1. Representing Molecules

2. Bonding Chemical bonds are forces that cause a group of atoms to behave as a unit. Bonds result from the tendency of a system to seek its lowest possible energy. Bond breaking always requires energy, and bond formation always releases energy.

3. Types of Bonds The type of bonding depends upon the nature of the atoms that are combined. A metal and a non-metal will form ionic bonds when electrons are transferred from the metal to the non-metal. The resulting attraction between oppositely charged ions forms a stable crystal.

4. Types of Bonds When metals bond with each other, the valence electrons are shared by the atoms in the entire crystal. The electrons are no longer associated with a specific nucleus, and are free to move throughout the sample.

5. Lewis Structures Lewis Structures, also known as Lewis dot diagrams, show how the valence electrons are arranged among the atoms in the molecule. For ionic compounds, it shows the end result when the metal loses its electrons to the non-metal.

6. Covalent compounds exist as discrete molecules, whereas ionic compounds consist of an aggregate of cations and anions.

7. Covalent Bonds When two (or more) non-metals form bonds, electrons are shared. The result is a covalent bond. Covalent bonds form because the attraction of electrons for the nuclei in the atoms is greater than the electron-electron repulsion or the nucleus-nucleus repulsion.

8. Types of Bonds There is usually an optimum bond length or internuclear distance where attractions between electrons and the nuclei are optimized and repulsions are minimized.

9. Covalent Bonding

10. Bond Energy Bond formation releases energy, and bond breaking requires energy.

11. Types of Covalent Bonds Atoms bonded together may share one, two or three pairs of electrons to make single, double or triple bonds. Double and triple bonds are stronger and shorter than single bonds between the same atoms.

12. Covalent Bonding When atoms of the same element form a covalent bond, the electrons are shared equally. Such a bond is called non-polar, or pure covalent. All homonuclear diatomic molecules contain non-polar bonds.

13. Electronegativity An electronegativity scale, developed by Linus Pauling (1901-1995), is used to predict the direction of the polarity of bonds. Electronegativity is a relative scale that reflects the ability of an atom to attract the electrons in a bond.

14. Electronegativity The small atoms in the upper right corner of the table, having high values of Z eff, also have high electronegativity values. Fluorine has the highest value at 4.0. The noble gases generally do not form compounds, and are not given electronegativity values.

15. Electronegativity The large metal atoms in the lower left hand corner of the periodic table have the lowest electronegativity values.

16. Electronegativity

17. Electronegativity Note that hydrogen has an electro-negativity value of 2.1, consistent with non-metals.

18. Covalent Bonding When atoms of different elements form a covalent bond, the electrons often are not shared equally. The electrons in the bond may spend, on average, more time on one of the atoms. This atom will have a slightly negative charge, indicated by the symbol δ –. The other atom will be slightly positive, indicated as δ +.

19. Covalent Bonding Covalent bonds with unequal sharing of the electrons in the bond are called polar bonds. An example is the molecule HF. The electrons spend more time on fluorine than on hydrogen. As a result, HF is a polar molecule. H F δ+δ+ δ-δ-

20. Dipole Moment Polar molecules have a dipole moment. This is a measure of the tendency of a molecule to line up in an electric field.

21. Dipole Moment Dipole moment, μ, depends upon the size of the partial charges and the distance between the charges. It is measured in Debye (D). A positive charge and a negative charge separated by 100 pm has a dipole moment of 4.80 D.

22. Polarity of Molecules Molecules with polar bonds may be polar, having a permanent dipole moment. Both the polarity of the bonds and the shape of the molecule must be considered. Once the Lewis structure (dot diagram) for the molecule has been determined, it is possible to predict the shape of the molecule and its polarity.

23. Polar Molecules

24. Lewis Structures – Covalent Molecules Once the Lewis structure has been obtained, the Valence Shell Electron Pair Repulsion approach can be used to predict the shape of the molecule and its polarity. CO 2 has polar carbon to oxygen bonds. Even though the bonds are polar, the molecule is non-polar because of its linear shape.

25. Polarity of Molecules Since CO 2 is a linear molecule, the dipoles cancel out, and the molecule is non-polar.

26. Polarity of Molecules Water also has polar bonds, and a bent shape. As a result, water is a polar molecule.

27. Polarity of Molecules The polarity of molecules has a profound effect on the properties and behavior of a substance. It will affect solubility, melting and boiling points, and other important aspects of molecular behavior.

28. The Continuum of Bond Types Metal and a non-metalBetween non-metals

29. Lewis Structures – Covalent Molecules In many covalent molecules, the non-metals share valence electrons. Electrons are shared so that each atom in the molecule has a full valence shell. H.. H H : H or H.. H or H - H Each hydrogen has access to two electrons, as does the noble gas helium.

30. Lewis Structures – Covalent Molecules For elements in period 2, the non-metals generally share enough valence electrons so that each atom obtains the same number of valence electrons as neon (a total of eight electrons). This may involve making multiple (double or triple) bonds between atoms.

31. Lewis Structures – Covalent Molecules Provide Lewis structures for elemental nitrogen, oxygen and fluorine.

32. Exceptions to the Octet “Rule” The elements B and Be sometimes form compounds with less than four electrons pairs on them. The are called electron deficient, and are often highly reactive. The elements in period 3 and below may accommodate more than four electron pairs.

33. Exceptions to the Octet “Rule” In addition, some molecules, such as NO or NO 2, have an odd number of electrons and do not obey the octet rule.

34. Practice Write Lewis dot diagrams for CO, NH 3 and CO 2 Write Lewis dot diagrams for CO, NH 3 and CO 2

35. Resonance Some molecules may have more than one valid Lewis structure. These structures differ in the placement of multiple bonds. In molecules with resonance, none of the Lewis structures accurately represents the true bonding in the molecule.

36. Resonance The molecule SO 2 has two resonance structures: [ O=S-O: ] ↔ [ :O-S=O] The molecule has two equivalent bonds between sulfur and oxygen. : : : : : : : : ::

37. Resonance The sulfur-oxygen bonds are identical- longer than double bonds, and shorter than single bonds. [ O=S-O: ] ↔ [ :O-S=O] The true structure of the molecule is in between the two Lewis structures drawn. : : : : : : : : ::

38. Formal Charges Formal charge is a way to keep track of the electrons in a covalent molecule. The formal charges can also be used to determine if one Lewis structure is more valid than another.

39. Formal Charges The formal charge on an atom is a comparison between the number of valence electrons on each atom and the number of electrons it has in the Lewis structure.

40. Formal Charges Consider the ion SCN -1. There are three valid Lewis structures for the ion. [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] Formal charges can be used to determine the major contributor(s) to the actual structure of the ion. ::: : : :

41. Formal Charges Divide the bonds in half and determine the number of electrons on each atom. [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] 6e 4e 6e 7e 4e 5e 5e 4e 7e 6e 4e 6e 7e 4e 5e 5e 4e 7e ::: : : :

42. Formal Charges Compare the number of electrons in the structure to the number of valence electrons. [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] 6e 4e 6e 7e 4e 5e 5e 4e 7e 6e 4e 6e 7e 4e 5e 5e 4e 7e 6e 4e 5e 6e 4e 5e 6e 4e 5e 6e 4e 5e 6e 4e 5e 6e 4e 5e ::: : : :

43. Formal Charges The net charge is the formal charge on each atom. 0 0 -1 -1 0 0 +1 0 -2 0 0 -1 -1 0 0 +1 0 -2 [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] 6e 4e 6e 7e 4e 5e 5e 4e 7e 6e 4e 6e 7e 4e 5e 5e 4e 7e 6e 4e 5e 6e 4e 5e 6e 4e 5e 6e 4e 5e 6e 4e 5e 6e 4e 5e ::: : : :

44. Formal Charges The net charge is the formal charge on each atom. 0 0 -1 -1 0 0 +1 0 -2 0 0 -1 -1 0 0 +1 0 -2 [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] The sum of the formal charges must equal the charge on the ion. ::: : : :

45. Formal Charges There are two rules used to determine the most likely Lewis structure(s). 1. Atoms try to achieve formal charges as close to zero as possible. 2. Any negative formal charges should reside on the most electronegative atoms.

46. Formal Charges The third Lewis structure is unlikely, due to the high formal charge on nitrogen. 0 0 -1 -1 0 0 +1 0 -2 0 0 -1 -1 0 0 +1 0 -2 [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] Since nitrogen is more electronegative than sulfur, the first structure should be the major contributor. ::: : : :

47. Formal Charges The actual molecule will be somewhere in between the first and second structures. 0 0 -1 -1 0 0 +1 0 -2 0 0 -1 -1 0 0 +1 0 -2 [:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:] The sulfur-carbon bond should be slightly longer than a double bond, and the carbon- nitrogen bond should be slightly shorter than a double bond. ::: : : :