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Chapter 9 and 10

Chapter 9 and 10. Bonding and Molecular Structure. Valence and Core Electrons. Valence Electrons – electrons that participate in bonding Core Electrons – electrons in an atom that do not participate in bonding Main Group Elements – s and p orbitals

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Chapter 9 and 10

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  1. Chapter 9 and 10 Bonding and Molecular Structure

  2. Valence and Core Electrons • Valence Electrons – electrons that participate in bonding • Core Electrons – electrons in an atom that do not participate in bonding • Main Group Elements – s and p orbitals • Transition Elements – ns and (n-1) d orbitals

  3. Lewis Dot Symbols for Elements • Element's symbol represents the atomic nucleus together with the core electrons. • Valence electrons are represented by dots that are placed around the symbol. • These symbols emphasize the ns2np2 “octet” that all noble gases except helium possess.

  4. Lewis Dot Symbols for Main Group Elements

  5. Chemical Bond Formation Ionic Bonds • Electrons are strongly displaced toward one atom and away from the other. • Generally involve metals from the left side of the periodic table interacting with nonmetals form the far right side.

  6. Lew Symbols and the Octet Rule for Ionic Compounds • The electron configuration of many substances after ion formation is that of an noble gas  octet rule. • Octet rule: Main-group elements gain, lose, or share in chemical bonding so that they attain a valence octet (eight electrons in an atoms valence shell).

  7. Example 1 • The electron configuration of each reactant in the formation of KCl gives: • K+ is that of [Ar] • Cl is also that of [Ar]. • The other electrons in the atom are not as important in determining the reactivity of that substance. • The octet rule is particularly important in compounds involving nonmetals.

  8. Energy in Ionic Bonding • When potassium and chlorine atoms approach each other we have: K(g) K+(g) + e Ei = +418 kJ Cl(g)+ e Cl(g) Eea = 349 kJ K(g)+Cl(g) K+(g) + Cl(g) E = + 69 kJ

  9. Driving Force of Ionic • Positive E = energy absorbed  energetically not allowed. • Driving force must be the formation of the crystalline solid. K+(g) + Cl(g)  KCl(s)

  10. Born-Haber Cycle and Lattice Energies • Overall energetics for the formation of crystalline solids can be determined from a Born-Haber cycle which accounts for all of the steps towards the formation of solid salts from the elements. For the formation of KCl from its elements we have:

  11. Born-Haber Cycle • Net energy change of 434 kJ/mol indicates energetically favored. • Energy for the fifth step is the negative of the • lattice energy: energy required to break ionic bonds and sublime (always positive).

  12. Example 2 Determine the lattice energy of BaCl2 if the heat of sublimation of Ba is 150.9 kJ/mol and the 1st and 2nd ionization energies are 502 and 966 kJ/mol, respectively. The heat for the synthesis of BaCl2(s) from its elements is 806.06 kJ/mol.

  13. LATTICE ENERGIES AND PERIODICITY • Lattice energy can also be determined from Coulomb’s law: • Directly proportional to charge on each ion. • Inversely proportional to size of compound (sum of ionic radii).

  14. Lattice Energies • Table (right) presents the lattice energies for alkali and alkaline earth ionic compounds. The lattice energies • decrease for compounds of a particular cation with atomic number of the anion. • decrease for compounds of a particular anion with atomic number of the cation.

  15. Covalent Bond Formation • Electrons involved in the bond are more or less evenly distributed between the atoms, and electrons are shared by two nuclei. • Covalent bonding generally occurs between nonmetals, elements that lie in the upper right corner of the periodic table.

  16. The Covalent Bond • Repulsive forces of the electrons offset by the attractive forces between the electrons and the two nuclei. • Bonds are characterized in terms of energy and bond distance

  17. Strengths of Covalent Bonds: • Bonds form because their formation produces lower energy state than when atoms are separated. • Breaking bonds increases the overall energy of the system. Energy for breaking bonds has a positive sign (negative means that energy is given off). H - H (g)  2H(g) DH = 436 kJ.

  18. Ionic vs Covalent Bond • Ionic compounds have high melting and boiling points and tend to be crystalline • Covalently bound compounds tend to have lower melting points since the attractive forces between the molecules are relatively weak.

  19. Lewis Structures • Lewis structures are used to indicate valence electrons and bonds

  20. Drawing Lewis Structures • Determine the arrangement of the atoms in the compound with respect to each other and draw a skeletal structure. • If the compound is binary, the first element written down is usually the central atom (hydrogen is an exception to this). • With a ternary compound (one with three kinds of elements) the middles atom in the formula is usually the central one.

  21. Lewis Structures Cont. • Determine the total number of valence electrons. • Subtract two electrons from the valence total for each bond in the skeletal structure.

  22. Lewis Structures Cont. • Determine how many electrons are required for each element to have a total of eight (there are several exceptions to this rule). • If a sufficient number of electrons are available, distribute the remaining electrons around the element symbols. • If a sufficient number is not available, add additional bonds making certain to subtract the electrons used. (Multiple bonds often occur among the atoms C, N, and O.

  23. Example 3 • Draw Lewis Electron Dot Structures for the following

  24. Lewis Structures and Resonance • Quantum theory indicates that any position is possible for an electron. • Equivalent electron positions often possible: • E.g. SO2 : O=S-O and :O-S=O. • Each structure equally likely. • the true form of the molecule is a hybrid of these and is called resonance and the hybrid form is called aresonance hybrid.

  25. Example 4 • Draw Lewis Electron Dot Structures for the following

  26. Exceptions to the Octet Rule • Although many molecules obey the octet rule, there are exceptions where the central atom has fewer or more than eight electrons. • Compounds in which an atom has fewer than eight valence electrons

  27. Compounds in which an atom has more than eight valence electrons • Generally, if a nonmetal is in the third period or greater it can accommodate as many as twelve electrons, if it is the central atom. • These elements have unfilled “d” subshells that can be used for bonding.

  28. Free Radicals • Molecules with an odd number of electrons

  29. Example 5 Draw Lewis Electron Dot Structures for the following XeF4, ICl3, and SF4

  30. FORMAL CHARGES • Formal Charge (of an atom in a Lewis formula) the hypothetical charge obtained by assuming that bonding electrons are equally shared between the two atoms involved in the bond. Lone pair electrons belong only to the atom to which they are bound. • Formal Charge = group number – number of lone pair electrons – ½ bonding electrons

  31. Example 6 Determine the formal charge on all elements: PCl3, PCl5, and HNO3.

  32. Formal charge (FC) allows the prediction of the more likely resonance structure. • To determine the more likely resonance structure: • FC should be as close to zero as possible. • Negative charge should reside on the most electronegative and positive charge on the least electronegative element.

  33. Example 7 Draw the resonance structures of H2SO4; determine the formal charge on each element and decide which is the most likely structure.

  34. Bond Properties • Bonds Order • Bond Length • Bond Energy • Bond Polarity

  35. Bond Order • The order of a bond is the number of bonding electrons pairs shared by two atoms in a molecule. • A fractional bond order is possible in molecules and ions having resonance structures

  36. Bond Length • Bond length is the distance between the nuclei of two bonded atoms. Bond length is determined in a large part by the size of the atoms. • Bond length becomes shorter as bond order increases

  37. Bond Dissociation Enthalpies • Bond dissociation energy, BE – the energy required to break one mole of a type of bond in an isolated molecule in the gas phase. • Useful for estimation of heat of unknown reactions.

  38. Bond energies are all positive • The energies are average bond energies • Bond energies are defined in terms of gaseous atoms of molecular fragments • Bond energies increase with bond order

  39. Hess’s law can be used with bond dissociation energies to estimate the enthalpy change of a reaction. • The breaking in a C – H bond would be C – H(g)  C(g) + H(g) H = BE = 410 kJ. • Sign always positive since energy must be supplied to break bond.

  40. Using Bond Dissociation Enthalpies • Estimate the heat of formation of H2O(g) from bond dissociation energies. Thus determine: • H2(g) + ½ O2(g)  H2O(g) = ?   H – H (g)  2H(g) H = BE = 436 kJ  ½ O=O  O(g) H = BE = 494/2 = 247 kJ 2H(g) + O(g)  H – O – H (g) H = 2BE = 2*459 kJ   H2(g) + ½ O2(g)  H2O(g) = 235 kJ Actual = 241.8 KJ • Can be determined by summing all the energies for the bonds broken and subtract from it the sum of the energies for the bonds formed.

  41. Example 8 Estimate the energy change for the chlorination of ethylene: • CH2=CH2(g) + Cl2(g) CH2ClCH2Cl

  42. Bond Polarity and Electronegativity • Electronegativities • increase from bottom to top of periodic table and • increase to a maximum towards the top right. • can provide an insight as to the type of bond that would be expected.

  43. Electronegativities

  44. Bond Polarity • Ionic bonds formed when displacement of electrons is essentially complete • Covalent bonds forms when no displacement of electrons occurs • Polar covalent forms when bond pair is not equally shared between two atoms and the electrons are displace toward one of the atoms from a point midway between them

  45. Polar Bonds With a polar bond exists between two atoms, a small charge on the atom due to that bond develops. + and  designates which is the positive and negative side respectively

  46. Example 9 Determine the relative polarities of HF, HCl, HBr and HI.

  47. Molecular Geometry and Directional Bonding • Atoms oriented in very well defined relative positions in the molecule. • Molecular Geometry = general shape of the molecule as determined by the relative positions of the atomic nuclei.

  48. Theories describing the structure and bonding of molecules are: • VSEPR = considers mostly electrostatics in determining the geometry of the molecule. • Valence Bond Theory = considers quantum mechanics and hybridization of atomic orbitals. • Molecular Orbital Theory = claims that upon bond formation new orbitals that are linear combinations of the atomic orbitals are formed.

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