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Bonding: General Concepts

Bonding: General Concepts. Chapter 8. Types of Chemical Bonds. Bond Energy is _____________________ _________________________________.

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Bonding: General Concepts

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  1. Bonding: General Concepts Chapter 8

  2. Types of Chemical Bonds • Bond Energy is _____________________ _________________________________. • Compounds always strive to be at the lowest energy orientation possible. They will bond or not bond depending on whether or not it is favorable from an energy perspective. • A higher value for bond energy indicates a stronger bond.

  3. Ionic Bonding • An ionic bond forms when electrons are _________ between atoms. When this happens both atoms become _____; one positive (____ of e-), the other negative (_____ of e-). • This type of bonding occurs between a _________ • and a _____________. • Elements bonded by an ionic bond form • ionic compounds.

  4. Coulomb’s Law E = • Describes _____________ between two particles. • _______ values indicate attraction, while _______ values indicate repulsive forces. • Ion pairs have lower energy (more favorable) when they are ______________________.

  5. Bond Length • Bond length is the distance between two ions that maximizes the favorable attractions and minimizes the amount of repulsive forces. • Favorable attractive forces: • Unfavorable attractive forces (repulsive forces):

  6. Covalent and Polar Covalent Bonds • A Covalent Bond occurs when two nuclei _________ electrons. • Occurs between two _____________ atoms. • Polar Covalent bonds occur when two nuclei _________________________. • The dipoles (+ and – centers) that are formed give the covalent bond more ionic character than normal. • Lower case delta (d) is used to indicate partial charge.

  7. Electronegativity • Follows the trend for electron affinity (synonymous), which means that it increases as we move ________ a group and _________ a period on the periodic table. • Electronegativity is __________________ _________________________________.

  8. Electronegativity affects bonds • If bonded atoms have a large difference in electronegativity (2.0 or greater), they are considered to have an _________ bond because _____________________________________. • If bonded atoms have a moderate difference in electronegativity (0.5-1.6), they are considered to have a _______-__________ bond because _____________________________________. • If bonded atoms have a negligible difference in electronegativity (below 0.5), they will be considered to have a ___________ bond because ______________________________________. • If bonded atoms have a difference in electronegativity between 1.6 and 2.0, their identities have to be considered. If a metal is involved, it will be deemed ionic. If 2 non-metals are bonded, it will be considered polar-covalent.

  9. Bond Polarity and Dipole Moments • Dipoles are formed in _________ bonds. • If the dipoles (a form of force) are aligned to point in exactly equal and opposite directions, they will _________________. • This makes the molecule non-polar. • The bond remains polar. • Molecules must be symmetrical for the cancellation to occur. Shapes that cancel are (draw them): 1. 2. 3.

  10. Assignment • Arrange the bonds in each of the following set in order of increasing polarity • C-F, O-F, Be-F • O-Cl, S-Br, C-P • C-S, B-F, N-O • Using only the periodic table as a guide, select A) the most electronegative element in group 6A B) The least electronegative element out of Al, Si, P C) The element in the group K, C, Zn, F that is most likely to form an ionic compound with Ba.

  11. Give three ions that are isoelectronic with argon. Place these ions in order of increasing size. • What two requirements must be satisfied for a molecule to be polar? • Rank the following bonds in order of increasing ionic character: • N-O, Ca-O, C-F, Br-Br, K-F

  12. Section 5:Formation of Binary Ionic Compounds • Lattice energy: the energy required to combine elements to form an ionic compound. (amount released when bond forms) • Lattice energy is negative. • Elements must start as a gas • The more exothermic (negative), the more likely the substance is to form spontaneously.

  13. Steps involved in forming ionic bonds from elements. • Vaporization of elements. (endo) Li(s) Li(g) F2(g)  2F(g) • Ionization of elements. (exo) Li(g) Li+(g) + e- F(g) + e-  F-(g) • Formation of solid by combination of ions. (very exo) Li+(g) + F-(g) LiF(s)

  14. Energy Diagram

  15. Partial Ionic Character of Covalent Bonds Any Compound That Conducts Electricity When Melted Is IONIC

  16. Models of Chemical Bonds • Models do not equal reality…they are merely something to help us visualize a concept near the truth. • Models are often wrong because they over-simplify.

  17. Covalent Bond Energy • Single bond = one pair shared electrons • Double bond = two pair shared electrons • Triple bond = three pair shared electrons • As more pairs of electrons are shared, the bond length shortens. (more orbitals have to overlap to allow the sharing to happen) • Single bonds usually contain the least amount of energy, while triple bonds usually contain the most…as bond length shortens, bond energy increases. Page 372 Tables 8.4 and 8.5

  18. Covalent Bond Energy Calculation • DHbond = sum energies required to break old bonds (positive signs) plus the sum of energies required to form new bonds (negative signs) • DHbond = S bonds broken – S bonds formed Use values in the tables on pg 372.

  19. Covalent Bond Energy Calculation Using the bond energies listed in Table 8.4, calculate the DH for the reaction of methane with chlorine and fluorine to give Freon-12 (CF2Cl2). CH4(g) + 2Cl2(g) + 2F2(g) CF2Cl2(g) + 2HF(g) + 2HCl(g)

  20. The VSEPR Model • Valence Shell Electron Pair Repulsion • A model of molecular structure based on the idea that ideal structures minimizes electron pair repulsions. • Draw and evaluate Lewis Structures bare naked electrons are the most repulsive!

  21. Molecular Geometry Models • We look at the molecular geometry of a single atom, not of an entire molecule. • Constituent groups are the things bonded to the atom under scrutiny. • Dashed lines represent a bond behind the plane of the paper; wedged lines represent a bond coming toward you (in front of the paper plane)

  22. Planar Geometry Linear 1-2 Constituents 0 Lone Pair Bond Angle: 180o Trigonal Planar 3 Constituents 0 Lone Pair Bond Angle: 120o Bent 2 Constituents 1 Lone Pair Bond Angle: <120o

  23. Tetrahedral and Derivatives Tetrahedral 4Constituents 0 Lone Pair Bond Angle: 109.5o Trigonal Pyramidal 3 Constituents 1 Lone Pair Bond Angle: 107.3o Bent 2 Constituents 2 Lone Pair Bond Angle: 104.5o

  24. Trigonal Bipyramidal and Derivatives Trigonal Bipyramidal 5 Constituents 0 Lone Pair Bond Angle: 90o, 120o See-Saw 4 Constituents 1 Lone Pair Bond Angle: <90o, <,120o, <180o T-Shaped 3 Constituents 2 Lone Pair Bond Angle: <90o, <180o Linear 2 Constituents 3 Lone Pair Bond Angle: 180o

  25. Octahedral and Derivatives Octahedral 6 Constituents 0 Lone Pair Bond Angle: 90o Square Pyramidal 5 Constituents 1 Lone Pair Bond Angle: <90o Square Planar 4 Constituents 2 Lone Pair Bond Angle: 90o

  26. Isomers • A set of compounds with the same chemical formula that exhibit decidedly different properties. • Structural Isomer = the same atoms are present in each molecule, but they are bonded to different things. • Stereoisomer = bonds are between same atoms, but have a different spatial relationship with each other.

  27. Concept Map ISOMERS STRUCTURAL STEREOISOMERS These are the ones we care about. Coordination Isomers Linkage Isomers Geometric Optical

  28. Trans Cis Geometric Isomers • Come in two forms: cis and trans • cis- prefix represents that similar constituent groups are on the SAME side. • trans- prefix represents that similar constituent groups are on OPPOSITE sides.

  29. More Complicated cis-, trans-

  30. Optical Isomers • Two forms of the molecule have different effects on plane polarized light. (one form creates destructive interference) • These isomers are non-super imposable mirror images. • Each isomer is referred to as an ENANTIOMER. • The chemical formula is said to have CHIRALITY.

  31. Hybridization of Orbitals • Natural orbitals overlap to form hybridized orbitals during bonding. • 2 types of bonding • Sigma (s) bonds = take equatorial positions, can be hybridized. • Pi (p) bonds = take axial positions, exist in non-hybridized orbitals • Possibilities for hybridization: sp3 sp2 sp dsp3 d2sp3

  32. How to become Hybrid Orbitals in which bonding electrons exist will become degenerate (equal energy). 2p ___ ___ ___ __ __ __ __ sp3 2s ___ By becoming equal energy, they also acquire same shape.

  33. sp3 Hybridization • Atoms with 4 constituent groups • Tetrahedral shapes

  34. sp2 hybridization • 3 constituent groups • Trigonal planar

  35. sp2 continues • One p orbital remains un-hybridized. This orbital has the ability to house lone pairs as well as form a double bond (pi bond!)

  36. sp hybridization • 2 constituents • Linear • Now 2 extra p remain un-hybridized. • 2 double bonds, or 1 triple bond can form.

  37. dsp3 and d2sp3 hybridization • dsp3 allows for 5 constituent groups • trigonal bipyramidal • d2sp3 allows for 6 constituent groups • octahedral

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