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Covalent Bonding

Learn about covalent bonding, bond polarity, molecular formulas, Lewis structures, VSEPR theory, molecular shapes, and more concepts related to covalent compounds. Discover how to draw Lewis structures and predict molecular geometries. Understand polar and nonpolar molecules along with their properties.

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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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