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Chemical Bonds

Chemical Bonds. All about the ……? 2d law of thermodynamics – entropy – natural state is toward minimum energy Balls roll downhill, hot things cool, students sleep in chemistry class All things get to lowest energy state possible This is the reason for chemical bonding

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Chemical Bonds

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  1. Chemical Bonds • All about the ……? • 2d law of thermodynamics – entropy – natural state is toward minimum energy • Balls roll downhill, hot things cool, students sleep in chemistry class • All things get to lowest energy state possible • This is the reason for chemical bonding • Elements want to reach the lowest, most compact energy state possible, which is the elusive “octet” • Full S and P sublevels – The Noble Gases • See Periodic Table

  2. Chemical Bonds

  3. Chemical Bonds • Exceptions to octet rule • Hydrogen – 1st 4or 5 elements, and inner transition and transition metals • But metals bond differently anyway- more later • Some atoms lose electrons to bond • Ionic bonds • Alkali/Alkali earth metals due to small number • With loss octet achieved (noble gas configuration) • Rules of thumb • <4 valence – lose • Low ionization energies • >4 valence – gain • High electron affinity

  4. Chemical Bonds • Ionic – atoms with great relative electronegativities – usually a metal and a nonmetal

  5. Chemical Bonds • Covalent - Both have strong electron affinities – share electrons – usually nonmetals

  6. Chemical Bonds • Metal – metal – metallic bonding • Only a few loosely held valence electrons • Low electron affinity • Electrons shared by many atoms in an “electron sea”

  7. Chemical Bonds • Bond Polarity & Bond Character • Polarity applies to covalent and ionic bonds • Tendency of an object to form 2 localized regions of opposite character • Earth – N & S magnetic poles – magnetically polarized • Magnets – Positive and Negative electrical poles where charges gather– electrically polarized • Bonds polar if electrons shared unequally between 2 atoms

  8. Chemical Bonds Electrons shared equally – nonpolar – H2 • Electrons shared unequally polar – HFl • Atom with greater electron affinity shares electron more making bond polar • See p.142 Ionic – electons shared hardly at all – effectively stolen

  9. Chemical Bonds • Polar bonds stronger due to charge – harder to break • Difference in strength is electronegativity • Atom’s ability to attract electrons to itself and form bonds

  10. Chemical Bonds • Range of polarity depends on difference in electronegativity • ∆En

  11. Chemical Bonds • Do Sect rvw ques. – p. 143

  12. Chemical Bonds • Covalent bonds • Non metals –hi electron affinity • Share electrons – maintain influence over • Share to attain octet (noble gas config) • Cl2 example – on board – dot notation • Shared electrons – bonding pair • Dense region of negative charge in between atoms where electrons are being shared – great attractive force to nucleus • Electrostatic force

  13. Chemical Bonds • Diatomic elements • Seven of them - memorize • H2, N2 O2 F2 Cl2 Br2 I2 • Mr. H. BrIClFON • Lewis Structures • 2 dimensional representations that show the bonds between different atoms • Similar to electron dot notation except a dash represents the bonding pair

  14. Chemical Bonds • Lewis structures – Hydrogen • Water example • Looking at the valence electrons in dot notation one can forecast what any given element needs to achieve stability • H? O? Cl? • Double Covalence – O2 • Triple Covalence – N2 • Triple Covalent Compound – Acetylene C2H2

  15. Chemical Bonds • Lewis structures – Guidelines • Illustrate Covalent Bonds ONLY • Electron sharing only occurs between Nonmetal-Nonmetal or Diatomic nonmetals • Nonmetals form covalent bonds to achieve octet (bonding and nonboding total) or 2 in case of H • Electrons normally shared in pairs (opposite spin) • Polyatomic molecules (3+ atoms) contain a central atom or atoms bonded to surrounding atoms. • Group 14 (4 valence elecs) more likely to be central • # Valence elecs for nonmetals = Gp# - 10 • # Covalent bonds nonmetals can participate • 8 – Valence e

  16. Chemical Bonds • Lewis structures – Drawing - H2CO • Write dot symbol for each • Determine # of valence available • Place element with most unpaired elecs in center (rem we said Gp 14 – C fam) • Place remaining symbols around central atom next to unpaired elecs • Make as many single bonds as possible • Create double/triple bonds as needed to reach octet • Check structure by counting valence elecs and make sure it matches available elecs

  17. Chemical Bonds • Polyatomic Ions • Atoms with hi elec affinity grab nearby electron to complete, creating covalently bonded anion • Ex. Cl- OH- • Lewis structures written same, but brackets put around ion and charge outside • Ex. NH4+ Ammonium cation • Common usage? NH4OH – ammonium hydro xide= aqueous ammonia or Windex • Addition of H cation to – surface of molecule (polar section) like H3O

  18. Chemical Bonds • Ionic bonding – NaCl • ∆en=? • Structure of ionic compounds • The formula unit – ratio of one ion to the other • No single molecules – a crystal lattice in the needed ratio to achieve octet

  19. Chemical Bonds • Calcium Fluoride • ∆en=? • Ca valence = ? • Fl valence = ? • So the formula unit will be: • CaFl2

  20. Chemical Bonds • Nature of ionic bonds • Crystal Lattice • Brittle crystal • Will not stretch or flatten, only shatter

  21. Chemical Bonds • Metallic bonds • Why no octet rule? • All have low en • All have 2-4 valence electrons • D or f sublevel must be filled before p can be so very difficult to achieve octet! • Thus, the mighty sea • Now – Sect rvw ques.

  22. Chemical Bonds • Properties of bonds • Covalent (? & ?) • Distinct molecules • Gases, liquids, solids with low melting points • Lack density, hardness, rigidity of metals/ionics • Usually poor conductors • Exceptions-network covalent substances Methane – CH4

  23. Chemical Bonds • Exceptions-network covalent substances • Covalently bonded in a continuous 3-D network Carbon Graphite Graphene Diamond

  24. Chemical Bonds • Exceptions-network covalent compounds • Silicates • Quartz • Hard, brittle crystals, high melting points, glassy luster, unusual electrical properties

  25. Chemical Bonds • Ionic Compounds • Electrostatically bonded • Strong, brittle,hard solids w/hi melting pts • Misalignment creates repulsive forces, causing them to fracture • Can be cleaved along flat surface • Poor conductors as solids – both elec&heat • Conductivity depends on mobility – these have none as solids • As liquids, or in solution – charged ions can freely move and are good conductors

  26. Chemical Bonds • Metallic bonds • Delocalized or free electrons holding together like a liquid glue causes all properties of metals • Hard, lustrous • Ductile • Malleable • Good conductors • Heat

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