boron and other group 3A elements may have either six or eight electrons in their outer shells and still be relatively stable.

27. Concept Check 5.3 Solution • To draw the Lewis structure, first draw skeletal structure of the molecule. • Count the number of valence electrons in BF3. B 3 valence electrons × 1B = 3 electrons F 7 valence electrons × 3F = 21 electrons 24 electrons total in structure

28. Concept Check 5.3 Solution 3.Put a single bond between each B and F atom. Place the electrons as dots to give octets to as many atoms as possible. In this case, there are three lone pairs of electrons on each F.

29. Concept Check 5.3 Solution 4.VSEPR can be applied to a correct Lewis structure to determine molecular geometry around the central atom. VSEPR is based on minimizing repulsions between electrons, either bonding or lone pairs, by maximizing the distance between them. BF3 has three bonding pairs of electrons and no lone pairs around B. A trigonal planar geometry best minimizes these repulsions.

30. Bond Polarity • The dots in covalently bonded Lewis structures appear to be equally shared but reality yields a less ideal picture. • Unequal electron sharing between two elements (polar bonds) can result in polar molecules. • Polar bonds arise whenever elements with different electron-attracting abilities form a bond. The symbols δ- and δ+ indicate partial negative and partial positive charges respectively. • Electronegativity is a measure of an atom’s ability to attract electrons in a bond.

31. Electronegativity Electronegativity increases as you move right across a row on the periodic table and decreases as you move down a column. A polar bond is analogous to a bar magnet; its uneven electron distribution results in a negative pole and a positive pole. Similarly, an entire molecule may be polar if uneven electron distribution within the molecule results in a negative pole and a positive pole.

32. Examples • Carbon monoxide, CO, is a polar molecule. • Carbon dioxide, CO2, contains polar bonds but is a nonpolar molecule. • VSEPR theory predicts a molecular shape for CO2 in which the polar bonds cancel each other.

33. Water • Since the water molecule is polar, attractions between water molecules tend to hold them together (high boiling point). • Molecules arrange themselves to maximize forces of attraction between them and minimize forces of repulsion (lower density when freezes). • Water is immiscible with nonpolar molecules.

34. Polar or Nonpolar? • Look at the Lewis structure and ask two questions. • Does the molecule contain polar bonds? • Do the polar bonds together give overall polarity to the molecule?

35. Polar or Nonpolar? • Look at the Lewis structure and ask two questions. • Does the molecule contain polar bonds? • Do the polar bonds together give overall polarity to the molecule? Yes, the C—F bond is polar. No, all of the individual C—F bond dipoles cancel each other; so overall, the molecule is nonpolar.

36. Concept Check 5.4 • Is BF3 a polar or nonpolar molecule

37. Concept Check 5.4 Solution • The first influence to consider is the polarity of the bonds. Are B—F bonds polar? Referring to Figure 5-4, using the difference between the electronegativity values for B and F we get: 4.0(F) – 2.0(B) = 2.0. • The geometry of the polar bonds determines whether the molecule is polar or not. Referring back to Concept Check 5.3, BF3 is trigonal planar. • Table 5.2 indicates that a trigonal planar structure with identical polar bonds is nonpolar, therefore, BF3 is nonpolar.

38. Chapter Summary • Molecular Concept • Ionic bonds • Covalent bonds • Lewis theory • Valence shell electron pair repulsion (VSEPR) • Polarity • Societal Impact • Life would be impossible without compounds. • Chemical bonding has allowed us to make products such as plastic. • The shapes of molecules determine their properties.