Bonding: General Concepts

1. Bonding: General Concepts AP Chemistry Unit 8 Author: BobCatChemistry

2. Types of Chemical Bonds

3. Ionic Bonds • Ionic Bonds are formed when an atom that loses electrons relatively easily reacts with an atom that has a high attraction for electrons. • Ionic Compounds results when a metal bonds with a nonmetal.

4. Bond Energy • Bond energy is the energy required to break a bond. • The energy of interaction between a pair of ions can be calculated using Coulomb’s law • r = the distance between the ions in nm. • Q1 and Q2 are the numerical ion charges. • E is in joules

5. Bond Energy • When the calculated energy between ions is negative, that indicates an attractive force. • A positive energy is a repulsive energy. • The distance where the energy is minimal is called the bond length.

6. Covalent Bonds • Covalent bonds form between molecules in which electrons are shared by nuclei. • The bonding electrons are typically positioned between the two positively charged nuclei.

7. Polar Covalent Bonds • Polar covalent bonds are an intermediate case in which the electrons are not completely transferred but form unequal sharing. • A δ- or δ+ is used to show a fractional or partial charge on a molecule with unequal sharing. This is called a dipole.

8. Electronegativity

9. Electronegativity Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. (electron love) • Relative electronegativities are determined by comparing the measured bond energy with the “expected” bond energy. • Measured in Paulings. After Linus Pauling the American scientist who won the Nobel Prizes for both chemistry and peace.

10. Electronegativity • Expected H-X bond energy=

11. Electronegativity • Electronegativity values generally increase going left to right across the periodic table and decrease going top to bottom.

12. Electronegativity and Bond type

13. Bond Polarity and Dipole

14. Dipoles and Dipole Moments A molecule that has a center of positive charge and a center of negative charge is said to be dipolar or to have a dipole moment. • An arrow is used to show this dipole moment by pointing to the negative charge and the tail at the positive charge.

15. Dipoles and Dipole Moments • Electrostatic potential diagram shows variation in charge. Red is the most electron rich region and blue is the most electron poor region.

16. Dipoles and Dipole Moments

17. Dipoles and Dipole Moments

18. Dipoles and Dipole Moments

19. Dipoles and Dipole Moments • Dipole moments are when opposing bond polarities don’t cancel out.

20. Dipoles and Dipole Moments

21. Example Problems • For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl2, SO3, CH4, H2S

22. HCl

23. Cl2

24. SO3

25. CH4

26. H2S

27. Ions: Electron Configurations and Sizes

28. Electron Configurations of Compounds • When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence electron configurations of both atoms. That is, both nonmetals attain noble gas electron configurations.

29. Electron Configurations of Compounds • When a nonmetal and a representative-group metal react to form a binary ionic compounds, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbitals of the metal are emptied. In this way both ions achieve noble gas electron configurations.

30. Predicting Ionic Formulas • To predict the formula of the ionic compound, we simply recognize that the chemical compounds are always electrically neutral. They have the same quantities of positive and negative charges.

31. Sizes of Ions Size of an ion generally follows the same trend as atomic radius. The big exception to this trend is where the metals become nonmetals and the ions switch charge.

32. Sizes of Ions A positive ion is formed by removing one or more electrons from a neutral atom, the resulting cation is smaller than the neutral atom. • Less electrons allow for less repulsions and the ion gets smaller.

33. Sizes of Ions An addition of electrons to a neutral atom produces an anion that is significantly larger than the neutral atom. • An addition of an electron causes additional repulsions around the atom and therefore its size increases.

34. Energy Effects in Binary Ionic Compounds

35. Lattice Energy Lattice energy is the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. • The lattice energy is often defined as the energy released when an ionic solid forms from its ions. • Lattice energy has a negative sign to show that the energy is released.

36. Lattice Energy Example Estimate the enthalpy of lithium fluoride and the changes of energy and lattice energy during formation: • Break down LiF into its standard state elements (use formation reaction): Li(s) + ½F2(g)  LiF(s) Li+(g) + F-(g) LiF(s)

37. Lattice Energy Example Li(s) + ½F2(g)  LiF(s) Li+(g) + F-(g)  LiF(s) • Use sublimation and evaporation reactions to get reactants into gas form (since lattice energy depends on gaseous state). Find the enthalpies to these reactions: Li(s) Li(g) 161 kJ/mol Li(g) + ½F2(g)  LiF(s)

38. Lattice Energy Example Li(g) + ½F2(g)  LiF(s) Li+(g) + F-(g)  LiF(s) • Ionize cation to form ions for bonding. Use Ionization energy for the enthalpy of the reaction. Li(g) Li+(g) + e- Ionization energy: 520 kJ/mol Li+(g) + ½F2(g)  LiF(s)

39. Lattice Energy Example Li+(g) + ½F2(g)  LiF(s) Li+(g) + F-(g)  LiF(s) • Dissociate diatomic gas to individual atoms: ½F2(g) F(g) ½ Bond dissociation energy of F-F = 154 kJ/ 2 = 77 kJ/mol Li+(g) + F(g)  LiF(s)

40. Lattice Energy Example Li+(g) + F(g)  LiF(s) Li+(g) + F-(g)  LiF(s) • Electron addition to fluorine is the electron affinity of fluorine: F(g) + e- F-(g) -328 kJ/mol Li+(g) + F-(g)  LiF(s)

41. Lattice Energy Example Li+(g) + F-(g)  LiF(s) Li+(g) + F-(g)  LiF(s) • Formation of solid lithium fluoride from the gaseous ions corresponds to its lattice energy: Li+(g) + F-(g)  LiF(s) -1047 kJ/mol

42. Lattice Energy Example The sum of these five processes yields the overall reaction and the sum of the individual energy changes gives the overall energy change and the enthalpy of formation: 161 kJ 520 kJ 77 kJ -328 kJ -1047 kJ Total = -617 kJ/mol • Li(s) Li(g) • Li(g) Li+(g) + e- • ½F2(g) F(g) • F(g) + e- F-(g) • Li+(g) + F-(g)  LiF(s)

43. Lattice Energy

44. Lattice Energy Lattice energy can be calculated with at form of Coulomb’s law: • Q is the charges on the ions and r is the shortest distance between the centers of the cations and anions. k is a constant that depends on the structure of the solid and the electron configurations of the ions.

45. Partial Ionic Character of Covalent Bonds

46. Bond Character Calculations of ionic character: Even compounds with the maximum possible electronegativity differences are not 100% ionic in the gas phase. Therefore the operational definition of ionic is any compound that conducts an electric current when melted will be classified as ionic.

47. Bond Character

48. The Covalent Chemical Bond

49. Chemical Bond Model A chemical bond can be viewed as forces that cause a group of atoms to behave as a unit. • Bonds result from the tendency of a system to seek its lowest possible energy. • Individual bonds act relatively independent.

50. Example • It takes 1652 kJ of energy required to break the bonds in 1 mole of methane. • 1652 kJ of energy is released when 1 mole of methane is formed from gaseous atoms. • Therefore, 1 mole of methane in gas phase has 1652 kJ lower energy than the total of the individual atoms. • One mole of methane is held together with 1652 kJ of energy. • Each of the four C-H bonds contains 413 kJ of energy.