2. Bonding – General Concepts

3. What is a Bond? A force that holds atoms together. We will look at it in terms of energy. –Bond energy - the energy required to break a bond. Why are compounds formed? –Because it gives the system the lowest energy.

4. Ionic Bonding An atom with a low ionization energy reacts with an atom with high electron affinity. The electron moves. Opposite charges hold the atoms together.

5. What about covalent compounds? The electrons in each atom are attracted to the nucleus of the other. The electrons repel each other, The nuclei repel each other. They reach a distance with the lowest possible energy. The distance between is the bond length.

6. Electronegativity The ability of an electron to attract shared electrons to itself. Pauling method Imaginary molecule HX Expected H-X energy = H-H energy + X-X energy 2  = (H-X) actual - (H-X) expected

7. Electronegativity   is known for almost every element Gives us relative electronegativities of all elements. Tends to increase left to right. decreases as you go down a group. Noble gases do not have values. Difference in electronegativity between atoms tells us how polar.

8. The ability of an atom in a molecule to attract shared electrons to itself. Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself.

9. Electronegativity difference Bond Type Zero Intermediate Large Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases

10. Ionic Bonds  Electrons are transferred  Electronegativity differences are generally greater than 1.7  The formation of ionic bonds is always exothermic!

11. Determination of Ionic Character Compounds are ionic if they conduct electricity in their molten state Electronegativity difference is not the final determination of ionic character

12. Q is the charge. r is the distance between the centers. If charges are opposite, E is negative –exothermic Same charge, positive E, requires energy to bring them together. –endothermic Coulomb’s Law

13. Size of ions Ion size increases down a group. Cations are smaller than the atoms they came from. Anions are larger. across a row they get smaller, and then suddenly larger. First half are cations. Second half are anions.

14. Periodic Trends Across the period nuclear charge increases so they get smaller. Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

15. Table of Ion Sizes

16. Size of Isoelectronic ions Iso - same Iso electronic ions have the same # of electrons Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 All have 10 electrons. All have the configuration 1s 2 2s 2 2p 6

17. Size of Isoelectronic ions Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

18. Ionic Compounds We mean the solid crystal. Ions align themselves to maximize attractions between opposite charges, and to minimize repulsion between like ions. Can stabilize ions that would be unstable as a gas. React to achieve noble gas configuration

20. List the following atoms in order of increasing ionization energy: Li, Na, C, O, F. 1.Li < Na < C < O < F 2.Na < Li < C < O < F 3.F < O < C < Li < Na 4.Na < Li < F < O < C 5.Na < Li < C < F < O

21. Sodium losing an electron is an ________ process and fluorine losing an electron is an _______ process. 1.endothermic, exothermic 2.exothermic, endothermic 3.endothermic, endothermic 4.exothermic, exothermic 5.more information needed

22. Which of the following statements is true about the ionization energy of Mg+? 1.It will be equal to the ionization energy of Li. 2.It will be equal to and opposite in sign to the electron affinity of Mg. 3.It will be equal to and opposite in sign to the electron affinity of Mg+. 4.It will be equal to and opposite in sign to the electron affinity of Mg2+. 5.none of the above

23. Choose the compound with the most ionic bond. 1.LiCl 2.KF 3.NaCl 4.LiF 5.KCl

24. In which pair do both compounds exhibit predominantly ionic bonding? 1.PCl 5 and HF 2.Na 2 SO 3 and BH 3 3.KI and O 3 4.NaF and H 2 O 5.RbCl and CaO

25. Which of the following arrangements is in order of increasing size? 1.Ga 3+ > Ca 2+ > K + > Cl – > S 2– 2.S 2– > Cl – > K + > Ca 2+ > Ga 3+ 3.Ga 3+ > S 2– > Ca 2+ > Cl – > K + 4.Ga 3+ > Ca 2+ > S 2– > Cl – > K + 5.Ga 3+ > Ca 2+ > S 2– > K + > Cl –

26. Which of the following species would be expected to have the lowest ionization energy? 1.F - 2.Ne 3.O 2- 4.Mg 2+ 5.Na +

27. Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

28. Forming Ionic Compounds Lattice energy - the energy associated with making a solid ionic compound from its gaseous ions. M + (g) + X - (g)  MX(s) This is the energy that “pays” for making ionic compounds. Energy is a state function so we can get from reactants to products in a round about way.

29. Calculating Lattice Energy Lattice Energy = k(Q 1 Q 2 / r) k is a constant that depends on the structure of the crystal. Q’s are charges. r is internuclear distance. Lattice energy is greater with more highly charged ions. This bigger lattice energy “pays” for the extra ionization energy. Also “pays” for unfavorable electron affinity.

30. Estimate  H f for Sodium Chloride Na(s) + ½ Cl 2 (g)  NaCl(s) Lattice Energy-786 kJ/mol Ionization Energy for Na495 kJ/mol Electron Affinity for Cl-349 kJ/mol Bond energy of Cl 2 239 kJ/mol Enthalpy of sublimation for Na109 kJ/mol Na(s)  Na(g) + 109 kJ Na(g)  Na + (g) + e - + 495 kJ ½ Cl 2 (g)  Cl(g) + ½(239 kJ) Cl(g) + e -  Cl - (g) - 349 kJ Na + (g) + Cl - (g)  NaCl(s) -786 kJ Na(s) + ½ Cl 2 (g)  NaCl(s) -412 kJ/mol

31. F-F- Cl - Br - I-I- Li + 1036853807757 Na + 923787747704 K+K+ 821715682649 Rb + 785689660630 Cs + 740659631604 Lattice Energies of Alkali Metals Halides (kJ/mol) OH - O 2- Na + 9002481 Mg 2+ 30063791 Al 3+ 562715,916 Lattice Energies of Salts of the OH - and O 2- Ions (kJ/mol)

32. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  Electrons are equally shared  Electronegativity difference of 0 to 0.3

33. Covalent Bonding Forces  Electron – electron repulsive forces repulsive forces  Proton – proton repulsive forces  Electron – proton attractive forces

34. Bond Length Diagram

35. The Octet Rule Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

36. Formation of Water by the Octet Rule

37. 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.

38. Shows how valence electrons are arranged among atoms in a molecule. 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. Reflects central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

39. C H H H Cl.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make carbon the central atom..

40. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons Ethene

41. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons Ethyne

42. Resonance Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. The actual structure is an average of the resonance structures. Benzene, C 6 H 6 The bond lengths in the ring are identical, and between those of single and double bonds.

43. Resonance Bond Length and Bond Energy Resonance bonds are shorter and stronger than single bonds. Resonance bonds are longer and weaker than double bonds.

44. Resonance in Ozone, O 3 Neither structure is correct. Oxygen bond lengths are identical, and intermediate to single and double bonds

45. Resonance in a carbonate ion: Resonance in an acetate ion: Resonance in Polyatomic Ions

46. Localized Electron Model Lewis structures are an application of the “Localized Electron Model” L.E.M. says: Electron pairs can be thought of as “belonging” to pairs of atoms when bonding Resonance points out a weakness in the Localized Electron Model.

47. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical

48. Fundamental Properties of Models  A model does not equal reality.  Models are oversimplifications, and are therefore often wrong.  Models become more complicated as they age.  We must understand the underlying assumptions in a model so that we don’t misuse it.

49. VSEPR – Valence Shell Electron Pair Repulsion X + E Overall Structure Forms 2 LinearAX 2 3 Trigonal PlanarAX 3, AX 2 E 4 TetrahedralAX 4, AX 3 E, AX 2 E 2 5 Trigonal bipyramidal AX 5, AX 4 E, AX 3 E 2, AX 2 E 3 6 OctahedralAX 6, AX 5 E, AX 4 E 2 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

50. VSEPR: Linear AX 2 CO 2

51. VSEPR: Trigonal Planar AX 3 AX 2 E BF 3 SnCl 2

52. VSEPR: Tetrahedral AX 4 AX 3 E AX 2 E 2 CCl 4 PCl 3 Cl 2 O

53. VSEPR: Trigonal Bi-pyramidal AX 5 AX 4 E AX 3 E 2 AX 2 E 3 PCl 5 SF 4 ClF 3 I3-I3-I3-I3-

54. VSEPR: Octahedral AX 6 AX 5 E AX 4 E 2 SF 6 ICl 4 - BrF 5