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Covalent Bonding Chapter 8

Covalent Bonding Chapter 8. Chemistry 2. Molecular Compounds 8.1. Molecules and Molecular Compounds 8.1. Covalent bond – SHARE e- Molecule – neutral group of atoms joined by covalent bond Diatomic = 2 atoms = O 2 Molecular Compound – compound composed of molecules = CO or CO 2

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Covalent Bonding Chapter 8

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  1. Covalent BondingChapter 8 Chemistry 2

  2. Molecular Compounds 8.1

  3. Molecules and Molecular Compounds 8.1 • Covalent bond – SHARE e- • Molecule – neutral group of atoms joined by covalent bond • Diatomic = 2 atoms = O2 • Molecular Compound – compound composed of molecules = CO or CO2 • Lower mp & bp than ionic compounds • Normally 2 or more nonmetals

  4. Molecular Formula 8.1 • Chemical formula • H20 or CO2 or O2 • 1 is omitted if there is only 1 atom • No structure or arrangement of atoms • Use diagrams

  5. The Nature of Covalent Bonding 8.2

  6. The Octet Rule in Covalent Bonding 8.2 • Share electrons to attain e- conf. of noble gas = 8 e- • Combination s of Group 4A, 5A, 6A & 7A likely to form covalent bonds

  7. The Electron Probability Distribution for the H2 Molecule 7

  8. Single Covalent Bonds 8.2 • 2 atoms held together by single pair of e- • 2 dots or H-H by each other represent • Structural Formula = H-H • Represent bonds and arrangement • Unshared pair = valence e- that is not shared

  9. Double or Triple Covalent Bonds 8.2 • Share 2 pairs or 3 pairs of e-

  10. Practice Problem • Write a lewis structure for CCl2F2 Step 1: Arrange Atoms (Carbon is “Central Atom” because is has the lowest group number and lowest electronegativity. Step 2: Determine total number of valence electrons 1 x C(4) + 2 x Cl(7) + 2 x F(7) = 32 Step 3: Draw in valence electrons Step 4: Draw single bonds in replace of 2 electrons between 2 atoms and subtract 2 e- for each single bond (4 x 2 = 8) so 32 – 8 = 24 remaining 11

  11. Coordinate Covalent Bonds 8.2 • Covalent bond in which one atom contributes both bonding e- • Molecular Formula = CO • Structural Formula = C O • Polyatomic Ion – tightly bound group of atoms that has a + or – charge and behaves as a unit • H+ attaches to NH3’s unshared e- • LOOK at page 225 SO3-2

  12. Bond Dissociation Energies 8.2 • E required to break the bond between 2 covalently bonded atoms • H + H  H2 = gives off large amount of heat • Product more stable than reactants • Big b.d.e. = strong covalent bond = normally unreactive • C-C = 347 kJ/mol • C C = 657 kJ/mol • C C = 908 kJ/mol

  13. Resonance • 2 or more possible e- dot structures • No back and forth changes actually occur • Just a way to vision • Drawing • Must adhere to octet rule • Sigma bonds not altered, pi and nonbonding e- are altered

  14. Exceptions to the Octet Rule • Can occur when odd number of valence e- • Atom requires less than octet of 8 e- • BF3-NH3 • Some expand octet to 10 or 12 (esp w/ P and S

  15. Bonding Theories 8.3

  16. Molecular Orbitals 8.3 • Orbitals overlap • REMEMBER: atomic orbitals are orbitals in s,p,d,f • Bonding orbital – molecular orbital that can be occupied by 2 e- of covalent bond • SIGMA BONDS σ • 2 atomic orbitals combine • Directly between 2 nuclei • Single bonds • p overlaps end to end • Pi Bonds π • 2nd bond of double bond, 2nd and 3rd bond of a triple bond (sigma is 1st of double) • Makes up 2 lobes • Tend to be weaker than sigma bonds • Orbital overlapping is less

  17. VSEPR Theory • Valence-shell electron pair repulsion theory = explains 3-D shape

  18. Hybrid Orbitals • Provided info about molecular bonding and molecular shape • Atomic orbitals mix to form same total number of equivalent hybrid orbitals • Single Bonds • CH4 = sp3

  19. http://www.chemguide.co.uk/atoms/bonding/covalent.html http://www.mikeblaber.org/oldwine/chm1045/notes/Geometry/Hybrid/Geom05.htm

  20. Hybrid Orbitals • Double Bonds

  21. Polar Bonds and Molecules 8.4

  22. Bond Polarity 8.4 • Nonpolar covalent bond – equally share electrons • Polar Covalent bond – unequal sharing • The more electronegative, the more strongly pulls on e- • Less electronegative atom = slightly δ+ charge • More electronegative atom = slightly δ- charge

  23. Use table 6.2 in Chapter 6 for electronegativity of elements • HCl • H = 2.1 • Cl = 3 • Electronegativity = .9 • Conceptual Problem Page 239 # 30-31

  24. Polar Molecules 8.4 • Often in a polar bond One end of molecule is slightly – and other end slightly + • Call DIPOLE • Ex: HCl

  25. Attractions Between Molecules 8.4 • Intermolecular attractions weaker than ionic or covalent bonds… but they are important!!! HOW? • Determine if solids, liquids, and gases • Surface tension • Van der Waals ForceS • Dipole Interactions – polar molecules attracted to one another • Similar to ionic but weaker • Dispersion Forces – caused by motion of e- • Temporarily attractive force that results when the e- in 2 adjacent atoms occupy positions that make them temporarily dipole • Weakest of all interactions • Occurs even in non-polar

  26. Hydrogen Bonds 8.4 • Attractive forces in which a hydrogen covalently bonded to a very electronegative atom is weakly bonded to unshared e- pair

  27. Intermolecular Attractions and Molecular Properties 8.4 • Physical properties = depends on type of bonding • Melting/boiling point = lower for covalent compared to ionic • Few covalent bonds have high mp • Most network solids (crystals) – solids in which all the atom s are covalently bonded • Ex: diamond • Each C is attracted to 4 other C’s

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