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 Q 1 and Q 2 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 H 2 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 H 2 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. 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. H F 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. The answer!! In increasing polarity Bond DifferenceBond type H – H0Covalent S – H 0.4 Polar covalent Cl – H0.9 Polar covalent O – H 1.4 Polar covalent F – H 1.9 Polar covalent POLARITYPOLARITY

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 CO 2 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. Show the direction of the bond polarities and indicate which ones have a dipole moment: HCl Cl 2 SO 3 CH 4 H 2 S Cl 2 has NO bond polarity because the e- are shared equally. Therefore, there is NO dipole moment.

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 O 2-, F -, Na +, Mg 2+, and Al 3+ 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 O 2- to Al 3+. 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 ionic charges 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 NH 4 Cl or Na 2 SO 4 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, triple and 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… H 2 (g) + F 2 (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): ∆H o = 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.

51. Localized electron model when applied tells us… 1.The description of the arrangement of the valence electrons (Lewis structures) 2.A prediction of the molecule’s geometry (VSEPR) 3.A description of the atomic orbital types used to share electrons or hold long pairs. (we’ll talk about this in Chapter 9)

52. 8.10 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.

53. Steps for writing/drawing Lewis Structures 1.Sum the valence electrons. (Add for – ions, subtract for + ions) 2.Use a pair of electrons to form a bond between each pair of bound atoms. 3.Atoms usually have noble gas configurations. Exceptions are Hydrogen (duet rule) and more we’ll talk about later.

54. 8.11 Exceptions to the Octet Rule We know that hydrogen is an exception as hydrogen’s outer most orbital is the 1s so when it is full it is 1s 2 therefore it follows the duet rule not the octet rule

55. Another e- poor exception Boron – Boron is electron deficient therefore it often only bonds with 6 electrons around it NOT 8. This is indicated often by its violent reactivity. Electron poor atoms are typically found in period 2 (this makes sense, because they don’t have 8 e- to begin with). They will NEVER exceed 8 since their 2s and 2p orbitals can’t exceed 8 B – 3e - 3F – 3x7e - 24e - FBF F 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24

56. Electron rich exceptions There are some occasions where an atom will exceed the octet rule. This is only observed for some elements in period 3 and beyond. Common ones are S and P, but there are more too! SF 6 S F F F F F F S – 6e - 6F – 42e - 48e - 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48

57. Should we try? Draw a Lewis structure for each of the following molecules: H 2 N 2 O 2 F 2 H 2 O NH 3 CO They get harder as you go CO 2 CH 3 OH *BeCl 2 C 2 H 6 O *PCl 5 NO 3 - *ICl 4 -

58. 8.12 Resonance Resonance is when there is more than one valid Lewis structure for a particular model Reality is that the actual structure is an average of the resonance structures.

59. Here’s one example OCO O OCO O OCO O 2-

60. However, Resonance is necessary because the LE model assumes that e- are localized between atoms, however, we know that electrons are really delocalized and they can move around the entire molecule. Resonance “compensates” for this defective assumption.

61. Odd electron molecules There are very FEW of them!!! Formed from nonmetals. The most common example is NO N – 5e - O – 6e - 11e - NO N O

62. Formal Charge How do we decide which Lewis Structure best describes the actual bonding in a molecule. One method is to estimate the charge on each atom in the various possible Lewis structures. We will use formal charge to do this.

63. Things to know: Formal charge is the difference between the # of valence e- on the free atom and the # of valence e- assigned to the atom in the molecule To use formal charge we need to know 2 things. 1.The number of valence electrons on the free neutral atom 2.The number of e- belonging to the atom in the molecule.