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Chapter 6 -Chemical Bonding

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Chapter 6 -Chemical Bonding

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  1. Chapter 6 -Chemical Bonding Batrachotoxin A steroid alkaloid derived from skin secretions of the Phyllobates and Dendrobates genera of South American poison-arrow frogs. It is one of the most potent venoms known.

  2. Bonds • Forces that hold groups of atoms together and make them function as a unit. • Ionic bonds– transfer of electrons • Covalent bonds– sharing of electrons • Polar Covalent bonds – unequal sharing of electrons that results in an unbalanced distribution of charge

  3. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994 e-

  4. Table of Electronegativities

  5. Shielding Effect • Electrons in the inner energy levels block the attraction of the nucleus for the valence electrons. • Shielding increases down a group. • This causes EN values to decrease.

  6. What is Nuclear Shielding? The nucleus (+) pulls the electron (-) close to the core. e- Inner electrons shield outer electrons from the nuclear pull. e- e- P+ P+ P+ P+ The further the electron is away from the nucleus, the weaker the nuclear pull. P+ P+ P+ P+ P+ e-

  7. ∆EN Values Range of EN Values 3.3 2.0 0.5 0 Polar Covalent Mostly Covalent Mostly Ionic Polar covalent- Electrons are shared, but unequally. There is some degree of ionic character in these bonds.

  8. Nonpolar Covalent Polar Covalent Ionic EN ∆ 1.7 0.3

  9. NaCl Sodium’s EN = 0.9 Chlorine’s EN = 3.0 3.0 - 0.9 = 2.1 therefore it is mostly ionic H2O Hydrogen’s EN = 2.1 0xygen’s EN = 3.5 3.5 - 2.1 = 1.4 therefore it is polar covalent Practice Problems Determine the bond type using EN values:

  10. Chapter 6 Chemical Bonding 6.2 Covalent Bonds

  11. Covalent Bonds

  12. Covalent Terms • Molecule: A neutral group of atoms that are held together by covalent bonds • Diatomic Molecule: A molecule containing only two atoms • Molecular Compound: A chemical compound whose simplest units are molecules • Chemical Formula: Indicates the relative numbers of atoms of each kind of a chemical compound by using atomic symbols and numerical subscripts • Molecular Formula: Shows the types and numbers of atoms combined in a single molecule of a molecular compound Example: BrINClHOF

  13. Bonding Forces • Electron – Electron • repulsive forces • Proton – Proton • repulsive forces • Electron – Proton • attractive forces

  14. Pure Covalent Bonding

  15. Bond Length Diagram

  16. Bond Energy • It is the energy required to break a bond. • It gives us information about the strength of a bonding interaction.

  17. Electron Dot Notation

  18. The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

  19. Hydrogen Chloride by the Octet Rule

  20. Formation of Water by the Octet Rule

  21. Comments About the Octet Rule • 2nd row elements C, N, O, F observe the octet rule. • 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. • 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. • When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

  22. Lewis Structures • Shows how valence electrons are arranged among atoms in a molecule. • Reflects central idea that stability of a compound relates to noble gas electron configuration.

  23. Completing a Lewis Structure CH3Cl • Make carbon the central atom • Add up available valence electrons: • C = 4, H = (3)(1), Cl = 7 Total = 14 • Join peripheral atoms to the central • atom with electron pairs. H .. .. H C Cl .. .. .. .. .. H • Complete octets on atoms other than hydrogen with remaining electrons

  24. How to Make a Lewis Dot Structure for a Molecule • Draw the Lewis Dot Diagram for each element • Place the atom with the lowest EN value in the center • Attach the rest of the atoms to the central • If lone electrons exist on adjacent atoms, pair them for multiple bonds.

  25. Multiple Covalent Bonds:Double Bonds • Two pairs of shared electrons • Higher bond energy and shorter bond length than single bonds

  26. Multiple Covalent Bonds:Triple bonds • Three pairs of shared electrons • Higher bond energy and shorter bond length than single or double bonds

  27. Resonance • Occurs when more than one valid Lewis structure can be written for a particular molecule. • These are resonance structures. • The actual structure is an average of • the resonance structures.

  28. Resonance in Ozone Neither structure is correct.

  29. Resonance in Polyatomic Ions Resonance in a carbonate ion: Resonance in an acetate ion:

  30. Covalent Network Compounds Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms

  31. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical

  32. Fundamental Properties of Models • A model does not equal reality. • Models are oversimplifications, and are therefore often wrong. • Models become more complicated as they age. • We must understand the underlying assumptions in a model so that we don’t misuse them.

  33. Chapter 6 Chemical Bonding 6.3 Ionic Bonding

  34. Ionic Bonds

  35. Ionic Bonds • Electrons are transferred • Electronegativity differences are generally greater than 1.7 • The formation of ionic bonds is always exothermic!

  36. Ionic Bonding • Formula Unit: The simplest collection of atoms from which an ionic compound's formula can be established • Lattice Energy: The energy released when one mole of an ionic crystalline compound is formed from gaseous ions Na+ (g) + Cl- (g)→ NaCl (s) + 787.5 kJ • Formation of ionic compounds is ALWAYS exothermic

  37. Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions. Formation of sodium chloride Na = 3s1 Cl = 3s23p5 Na+ = 2s22p6 Cl- = 3s23p6

  38. Polyatomic Ions • A charged group of covalently bonded atoms • Creation of octets results in an excess or deficit of electrons

  39. A Comparison of Ionic and Molecular Compounds

  40. Chapter 6 Chemical Bonding 6.4 Metallic Bonding

  41. The Metallic Bond Model • The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons • Electron Delocalization in Metals • Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move freely throughout the metal • The valence electrons do not belong to any one atom!

  42. Swim in the Sea of Valence Electrons… • In metals, the valence electrons are held loosely. The vacant p & d orbitals overlap. • Metal atoms DO NOT lose their valence electrons in metallic bonding, rather they release them into a “Sea of Electrons” • Although the atoms are bonded together, they are not bonded to any one particular atom, it is more like a large network.

  43. Sea of Valence Electrons…

  44. Bonding Between Metals • Results in an interaction that hold metal atoms together, however, it is not called a compound. • Special Properties result from this interaction: • Malleable- pounded/ rolled into sheets (aluminum foil) • Ductile- Drawn into wire. (copper wires) • Conductivity- the flow of electrons • Luster- Shiny-The narrow range of energy differences between orbitals allows electrons to be easily excited, and emit light upon returning to a lower energy level

  45. Metallic Properties • Metals are good conductors of heat and light • Metals have luster (shiny) • The narrow range of energy differences between orbitals allows electrons to be easily excited, and emit light upon returning to a lower energy level • Metals are Malleable- can be hammered into thin sheets • Metals are ductile- ability to be drawn into wire • Metallic bonding is the same in all directions, so metals tend not to be brittle

  46. Metallic Bond Strength • Heat of Vaporization • The ease with which atoms in a metallic solid can be separated from one another into individual gaseous atoms is related to bond strength

  47. Chapter 6 Chemical Bonding 6.5 Molecular Geometry

  48. VSEPR Model (Valence Shell Electron Pair Repulsion) • The structure around a given atom is determined principally by minimizing electron pair repulsions. • The model for predicting molecular shapes • VSEPR Theory:Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible

  49. VSEPR and Unshared Electron Pairs • Unshared pairs take up positions in the geometry of molecules just as atoms do • Unshared pairs have a relatively greater effect on geometry than do atoms • Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs • Lone pairs do not cause distortion when bond angles are 120° or greater