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Covalent Bonding

Covalent Bonding. B etween nonmetal atoms Share valence electrons between atoms Electron clouds overlap Ex: H 2 O CH 4 NH 3 CO 2. Electronegativity. Difference in EN smaller than in ionics and is usually < 1.7 Ex: HCl H = 2.2 Cl = 3.2 Difference = 1.0.

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Covalent Bonding

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  1. Covalent Bonding • Between nonmetal atoms • Share valence electrons between atoms • Electron clouds overlap Ex: H2O CH4 NH3 CO2

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

  3. Bond Polarity Polar Bonds: • Have 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)

  5. Comparing Bond Types

  6. 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 3 pair here = 6 electrons!

  7. Molecular Formulas • Covalent compounds are molecules. • Made up of all nonmetals. Molecular formulas: show actual number of atoms of each element present Ex: H2O 2 hydrogen atoms, 1 oxygen atom

  8. Ionic Compounds (SALTS) formula unit indicates simplest ratio of ions in the crystal structure • Covalent Compounds (MOLECULES) molecular formula indicates actual number of atoms present in molecule NaCl = 1:1 ion ratio CH4 = 5 atoms in molecule Formula Units vs. Molecules 7 minutes https://www.youtube.com/watch?v=dHWqJeSs8ms

  9. Structural Formulas of Molecules • Show how the atoms are bonded together • Use “lines” to show covalent bonds • Dots show free electron pairs

  10. Empirical Formulas • Show simplest whole number ratio of atoms or ions in the compound. Ex: MgCl2 1 : 2 ion ratio Al2(SO4)3 2 : 3 ion ratio NOTE: • All ionic compounds have empirical formulas

  11. You can simplify some molecular formulas to make them empirical ratios Ex: C6H12O6 Simplest ratio of atoms ___________ C6H6 Simplest ratio of atoms ___________ CO2 Simplest ratio of atoms ___________

  12. Naming Binary Covalent Compounds Prefix system indicates number of atoms Add “-ide” ending

  13. Examples of Naming Covalents • Note: use “mono” prefix if only one of 2nd element.

  14. Drawing Covalent Molecules Lewis Structures

  15. STEPS TO DRAW MOLECULE • Count total valence e- in the molecule • Draw molecule with single bonds between atoms then subtract these e- from total • Evenly distribute remaining e- in pairs to all atoms in molecule that still need e- • Check to see if all obey octet rule (Hydrogen is 2 e-) • If deficient, shift over free e- pairs to make double or triple bonds as needed.

  16. Draw NH3 • Draw H2O • Draw CH4

  17. Crash Course Chemistry: Lewis Structures: (11 minutes) http://www.youtube.com/watch?v=a8LF7JEb0IA

  18. Drawing Polyatomic Ions • Covalently bonded atoms with a group charge • 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

  19. VSEPR and Molecular Shape • Assume valence electrons repel each other. • Molecule adopts 3D geometry that minimizes this repulsion. • Valence Shell Electron Pair Repulsion (VSEPR) theory.

  20. Predicting Molecular Geometries • Draw Lewis structure • Count total number of electron pairs around the central atom (both shared and unshared) • Arrange electron pairs to minimize e-repulsion • Multiple bounds count as one bonding pair

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

  22. Is a Molecule Polar? • If the centers of negative and positive charge do not coincide, then the molecule is polar.

  23. Look for Symmetry • Polar Molecules: • Have polar bonds and are not symmetrical • Positive & negative “partial charges” don’t overlap • They have a “dipole moment” • Nonpolar Molecules • Have nonpolar bonds OR • Have polar bonds and are symmetrical • Centers of positive & negative charge overlap

  24. Example: In CO2, the polarity of each C-O bond is cancelled because the molecule is linear. In H2O, the polar H-O bonds do not cancel because the molecule is bent.

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

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

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

  28. All pyramids are asymmetrical. • These molecules are always POLAR!

  29. 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

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

  31. Linear: 2 or 3 atoms in a line • Can be polar or nonpolar depending on type of bonding and symmetry • EX: Cl2, O2, N2, HCl, CO2

  32. Ex: Linear Symmetrical Molecules

  33. Properties of Covalent CompoundsNote: These properties can vary depending on if molecule is polar or not!

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

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

  36. All these nonpolar diatomics have really low MP/BP

  37. Solubility “Like Dissolves Like” • Polar Molecules dissolve in polar solvents as they are attracted to them • like H2O, CHCl3, NH3 etc. • Non-polar Molecules dissolve in non-polar solvents • like hexane, CCl4

  38. Oil and water don’t mix! • How does soap work?

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

  40. 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)

  41. Other Types of Covalent Bonds

  42. 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

  43. 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

  44. Ex: Forming Ammonium Ion

  45. 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.

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

  47. Bonding in Pure Metals Hunting the elements: https://www.youtube.com/watch?v=Yu7Sq2zjBDc

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