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Understanding Covalent Bonding and Molecular Compounds: Chapter 6.2 Overview

This chapter explores the fundamentals of covalent bonding, focusing on molecular compounds. It defines molecules as neutral groups of atoms bonded by covalent bonds and molecular formulas as representations of these compounds. Key concepts include the formation of covalent bonds, bond length, bond energy, and the octet rule, essential for understanding electron sharing between atoms. The chapter also highlights the significance of Lewis structures and resonance structures in depicting molecular bonding, alongside examples of diatomic and complex molecules like ammonia, formaldehyde, and ozone.

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Understanding Covalent Bonding and Molecular Compounds: Chapter 6.2 Overview

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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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