Chapter 8 Bonding

1. Chapter 8 Bonding

2. 8.1 Types of Bonds • There are lots of experiments we can do to determine the nature of materials • Melting point • Conductivity • Solubility • Charge distribution in an electric field • Bond Energy – The energy required to break a bond • Tells us the strength of a bonds interactions.

3. Ionic Compounds • We all know that ionic bonding is a result of electrostatic attractions of oppositely charged ions. • And ionic compounds are formed when a nonmetal and a metal react.

4. Coulomb’s law • Coulomb’s law is used to measure the energy of interaction between a pair of ions. • Where E is energy in joules, r is the distance between the ion centers in nm, and Q1 and Q2 are the charges of the ions.

5. Coulomb’s law in use • Let’s look at salt with ions at the given distance: • Notice the E is negative (indicating an attractive force) which means the ion pair has LOWER energy than the separated ions.

6. Repulsive forces • Let’s look at H-H bonds. When will a H2 molecule be favored • A bond will form if the energy of the aggregate (whole created from its parts) is lower than that of the separated ions!!

7. An energy profile

8. Bond length • The distance where the energy is MINIMAL is called the bond length. • The type of bond we seen in an H2 molecule is called a covalent bond. (where electrons are shared) A mutual attraction of the two nuclei for the shared electrons. \

9. Polar Covalent Bonds • In between the two extreme types of bonding (ionic and covalent) is an intermediate case; called polar covalent bonds. • These bonds are between atoms that are not so different that electrons are completely transferred but are different enough that there is unequal sharing.

10. F H H—F • Look at the HF molecule below. The symbol δ (lowercase delta) indicates a fractional charge. • This indicates that fluorine has a stronger attraction for the shared e- than hydrogen does. electron rich region electron poor region

11. 8.2 Electronegativity • Electronegativity – is the ability of an atom in a molecule to attract shared electrons to itself. • Linus Pauling’s model is what we used to assign a value to electronegativity. • The trend is that electronegativity generally increases across a period and decreases down a group.

13. Relationship between electronegativity and bond type.

14. Using electronegativity to determine bond polarity • Using electronegativity values, arrange the following bonds in order of increasing polarity: H – H, O – H, Cl – H, S – H, and F – H. Remember bond polarity is the difference of the electronegativities of the atoms forming the bond

15. P O L A R I T Y The answer!!In increasing polarity Bond Difference Bond type H – H 0 Covalent S – H 0.4 Polar covalent Cl – H 0.9 Polar covalent O – H 1.4 Polar covalent F – H 1.9 Polar covalent

16. 8.3 Bond Polarity & Dipole moments When a molecule has a center of positive charge and center of negative charge it is said to be dipolar or have a dipole moment. The dipolar character of a molecule is often represented by an arrow that points toward the negative charge center and the tail indicates the positive center of charge.

17. Another way to look at it… Electrostatic potential maps also show the charge distribution of a molar molecule. Red shows the electron rich area and the violet shows the electron poor region. Any diatomic molecule with a polar bond will show a molecular dipole moment

18. Sometimes… Polyatomic molecules exhibit dipolar behavior. And in an electric field, water acts as if it has 2 centers of charge.

19. Other times… Individual bond polarities are arranged in such a way that they “cancel” each other out. As seen below in the CO2 molecule. Polar bonds but NO dipole moment. There are many more of these cases!!

20. Types of Molecules with Polar Bonds but No Resulting Dipole Moment

21. Try these. Cl2 has NO bond polarity because the e- are shared equally. Therefore, there is NO dipole moment. Show the direction of the bond polarities and indicate which ones have a dipole moment: HCl Cl2 SO3 CH4 H2S

22. 8.4 Ions: Electron Configuration & Sizes • Generalizations: • When 2 nonmetals react to form a covalent bond, they share e- so that the valence electron configuration of both atoms is complete. (i.e. both nonmetals have attained noble gas electron configurations) • When a nonmetal and a representative metal react to form a binary ionic compound, ions form so that the valence e- configuration of the nonmetal achieves the e- config of the next noble gas and the valence orbitals of the metal are emptied. (both ions attain noble gas e- configurations)

23. Predicting formulas of ionic compounds When the term ionic compound is used, typically, the state of matter being referred to is a SOLID. The + and – ions are packed together in such a way so that the + + and - - repulsions are minimized and the + - attractions are maximized.

24. Sizes of Ions The size of an ion plays an important role in the stability of an ionic solid and the properties of the ions in solution. Most of the time ionic radii are determined by measuring the distances between ions in ionic compounds… BUT this assumes how the distance is divided between the two ions.

25. Controversy This method of determining ion size creates a controversy…sooo….we will concentrate on the trends! Consider relative ion size and the size of the parent atom. + ions are formed by removing an e- so the resulting cation is SMALLER than its parent atom. And the opposite is true for anions.

26. Depends on parents Remember the generalizations we talked about… Size depends on the “parent’s” position in the periodic table. A given period has both elements that give up and gain electrons. So some elements get larger (to emulate the next noble gas) and some elements get smaller (to “look like” the previous noble gas)

27. Isoelectronic ions. • Isoelectronic ions are ions that contain the same number of electrons. • For example O2-, F-, Na+, Mg2+, and Al3+ each have 10 electrons ([Ne]) so the amount of repulsions should be the same. • Let’s look at how their sizes vary.

28. What other factor should we look at for size? • When looking at isoelectronic ions, the other factor to consider is the number of protons in the nucleus. The number of protons increases from 8 to 13 from O2- to Al3+. • So as the number of protons increases, the ATTRACTION to the 10 electrons is GREATER and it causes the ions to become smaller.

29. 8.5 Energy Effect in Binary Ionic Compounds • We know that metals and nonmetals react by transferring e- and that the result is a solid ionic compound due to the oppositely charged ions having a lower energy than the original elements. • As always ENERGY tells us how strongly the ions are attracted to each other. In a solid ionic compound the energy is called lattice energy.

30. Lattice energy defined • Lattice energy is defined as the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. • It is the energy released when an ionic solid forms from its ions! • Remember the sign is determined from the system’s P.O.V! Exo = negative: since the energy is leaving the system.

31. Let’s look at the energy changes involved with the formation of this ionic solid • Sublimation – solid to gas phase • Ionization of Li • Dissociation of fluorine • Formation of F- ions • Formation of solid LiF from gaseous ions

32. This is an energy diagram!Notice the exothermic reaction…more energy is released than absorbed in the process

33. Lattice energy calculations • This leads us to a modified version of Coulomb’s law: • Where k is a proportionality constant that depends on the structure of the solid and the electron configurations of the ions and r is the shortest distance between the ions.

34. Looking at the formula, can you see that the process becomes MORE exothermic as the ioniccharges increase AND as the distance between the ions decreases • Let’s look at how charges affect energy by comparing MgO and NaF energy diagrams.

36. 8.6 Partial Ionic Character of Covalent Bonds • Let’s think of polar covalent bonds and ionic. We understand these bonds to be between atoms that have different electronegativities. • They either share e- unequally or transfer one or more e-. So how do we tell if it is polar covalent or ionic? • The REAL answer is there are NO completely ionic compounds. Let’s look.

38. A complication • The previous slide defies the idea that we know many of those compounds (that are above 50%) as ionic solid. • We must consider that the compounds in the table are in the gas phase. And that these results can not be assumed to also apply to the solid phase.

39. Another complication • Another complication in identifying ionic compounds is that many contain polyatomic ions. • Polyatomic ions are held together by covalent bonds • So calling NH4Cl or Na2SO4 ionic is ambiguous.

40. What will we call ionic? • So from now on…we define ionic compound as any compound that conducts an electric current when melted.

41. 8.7 The Covalent Chemical Bond: A Model • Bonding is a model proposed to explain molecular stability. • The bond concept is a human invention to provide a method for dividing up the energy evolved when a stable molecule is formed from its component atoms.

42. Sensible… • It is physically sensible and makes sense that atoms can form stable groups by sharing e- since shared e- give a lower energy state because they are simultaneously attracted by 2 nuclei.

43. 8.8 Covalent Bond Energies and Chemical Reactions • From the stepwise decomposition of methane it is determined that the energy to break a C—H bond varies in a nonsystematic way. This also shows that bond strength varies significantly with its environment. • So bond energy as an average is what is useful to chemists and given in a table format.

45. Single, double, tripleand I don’t mean baseball • Single bonds – 1 pair of e-shared. • Double bonds – 2 pair of e- shared. • Triple bonds – 3 pair of e-shared. • There is a relationship between bond length and # of shared e-. The more e- shared, the shorter the bond.

46. Bond Energy and Enthalpy • Using bond energy, the energies of reactions can be approximated. • Breaking bonds requires energy so the sign is + • Formation of bonds releases energy so the sign is -

47. The “formula” • ∆H = sum of the energies required to break old bonds plus the sum of the energies released in the formation of new bonds. energy required energy released ∑ is the sum of the terms and D represents the bond energy per mole of bonds. (D is always positive)

48. Let’s try with… H2(g) + F2(g)  2HF(g) We need to break 1) H—H bond 1) F—F bond We need to form 1) H—F bond

49. If we compare this with ∆H from standard enthalpy for HF (-271kJ/mol): ∆Ho = 2mol • -271kJ/mol = -542 kJ So bond energies work well to find ∆H

50. 8.9 The localized Electron Bonding Model • A “localized electron” assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. • Lone pairs are the e- that are localized on the atom (or in space) • Bonding pairs are those e- between the atoms.