How to Use This Presentation

1. How to Use This Presentation • To View the presentation as a slideshow with effects select “View” on the menu bar and click on “Slide Show.” • To advance through the presentation, click the right-arrow key or the space bar. • From the resources slide, click on any resource to see a presentation for that resource. • From the Chapter menu screen click on any lesson to go directly to that lesson’s presentation. • You may exit the slide show at any time by pressing the Esc key.

2. Resources Chapter Presentation Bellringer Transparencies Sample Problems Visual Concepts Standardized Test Prep

3. Covalent Compounds Chapter 6 Table of Contents Section 1Covalent Bonds Section 2Drawing and Naming Molecules Section 3Molecular Shapes

4. Chapter 6 Section1 Covalent Bonds Bellringer • Make a list of the elements that form ionic bonds. Note that most ionic bonds contain a metal and a nonmetal.

5. Chapter 6 Section1 Covalent Bonds Objectives • Explain the role and location of electrons in a covalent bond. • Describe the change in energy and stability that takes place as a covalent bond forms. • Distinguish between nonpolar and polar covalent bonds based on electronegativity differences.

6. Chapter 6 Section1 Covalent Bonds Objectives, continued • Compare the physical properties of substances that have different bond types, and relate bond types to electronegativity differences. • Ignore pgs.192,193

7. Section1 Covalent Bonds Chapter 6 Sharing Electrons • When an ionic bond forms, electrons are rearrangedand are transferred from one atom to another to form charged ions. • In another kind of change involving electrons, the neutral atoms share electrons.

8. Section1 Covalent Bonds Chapter 6 Sharing Electrons, continued Forming Molecular Orbitals • A covalent bond isa bond formed when atoms share one or more pairs of electrons. • The shared electrons move within a space called a molecular orbital. • Amolecular orbital is the region of high probability that is occupied by an individual electron as it travels with a wavelike motion in the three-dimensional space around one of two or more associated nuclei.

9. Formation of a Covalent Bond Chapter 6

10. Visual Concepts Chapter 6 Chemical Bond

11. Section1 Covalent Bonds Chapter 6 Energy and Stability Energy Is Released When Atoms Form a Covalent Bond

12. Section1 Covalent Bonds Chapter 6 Energy and Stability, continued Potential Energy Determines Bond Length • When two bonded hydrogen atoms are at their lowest potential energy, the distance between them is 75 pm. • The bond length isthe distance between two bonded atoms at their minimum potential energy. • However, the two nuclei in a covalent bond vibrate back and forth. The bond length is thus the average distance between the two nuclei.

13. Visual Concepts Chapter 6 Bond Length

14. Section1 Covalent Bonds Chapter 6 Energy and Stability, continued Bonded Atoms Vibrate, and Bonds Vary in Strength • The bond length is the average distance between two nuclei in a covalent bond. • At a bond length of 75 pm, the potential energy of H2 is –436 kJ/mol. • Thus 436 kJ of energy must be supplied to break the bonds in 1 mol of H2 molecules. • The energy required to break a bond between two atoms is the bond energy. • Bonds that have the higher bond energies (stronger bonds) have the shorter bond lengths.

15. Visual Concepts Chapter 6 Bond Energy

16. Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding • In covalent bonds between two different atoms, the atoms often have different attractions for shared electrons. • Electronegativity values are a useful tool to predict what kind of bond will form.

17. Visual Concepts Chapter 6 Electronegativity

18. Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally • When the electronegativity values of two bonding atoms are similar, bonding electrons are shared equally. • A covalent bond in which the bonding electrons in the molecular orbital are shared equally is a nonpolar covalent bond.

19. Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally, continued • When the electronegativity values of two bonding atomsare different, bonding electrons are shared unequally. • A covalent bond in which the bonding electrons in the molecular orbital are shared unequally is a polar covalent bond.

20. Predicting Bond Character from Electronegativity Differences Chapter 6

21. Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends • In a polar covalent bond, the ends of the bond have opposite partial charges. • A molecule in which one end has a partial positive charge and the other end has a partial negative charge is called a dipole. • In a polar covalent bond, the shared pair of electrons is not transferred completely. Instead, it is more likely to be found near the more electronegative atom.

22. Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends, continued • The symbol is used to mean partial. • + is used to show a partial positive charge • – is used to show a partial negative charge charge • example: H+F– • Because the F atom has a partial negative charge, the electron pair is more likely to be found nearer to the fluorine atom

23. Visual Concepts Chapter 6 Comparing Polar and Nonpolar Covalent Bonds

24. Chapter 6 Section1 Covalent Bonds Polarity Is Related to Bond Strength • In general, the greater the electronegativitydifference, the greater the polarity and the stronger the bond.

25. Section1 Covalent Bonds Chapter 6 Electronegativity and Bond Types • Differences in electronegativity values provide one model that can tell you which type of bond two atoms will form. • Another general rule states: • A covalent bond forms between two nonmetals. • An ionic bond forms between a nonmetal and a metal.

26. Section1 Covalent Bonds Chapter 6 Properties of Substances Depend on Bond Type • The type of bond that forms (metallic, ionic, or covalent) determines the properties of the substance. • The difference in the strength of attraction between the basic units of ionic and covalent substances causes these types of substances to have different properties.

27. Properties of Substances with Metallic, Ionic, and Covalent Bonds Chapter 6

28. Chapter 6 Section 6.1 review, pg 198 1. Describe the attractive forces and repulsive forces that exist between two atoms as the atoms move closer together. The positive nucleus of each atom attracts the electrons of the other atom. At the same time, the nuclei repel each other, as do the electron clouds. 3. In what two ways can two atoms share electrons when forming a covalent bond? The two atoms may share electrons equally, forming a nonpolar covalent bond, or unequally, forming a polar covalent bond.

29. Chapter 6 5. How are the partial charges shown in a polar covalent molecule? The symbol is written as a superscript on the element with the partial positive charge. The symbol is written as a superscript on the element with the partial negative charge. 6. What information can be obtained by knowing the electronegativity differences between two elements? Knowing the electronegativity difference suggests what type of bond will form between the atoms of the two elements.

30. Chapter 6 7. Why do molecular compounds have low melting points and low boiling points relative to ionic substances? The attractive forces between individual molecules are weak, accounting for the low melting point of molecular compounds.

31. Section2 Drawing and Naming Molecules Chapter 6 Bellringer • Classify the following compounds according to the type of bonds they contain: • NO • CO • HF • NaCl • HBr • NaI

32. Section2 Drawing and Naming Molecules Chapter 6 Objectives • Draw Lewis structures to show the arrangement of valence electrons among atoms in molecules and polyatomic ions. • Explain the differences between single, double, and triple covalent bonds. • Draw resonance structures for simple molecules and polyatomic ions, and recognize when they are required.

33. Section2 Drawing and Naming Molecules Chapter 6 Objectives, continued • Name binary inorganic covalent compounds by using prefixes, roots, and suffixes.

34. Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures • Valence electrons are the electrons in the outermost energy level of an atom. • A Lewis structure is a structural formula in which valence electrons are represented by dots. • In Lewis structures, dot pairs or dashes between two atomic symbols represent pairs in covalent bonds.

35. Visual Concepts Chapter 6 Valence Electrons

36. Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons • As you go from element to element across a period, you add a dot to each side of the element’s symbol.

37. Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • You do not begin to pair dots until all four sides of the element’s symbol have a dot.

38. Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • An element with an octet of valence electrons has a stable configuration. • The tendency of bonded atoms to have octets of valence electrons is called the octet rule.

39. Visual Concepts Chapter 6 The Octet Rule

40. Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • When two chlorine atoms form a covalent bond, each atom contributes one electron to a shared pair. • An unshared pair, or a lone pair, is a nonbonding pair of electrons in the valence shell of an atom.

41. Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • A single bond is a covalent bond in which two atoms share one pair of electrons • The electrons can pair in any order. However, any unpaired electrons are usually filled in to show how they will form a covalent bond.

42. Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Draw a Lewis structure for CH3I.

43. Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Solution Draw each atom’s Lewis structure, and count the total number of valence electrons. number of dots: 14 Arrange the Lewis structure so that carbon is the central atom.

44. Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Solution, continued Distribute one bonding pair of electrons between each of the bonded atoms. Then, distribute the remaining electrons, in pairs, around the remaining atoms to form an octet for each atom. Change each pair of dots that represents a shared pair of electrons to a long dash.

45. Chapter 6.2 Sample Problem A,practice pg.202 1)Draw the Lewis structures for H2S, CH2Cl2, NH3, and C2H6.

46. Chapter 6.2 Sample Problem A,practice pg.202 2)Draw the Lewis structures for methanol, CH3OH.

47. Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Draw a Lewis structure for the sulfate ion,

48. Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Solution Count electrons for all atoms. Add two additional electrons to account for the 2− charge on the ion. number of dots: 30 + 2 = 32 Distribute the 32 dots so that there are 8 dots around each atom.

49. Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Solution, continued Change each bonding pair to a long dash. Place brackets around the ion and a 2 charge outside the bracket to show that the charge is spread out over the entire ion.

50. Chapter 6.2 Sample Problem B, practice pg.203 1)Draw a Lewis structure for ClO3- 2) Draw the Lewis structure for the hydronium ion, H3O +