1. Covalent Bonding Section 4.2

2. Covalent Bonds Sharing Electrons –Covalent bonds form when atoms share one or more pairs of electrons nucleus of each atom is attracted to electron cloud of other atom neither atom removes an electron from the other

3. Covalent Bonding

4. Covalent Bonds space where electrons move is called molecular orbital made when atomic orbitals overlap

5. Molecules (Space-filling Models)

6. Covalent Bonds Energy and Stability –Noble gases are stable (full octet) (low P.E.) –Other elements are not stable (high P.E.) covalent bonding decreases potential energy because each atom achieves an electron configuration like a noble gas

7. Covalent Bonds Energy and Stability because P.E. decreases when atoms bond, energy is released –atoms lose P.E. when they bond –loss of P.E. implies higher stability

9. Covalent Bonds Energy and Stability potential energy determines bond length at minimum P.E., distance between two bonded atoms' nuclei is called bond length –bonded atoms vibrate –therefore, bond length is an average length

10. Covalent Bonds bonds vary in strength bond energy is the amount of energy required to break the bonds in 1 mol of a chemical compound bond energy predicts reactivity bond energy is equal to loss of P.E. during formation

11. Multiple Bonds Consider 2 carbon atoms joined by a single, double, or triple bond StructureBond length (pm) Bond strength (kJ mol -1 C ̶ C 154364 C = C134602 C ≡ C120835

12. Multiple Bonds The shorter the bond, the greater the bond strength. The more bonds between two covalently bonded atoms, the shorter the bond Length: triple bond < double bond < single bond Strength: single bond < double bond < triple bond

13. Another Example Consider the COOH group which is attached to certain acids The C is single bonded to one O and double bonded to the other StructureBond length (pm) Bond strength (kJ mol -1 C ̶ O 143358 C = O120799

14. Covalent Bonds Electronegativity –Atoms share electrons equally or unequally nonpolar covalent bond: bonding electrons shared equally (when diff in EN value is zero) polar covalent bond: shared electrons more likely to be found around more electronegative atom (when diff. in EN value is between 0 and less than 1.8)

15. Covalent Bonds Differences in electronegativity can be used to predict type of bond (but boundaries are arbitrary) Subtract the electronegativities of the atoms in the bond to determine the difference

16. Bond Types

17. Practice: Calculate the bond type N and H F and F Ca and Cl C and O Polar Non-polar Ionic Polar

18. Covalent Bonds Polar molecules have positive and negative ends (poles) such molecules are called dipoles δ (“delta”) means “partial” in math and science positive end is designated as δ + negative end is designated asδ - example: H δ+ F δ-

19. Drawing and Naming Lewis Electron-Dot Structures –Lewis structures represent valence electrons with dots position of electrons is symbolic (not literal) shows only the valence electrons of an atom dots around the atomic symbol represent electrons

20. Lewis Structures of Second- Period Elements

21. Drawing Lewis Structures for Compounds Refer to extra notes on moodle

22. Coordinate Covalent (Dative) Bonds resulting bond when a shared pair of electrons originates from one atom Examples: CO (O contributes 2 electrons for one of the bonds), NH 4 + (N contributes 2), H 3 O + (O contributes 2)

23. Drawing and Naming Polyatomic Ions: be sure to add or subtract the electrons use brackets [] to show overall charge Examples: NH 4 + and OH -

24. Naming Covalent Compounds

25. MOLECULAR GEOMETRY VSEPR V alence S hell E lectron P air R epulsion theory.V alence S hell E lectron P air R epulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs.Most important factor in determining geometry is relative repulsion between electron pairs. Molecules adopt the shape that minimizes the electron pair repulsions.

26. Molecular Shapes Lewis structures show which atoms are connected where, and by how many bonds, but they don't properly show the 3-D shape of the molecules To find the actual shape, first draw the Lewis structure, and then apply the VESPR Theory

27. Molecular Shapes Determining Molecular Shapes –Three-dimensional shape helps determine physical and chemical properties –valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes based on the idea that electrons repel one another

28. VSEPR Rules To apply VSEPR theory: 1: Draw the Lewis structure of the molecule and identify the central atom 2: Count the number of negative centers (lone and bonding pairs) surrounding the central atom. 3: Predict molecular shape by assuming that electrons orient themselves so they are as far away from one another as possible

29. VSEPR Shape Predictor Table

31. Bond Angles Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected:

32. Bond Angles Methane, CH 4, has a perfect tetrahedral bond angle of about 109°, while the N-H bond angle of ammonia, NH 3, is slightly less at 107°:

33. Bond Angles The oxygen of water has two bonded electron pairs and two non-bonded "lone" electron pairs. The lone pairs push the bond down so it is smaller The geometry is defined by the relationship between the H-O-H atoms and water is said to be "bent" or "angular" shape of 104.5°.

34. ShapeBond AngleExample linear180CO 2 and HCN Trigonal planar120BF 3 and SO 3 V-shaped117SO 2 Tetrahedral109CH 4 and CCl 4 Trigonal pyramidal 107NH 3 and NF 3 Bent linear104.5H 2 O and H 2 S Table of Common Shapes, Angles and Examples

35. Simple Shapes

36. Trigonal Planar

37. Tetrahedral

38. Bent

39. Molecular Shapes Determining Molecular Shapes and angles. –Let’s try some. CO NCl 3 BF 3 CH 4 H 2 O SO 2

40. Polarity of Molecules Two atoms: bond polarity is the molecular polarity More than 2 atoms: the geometry of the molecule must be considered If the bonds are non-polar, the molecular is non-polar Some molecules with polar bonds can be non-polar

41. More Sometimes the partial charges cancel each other out because they are directly opposite each other Consider CO 2 and CCl 4 The symmetrical distribution of the bonds leads to cancellation of the charges

42. Polar Bonds

43. Giant Covalent Lattices Consist of a 3-D lattice of covalently bonded atoms Atoms can be either all of the same type as in silicon and carbon, or of two different elements such as silicon and dioxide

44. Allotropes of Carbon Allotropes are two or more crystalline forms of the same element in which the atoms are bonded differently Carbon has 3 allotropes –Diamond, graphite, C 60 (fullerene)

45. Diamond Each C atom is tetrahedrally bonded to 4 other C atoms by single covalent bonds (giant covalent lattice) Very rigid 3-D network with bond angles of 109.5° Very hard, poor conductor, very dense

46. Graphite Each C atom is covalently bonded to 3 other C atoms (giant covalent lattice) 2-D network is formed consisting of hexagonal rings of carbon atoms Crystal structure is composed of many layers of rings Very soft and slippery, good conductor along the plane of the layers, not as dense as diamond

47. Fullerene C atoms are arranged into the shape of a soccer ball Has 60 corners and 32 faces Bonding is a series of single and double bonds Soft, very poor conductor, less dense than graphite

48. DiamondGraphiteFullerene

49. Silicon Dioxide Exists as a giant covalent structure Most common form is quartz Structure is similar to diamond Hard, transparent, high melting and boiling point Impure form of silicon dioxide is sand

50. Silicon Dioxide

51. Silicon Has a giant covalent structure similar to diamond Less hard than diamond owing to the larger size of the silicon atoms (longer and weaker bonds) Is an insulator, but can be made to conduct electricity by adding small amounts of other atoms

52. Silicon structure is like the diamond structure