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Chemical Bonding

Chemical Bonding

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Chemical Bonding

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  1. Chemical Bonding Chapters 7-8 General Chemistry

  2. Objectives • Explain how atoms combine to form compounds through both ionic and covalent bonding. • Draw Lewis dot structures for simple molecules. • Relate electronegativity and ionization energy to the type of bonding an element is likely to undergo. • Predict the geometry of simple molecules and their polarity (valence shell electron pair repulsion).

  3. Valence Electrons • Valence electrons are the number of electrons in highest occupied energy level of an atom • The s and p electrons in the outer energy level • Fluorine [He] 2s2 2p5 = 7 valence e- • The electrons responsible for the chemical properties of atoms are those in the outer energy level • Core electrons -those in the energy levels below the outer energy level 2s2 2p5

  4. e- Configuration and Valence e-

  5. Lewis Dot (Electron Dot) Diagrams • Lewis Dot (electron dot) diagrams show valence e- as dots around symbol of element X

  6. Lewis Dot Diagrams of Selected Elements

  7. Octet Rule • The octet rule: atoms of elements gain, lose or share e- so that each atom has a full outermost energy level • Want to achieve the e- configuration of a noble gas • Why named “octet”? • Exceptions?

  8. Chemical Bonding • When atoms bond, the valence electrons are redistributed to make the atom more stable • Ionic bonding: results from the electrical attraction between large numbers of cations and anions • Covalent bonding: results from the sharing of electrons between two atoms

  9. Ionic Bonding

  10. Remember Ions ? • Ions: charged atoms • Cations: positively charged atoms • Metals, like sodium, tend to lose electrons to create a noble gas configuration (cations) • Anions: negatively charged atoms • Nonmetals, like chlorine, tend to gain electrons to create a noble gas configuration (anions)

  11. Ionic Bonds • Formed between metal and nonmetal atoms • Anions and cations are held together by opposite charges • The bond is formed through the transfer of electrons • Ionic compounds are called salts • Simplest ratio is called the formula unit • Example: Na+ will bond with Cl- to make sodium chloride, NaCl

  12. Electronegativity • Electronegativity: reflects an atom’s ability to attract electrons in a chemical bond • Metals generally have low electronegativity • Nonmetals generally have high electronegativity

  13. How Determine if Ionic? • Ionic bonds form between 2 atoms with difference in electronegativity of 2.0 or greater

  14. Properties of Ionic Compounds • Conduct electricity in aqueous form • are electrolytes • High melting and boiling points • Usually solids at room temperature • Have crystalline shape • Example: sodium chloride (table salt)

  15. Lattice Energy • The strength of an ionic bond compared to another ionic bond is determined by the lattice energy • Lattice energy is the energy released when one mole of an ionic compound is formed from gaseous ions • Examples: • NaCl -787.5 kJ/mol (weaker bond) • MgO -3760 kJ/mol (stronger bond)

  16. Crystalline structure

  17. Ionic Bonding Lewis Dot Diagrams Na Cl

  18. Ionic Bonding Lewis Dot Diagrams Na+ Cl-

  19. Ionic Bonding Lewis Dot Diagrams • All the electrons must be accounted for! Ca P

  20. Ionic Bonding Lewis Dot Diagrams Ca P

  21. Ionic Bonding Lewis Dot Diagrams Ca2+ P

  22. Ionic Bonding Lewis Dot Diagrams Ca2+ P Ca

  23. Ionic Bonding Lewis Dot Diagrams Ca2+ P 3- Ca

  24. Ionic Bonding Lewis Dot Diagrams Ca2+ P 3- Ca P

  25. Ionic Bonding Lewis Dot Diagrams Ca2+ P 3- Ca2+ P

  26. Ionic Bonding Lewis Dot Diagrams Ca Ca2+ P 3- Ca2+ P

  27. Ionic Bonding Lewis Dot Diagrams Ca Ca2+ P 3- Ca2+ P

  28. Ionic Bonding Lewis Dot Diagrams Ca2+ Ca2+ P 3- Ca2+ P 3-

  29. Ionic Bonding Lewis Dot Diagrams = Ca3P2 Formula Unit

  30. Metallic Bonding

  31. + + + + + + + + + + + + Metallic Bonds • Metallic bonding is the bonding that results from the attraction between metal atoms and the surrounding sea of electrons. • Bond between two metal atoms

  32. Sea of Electrons • Metals hold on to their valence electrons very weakly. • Think of them as positive ions (cations) floating in a sea of electrons • Electrons are free to move through the solid. • Metals conduct electricity.

  33. Covalent Bonding

  34. Covalent Bonds • Two nonmetals share electrons to achieve full octet of electrons • By sharing, both atoms get to count the electrons toward a noble gas configuration. • Form molecules - compounds that are bonded covalently

  35. Examples of Molecules

  36. How determine if covalent? • Covalent bonds form between 2 atoms with difference in electronegativity of less than 2

  37. Properties of Covalent Compounds • Do not conduct electricity in aqueous solution • Are non-electrolytes • Relatively low melting and boiling points • Can be gasses, liquids or solids @ room temp • Examples: sugar, wax, carbon dioxide

  38. Comparison of MP, BP in Ionic and Covalent Compounds

  39. Bond Energy • The strength of a covalent bond compared to another covalent bond is determined by the bond energy • Bond Energy: energy required to break a chemical bond and form neutral, isolated atoms • Stronger covalent bonds have a higher bond energy

  40. Bond Energy and Bond Length

  41. Bond Length • Bond Length: the average distance between two bonded atoms • The longer the bond, the smaller the bond energy (the weaker the bond) • The shorter the bond, the larger the bond energy (the stronger the bond)

  42. Single covalent Double covalent Triple covalent Share 2 e- (one pair) Share 4 e- (two pairs) Share 6 e- (three pairs) Types of Covalent Bonds

  43. F Covalent bonding • Fluorine has seven valence electrons

  44. F F Covalent bonding • Fluorine has seven valence electrons • A second F atom also has seven • By sharing electrons…

  45. Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons… • …both end with full orbitals F F 8 Valence electrons

  46. Bonding and Nonbonding Electrons • Bonding (shared) electrons are involved in a chemical bond • Nonbonding (unshared or lone pair) electrons are not involved in bonding and belong exclusively to one atom Nonbonding electrons Bonding electrons

  47. Diatomic Elements • Seven pure elements that exist as pairs in nature • Are covalently bonded • H2 N2 O2 F2 Cl2 Br2 I2 • Ways to remember: • Br I N Cl H O F • H, N, O, Halogens

  48. Polarity

  49. Bond Polarity • Atoms of elements do not always share electrons equally • Polar covalent bond: unequal sharing of electrons (dif electroneg 0.5 – 1.9) • Nonpolar covalent bond: equal sharing of electrons (dif electroneg 0.0-0.4)