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Covalent Bonding Chapter 6.2

Covalent Bonding Chapter 6.2. Molecular Compounds . molecule: neutral group of atoms held together by covalent bonds molecular compound: compound whose simplest unit is a molecule. Formulas. chemical formula: tells the number of each type of atom in a compound

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Covalent Bonding Chapter 6.2

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  1. Covalent Bonding Chapter 6.2

  2. Molecular Compounds • molecule: neutral group of atoms held together by covalent bonds • molecular compound: compound whose simplest unit is a molecule

  3. Formulas • chemical formula: tells the number of each type of atom in a compound • molecular formula: tells the number of each type of atom in a molecular compound • ex. H2O, Cl2, C6H12O2

  4. Molecular Compounds • diatomic molecule: a molecules containing only 2 atoms • usually refers to 2 of the same atoms • ex: O2, Br2, F2, etc. • 7+1 rule

  5. Formation of Covalent Bond

  6. Formation of Covalent Bond • approaching nuclei and electron clouds are attracted to each other to create a decrease in PE • two nuclei and two electron clouds repel each other creating an increase in PE

  7. Formation of Covalent Bond • a distance between the nuclei is reached in which • repulsion and attraction forces are equal • potential energy is at the lowest point possible • at the bottom of the curve on PE graph

  8. Covalent Bonds • Bond Length • distance between two bonded atoms at their lowest PE • average distance since there are some vibrations • measured in pm (1012 pm = 1 m) • stronger the bond, shorter the bond

  9. Covalent Bonds • Bond Energy • energy is released when atoms become because they have lower PE • the same amount of energy must be used to break the bond and form neutral isolated atoms • stronger bond, higher bond energy • average since varies a small amount based on atoms in entire molecule • in kJ/mol

  10. Octet Rule • representative elements can “fill” their outer energy level by sharing electrons in covalent bonds • Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons • Duet Rule- applies to H and He

  11. Octet Rule • Less than 8: • Boron: 6 in outer energy level • More than 8: • anything in 3rd period or heavier • because may use the empty d orbital • ex: S, P, I

  12. Electron Dot Diagrams • a way to show electron configuration • identifies the number and pairing of valence electrons to show how bonding will occur • write the noble gas notation • identify the number of valence • identify how many are paired and how many are alone • do not go by Figure 6-10

  13. Example • Nitrogen • 1s2 2s2 2p3 • 5 valence • 2 are paired • 3 are alone • Sulfur • 1s2 2s2 2p6 3s2 3p4 • 6 valence • 4 paired (2 pairs) • 2 are alone N

  14. Lewis Structures • like dot diagrams but for entire molecules • atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons • unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol • bonding electrons: written in between 2 atoms as a dash

  15. Types of Bonds • single- sharing of one pair of electrons • weakest, longest • double- sharing of 2 pairs of electrons • stronger and shorter • triple- sharing of 3 pairs of electrons • strongest and shortest • multiple bonds include double and triple bonds

  16. Drawing Lewis Structures • find the number of valence electrons in each atom and add them up • draw the atoms next to each other in the way they will bond • add one bonding pair between each connected atoms • add the rest of the electrons until all have 8 (consider exceptions to octet rule)

  17. Example 1 H H C Cl H • CH3Cl • methyl chloride • C: 4 x 1 = 4 • H: 1 x 3 = 3 • Cl: 7 x 1 = 7 • total = 14 electrons • carbon is central duet octet duet octet duet H H C Cl H

  18. Example 2 • NH3 • ammonia • N: 5 x 1 = 5 • H: 1 x 3 = 3 • total = 8 • N is central H N H H

  19. Example 3 • N2 • nitrogen gas • N: 5 x 2 = 10 • 10 electrons N N N N

  20. Example 4 • CH2O • formaldehyde • C: 4 x 1 = 4 • H: 1 x 2 = 2 • O: 1 x 6 = 6 • total = 12 • C is central H C H O

  21. O O O Example 5 • O3 • ozone • O: 6 x 3 = 18 • two completely equal arrangements • the real structure is an average of these two • where each bond is sharing 3 electrons instead of 4 or 2 O O O

  22. O O O O O O Resonance Structures • resonance – bonding between atoms that cannot be represented in on Lewis structure • show all possible structures with double-ended arrow in between to show that electrons are delocalized

  23. Example 6 • NO31- • N: 5 x 1 = 5 • O: 6 x 3 = 18 • total = 23 + 1 = 24

  24. Covalent Network Bonding • a different type of covalent bonding • not specific molecules • lots of nonmetal atoms covalently bonded together in a network in all directions • example: • diamond • silicon dioxide • graphite

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