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Ch10lecture 150104202336-conversion-gate02
- 1. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 1 of 48 PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY TENTH EDITION GENERAL CHEMISTRY Principles and Modern Applications PETRUCCI HERRING MADURA BISSONNETTE Chemical Bonding I: Basic Concepts 10
- 2. Chemical Bonding I: Basic Concepts Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 2 of 48
- 3. 10-1 Lewis Theory: An Overview 1. Valence e- play a fundamental role in chemical bonding. 2. e- transfer leads to ionic bonds. 3. Sharing of e- leads to a covalent bond. 4. e- are transferred or shared to give each atom a noble gas configuration, the octet. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 3 of 48 Gilbert Newton Lewis (1875-1946)
- 4. Lewis Symbols and Lewis Structures Copyright © 2011 Pearson Canada Inc.Slide 4 of 48 General Chemistry: Chapter 10 A chemical symbol represents the nucleus and the core e- . Dots around the symbol represent valence e- . Si • • • • N •• • • • P •• • • • As •• • • • Sb •• • • • Bi •• • • • •• Al• • • Se• • • •• Ar •• •• •• I • •• •• ••
- 5. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 5 of 48 Lewis structure
- 6. 10-2 Covalent Bonding: An Introduction Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 6 of 48
- 7. Coordinate Covalent Bonds Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 7 of 48 FIGURE 10-2 Formation of the ammonium ion, NH4 + HN •• H H H Cl N H H H H + •• Cl •• •• •• -
- 8. Multiple Covalent Bonds Copyright © 2011 Pearson Canada Inc.Slide 8 of 48 General Chemistry: Chapter 10 C• • • • O• • •• • • O• • •• • • CO O • •• ••• •• •• •• • CO O •• ••• •• •• •• CO O •• •• •• ••
- 9. Multiple Covalent Bonds Copyright © 2011 Pearson Canada Inc.Slide 9 of 48 General Chemistry: Chapter 10 N• ••• • N N • •• • • •• • N• • •• • N N • •• •• • N N •• ••
- 10. Paramagnetism of Oxygen Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 10 of 48
- 11. 10-3 Polar Covalent Bonds and Electrostatic Potential Maps Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 11 of 48 FIGURE 10-4 Determination of the electrostatic potential map for ammonia
- 12. The electrostatic potential maps for sodium chloride, hydrogen chloride and chlorine FIGURE 10-5 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 12 of 48
- 13. Electronegativity Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 13 of 48 FIGURE 10-6 Electronegativities of the elements
- 14. Percent ionic character of a chemical bond as a function of electronegativity difference FIGURE 10-7 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 14 of 48
- 15. 10-4 Writing Lewis Structures • All the valence e- of atoms must appear. • Usually, the e- are paired. • Usually, each atom requires an octet. • H only requires 2 e- . • Multiple bonds may be needed. • Readily formed by C, N, O, S, and P. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 15 of 48
- 16. Skeletal Structure Copyright © 2011 Pearson Canada Inc.Slide 16 of 48 General Chemistry: Chapter 10 Identify central and terminal atoms. C H H H HC H H O
- 17. Skeletal Structure Copyright © 2011 Pearson Canada Inc.Slide 17 of 48 General Chemistry: Chapter 10 • Hydrogen atoms are always terminal atoms. • Central atoms are generally those with the lowest electronegativity. • Carbon atoms are always central atoms. • Generally structures are compact and symmetrical.
- 18. Summary scheme for drawing Lewis Structures FIGURE 10-8 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 18 of 48
- 19. Formal Charge Copyright © 2011 Pearson Canada Inc.Slide 19 of 48 General Chemistry: Chapter 10 FC = #valence e- - #lone pair e- - #bond pair e-2 1 FC(O) = 6 - 4 – (4) = 0 2 1 FC(N) = 5 - 0 – (8) = +1 2 1 •• •• •• •• O=N=O +
- 20. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 20 of 48 Formal Charge of an Alternative Lewis Structure •• •• •• •• •• •• O—N—O FC(O≡) = 6 - 2 – (6) = +1 2 1 FC(N) = 5 - 0 – (8) = +1 2 1 FC(O—) = 6 - 6 – (2) = -1 2 1 •• O N O •• •• •• + + -
- 21. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 21 of 48 • Sum of FC is the overall charge. • FC should be as small as possible. • Negative FC usually on most electronegative elements. • FC of same sign on adjacent atoms is unlikely. General Rules for Formal Charge + •• O≡N—O •• •• •• -+
- 22. 10-5 Resonance Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 22 of 48 O O O O O O •••• •• •• •• •• •• •• •• •• •• ••+ +- - Electrostatic potential map of ozone O O O •• •• •• •• ••+ -½-½
- 23. 10-6 Exceptions to the Octet Rule Odd-ElectronSpecies Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 23 of 48 •• •• •• • H—C—H H • • O—H •• •• N=O
- 24. Incomplete Octets Copyright © 2011 Pearson Canada Inc.Slide 24 of 48 General Chemistry: Chapter 10 B F FF •• •• •• •• •• •• •• •• •• B F FF - •• •• •• •• •• •• + •• •• •• •• •• B F FF •• •• •• •• •• •• •• - + •• ••
- 25. Expanded octets Copyright © 2011 Pearson Canada Inc.Slide 25 of 48 General Chemistry: Chapter 10 •• •• •• P Cl ClCl •• •• •• •• •• •• •• P Cl Cl •• •• Cl •• •• •• •• •• •• •• •• •• Cl •• •• •• Cl•• S F F •• •• F •• •• •• •• •• •• •• •• •• F •• •• •• F•• F •• •• ••
- 26. Expanded Valence Shells Copyright © 2011 Pearson Canada Inc.Slide 26 of 48 General Chemistry: Chapter 10
- 27. 10-7 The Shapes of Molecules Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 27 of 48 FIGURE 10-9 Geometric shape of a molecule
- 28. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 28 of 48 Bond length – distance between nuclei. Bond angle – angle between adjacent bonds. VSEPR Theory Electron pairs repel each other whether they are in chemical bonds (bond pairs) or unshared (lone pairs). Electron pairs assume orientations about an atom to minimize repulsions. Electron group geometry – distribution of e- pairs. Molecular geometry – distribution of nuclei.
- 29. Balloon analogy to valence-shell electron-pair repulsion FIGURE 10-10 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 29 of 48
- 30. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 30 of 48 Methane, Ammonia and Water
- 31. Several electron-group geometries illustrated FIGURE 10-12 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 31 of 48
- 32. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 32 of 48 Table 10.1 Molecular Geometry as a Function of Electron Group Geometry
- 33. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 33 of 48
- 34. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 34 of 48
- 35. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 35 of 48
- 36. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 36 of 48
- 37. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 37 of 48 Replace this part of table 10.1 this format is not present in the text. Parent table is on page 425 and 426
- 38. Applying VSEPR Theory Copyright © 2011 Pearson Canada Inc.Slide 38 of 48 General Chemistry: Chapter 10 1. Draw a plausible Lewis structure. 2. Determine the number of e- groups and identify them as bond or lone pairs. 3. Establish the e- group geometry. 4. Determine the molecular geometry.
- 39. Structures with Multiple Covalent Bonds Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 39 of 48 •• •• •• •• S OO - •• +•• •• •• •• •• S OO - •• •• +•• •• •• •• •• S OO •• •• Molecules with More Than One Central Atom The geometric distribution of terminal atoms around each central atom must be determined and the results then combined into a single description of the molecular shape. See Example 10-12.
- 40. Molecular Shapes and Dipole Moments Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 40 of 48 FIGURE 10-14 Polar molecules in an electric field
- 41. Molecular shapes and dipole Moments FIGURE 10-15 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 41 of 48
- 42. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 42 of 48 10-8 Bond Order and Bond Length Bond Order Single bond, bond order = 1 Double bond, bond order = 2 Triple bond, bond order = 3 Bond Length The distance between the centers of two atoms joined by a covalent bond. the length of the covalent bond between two atoms can be approximated as the sum of the covalent radii of the two atoms.
- 43. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 43 of 48
- 44. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 44 of 48 •• •• •• C O OO - •• •• •• •• •• - •• •• •• C O OO - •• •• •• •• •• - •• •• •• C O OO - •• •• •• •• •• - average bond order = (1 + 1 + 2) bonds (3) structures = 4 3 = 1 ⅓ Average Bond Order
- 45. 10-9 Bond Energies Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 45 of 48 FIGURE 10-16 Some bond energies compared
- 46. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 46 of 48
- 47. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 47 of 48 Calculating an Enthalpy of Reaction from Bond Energies. ΔHrxn = Σ ΔH(bond breakage) + ΣΔH(bond formation) ≈ Σ BE(reactants) - Σ BE(products)
- 48. End of Chapter Questions Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 10Slide 48 of 48 c •Testing your decisions: –If you get an error or a nonsense result, then climb back to an intersection where you KNOW you were correct, and take another route. d? Answer c e f a b e?!#
Editor's Notes
- In this chapter, you will learn how to deduce and write chemical formulas and how to use the information incorporated into chemical formulas. The chapter ends with an overview of the relationship between names and formulas—chemical nomenclature.
- A scanning tunneling microscope image of 48 iron atoms adsorbed onto a surface of copper atoms. The iron atoms were moved into position with the tip of the scanning tunneling microscope in order to create a barrier that forced some electrons of the copper atoms into a quantum state seen here as circular rings of electron density. The colors are from the computer rendering of the image. In this chapter we discuss the periodic table and the properties of atoms and ions. In this chapter, we will use the table as a backdrop for a discussion of some properties of elements, including atomic radii, ionization energies, and electron affinities. These atomic properties also arise in the discussion of chemical bonding in the following two chapters, and the periodic table itself will be our indispensable guide throughout much of the remainder of the text.
- Lewis’s contribution to the study of chemical bonding is evident throughout this text. Equally important, however, was his pioneering introduction of thermodynamics into chemistry.
- Note double headed arrow for two electron movement The H atom of HCl leaves its electron with the Cl atom and, as H+ attaches itself to the NH3 molecule through the lone-pair electrons on the N atom. The ions NH4+ and Cl- are formed.
- The electrostatic potential at any point on the charge density surface of a molecule is defined as the change in energy that occurs when a unit positive charge is brought to this point, starting from another point that is infinitely far removed from the molecule. The surface encompassing the ammonia molecule is analogous to the 95% surface of electron charge density for atomic orbitals discussed in Chapter 8. The electrostatic potential map gives information about the distribution of electron charge within this surface.
- The dark red and dark blue on the electrostatic potential map correspond to the extremes of the electrostatic potential, negative to positive, for the particular molecule for which the map is calculated. To get a reliable comparison of different molecules, the values of the extremes in electrostatic potential (in KJ mol-1) must be the same for all of the molecules compared. In the maps shown here the range is -250 to 750 kJ mol-1.
- As a general rule, electronegativities decrease from top to bottom in a group and increase from left to right in a period of elements. The values are from L. Pauling, The Nature of the Chemical Bond, 3rd ed., Cornell University, Ithaca, NY, 1960, page 93. Values may be somewhat different when based on other electronegativity scales.
- H can only accommodate two electrons H and O are common exceptions to rule 2 Organic compounds are not compact nor symmetrical.
- The formal charge on an atom in a Lewis structure is the number of valence e- in the free atom minus the number of e- assigned to that atom in the Lewis structure. This is essentially Example 10-7
- Many Lewis structures may be written for a given structure.. Ozone has two good possibilities, but neither gives the correct structure that has two equivalent O-O bonds. All the atoms in an ozone molecule have the same electronegativity and yet the distribution of electron density is nonuniform. The reason will become apparent when we describe in a more sophisticated way the bonding in this molecule.
- H
- It is not clear which is the more correct representation.
- To establish the shape of the triatomic H2O molecule shown here, we need to determine the distances between the nuclei of the bonded atoms and the angle between adjacent bonds. In H2O the bond lengths d1= d2 = 95.8 pm and the bond angle is α= 104.45°.
- When two elongated balloons are twisted together, they separate into four lobes. To minimize interferences, the lobes spread out into a tetrahedral pattern. (A regular tetrahedron has four faces, each an equilateral triangle.) The lobes are analogous to valence-shell electron pairs.
- The electron-group geometries pictured are trigonal-planar (orange), tetrahedral (gray), trigonal bipyramidal (pink), and octahedral (yellow). The atoms at the ends of the balloons are not shown and are not important in this model.
- In a multiple covalent bond, all electrons in the bond are confined to the region between the bonded atoms, and together constitute one group of electrons. Let us test this idea by predicting the molecular geometry of sulfur dioxide. S is the central atom, and the total number of valence electrons is 3 x 6 = 18. Because we count the electrons in the double covalent bond as one group, the electron-group geometry around the central S atom is that of three electron groups—trigonal-planar. Of the three electron groups, two are bonding groups and one is a lone pair. This is the case of AX2E (see Table 10.1). The molecular shape is angular, or bent, with an expected bond angle of 120° (The measured bond angle in SO2 is 119°.)
- The device pictured is called an electrical condenser (or capacitor). It consists of a pair of electrodes separated by a medium that does not conduct electricity but consists of polar molecules. (a) When the field is off, the molecules orient randomly. (b) When the electric field is turned on, the polar molecules orient in the field between the charged plates so that the negative ends of the molecules are toward the positive plate and vice versa.
- HCl is a polar molecule
- An interesting situation arises for molecules in which resonance is present. In such molecules, fractional bond orders are possible. Consider, for example, the carbonate anion, The CO bond distance in the carbonate anion is 129 pm, which is intermediate between a single bond (143 pm) and a double bond (120 pm), as we might expect for a fractional bond order.
- The same quantity of energy, 453.93 kJ/mol, is required to break all H-H bonds. In H2O more energy is required to break the first bond (498.7 kJ/mol) than to break the second (428.0 kJ/mol) The second bond broken is that in the •OH radical. The O-H bond energy in H2O is the average of the two values: 463.4 kJ/mol.
- You can use bond energies in exactly the same way you can use enthalpies of formation. Enthalpy of formation is more accurately known and bond energy is usually an average (thus the approximation sign), but it can be used effectively if formation data is unavailable.