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Chemical BONDING

Chemical BONDING

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Chemical BONDING

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  1. ChemicalBONDING

  2. Chemical Bond • A bond results from the attraction of nuclei for electrons • All atoms trying to achieve a stable octet • IN OTHER WORDS • the p+ in one nucleus are attracted to the e- of another atom • Electronegativity

  3. You complete me. What did the atom of fluorine say to the atom of sodium?

  4. Two Major Types of Bonding • Ionic Bonding • forms ionic compounds • transfer of e- • Covalent Bonding • forms molecules • sharing e-

  5. One minor type of bonding • Metallic bonding • Occurs between like atoms of a metal in the free state • Valence e- are mobile (move freely among all metal atoms) • Positive ions in a sea of electrons • Metallic characteristics • High mp temps, ductile, malleable, shiny • Hard substances • Good conductors of heat and electricity as (s) and (l)

  6. It’s the mobile electrons that enable me-tals to conduct electricity!!!!!!

  7. IONic Bonding • electrons are transferred between valence shells of atoms • ionic compounds are made of ions NOT MOLECULES • ionic compounds are calledSaltsorCrystals

  8. IONic bonding • Always formed between metals and non-metals [METALS ]+ [NON-METALS ]- Lost e- Gained e-

  9. IONic Bonding • Electronegativity difference > 2.0 • Look up e-neg of the atoms in the bond and subtract NaCl CaCl2 • Compoundswithpolyatomic ions NaNO3

  10. Properties of Ionic Compounds SALTS Crystals • hard solid @ 22oC • high mp temperatures • nonconductors of electricity in solid phase • good conductors in liquid phase or dissolved in water (aq)

  11. Covalent Bonding molecules • Pairs of e- are shared between non-metal atoms • electronegativity difference < 2.0 • forms polyatomic ions

  12. Properties of Molecular Substances Covalent bonding • Low m.p. temp and b.p. temps • relatively soft solidsas compared to ionic compounds • nonconductors of electricity in any phase

  13. NO2 sodium hydride Hg H2S sulfate NH4+ Aluminum phosphate KH KCl HF Covalent, Ionic, metallic bonding? • CO • Co Also study your characteristics!

  14. Drawing ionic compounds using Lewis Dot Structures • Symbol represents the KERNEL of the atom (nucleus and inner e-) • dots represent valence e-

  15. How did we get here? NaCl • This is the finished Lewis Dot Structure [Na]+ [ Cl ]-

  16. Step 1 afterchecking that it is IONIC • Determine which atom will be the +ion • Determine which atom will be the - ion • Step 2 • Write the symbol for the + ion first. • NO DOTS • Draw the e- dot diagram for the – ion • COMPLETE outer shell • Step 3 • Enclose both in brackets and show each charge

  17. Draw the Lewis Diagrams • LiF • MgO • CaCl2 • K2S

  18. Drawing molecules using Lewis Dot Structures • Symbol represents the KERNEL of the atom (nucleus and inner e-) • dots represent valence e-

  19. Always remember atoms are trying to complete their outer shell! The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four

  20. How did we get here? Methane CH4 • This is the finished Lewis dot structure

  21. Step 1 • count total valence e- involved • Step 2 • connect the central atom (usually the first in the formula) to the others with single bonds • Step 3 • complete valence shells of outer atoms • Step 4 • add any extra e- to central atom IF the central atom has 8 valence e- surrounding it . . YOU’RE DONE!

  22. Sometimes . . . • You only have two atoms, so there is no central atom, but follow the same rules. • Check& Share to make sure all the atoms are “happy”. Cl2 Br2 H2 O2 N2 HCl

  23. DOUBLE bond • atoms that share two e- pairs (4 e-) O O • TRIPLE bond • atoms that share three e- pairs (6 e-) N N

  24. Draw Lewis Dot Structures You may represent valence electrons from different atoms with the following symbols x, , CO2 NH3

  25. Draw the Lewis Dot Diagram for polyatomic ions • Count all valence e- needed for covalent bonding • Add or subtract other electrons based on the charge REMEMBER! A positive charge means it LOST electrons!!!!!

  26. Draw Polyatomics • Ammonium • Sulfate

  27. Types of CovalentBonds • NON-Polar bonds • Electrons shared evenly in the bond • E-neg difference is zero Between identical atoms Diatomic molecules

  28. Types of Covalent Bonds Polar bond • Electrons unevenly shared • E-neg difference greater than zero but less than 2.0 closer to 2.0 more polar more “ionic character”

  29. H H C H H non-polar MOLECULES • Sometimes the bonds within a molecule are polar and yet the molecule is non-polar because its shape is symmetrical. Draw Lewis dot first and see if equal on all sides

  30. Polar molecules (a.k.a. Dipoles) • Not equal on all sides • Polar bond between 2 atoms makes a polar molecule • asymmetrical shape of molecule

  31. H Cl + -

  32. O Water is asymmetrical + + H H -

  33. H H H H O Water is a bent molecule

  34. Making sense of the polar non-polar thing BONDS Non-polar Polar Identical Different MOLECULES Non-polar Polar Symmetrical Asymmetrical

  35. IONIC bonds …. Ionic bonds are so polar that the electrons are not shared but transferred between atoms forming ions!!!!!!

  36. VSEPR Theory • Valence Shell Electron Pair Repulsion Theory • Electron pairs orient themselves in order to minimize repulsive forces. C. Johannesson

  37. Lone pairs repel more strongly than bonding pairs!!! VSEPR Theory • Types of e- Pairs • Bonding pairs - form bonds • Lone pairs - nonbonding e- C. Johannesson

  38. 4 Shapes of molecules

  39. 1. Linear (straight line) Ball and stick model Space filling model

  40. 2. Bent Ball and stick model Space filling model

  41. 3.Trigonal pyramid Ball and stick model Space filling model

  42. 4.Tetrahedral Ball and stick model Space filling model

  43. Intermolecular attractions • Attractions between molecules • van der Waals forces • Weak attractive forces between non-polar molecules • Hydrogen “bonding” • Strong attraction between special polar molecules

  44. van der Waals • Non-polar molecules can exist in liquid and solid phases because van der Waals forces keep the molecules attracted to each other • Exist between CO2, CH4, CCl4, CF4, diatomics and monoatomics

  45. van der Waals periodicity • increase with molecular mass. • increase with closer distance between molecules • Decreases when particles are farther away

  46. H “bond” Hydrogen “Bonding” • Strong polar attraction • Like magnets • Occurs ONLY between H of one molecule and N, O, F of another

  47. His shared between 2 atoms of OXYGEN or 2 atoms of NITROGEN or 2 atoms of FLUORINE Of 2 different molecules

  48. Why does H “bonding” occur? • Nitrogen, Oxygen and Fluorine • small atoms with strong nuclear charges • powerful atoms • very high electronegativities

  49. Intermolecular forces dictate chemical properties • Strong intermolecular forces cause high b.p., m.p. and slow evaporation (low vapor pressure) of a substance.