Chapter 6: Covalent Compounds

1. Chapter 6: Covalent Compounds

2. 6.1 Covalent Compounds

3. What is a covalent bond? • Sharing of electrons between atoms • Electron orbitals overlap to form molecular orbitals • Molecular orbitals are the region where the shared electrons move • Covalent bonds are also called molecular bonds.

4. How is a molecular bond different from an ionic bond? • Ionic Bond: Transfer of electrons • Covalent or Molecular Bond: Sharing of electrons

5. Bond Length • Average distance between nuclei of two atoms • Not Fixed: atoms vibrate; bond acts like a spring

6. Bond Energy • Energy needed to break a bond • As bond energy goes up, bond length goes down (inverse relationship) H Br H F 141.1 pm 91.7 pm Bond energy = 366 kJ/mol Bond energy = 570 kJ/mol

8. Electronegativity • Difference in electronegativities can be used to determine bond type

9. Predicting Bond Type from Electronegativity Difference • Nonpolar Covalent: Electronegativity Difference < 0.5 • Polar Covalent: Electronegativity Difference ≥ 0.5 and ≤ 2.1 • Ionic: Electronegativity Difference > 2.1

10. Types of Covalent Bonds 1. Nonpolar Covalent Bond • Bonding electrons shared evenly • Difference in electronegativity is less than 0.5 • Examples: C – H H – H

11. Types of Covalent Bonds 2. Polar Covalent Bond • Bonding electrons not shared evenly • Difference in electronegativity is 0.5 to 2.1 • Tend to be stronger than nonpolar covalent bonds • Electrons tend to be closer to more electronegative atoms • Example: H – F

12. Polar Covalent Bonds • Polar covalent bonds have a dipole. • Dipole: one end has a partial positive charge and the other end has a partial negative charge. • δ+ is used to indicate a partial positive charge • δ- is used to indicate a partial negative charge • Examples:

13. Ionic Bonds • Metal + Nonmetal • Electronegativity difference is greater than 2.1 • Example: Li – F

14. Summary • Nonpolar Covalent: Electronegativity Difference < 0.5 • Polar Covalent: Electronegativity Difference ≥ 0.5 and ≤ 2.1 • Ionic: Electronegativity Difference > 2.1

15. Determine if the following compounds are ionic, polar covalent, or nonpolar covalent: • Na – F • Ca – O • C – Cl • Al – Cl • Br – Br

16. Determine if the following compounds are ionic, polar covalent, or nonpolar covalent: • Na – F = ionic (electronegativity difference = 3.1) 2. Ca – O = ionic (electronegativity difference = 2.4) 3. C– Cl = polar covalent (electronegativity difference = 0.6) 4. Al – Cl = polar covalent (electronegativity difference = 1.6) • Br – Br = nonpolar covalent (electronegativity difference = 0)

17. Properties of Covalent Compounds • Low melting and boiling points • Soft and squishy • Tend to be flammable • Do NOT conduct electricity in water • Are NOT very soluble in water

18. 6.2 Drawing Covalent Compounds: Lewis Structures

19. Gilbert N. Lewis (1875 – 1946) • Gilbert N. Lewis created Lewis structures • Lewis structures (electron-dot diagrams) • Structural formula for drawing molecules • Dots = electrons • Lines = electron pairs/bonds • Only valence electrons are shown

20. Lewis Structures • These are examples of Lewis structures H Be C O

21. 1. Lewis Structures of Atoms • A Lewis diagram consists of the element symbol in the center and dots going around it representing the valance electrons. • The dots representing the valance electrons must be evenly distributed around all 4 sides of the symbol. • Group 2 elements and Helium are the exception to the rule of even distribution.

22. When placing dots around the element symbol, you should start on the left side and go clock wise until you run out of electrons. • Nitrogen has 5 Valence electrons N N N N N

23. Practice Questions Draw a dot diagram of the following elements: • Potassium • Boron • Chlorine • Neon • Lead • Oxygen

24. K B Cl Ne Pb O

25. Group 2 and Helium -Group 2 and Helium are the exceptions because their valence electrons are in the first shell. -Therefore the 2 valence electrons for Helium and Group 2 elements need to be put on the same side of the element symbol. He

26. Why do Lewis structures help us? • These diagrams allow us to understand how atoms will bind based on their valance electron configuration.

27. Bonding of atoms • When you have unpaired electrons on a side of an atom that atom will want to bond/share electrons with other atoms • This is because all atoms want a full valance shell • Oxygen has two unpaired electrons • Hydrogen has one free electron so 2 H can bond with O, this make H O 2 H O H O H

28. 2. Drawing Lewis Structures of Compounds • Step 1: Count the total number of valence electrons in the compound. Example: CH3I

29. Step 2: Draw a Lewis structure for each atom in the compound. Example: CH3I 14 valence electrons

30. Step 3: Arrange the atoms • Halogen and Hydrogen atoms often bond to only 1 other atom and are usually at the end of the molecule • Carbon is often in the center of the molecule • With the exception of carbon, the atom with the lowest electronegativity is often the central atom Example: CH3I 14 valence electrons

31. Step 4: Distribute the dots • Distribute the electron dots so that each atom (except H, Be, and B) satisfy the octet rule Example: CH3I 14 valence electrons

32. Step 5: Draw the bonds • Change each pair of dots that represent a shared pair of electrons to a long dash. Example: CH3I 14 valence electrons

33. Step 6: Verify the structure • Count the electrons in your structure

34. Practice 1.) H2S 2.) NH3 3.) BF3

35. More Practice 4.) C2H6 5.) CH2Cl2 6.) CH3OH 7.) SCl2 8.) CHF3

36. 3. Lewis Structures with Polyatomic Ions

37. More Practice

38. 3. Lewis Structures of Compounds with Multiple Bonds C, O, N often form double bonds by sharing 2 pairs of electrons. C, N may even form triple bonds by sharing 3 pairs of electrons.

39. Practice 1.) O2 2.) C2H4 3.) CO2 4.) N2

40. More Practice 5.) CO 6.) C2H2 7.) HCN 8.) HC2Cl 9.) N2F2

41. 4. Resonance Structures -When a molecule has two or more possible Lewis structures, the two structures are called resonance structures. -Place double-headed arrows between the structure to show that the actual molecule is an average of the two possible states.

42. Practice 1.) SO2 2.) O3 3.) NO2-

43. 6.3 Naming Covalent Molecules

44. Writing Formulas Covalent Compounds

45. Covalent Compounds • Formed from 2 nonmetals or 1 metalloid and 1 nonmetal.

47. Writing Formulas for Covalent Compounds • Write the symbol for the cation. • From the prefix, write the number of atoms as a subscript behind the symbol. • Write the symbol of the anion. • From the prefix, write the number of atoms as a subscript behind the symbol. • DO NOT use charges.

48. Examples • Carbon Monoxide = CO • Tetraphosphorus decaoxide = P4O10 • Carbon dioxide • Sulfur dioxide

49. Examples • Carbon Monoxide = CO • Tetraphosphorus decoxide = P4O10 • Carbon dioxide = CO2 • Sulfur dioxide = SO2

50. Naming Compounds Covalent Compounds