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Chapter 8

Chapter 8. Covalent Compounds. 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. Covalent Bonding. 8.1 Covalent Bonds.

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Chapter 8

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  1. Chapter 8 Covalent Compounds

  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. 8.1 Covalent Bonds • Sharing Electrons • Covalent bonds • space where electrons move is called molecular orbital • made when atomic orbitals overlap

  5. Molecules

  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 electron configuration like noble gas

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

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

  9. Covalent Bonds • Energy and Stability • 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

  10. Bond Energies and Lengths

  11. Covalent Bonds • Electronegativity • Atoms share electrons equally or unequally • nonpolar covalent bond: bonding electrons shared equally • polar covalent bond: shared electrons more likely to be found around more electronegative atom

  12. Covalent Bonds • Electronegativity • Atoms share electrons equally or unequally • difference in electronegativity can be used to predict type of bond (but boundaries are arbitrary)

  13. Bond Types

  14. Covalent Bonds • Electronegativity • Polar molecules have positive and negative ends • such molecules called dipoles •  (“delta”) means partial in math and science • positive end—+ • negative end—- • example: H+F-

  15. Electronegativity Difference for Hydrogen Halides

  16. Covalent Bonds • Electronegativity • Polarity is related to bond strength • greater electronegativity means • greater polarity means • greater bond strength

  17. Covalent Bonds • Electronegativity • Bond type determines properties of substances • metallic bonds: electrons can move from one atom to another—good conductors • ionic bonds: hard and difficult to break apart • covalent bonds: low melting/boiling points

  18. Properties of Substances with Different Types of Bonds

  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 atomic symbol represent electrons

  20. Lewis Structures of Second-Period Elements

  21. Drawing and Naming • Lewis Electron-Dot Structures • Cl2 • HCl

  22. Drawing and Naming • Lewis Electron-Dot Structures • Drawing • 1. Gather information • draw Lewis structure for each atom in compound; place one electron on each side before pairing • determine total number of valence electrons

  23. Drawing and Naming • Lewis Electron-Dot Structures • Drawing • 2. Arrange atoms • arrange structure to show bonding • halogens and hydrogen usually make one bond at end of molecule • carbon usually in center

  24. Drawing and Naming • Lewis Electron-Dot Structures • Drawing • 3. Distribute the dots so that each atom satisfies octet rule (except H, Be, B) • 4. Draw the bonds as long dashes • 5. Verify the structure by counting number of valence electrons

  25. Drawing and Naming • Lewis Electron-Dot Structures • Polyatomic Ions • use brackets [] to show overall charge • example:

  26. Drawing and Naming • Lewis Electron-Dot Structures • Multiple Bonds • sharing two pairs of electrons is a double bond • sharing three pairs of electrons makes triple bonds • example:

  27. Drawing and Naming • Lewis Electron-Dot Structures • Resonance Structures • sometimes, multiple structures are possible • show all possibilities • example:

  28. 6.2 Drawing and Naming • Naming Covalent Compounds • First name: name of first element in formula • usually least electronegative • requires a prefix if more than one of them • Second name: ends in –ide • requires a prefix if more than one of them

  29. Naming Covalent Compounds

  30. 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 idea that electrons repel one another

  31. Molecular Shapes

  32. Molecular Shapes • Lewis structures show which atoms are connected where, and by how many bonds, but they don't properly show 3-D shapes of molecules. To find the actual shape of a molecule, first draw the Lewis structure, and then use VSEPR Theory.

  33. MOLECULAR GEOMETRY VSEPR • Valence Shell Electron Pair Repulsion theory. • Most important factor in determining geometry is relative repulsion between electron pairs.

  34. MOLECULAR GEOMETRY Molecule adopts the shape that minimizes the electron pair repulsions.

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

  36. VSEPR Shape Predictor Table -

  37. VSEPR Shape Predictor Table -

  38. 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 centre so that bond angles are slightly smaller than expected:

  39. Bond Angles • Methane, CH4, has a perfect tetrahedral bond angle of 109° 28' (109.47°), while the H-N-H bond angle of ammonia, NH3, is slightly less at 107°:

  40. Bond Angles • The oxygen of water has two bonded electron pairs and two nonbonded "lone" electron pairs giving a total VSEPR coordination number of 4. • But 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°.

  41. Molecular Shapes • Determining Molecular Shapes and angles. • Let’s try some. • CO • CO2 • BF3 • CH4 • SnCl2 • SO2

  42. Simple Shapes

  43. Trigonal Planar

  44. Tetrahedral

  45. Bent

  46. Molecular Shapes • Molecular Shape Affects Properties • Shape affects polarity • compare CO2 and H2O • polarity affects properties (such as boiling point) due to attractions between molecules

  47. Polar Bonds

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