1 / 45

Covalent Bonding

Covalent Bonding. Happens between nonmetals Share valence electrons between atoms Electron clouds overlap Ex: H 2 O CH 4 C 6 H 12 O 6 NH 3 CO 2. Electronegativity. Difference in EN is smaller than in ionics and is usually < 1.7 Ex: HCl H = 2.2 Cl = 3.2

clint
Télécharger la présentation

Covalent Bonding

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Covalent Bonding • Happens between nonmetals • Share valence electrons between atoms • Electron clouds overlap Ex: H2O CH4 C6H12O6 NH3 CO2

  2. Electronegativity • Difference in EN is smaller than in ionics and is usually < 1.7 • Ex: HCl H = 2.2 Cl = 3.2 Difference = 1.0

  3. Bond Polarity • Polar Bond: when there is a difference in EN values. (unequal sharing) • Ex: H Cl EN=2.2 EN=3.2

  4. NonPolar Bond: no difference in EN values. (equal sharing) Ex: O2, N2, Cl2, H2 (all the diatomics!)

  5. Single, Double, Triple Bonds • Atoms can share single double or triple bonds between them. • Each bond represents a shared pair of electrons. • http://youtu.be/1wpDicW_MQQ

  6. Naming Covalent Compounds • Use prefix system to indicate the number of atoms of each element present. • Mono • Di • Tri • Tetra • Penta • Hexa • Hepta • Octa • Nona • Deca

  7. Molecular Formulas • Covalent compounds are molecules. • Molecular formulas: show actual number of atoms of each element present in compound Ex: H2O 2 hydrogen atoms and 1 oxygen

  8. Structural Formulas • Show how the atoms are bonded together in a covalent molecule. • Use “lines” to show covalent bonds

  9. Empirical Formulas Empirical formulas: • Show simplest whole number ratio of atoms or ions in the compound. • Ionic compounds are “salts” or ionic crystals. • All ionic compounds have empirical formulas • Ex: MgCl2 1 : 2 ion ratio

  10. You can simplify some molecular formulas to make them empirical ratios Ex: C6H12O6 Simplest ratio of atoms CH2O

  11. Molecule vs. Ionic Crystal CH4 = 5 atoms in molecule NaCl = 1:1 ion ratio

  12. Properties of Covalent Compounds

  13. Melting Point • Lower than Ionics • To melt, you are only separating the weak bonds between molecules (not within).

  14. Melting Point • Polar Molecules (dipoles/”mini-magnets”): • Have higher melting points than non-polars because they are harder to separate.

  15. Solubility • Polar Molecules dissolve in Polar Solvents (like H2O, CHCl3, NH3 etc.) • Non-polar Molecules dissolve in Non-Polar Solvents (like hexane, CCl4) “Like Dissolves Like”

  16. Conductivity • Covalent Molecules do not conduct well as they do not form ions. • Except Acids!!!! • Acids are covalently bonded but in water (aqueous) they will ionize and conduct current. • (Acids are not on this test)

  17. Decompose • If the heat gets high enough covalent compounds will break down and decompose. • (Remember the lab, sugar melted first, then it burned and turned into black carbon)

  18. Shapes of Molecules And how do you draw them?

  19. VSEPR • Valence Shell Electron Pair Repulsion • Valence electrons will orient themselves around the “central” atom to be as far apart from each other as possible. • This influences the “shape” of the molecule.

  20. Polarity of a Molecule • Polar Molecules: • Have polar bonds and are not symmetrical • Nonpolar Molecules • Have nonpolar bonds OR • Have polar bonds and are symmetrical

  21. Regents Shapes to Know: • Tetrahedral • Pyramidal • Bent • Linear

  22. Honors Shapes to Know: See Honors Packet • Linear • Trigonal Planar • Bent (with 3 e- regions and with 4) • Trigonal Pyramidal • Tetrahedral • Trigonal Bipyramidal • SeeSaw • T-Shaped • Linear • Octahedral • Square Pyramidal • Square Planar • VSEPR SHAPES FOR HONORS • http://youtu.be/FjjhUI4wFTE • Example of Shapes (honors) • http://youtu.be/i3FCHVlSZc4

  23. Tetrahedral • Has 4 atoms bonded (no free pairs)

  24. Symmetry? Depends on what atoms are attached. Can be polar (asymmetrical) or nonpolar (symmetrical)

  25. Pyramidal • Three atoms bonded (one free pair)

  26. Symmetry? • All pyramids are asymmetrical. • These molecules are always POLAR!

  27. Bent • Two atoms attached (2 free pair) The 2 free pair make it bent and not linear. These are always asymmetrical so are always polar. H2O

  28. Hey, Water is Polar!!!!! Never forget this!!!

  29. Linear • 2 atoms in molecule • Ex: Cl Cl (nonpolar bond, sym., nonpolar molecule) H Br (polar bond, asym., polar molecule) • 3 atoms in molecule

  30. Drawing Polyatomic Ions • Covalently bonded atoms with a group charge • Follow same steps for drawing covalents. • Add or subtract electrons from total valence depending on charge. • Draw brackets around ion and indicate charge. • Ex: (SO4) -2 6 + 4(6) + 2 = 32 electrons

  31. Remember!! • If an ionic compound contains a polyatomic ion it contains both ionic and covalent bonds!!

  32. Coordinate Covalent Bonding • Covalent bond in which one of the bonding atoms donates both of the electrons to the bond. • The other atom donates nothing. Ex: Forming Hydronium Ion

  33. To form this type of bond you must have: • A molecule with a free pair of electrons • Something that needs to gain 2 electrons H+1

  34. Ex: Forming Ammonium Ion

  35. Exceptions to the Octet Rule (Honors) • Less than Octet • Ex: BeH2 • Odd Number of Valence Electrons • Ex: NO • More than Octet • Ex: SF6

  36. Resonance(Honors) • Electrons “resonate” between multiple bonding sites. • Adds strength to all the bonds. • Check if it is possible to draw double or triple bonds at more than one bonding site.

  37. Sigma and Pi Bonds (Honors) • Sigma Bond: • End to end overlap of orbitals • All single bonds are sigma • One of the bonds in a multiple bond is a sigma • Pi Bond: • Side to side overlap of orbitals. • Found in double or triple bonds.

  38. Network Solids • Giant network of covalently bonded atoms. • Large macromolecules • Extremely strong structures • Unusually high M.P. • Do not dissolve Diamonds are a giant network of carbon atoms.

  39. Ex: C (s) (graphite, diamond, buckyball), SiO2 (quartz), GeO2 “Buckyball”

  40. Metallic Bonding • Pure metals or alloys. • Ex: Mg, Fe, Brass, Au, Ni, Cu • “Delocalized” valence electrons move about between all the metal atoms.

  41. Metallic Bonding Properties • Conducts heat and electricity very well Conducts as a solid too! • Does not dissolve in solvents • Malleable and Ductile • Relatively high melting point. • Higher MP than covalents. • Similar MP to most ionics

  42. Metallic Bonding • http://www.drkstreet.com/resources/metallic-bonding-animation.swf

  43. Ionic vs. Covalent Bonding • http://youtu.be/QqjcCvzWwww • http://youtu.be/yjge1WdCFPs • Electronegativity • http://youtu.be/Kj3o0XvhVqQ • Songs • It’s a chemical bond baby” • http://youtu.be/wWUYHHo-zB0 • Dancin Queen (Ionic/Covalent Bonds) • http://youtu.be/BCYrNU-7SfA • Isn’t it Ionic • http://youtu.be/rwRtfrgJL5E

More Related