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BONDING

BONDING. TOPIC 4. Terms. Covalent Bonding. Bonds Breaking them takes energy Making them gives off energy. Exothermic More energy is given off than put in Endothermic More energy is absorbed than given off Intra molecular Forces Forces within molecules (ionic, covalent and metallic)

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BONDING

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  1. BONDING TOPIC 4

  2. Terms Covalent Bonding • Bonds • Breaking them takes energy • Making them gives off energy

  3. Exothermic • More energy is given off than put in • Endothermic • More energy is absorbed than given off • Intramolecular Forces • Forces within molecules (ionic, covalent and metallic) • Intermolecular Forces (IMF) • Forces between particles

  4. Ionic Bonding + - Less e- = Less e- repulsion More e- = e- more repulsion. Metal: K Non-Metal: Cl • If the electronegative difference between the atoms involved is =>1.8 • There are always exceptions to this rule! • Will conduct electricity in its molten or aqueous state (This test proves ionic)

  5. Intramolecular Forces Drawing Ionic Bonding Lewis Dot Diagram - + X Electrons are in pairs Na Cl Special Note: The ionic bond is the electrostaticattractionbetween oppositely charged ions! Ionic Bonding • Just use the valence shell • Be sure to include square brackets and charge after electron exchange.

  6. Lewis Diagrams Combine C Al Fe Mg Be Cl F Cl O Br Lewis Dot diagrams us the atoms valance shell electrons

  7. Decomposition Intramolecular Forces 2Na+(aq) + 2Cl-(aq)  2Na(s) + Cl2(g) + + + + + + - - CATHODE (-) ANODE (+) + - + + + - - + + Conductivity is FINITE NaCl • When in molten or aqueous state, ionic substances WILL conduct electricity, by the movement of (+) and (-) ions. • This is different from how METALS conduct electricity!

  8. Ionic Compounds Giant Ionic Lattices Force Cation Anion + - + - + - + - + - + - + - + - + - + - + - + - + Like charges repel Metal: K Non-Metal: Cl • No bonds are made!!! • Static attractions holds them together. (opposites attract) • When a force is applied, ionic compounds will make a clean break. • Physical characteristics • Hard and brittle • Solid doesn’t conduct Electricity • More soluble in water than other solvents • High MP and BP

  9. Cubic or Isometric Giant Ionic Lattices Table Salt NaCl

  10. Tetragonal Giant Ionic Lattices Cassiterite SnO2

  11. Orthorhombic Giant Ionic Lattices Agagonite CaCO3 • Also found in mollusk shells and coral

  12. Hexagonal Giant Ionic Lattices Beryl Be3Al2(SiO 3)6

  13. Trigonal Giant Ionic Lattices Quartz SiO2

  14. Ionic Bonding Giant Ionic Lattices Beryl Be3Al2(SiO 3)6

  15. Triclinic Giant Ionic Lattices Copper(II) Sulfate CuSO4

  16. Intramolecular Forces Transition Metals Fe2+ Cu+ Fe3+ Cu2+ Iron(II) Oxide Copper(I) Oxide Iron (III) Oxide Copper(II) Oxide Multiple Ions • Transition metals can have multiple ions. • Ones you should know.

  17. Ions Reminder SO4-2 PO4-3 NO3- CO3-2 OH- HCO3- NH4+ Polyatomic Ions • Be sure to review your polyatomic ions!!!

  18. Covalent Bonding Topic 4

  19. Intramolecular Forces Covalent Bonding 2.1 3.0 X Differences |3-2.1| =0.9 H Cl Special Note: The covalent bond is the electrostaticattractionbetween pairs of e- and positively charged nuclei! COVALENT BONDING • If the electronegative difference between the atoms involved is <1.8 • Will NOT conduct electricity • Electrons are shared

  20. Questions Review Na Ca + CO3 Cl + Li + O + Na SO3 NO3 K + For ionic compounds to form the valance shells of both metal and non-metal must be full!! What is the chemical formula? What is the names for each?

  21. Intramolecular Forces Covalent Bonding H H X X C Cl H H H X H X H H COVALENT BONDING • Structural formula • Lewis structure

  22. Lewis Structures Intramolecular Forces H2O 1 1 6 H H Hydrogen can only hold 2e- remaining must be paired on Oxygen O - 4 = 4 8 COVALENT BONDING • 1) Sum all valence e- • 2) Subtract 2e- for every bond • 3) Place e- around periphery atoms to form octets. The remaining around central atom • 4) All atoms MUST be paired!!!!!!

  23. Intramolecular Forces Lewis Structures Covalent Bonding HL: PCl5, PCl4+, PCl6- and XeF4 • Draw the following Lewis structures • H2 Cl2 • O2 N2 • HCN C2H6 • C2H4 C2H2

  24. Special Lewis Structures Intramolecular Forces + + Lone pair of e- H N H H Electrophile H Covalent Bonding • Coordinate or dative covalent bonds • When both e- are shared from the same atom. (Not one from each as before) • Occurs when a non bonding e- pair donates an e- to an e- deficient atom.

  25. Intramolecular Forces Special Lewis Structures Covalent Bonding • Draw the following Lewis structures • CO • H3O+

  26. Length, Strength & Hybrid Resonance Intramolecular Forces O 2- 2- 2- O O C O O C Don’t forget to show the e- pairs!! C O O O O CO32- • More bonds = more strength & shorter bonds • Resonance structures • Bond length is longer than a double bond but shorter than a single bond

  27. Length & Strength Intramolecular Forces Ethene Carboxylic Acid Ethyne O H H C R = Functional Group C C H C C H R OH H H CO32- • Compare the two molecules • Ethyne has stronger and shorter bonds • C=O bond is stronger and shorter due to Oxygen being more electronegative

  28. Bond Polarity Intramolecular Forces δ+ δ- Dipole Moment H X Cl Covalent Bonding • Non-Metals are fighting for e- • Atom with larger electronegativity will hold the e- closer to itself. • Atoms become slightly charged.

  29. Exceptions to the Octet Rule Intramolecular Forces F B F F Covalent Bonding • BF3 • Actual structure: Boron is e- deficient • This is known because of its reactivity towards electron rich molecules such as NH3 • CNOF all obey the octet rule.

  30. Intramolecular Forces Formal Charge Covalent Bonding • SO42- • Single bonds (8 e- around S) • Double bonds (12 e- around S) • Formal Charge = (# valence e- on free atom) – (# valence e- assigned to the atom in the molecule) • (Valence e-)assigned = (# lone pair e-) + ½ (# of shared e-) • 1) Molecules attempt to achieve Formal Charge as close to 0 as possible. • 2) Any negative Formal charge will reside on most electronegative atom.

  31. VSEPR (shape) Intramolecular Forces 3 Pairs of e- 120o 2- O F C B 2 Pairs of e- 180o C O O O O F F Covalent Bonding • VSEPR (Valence Shell Electron Pair Repulsion) • Paired e- attempt to get as far away from each other as possible. • Multiple bonds still count only as 1 pair!!

  32. VSEPR Intramolecular Forces 4 Pairs of e- 109.5o H Lone pair 107o Lone pair 104.5o O N C H H H H H H H H Covalent Bonding • Tetrahedral • Lone pair e- have increased charge density and require more room • More repulsion from lone pair will decrease bond angle.

  33. Intramolecular Forces Home Work Covalent Bonding • Predict the shape AND bond angles • H2S PbCl4 H2CO SO2 • NO3- PH3 NO2- • NH2- POCl3 CO2

  34. HL VSEPR

  35. HL VSEPR

  36. Expanded Valance Shell (14.1) • Molecules with more than 8 electrons • Electron promotion:

  37. Dipole Moment Molecule Polarity (4.2.6) 2δ- Cl Cl Cl δ- H δ- Non Polar O δ+ C C H H H H δ+ δ+ δ+ H H H Covalent Bonding • Polarity effects state change (physical change) • Unequal sharing causes a dipole moment to form • Q: Why is BF3 non-polar whereas PF3 is polar?

  38. Hybridization (14.2.2) • Sigma bond: σ (single bond) • Axial overlap of orbital’s 1s1 2px2 py2 pz2 H Cl

  39. Hybridization (14.2) • Sigma bond: σ (single bond) • Axial overlap of orbital’s Cl Cl

  40. Hybridization (14.2) • Pi bond: π(Double bond, one σ bond) • Parallel overlap of orbital’s O O N N

  41. Hybridization (14.2.3) • Hybridization electron promotion • New Orbital sp3 2s2 2px2 py2 pz2 Excited State Ground State C 4 Equal orbital`s capable of holding a maximum of 2 electrons each

  42. Hybridization (14.2) • How to determine Hybridized orbital`s • Look at the shape

  43. Carbon Allotropes C C C C C Giant Covalent • 1) Diamond (Tetrahedron, localized e-) • Very hard and does not conduct electricity • 2) Fullerenes (C60) Hexagonal and pentagonal rings • Nanotubes

  44. Allotropes Carbon C C C C Weak Pi Bonds C C HL: sp hybrid Delocalized electrons able to move Giant Covalent • 3) Graphite (Planar, delocalized e-) • Weak pi bonding between sheets cause it to conduct electricity and be slippery. • Bonds are shorter than a tetrahedral due to the pi bonding

  45. Benzene (14.3) Pi bonds overlap allowing for electrons to be delocalized over the entire molecule. C C C C C C C6H6 • Planar, delocalized e- • Regular bonding would predict an alternating double bond (Resonance structure) • Hybrid theory shows sp2 configuration

  46. Intramolecular Forces Silicon Si Si Si Si Si Si Si Si Si Si Silicon Tetrahedron Configuration Similar to diamond

  47. Intramolecular Forces Silicon & Silicon dioxide SiO2 but based on a network of SiO4 O Si O O O Quartz • Single bonds formed between Oxygen to satisfy the octet. • HL: Less overlap in the P-sub orbital due to atomic size difference therefore Pi bonds do not form.

  48. Metallic Bonding Topic 4

  49. Metallic Bonding Intramolecular Forces + + + + - + + + + - - - Sea of electrons + + + + + - - + + + + - - Conductivity is INFINITE Metallic Bonding • In solid state • Outer e- are delocalized and free to move about • Bond is a result of electrostatic attraction between Fixed positive metal ions and delocalized e-

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