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Chemical Bonding: Basic Concepts

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  1. Chemical Bonding:Basic Concepts Chapter 6

  2. Why do atoms form bonds? To become more stable

  3. What is a bond? • A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

  4. Bonding Forces Electron – electron repulsive forces Proton – proton repulsive forces Electron – proton attractive forces

  5. Three types of bonds

  6. What type of bond will form? • Type of bond is dependent on differences in electronegativity

  7. Remember- Electronegativity is … A measure of the ability of an atom in a chemical compound to attract electrons

  8. Rule of thumb • Metal and non-metal usually form an ionic bond • Non-metal and non-metal from a covalent bond • Metals come together to form metallic bonds

  9. See Electronegativity Table Pg. 161 What type of bond? Chlorine & Calcium 3.0 – 1.0 = 2.0 Ionic Chlorine Covalent Chlorine & Oxygen 3.5 – 3.0 = 0.5 Oxygen Covalent Bromine & Chlorine 3.0 – 2.8 = 0.2 Chlorine

  10. Pause for a CauseCreate a “PFC” page in your notebook to use for this unit. PFC #1 Determine what type of bond is present in each of the following • Magnesium chloride • Water (H2O) • Chlorine gas (Cl2) • Tungsten shavings Check your answers:

  11. Properties of each type of bond

  12. Covalent Bond A chemical bond that results in the sharing of electrons between two atoms

  13. Characteristics of Covalent Bonds a) Usually occur between atoms of nonmetals b) Result in a particle called a molecule c) Have low melting points (therefore are not solids at room temperature) d) Do not usually conduct electricity.

  14. Ionic Bond • The chemical bond resulting from the electrostatic attraction between positive and negative ions.

  15. Electrons are transferred • Opposite charges attract (electrostatic force)

  16. Ionic substances have predictable characteristics. a) Have extremely high melting points b) Tend to be soluble in water. c) Usually form a crystalline lattice structure as a solid. d) Are good conductors of electricity in the molten state.

  17. Metallic BondingBonding that occurs between atoms of a metal due to delocalized electrons forming an electron sea

  18. Characteristics of Metallic Bonding • The highly mobile “sea of electrons” in metals account for its high conductivity. • Metals are better conductors of heat and electricity than ionic and molecular compounds. • Metals are also malleable, ductile and shiny.

  19. Pause for a Cause #2 Tell what kind of bonds I have… Check your answers Ionic Metallic Covalent Let’s hear some of yours! • Heat me up and I can conduct electricity. • My electrons flow like fish in the sea. • My bonding partner and I are very dull (not shiny). • Write your own riddle about one type of bond.

  20. How strong are bonds? Ionic (Lattice energy) Covalent (bond strength) Dependent on two factors Bond type (single/double/triple) Size (Must remember your size trends) Shorter bonds are stronger Longer bonds are weaker Dependent on two factors • Charge • Size

  21. lattice energy cmpd MgF2 2957 Q= +2,-1 MgO 3938 Q= +2,-2 LiF 1036 LiCl 853 Electrostatic (Lattice) Energy Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. r F < r Cl

  22. Shorter bonds are stronger bonds Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond 9.4

  23. Pause for a Cause #3 Label each set as covelent or ionic and circle the substance with the greatest lattice or bond energy Check your answers • MgCl2 or MgO • H2O or H2S (both have single bonds) • N2 (triple bond) or H2 (single bond) • NaCl or KCl

  24. F H F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e- poor e- rich d+ d- 9.5

  25. Lewis Structures

  26. 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.

  27. - - - - + Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6

  28. Why should two atoms share electrons? + 8e- 8e- 7e- 7e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F2

  29. Writing Lewis Structures • Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. • Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. • Complete octets for as many atoms as possible, starting with the outer atoms. • If all octets are not filled, create double or triple bonds within the structure (with C, S, P, O or N) 9.6

  30. F N F F Write the Lewis structure of nitrogen trifluoride (NF3). 9.6

  31. 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24 Write the Lewis structure of the carbonate ion (CO32-). 9.6

  32. What are the resonance structures of the carbonate (CO32-) ion? - - + + O O O O O O O O O C C C O O O - - - - O O O - - A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. 9.8

  33. Pause for a Cause #5 Draw Lewis structures for the following molecules 1. HBr 2. CF4 3. NH2Cl 4. SiCl3Br 5. NH3 6. H2S

  34. BF3 SF4 Violations of the Octet Rule • B and Sn tends to form bonds with 6 electrons. • 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. • Be tends to form bonds with 4 electrons.

  35. H H C O H C O H formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 Two possible skeletal structures of formaldehyde (CH2O) An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. 9.7

  36. H C O H total number of nonbonding electrons formal charge on an atom in a Lewis structure total number of bonding pairs + = -1 +1 formal charge on C = formal charge on O = 9.7

  37. C – 4 e- H O – 6 e- C O H 2H – 2x1 e- 12 e- formal charge on an atom in a Lewis structure total number of total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 single bonds (2x2) = 4 2 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 0 0 formal charge on C = 4 - 0- ½ x 8 = 0 formal charge on O = 6 - 4- ½ x 4 = 0 9.7

  38. Which is the most likely Lewis structure for CH2O? H C O H H C O H -1 +1 0 0 Formal Charge and Lewis Structures • For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. • Lewis structures with large formal charges are less plausible than those with small formal charges. • Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. 9.7

  39. What shape do molecules have? Determined by Valence Shell Electron Pair Repulsion Theory VSEPR Theory- electrostatic repulsion between the valence-level pairs surrounding an atom causes these pairs to be oriented as far apart as possible. What does it mean?

  40. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class linear linear B B Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB2 2 0 10.1

  41. 0 lone pairs on central atom Cl Be Cl 2 atoms bonded to central atom 10.1

  42. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar Arrangement ofelectron pairs Molecular Geometry Class VSEPR AB2 2 0 linear linear AB3 3 0 10.1

  43. 10.1

  44. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class tetrahedral tetrahedral VSEPR AB2 2 0 linear linear AB4 4 0 10.1

  45. 10.1

  46. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class trigonal bipyramidal trigonal bipyramidal VSEPR AB2 2 0 linear linear tetrahedral tetrahedral AB4 4 0 AB5 5 0 10.1

  47. 10.1

  48. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class trigonal bipyramidal trigonal bipyramidal AB5 5 0 octahedral octahedral VSEPR AB2 2 0 linear linear tetrahedral tetrahedral AB4 4 0 AB6 6 0 10.1

  49. 10.1

  50. 10.1