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Bonding – General Concepts

Bonding – General Concepts. 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. Ionic Bonding.

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Bonding – General Concepts

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  1. Bonding – General Concepts

  2. 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.

  3. 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.

  4. 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 • D = (H-X) actual - (H-X)expected

  5. Electronegativity • D 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. • Most Noble gases do not have values. • Difference in electronegativity between atoms tells us how polar.

  6. Electronegativity: The ability of anatom in a molecule to attract shared electrons to itself.

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

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

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

  10. Coulomb’s Law • 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

  11. 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.

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

  13. Table of Ion Sizes

  14. 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 1s22s22p6

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

  16. 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

  17. List the following atoms in order of increasing ionization energy: Li, Na, C, O, F. • Li < Na < C < O < F • Na < Li < C < O < F • F < O < C < Li < Na • Na < Li < F < O < C • Na < Li < C < F < O

  18. Sodium losing an electron is an ________ process and fluorine losing an electron is an _______ process. • endothermic, exothermic • exothermic, endothermic • endothermic, endothermic • exothermic, exothermic • more information needed

  19. Which of the following statements is true about the ionization energy of Mg+? • It will be equal to the ionization energy of Li. • It will be equal to and opposite in sign to the electron affinity of Mg. • It will be equal to and opposite in sign to the electron affinity of Mg+. • It will be equal to and opposite in sign to the electron affinity of Mg2+. • none of the above

  20. Choose the compound with the most ionic bond. • LiCl • KF • NaCl • LiF • KCl

  21. In which pair do both compounds exhibit predominantly ionic bonding? • PCl5 and HF • Na2SO3 and BH3 • KI and O3 • NaF and H2O • RbCl and CaO

  22. Which of the following arrangements is in order of increasing size? • Ga3+ > Ca2+ > K+ > Cl– > S2– • S2– > Cl– > K+ > Ca2+ > Ga3+ • Ga3+ > S2– > Ca2+ > Cl– > K+ • Ga3+ > Ca2+ > S2– > Cl– > K+ • Ga3+ > Ca2+ > S2– > K+ > Cl–

  23. Which of the following species would be expected to have the lowest ionization energy? • F- • Ne • O2- • Mg2+ • Na+

  24. 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.

  25. 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.

  26. Calculating Lattice Energy • Lattice Energy = k(Q1Q2 / 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.

  27. Estimate Hf for Sodium Chloride Na(s) + ½ Cl2(g)  NaCl(s) Na+(g) + Cl-(g)  NaCl(s) -786 kJ Na(g)  Na+(g) + e- + 495 kJ ½ Cl2(g)  Cl(g) + ½(239 kJ) Cl(g) + e- Cl-(g) - 349 kJ Na(s)  Na(g) + 109 kJ Na(s) + ½ Cl2(g)  NaCl(s) -412 kJ/mol

  28. Lattice Energies of Alkali Metals Halides (kJ/mol) Lattice Energies of Salts of the OH- and O2- Ions (kJ/mol)

  29. 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.

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

  31. Covalent Bonding Forces • Electron – electron • repulsive forces • Proton – proton • repulsive forces • Electron – proton • attractive forces

  32. Bond Length Diagram

  33. 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

  34. Formation of Water by the Octet Rule

  35. 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.

  36. Lewis Structures • 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.

  37. Completing a Lewis Structure -CH3Cl • Make carbon the central atom • Add up available valence electrons: • C = 4, H = (3)(1), Cl = 7 Total = 14 • Join peripheral atoms • to the central atom • with electron pairs. H .. .. .. H .. • Complete octets on • atoms other than • hydrogen with remaining • electrons C .. Cl .. .. H

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

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

  40. Resonance • Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. Benzene, C6H6 • The actual structure is an average of the resonance • structures. • The bond lengths in the ring are identical, and • between those of single and double bonds.

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

  42. Resonance in Ozone, O3 Neither structure is correct. Oxygen bond lengths are identical, and intermediate to single and double bonds

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

  44. 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.

  45. 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

  46. 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.

  47. VSEPR – Valence Shell Electron Pair Repulsion A = central atom X =atoms bonded to A E = nonbonding electron pairs on A

  48. VSEPR: Linear AX2 CO2

  49. VSEPR: Trigonal Planar AX3 BF3 AX2E SnCl2

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