Chapters 9 and 10

1. Chapters 9 and 10 Chemical Bonding I and II

2. Metallic Bonds • Attraction between a lattice (pattern) of positive ions and delocalized electrons • “sea of electrons” • Malleable – metals can be reshaped without breaking because of the lattice structure • High Conductivity – electrons are free to move around which means metals are good conductors • What is a metallic bond?

3. Ionic Bond • An electrostatic force that holds ions together in an ionic compound • Opposite charges attract • Like charges repel • Two ways: • Periodic position • Difference in electronegativity • What is an ionic bond? • How do we determine whether a bond is ionic or not?

4. Ionic Bond • Metals are always cations • Nonmetals can be anions • A metal-nonmetal bond is called an ionic bond • Ex. NaCl • O-H • Ca-Br • Be-Br • How does periodic position determine whether a bond will be ionic? • Are the following bonds ionic, yes or no?

5. Ionic Bond • Determine the difference between the electronegativity values • Generally, if the difference is above 1.7 (book uses 2.0), the bond is considered ionic • Please understand that there is no definite cut off point • O-H • Ca-Br • Be-Br • How does the difference in EN work for ionic bonds? • Are the following bonds ionic based on the their EN values, Yes or no?

6. Electronegativity Table

7. Lewis Dot Diagrams (Symbols) • Responsible for how an element reacts • Shows the element symbol with only the number of valence electrons (dots) • NOTE: Remember how we placed them before • Ex. Nitrogen • Why are valence electrons important? • How do Lewis Dot Structures emphasize this?

8. Lewis Dot Diagrams Practice • Determine the Lewis Dot Structures for Carbon, Magnesium and Phosphorus.

9. Formation of an Ionic Bond Practice • Using Lewis Dot Structures, show the formation of sodium chloride (NaCl). • Using the chips, show all of the reactants with their valence electrons: Na + Cl • Using the chips, show the product, sodium chloride. • Note: It will still look like separate ions. • What happened to the electron? • IT WAS TRANSFERRED! • The positive and negative ions attract each other which creates the bond!

10. Formation of an Ionic Bond Practice • Using Lewis Dot Structures, show the formation of magnesium nitride (Mg3N2). • Using the chips, show all of the reactants with their valence electrons: 3Mg + 2N  • Using the chips, show the product, magnesium nitride. • Note: It will still look like separate ions.

11. Formation of an Ionic Bond Practice • Using Lewis Dot Structures, show the formation of aluminum oxide (Al2O3). • Using the chips, show all of the reactants with their valence electrons. • Using the chips, show the product, aluminum oxide.

12. Ionic Compounds • Numerous ionic bonds holding ions together through static electric charge • At STP, all ionic compounds are solid • Lattice shape, ions stacked on top of each other in a pattern • What is the physical structure of an ionic compound?

13. Covalent Bonds • A bond where at least two electrons are shared • Usually, between two or more nonmetals (or metalloids) • Each shared electron is attracted to the nucleus of each atom in the bond • What are covalent bonds?

14. Covalent Bonds • EACH shared pair is represented by a single line • Lone pairs of electrons NEED to be shown (can be left as dots or a line) • Ex. F2 • Due to sharing, all elements in the compound have a full valence shell • How do we represent covalent bonds using Lewis structures? • How does this support the octet rule?

15. Multiple Covalent Bonds • More lines! • As long as the octet rule is obeyed multiple pairs can be shared! • One shared pair – single bond • Two shared pairs – double bond • Three shared pairs – triple bond • How do we represent multiple covalent bonds using Lewis structures?

16. Writing Lewis Structures (pg. 343) • Using the chemical symbols place bonded atoms next to each other. • Generally, the least electronegative is in the center. • Count the number of valence electrons present. • Begin drawing single covalent bonds connecting the central atom to other atoms until the octet rule is satisfied. Do not forget to show lone pairs! • If the octet rule is not satisfied for the central atom, begin adding double and triple bonds using lone pairs.

17. Modeling – Multiple Covalent Bonds • Using Lewis Dot Structures, show the formation of carbon dioxide (CO2). • Using the chips, show all of the reactants with their valence electrons. C + 2O  • Using the chips, show the product, carbon dioxide. • NOTE: You are going to need to share two pairs to each oxygen. • Remove each shared pair and replace it with a line.

18. Modeling – Multiple Covalent Bonds • Using Lewis Dot Structures, show the formation of diatomic nitrogen (N2). • Using the chips, show all of the reactants with their valence electrons. 2N  • Using the chips, show the product, diatomic nitrogen. • Remove each shared pair and replace it with a line.

19. Covalent Bonding Practice • On your whiteboard generate the Lewis structures for the following covalent compounds? • Make sure to show lone pairs and multiple bonds! • Nitrogen Triflouride (NF3) • Water (H2O) • Oxygen (O2) • Methane (CH4)

20. Ionic vs. Covalent • Covalent: • Weak bonds • Usually gasses and liquids • Insoluble in water • Non-electrolytes • Ionic: • Strong electrostatic bonds • Usually solids • Soluble in water • Good electrolytes • How are the physical properties of ionic compounds different from covalent?

21. Length Does Matter • Bond length is the distance between two bonded atoms • Single bonds are longer and more unstable • Triple bonds are short but are more stable • How does the type of bond affect the length of the bond?

22. Bond Dissociation Energy (BDE) • Bond dissociation energy • Also called bond enthalpy • Energy change required to “break” a specific chemical bond • Due to other forces, must be calculated when compounds are gaseous • Molar quantity • Ex. Cl2 • Cl2 Cl + Cl • ∆H = 242.7 kJ • How do we measure how stable (strong) a bond is?

23. Bond Dissociation Energy • Due to other molecular forces in a compound, we simplify and use the average BDE • Notice the increase in stability (strength) as the number of bonds is increased • These average bond energies can be used to determine enthalpies of reaction (∆Hrxn) • How do we determine BDE when dealing with polyatomic compounds?

24. Bond Dissociation Energy • Determine the enthalpy of the reaction: H2(g) + F2(g)  2HF(g)

25. Bond Dissociation Energy • Determine the enthalpy of the reaction: H2(g) + F2(g)  2HF(g)

26. Bond Dissociation Energy • Determine the enthalpy of the reaction: H2(g) + F2(g)  2HF(g)

27. Bond Dissociation Energy • Determine the enthalpy of the reaction: H2(g) + F2(g)  2HF(g) • ∆Hrxn = ΣBEreactants– ΣBEproducts • Can also be break minus make • ∆Hrxn = 593.3 kJ – 1136.4 kJ • ∆Hrxn = -543.1 kJ • Exothermic Reaction • This method is NOT as accurate as the Direct/Indirect (Hess) Methods in Chapter 6 primarily because we are using averages!

28. Bond Dissociation Energy Practice • Determine the enthalpy of the reaction: 3F2(g) + NH3(g)  3HF(g) + NF3(g) • Answer = -871 kJ

29. Bond Dissociation Energy Practice N-F 343 kJ/mol

30. Bond Dissociation Energy Practice • Determine the enthalpy of the reaction: 2H2(g) + O2(g)  2H2O(g) • Answer = -871 kJ

31. Bond Dissociation Energy Practice • Determine the enthalpy of the reaction: H2(g) + C2H4(g)  C2H6(g) • Answer = -871 kJ

32. VSEPR • Need to use the VSEPR (Valence Shell Electron Pair Repulsion) model • Based on: • How many atoms are bonded to the central atom • How many lone pairs are on the central atom • Single, double and triple bonds can be treated as if they were the same • How do we determine molecular shapes?

33. VSEPR • What are the possible VSEPR geometries?

34. How many atoms and lone pairs are connected to the central atom? How many lone pairs are on the atom? 1 4 Linear – 180° 0 of 4 1 of 4 2 of 4 3 of 4 Tetrahedral – 109.5° Bent – 104.5° How many lone pairs are on the atom? 2 Trigonal Pyramidal – 107.5° 0 of 2 1 of 2 Linear – 180° 3 Linear – 180° Linear – 180° How many lone pairs are on the atom? VSEPR Flow Chart 0 of 3 1 of 3 2 of 3 Trigonal Planar – 120° Bent – 118° Linear – 180°

35. VSEPR Practice • Please build the following molecule, draw it, and determine its shape and bond angle. • Methane (CH4) – black • Tetrahedral • Ammonia (NH3) – dark blue • Trigonal pyramidal • Water (H2O) – red • Bent • Carbon Dioxide (CO2) – black/red • Linear • Ammonium (NH4+) – dark blue • Tetrahedral • Hydronium (H3O+) – red • Trigonal pyramidal

36. VSEPR Anomalies • Covalent (molecular) compounds that have a net overall charge • Ammonium (NH4+) • Hydronium (H3O+) • It has partial charges within the molecule that move back and forth • This is called resonance! • What are coordinate covalent compounds? • What is peculiar about sulfur dioxide (SO2)?

37. Octet Rule Exceptions • Incomplete Octet • Be will make only two, usually covalent, bonds with its 2 VE • B will make 3 bonds with its 3 VE • Odd-Electron Molecules • Molecules with odd numbers of electrons will never be able to obtain 8 VE, such as NO and NO2 • Resonance occurs • The Expanded Octet • Elements in and beyond the third period can have more than 8 VE • Transition elements use the 18 electron rule because you move into the d and f orbitals • What are some exceptions to the octet rule?

38. Allotropes • What is an allotrope? • An element whose bonds can be rearranged to form different structures and patterns • Each structure, however, has its own distinct properties This photo shows waxy white phosphorus (yellow cut), red phosphorus, violet phosphorus and black phosphorus. The allotropes of phosphorus have markedly different properties from each other.

39. Allotropes of Carbon • Graphite • Bond angle = 120 • Multi-layered interlocking hexagon pattern • Layers can slide past one another making graphite relatively soft • Electrons move between layers allowing it to conduct electricity • 1 layer  Graphene • Compare and contrast the three main allotropes of carbon.

40. Graphene

41. Allotropes of Carbon • Diamond • Bond angle = 109.5 • All four of the outer electrons on the carbon atom form bonds with other carbon atoms so there are no mobile electrons • Diamond does not conduct electricity • Rigid structure, incredibly hard to break • Compare and contrast the three main allotropes of carbon.

42. Allotropes of Carbon • Fullerenes • Two hollow shapes: • Sphere – Buckeyball • Cylinder – Nanotube • Only known structures of carbon that can be dissolved • Much of their properties are still unknown and are being tested • Compare and contrast the three main allotropes of carbon.

43. Allotropes of Silicon • Amorphous (brown) and Crystalline (“silverish”) • Bond angles = 109.5 • Both have similar structures to diamond, but amorphous silicon has some disconnected bonds allowing it to be more flexible • Compare and contrast the two main allotropes of silicon.

44. Silicon Dioxide • Usually called silica • Most common compound found in the Earths crust • Many shapes exist such as tetrahedral and helical • Many forms exist such as quartz and sand • Used in electronics, glass production, biomedical research, optics, space shuttle and satellites to name a few… • What are the special properties of silicon dioxide?

45. Polarity • Describes how the electrons are shared within the bond • “tug of war” • Based on EN differences • Ionic (extremely polar!) • Difference in EN > 1.7 (or 2.0) • Polar covalent • Difference in EN  ~ 0 – 1.7 ( or 2.0) • Nonpolar covalent • Difference in EN ~ 0 • What is polarity? • What are the different types of bonds?

46. Polarity • Dipole moment • Modified arrow that points towards the more EN atom in the bond • Ex. H–F • How do we visually represent polarity within a bond?

47. Polarity • Electron density is shifted so the atoms become partially charged • Use lower case delta to represent a partial charge • Ex. H–F • What effect does a polar bond have on the charges of the atoms?

48. Polarity • Please determine the difference in EN, determine the type of bond and draw a dipole moment with partial charges if appropriate. (Data Booklet – Table 7) • O-H • C-H • F-F • S-O • F-Li

49. Polarity • Determine the direction and relative strength of every dipole moment in the molecule and then we can treat them like vectors • They can… • combine with each other • cancel each other out • If a molecule has polar bonds, how do we determine the polarity of the molecule?

50. Polarity • Two cases: • Symmetrical – dipole moments cancel so the molecule is nonpolar • Ex. BF3 • Asymmetrical – dipole moments do not cancel and you get a resultant dipole moment • Ex. H2O • How do we determine the polarity of a molecule with polar bonds?