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Chapter 12

Chapter 12. Chemical Bonds. Types of Bonds. Ionic bond: one that holds an ionic cmpd together Relatively strong Results in high melting/boiling points Covalent bond: one that holds a covalent cmpd together Not as strong Lower melting/boiling points

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Chapter 12

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  1. Chapter 12 Chemical Bonds

  2. Types of Bonds • Ionic bond: one that holds an ionic cmpd together • Relatively strong • Results in high melting/boiling points • Covalent bond: one that holds a covalent cmpd together • Not as strong • Lower melting/boiling points • 2 types: polar covalent and nonpolar covalent

  3. In an ionic cmpd, the anion is always larger while the cation fits within the gaps of the anions making a strong crystal lattice.

  4. Polar Covalent Bonds • Type of covalent bond • Electrons are shared unequally • One element has a high affinity for gaining the electron thus forming a slight – and the other element is slightly +

  5. Nonpolar Covalent Bonds • Type of covalent bond • Electrons are shared equally • Usually occurs between 2 identical elements or symmetrical molecules

  6. Electronegativity • EN-ability of an atom in a molecule to attract shared electrons to itself • Can help determine the type of bond as well • Generally a difference of 0=nonpolar covalent bond, 0.1-2.0-polar covalent bond, above 2.0=ionic bond

  7. Dipole Movements • All polar molecules have a dipole movement • Dipole movement-something that exhibits a + and – charge • An arrow represents the direction of the dipole movement • The more EN element has the arrow side and the least EN has a plus end

  8. Water’s polarity is what allows water to attract both + and – ions, allowing many things to dissolve in water. • Also causes water molecules to adhere to each other strongly (reason why a lot of energy is needed to evaporate water)

  9. Examples • Page 404 Ex. 12.1 • Page 406 # 4 and 6

  10. Lewis Structures • A representation that shows how valence electrons bond. • Hydrogen is the only element that follows the duet rule-needs 2 electrons to be stable • Other elements follow the octet rule-need 8 electrons to be stable • Bonding electrons-those involved in bonding • Lone pair electrons-those not involved in bonding

  11. Examples • H2 • H2O • PH3 • CCl4

  12. Multiple Bonds • Single bond-sharing of 2 electrons • Double bond-sharing of 4 electrons • Triple bond-sharing of 6 electrons • Ex: HCN

  13. Resonance & Formal Charge • Some elements exhibit multiple Lewis structures like carbon dioxide. See if you can draw all 3.

  14. Resonance structures-when more than 1 valid Lewis Structure can be drawn • Use arrows to show resonance • How do we know which is right? • Want the formal charge on all elements to be as close to zero as possible • # of valence – nonbonding – (bonding / 2)

  15. Some resonance structures will not require formal charge as they are mirror images of each other • Ex: ozone(O3), SO3 • What is the formal charge on each element in the above?

  16. Polyatomic Ions • Add/subtract electrons as needed • Put the structure in brackets with the charge outside • Check for resonance • Ex: NO2 -1, CO3-2

  17. Exceptions to the rule • Boron is happy with 6 electrons • Anything to the right and below phosphorus can technically have expanded octets. We will work mostly with phosphorus and sulfur. • Ex: BF3, SF6, PCl5

  18. Molecular Structure • 3D arrangement of the valence electrons (both bonded and non) • VSEPR theory-valence shell electron pair repulsion theory • The structure around a given atom is determined by minimizing repulsions between electrons pairs. • AKA: electrons spread out as far away from each other as possible

  19. Electron pair geometry vs. molecular geometry • Electron pair includes all electrons including bonded and non • Molecular refers only to the bonding electrons

  20. http://intro.chem.okstate.edu/1314f00/lecture/chapter10/vsepr.htmlhttp://intro.chem.okstate.edu/1314f00/lecture/chapter10/vsepr.html • Main types/Subtypes with bond angles • Linear- 180 • Trigonal Planar -120 • Bent- under 118 • Tetrahedral- 109.5 • Trigonal pyramidal- 107 • Bent- 105 • Trigonalbipyramidal- 90,120,180 • Octahedral- 90 • Square Planar- 90

  21. Linear-Electron • Linear-Molecular • EX: CO2 • Bond angle of 180

  22. Trigonal Planar-Electron • Trigonal Planar-Molecular • Ex: BF3 • Bond angle of 120 • Bent-Molecular • Ozone • Bond angle of 118

  23. Tetrahedral-Electron • Tetrahedral-molecular • Ex: Methane, CH4 • Bond angles of 109.5 • Trigonal Pyramidal-molecular • Ex: NH3 • Bond angle of 107 • Bent-molecular • Ex: Water • Bond angle of 105

  24. TrigonalBipyramidal-Electron • Trigonalbipyramidal-molecular • Ex: PCl5 • Bond angles of 90, 120, 180

  25. Octahedral-Electron • Octahedral-molecular • Ex: SF6 • Bond angle of 90 • Square planar-molecular • Ex: XeF4 • Bond angle of 90

  26. In-Class Practice • Provide bond angles for all • Page 429 honors book Ex. 12.6 • Others: nitrite ion, sulfur trioxide, phosphorus pentafluoride, SiF6-2

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