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Chemical Bonding

Chemical Bonding. Chapter 6. Chemical Bonding & Structure. Molecular bonding and structure play the central role in determining the course of chemical reactions. Bonds. Forces that hold groups of atoms together and make them function as a unit. Bond Energy.

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Chemical Bonding

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  1. Chemical Bonding • Chapter 6

  2. Chemical Bonding & Structure • Molecular bonding and structure play the central role in determining the course of chemical reactions.

  3. Bonds • Forces that hold groups of atoms together and make them function as a unit.

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

  5. Bond Energies • Bond breaking requires energy (endothermic). • Bond formation releases energy (exothermic).

  6. Chemical Bonds • Chemical Bond • Ionic Covalent • Cation Anion Molecule

  7. Ionic Bonds • Formed from electrostatic attractions of closely packed, oppositely charged ions. • Formed when an atom that easily loses electrons (metal) reacts with one that has a high electron affinity(nonmetal). • 2Na(s) +Cl2(g) ----> 2Na+(aq) + 2Cl-(aq)

  8. Figure 11.8: The structure of lithium fluoride

  9. Figure 11.1: The formation of a bond between two hydrogen atoms

  10. Covalent Bonding • Covalent bonds are formed by sharing electrons between nuclei. • H.+ .H ----> H-H 2 hydrogen atoms hydrogen molecule

  11. Types of Covalent Bonds • Polar covalent bond -- covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other. A dipole moment exists. HOH, HCl, & CO • Nonpolar covalent bond -- covalent bond in which the electrons are shared equally between both atoms. No dipole moment exists. CO2, CH4, & Cl2

  12. Electronegativity • The ability of an atom in a molecule to attract shared electrons to itself. • As electronegativity increases, the attraction for electrons increases. Fluorine has the highest value at 4.0 and cesium and francium are lowest at 0.7.

  13. Pauling Electronegativity Values

  14. Electronegativity values for selected elements. See Figure 11.3 on page 334 in Zumdahl.

  15. Percent Ionic Character where xA is the larger electronegativity and xB is the smaller value. Watch significant figures!!! Ionic Bond % IC > 50 % Polar Covalent % IC 5 - 50 % Nonpolar Covalent % IC < 5 %

  16. Percent Ionic Character What type of bonding & % ionic character does KCl have? Ionic

  17. Percent Ionic Character What type of bonding & % ionic character does HOH have? Polar covalent

  18. Percent Ionic Character What type of bonding & % ionic character does N2 have? Nonpolar covalent

  19. Three Possible Types of Bonds Nonpolar Covalent (Electrons equally shared.) Polar Covalent (Electrons shared unequally.) Ionic (Electrons are transferred.)

  20. Figure 11.2: Probability representations of the electron sharing in HF

  21. Polarity • A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. partial positive charge partial negative charge

  22. The Effect of an electric field on hydrogen fluoride molecules.

  23. Dipole Moment for the water molecule.

  24. Polar Water Molecule • The polarity of water allows it to dissolve ionic materials which are essential for life. • The polarity of the water molecule allows water molecules to attract each other strongly (hydrogen bonds). Because of this fact water remains as a liquid at room temperatures and allows the existence of life as we know it.

  25. Dipole moment for the ammonia molecule.

  26. Nonpolar molecule--zero dipole moment.

  27. Achieving Noble Gas Electron Configurations (NGEC) • Two nonmetals react: They share electrons to achieve NGEC. • A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

  28. Noble Gas Configuration • When a Group I, II, or III metal reacts with a nonmetal to form a binary ionic compound, the nonmetal gains electrons to obtain the configuration of the next noble gas. The metal loses electrons to gain the configuration of the previous noble gas. • Na ----> Na+ + e- configuration of Ne • Cl + e- ----> Cl- configuration of Ar

  29. Noble Gas ConfigurationContinued • When two nonmetals react to form a covalent bond, they share electrons to form noble gas configurations for both. • H.+ ----> H-- • Hydrogen gains the noble gas configuration of helium, while Chlorine gains the configuration of Argon.

  30. Anion Size • Anions are always larger than the parent atom because they have added electrons which repel each other. As well, the number of protons is less than the number of electrons so they are not held as tightly.

  31. Cation Size • Cations are always smaller than the parent atom because they have lost an entire electron shell. As well, the number of protons is greater than the number of electrons so the electrons are held tighter.

  32. Relative sizes of some ions and their parent atoms.

  33. Lewis Structure • 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. • Developed by G.N. Lewis in 1902.

  34. Lewis Structures • Na. sodium atom [Na]+ sodium ion • sulfur atom [ ] sulfide ion

  35. Covalent Compounds Lewis Structures Ionic Compounds In ionic compounds, electrons are transferred and ions are formed. In covalent compounds, electrons are shared to form a molecule. Potassium gains the stability of argon, bromine of krypton, and fluorine of neon.

  36. Lone Pairs & Bonding Pairs Electrons shared between atoms are bonding pairs. Electrons that are not involved in bonding are called lone pairs. Each fluorine has three lone pairs and one bonding pair shared between them.

  37. Octet • Neon does not form bonds because it has a full outer shell of electrons--an octet. An octet is four pairs of electrons and represents extra stability for atoms and ions.

  38. Rules for Writing Lewis Structures • Sum the valence electrons from all the atoms. • Use a pair of electrons to form a bond between each pair of bound atoms. • Arrange remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second-row elements.

  39. Lewis Structures • NO+ • 5 e- + 6 e- - 1 e- = 10 e- • [:NO:]+ • Each atom has an octet and is satisfied.

  40. Single, Double, & Triple Bonds Single bonds -- one shared pair of electrons. Double bonds -- two shared pairs of electrons. Triple bonds -- three shared pairs of electrons.

  41. Bond Strength = Triple > Double > Single • For bonds between same atoms, CN > C=N > C—N • Though Double not 2x the strength of Single and Triple not 3x the strength of Single • Bond Length = Single > Double > Triple • For bonds between same atoms, C—N > C=N > CN

  42. 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.

  43. Electron Deficient Molecules • Beryllium chloride -- BeCl2 -- is electron deficient with four electrons. It forms a linear molecule. • Boron trifluoride -- BF3 -- is electron deficient with six electrons. It forms a trigonal planar molecule. • See page 351 for the reaction between boron trifluoride and ammonia.

  44. Four Failures of Lewis Structures • Lewis Structures cannot adequately explain: • 1. electron-deficient molecules. • 2. the paramagnetism of oxygen and other similar substances. • 3. odd-electron molecules. • 4. resonance.

  45. Odd-Electron Molecules • NO2 • contains 17 electrons. • cannot satisfy the octet rule. • a more sophisticated model is needed-the molecular orbital model.

  46. 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 called a resonance hybrid. See the resonance structures for the nitrate ion on page 362 in Zumdahl.

  47. O S O O S O •• •• •• •• •• •• • • •• •• • • •• •• • • • • • • • • •• •• Resonance • Resonance structures have Lone Pairs and Multiple Bonds in different positions. • The actual molecule is a combination of all the resonance forms – it does not resonate between the two forms, though we often draw it that way!

  48. Stereochemistry The study of the three-dimensional arrangement (molecular structure) of atoms or groups of atoms within molecules and the properties which follow such arrangement.

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