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Unit 5

Unit 5. Bonding processes. Isolated atoms. Rarely occur in nature Only the noble gases consist of individual, nonreactive atoms Atoms tend to combine with one another in various ways to form more complex structural units. Bonding processes. Why would elements bond?

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Unit 5

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  1. Unit 5 Bonding processes

  2. Isolated atoms • Rarely occur in nature • Only the noble gases consist of individual, nonreactive atoms • Atoms tend to combine with one another in various ways to form more complex structural units

  3. Bonding processes • Why would elements bond? • Elements combine in specific ratios to gain a noble gas configuration • How many bonds they make depends on their oxidation number (charge for elements). • Elements combine in specific ratios to have 8 valence electrons – called the octet rule

  4. H2 H H Cl2 Cl Cl HCl H Cl Octet Rule • In general - atoms combine to share electrons and fill s & p sublevels.

  5. Exceptions to the octet rule • We don’t want 8!! • H, He, Li, Be • Incomplete Octet (rare) • B – will only bond with 3 H • How many electrons are present? • Odd # Electrons • NO, and NO2 • Expanded Octet (lone pair splits) • SF6, PCl5, XeF6

  6. Bond types • Energy is always released when a covalent bond forms between atoms (exothermic) • It always takes energy to break a bond (endothermic) • Atoms are constantly rearranging to form new bonds • The new product is lower in energy and more stable than the reactants

  7. 3 types of bonding compounds • Ionic (metal and a nonmetal) • Covalent (nonmetal and a nonmetal) • Metallic (metal + metal)

  8. Ionic Compounds • Ions – atoms or groups of atoms that have an electrical charge • An atom becomes an ion by losing or gaining electrons • Metal and a nonmetal • Involves the transferring of electrons • Held together by strong electrical forces between oppositely charged ions (ion – ion force the strongest of all) – called ionic bonds • Tend to be solids at room temperature • Have relatively high melting points • Show the simplest ratio between cation and anion

  9. Ionic bonds • Ionic bonds are strong attractive forces that exist between cations and anions in an ionic compound • Exist as large-ordered, three-dimensional network of positively and negatively charged ions called an ionic lattice • Ions are arranged in such a way that positive ions are adjacent to negative ions • Bonds are very rigid and sturdy

  10. H+ H+ H+ H+ H+ H+ - - - - - - Ionic Bond Type • Two atoms approach each other the nuclides attract the outer electrons • One element “steals” the electrons from the other

  11. Ionic Nomenclature • Know rules for naming the 6 types of ionic bonds – both IUPAC (stock) and Classical • Binary • Polyatomic • Multi-valent • Hydrates • Acids • Bases

  12. Covalent Compounds • Also called molecular compounds • Nonmetal and a nonmetal • NOT composed of ions • Involves the sharing of electrons to for an octet • Shown by dashes to indicate shared pairs of e-

  13. Covalent bonds • Tend to occur between nonmetals, and exist in “molecules” • They share pairs of electrons to achieve stability • They’re called molecular compounds • The bonds are flexible – they can vibrate (spring-like action)

  14. - - - - - - - - - - - - - - H+ H+ H+ Cl+ - - Cl+ Cl+ - - - Forming an Covalent Bond The sharing results is an octet for each atom from 2 electrons being shared by 2 atoms! The difference in attraction is not great enough for transferring. The nuclides share electrons to still have an octet of valance electrons • Hydrogen and Chlorine

  15. Metallic Bonds • Metal atoms have a low attraction for valance e- • Valance electrons can be shared by all the atoms • Valance electrons move between metal ions

  16. More Metal Bond • - Metals are good conductors of heat and electricity. This is directly due to the mobility of the electrons. • - The "cement" effect of the electrons determines the hardness of the metal.  Some metals are harder than others; the strength of the "cement" varies from metal to metal. • -  Metals are lustrous.  This is due to the uniform way that the valence electrons of the metal absorb and re-emit light energy. • - Metals are malleable (can be flattened) and ductile (can be drawn into wires) because of the way the metal cations and electrons can "flow" around each other, without breaking the crystal structure. • *Metallic bonds are best characterized by the phrase "a sea of electrons"* • http://yteach.co.uk/page.php/resources/view_all?id=bond_charge_electrical_ion_ionic_electrovalent_covalent_multiple_coordinate_metalic_t_page_27&from=search

  17. M+ e- M+ e- e- M+ e- M+ M+ e- M+ e- e- M+ M+ e- Even more about metals • Cations are surrounded by delocalized outer electrons • Ductile • Malleability • Wide range of melting points depends on valance electrons • Insoluble in water

  18. Polyatomic Compounds • Polyatomic ions are like molecules but with a charge • The group of atoms, together, have extra or too few electrons • Electrons have been gained or lost to have octets • The bond between the atoms are covalent • The bond between particles are ionic

  19. Practice identifying the types of bonds • Get out a separate sheet of paper and number 1-20 • Write down the following compounds, name them, and classify the following compounds as ionic, covalent, or polyatomic • 1. CaCl2 • 2. CO2 • 3. H2O • 4. BaSO4

  20. 5. K2O • 6. NaF • 7. Na2CO3 • 8. CH4 • 9. SO3 • 10. LiBr • 11. MgO • 12. NH4Cl

  21. 13. HCl • 14. KI • 15. NaOH • 16. NO2 • 17. AlPO4 • 18. FeCl3 • 19. P2O5 • 20. N2O3

  22. Lewis Dot Structures • Used to show the number of valence electrons an element has • Also used to show bonding between two or more atoms • The e- are either transferred or shared to form an octet • Modifications are made when electrons are shared in pairs

  23. Electron Dot Structures • Count Valance electrons • If polyatomic • add extra electrons for anions • remove electrons for cations • Arrange Element closes to group 14 in center • Place other elements on each side • H and Halogens on outside

  24. Drawing Lewis Dot Structures • Ionic compounds – only dots are used; electrons are transferred not shared • Covalent compounds – dots and lines are used; electrons can be shared

  25. Ionic compounds Lewis Dots • Step One: Write the correct formula for the compound • Example: Al + F

  26. Step Two: List each element separate, making sure to evenly distribute the anion and cation with each other • Example:AlF3

  27. Step Three: Draw the Lewis Dot diagrams of each of the individual elements • Example: F • F Al F

  28. Step Four: Use arrows to show how the electrons move from the cation(s) to the anion(s) • Example: F • F Al F

  29. End result: The cation(s) do not have any valence electrons, while the anion(s) have complete shell of valence electrons • Example: F • F Al F

  30. Drawing Ionic compounds • Na + Cl  • Li + S  • Ca + Cl 

  31. Homework • Write the Lewis dot diagrams for the following compounds: • Magnesium sulfide • Calcium oxide • Copper (II) oxide • Copper (III) fluoride • Iron (I) bromide

  32. Covalent Bonds – more than one type • Single Bond • Only 2 electrons are shared • Cl - Cl • Double Bond • 4 electrons are shared • O = O • Triple Bond • 6 electrons shared • N N

  33. Different ways to write covalent (molecular) substances • Molecular formulas • The number of atoms of each element is shown with a subscript • Structural formulas • Show bonding pattern within the molecule

  34. Structural formulas – review of writing Lewis dot for covalent compounds • Rules for writing Lewis Dot structures. • Use the periodic table to determine how many valence electrons each element has and how many bonds each element is probably going to want to make Example: CO2

  35. Step Two: • Count the total number of valence electrons – write this number down. This is the number of electrons your drawing must have. Example: CO2 • C = • O = • O =

  36. Write down the symbols for the elements with them connected by a single line • Use a dash to show the sharing of electrons between the elements Example: CO2

  37. Add remaining electron pairs around each element • Adjust the lone pairs to satisfy the octet rule • You can have a double or triple bond Example: CO2 O - C - O

  38. Cl Cl C Cl Cl Examples 4 + 4*7 = 32 • Count total Valance e- • CCl4 • Place C in center and Cl at sides • 2 electrons per single bonds • Add e- pairs on outside • End with all elements having an octet

  39. Lewis Dot Continued • H and Cl • O and O • N and N

  40. O C O O C O Building Molecules • Fill octets • Double or triple bonds? O C O OR

  41. Condensed structural formula • Shows the bonding pattern in the molecule and highlights the presence of a reactive group of atoms within the molecule

  42. Review • What is the difference between an ionic, covalent, and polyatomic compound? • Write the Lewis dot for the following: • NH3 • SrI2

  43. Resonance structures • Valid dot diagrams that differ from one another only in the arrangement of electrons are called resonance forms • The actual molecule is considered by chemists to be some combination of the different forms, with all of the forms not necessarily equally weighted

  44. O S O O S O O S O O S O Resonates Resonance • More than one way to make double or triple bonds - SO2 Same bond lengths and same bond strengths!

  45. Formal Charge • What if more than one possible structure? • Find the formal charge • Draw possible Lewis structure • F.C. = EV - (EU + 1/2 EB) • EV = # valance electrons • EU = # unbounded electrons • EB = # bonded electrons http://highered.mcgraw-hill.com/olcweb/cgi/pluginpop.cgi?it=swf::100%::100%::/sites/dl/free/0072512644/117354/05_Formal_Charge_Calculations.swf::Formal Charge Calculations

  46. F F F F B B Be F F F F Be F F Formal Charge • Most likely structure has • F.C. close to zero • Negative F.C. on higher electronegative atoms • F.C. add up to Zero or Ion Charge!!!! BF3 BeF2

  47. HCN N2O H C N H C N N N O N O N H N C Formal Charge

  48. Polyatomics • Groups of atoms that have a charge. This tells us that extra electrons have been gained or lost to fulfill the octet rule of elements • Example – draw the Lewis dot for NO3

  49. Shapes of molecules • VSEPR Theory • Valence-Shell Electron-Pair repulsion theory • Electrons are negatively charged and therefore repel each other • Distances between electron pairs are maximized

  50. VSEPR Theory • Electron-pair geometry (parent structure) • Considers all sets of bonding and nonbonding electron pairs • Looks at how many pairs are around the central atom • Molecular geometry (molecular shape) • From the electron-pair geometry, determine the arrangement of the atoms in the molecule • Look to see what’s happening to the e-

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