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3. Chapter Nine Chemical Bonds

4. Bonding Review • Lewis dot structures: the symbol represents the nucleus and core electrons and dots represent the valence electrons

5. Example 9.1 Give Lewis symbols for magnesium, silicon, and phosphorus.

6. Bonding Review • Octet Rule – Eight is great, except for hydrogen and helium 2 will do!

7. Bonding Review • Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic compounds. • Chemical bonds are electrostatic forces; they reflect a balance in the forces of attraction and repulsion between electrically charged particles.

8. In this unit… • We will look at intramolecular forces and intermolecular forces: • Intra- means “within” a molecule • Examples are ionic, covalent, and metallic bonding • Inter- means “between” molecules • Examples are hydrogen bonding, dipole interactions, and dispersion forces • First let’s look at intramolecular forces.

9. Bonding Review • Ionic bonding – a bond between a metal and a nonmetal where one atom gives an electron and one takes an electron.

10. Bonding Review • When atoms lose or gain electrons, they may acquire a noble gas configuration, but do not become noble gases. • Because the two ions formed in a reaction between a metal and a nonmetal have opposite charges, they are strongly attracted to one another and form an ion pair. • The net attractive electrostatic forces that hold the cations and anions together are ionic bonds. • The highly ordered solid collection of ions is called an ionic crystal.

11. Bonding Review Na donates an electron to Cl … … and opposites attract. Sodium reacts violently in chlorine gas.

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15. Using Lewis Symbolsto Represent Ionic Bonding • Lewis symbols can be used to represent ionic bonding between nonmetals and: the s-block metals, some p-block metals, and a few d-block metals. • Instead of using complete electron configurations to represent the loss and gain of electrons, Lewis symbols can be used.

16. Example 9.2 Use Lewis symbols to show the formation of ionic bonds between magnesium and nitrogen. What are the name and formula of the compound that results?

17. Try it Out! • Use Lewis symbols to show the formation of ionic bonds between sodium and phosphorus. What are the name and formula of the compound that results?

18. Lattice Energy • Energy released when positive and negative ions form crystal lattice due to their attraction for each other. • Stability of ionic compounds (high melting point, brittle) because of large lattice energy

19. Coulomb’s Law • Describes the electrostatic interaction between charged particles. • Says that the force of attraction or repulsion between two point charges is directly proportional to the product of magnitude of each charge and indirectly proportional to the square of distance between them

20. Coulomb’s Law • Simply put: • The higher the charges on ions, the greater the lattice energy. • The higher the distance between ions, the smaller the lattice energy. • If the charges are opposite in sign the forces are attractive, if they are the same sign the forces are repulsive.

21. Examples • Which has more lattice energy: LiBr or CaO? • CaO because they both have a charge of 2 while Li and Br have a charge of 1. • Which has more lattice energy: NaCl or CsCl? • NaCl because all ions have the same charge, but Cs is much larger. Therefore the distance between atoms in CsCl is greater and the lattice energy is smaller.

22. Sample AP Exam Question • The melting point of MgO is higher than that of NaF. Explanations for this observation include which of the following? • Mg+2 is more positively charged than Na+1 • O-2 is more negatively charged than F-1 • The O-2 is smaller than the F-1 ion. • II only • I and II only • I and III only • II and III only • I, II, and III B

23. Bonding Review • Covalent bond – a bond between two nonmetals where electrons are shared.

24. Bonding Review • Single covalent bond – a pair of electrons is shared between two atoms (one dash/line) Nonbonding electrons/lone pairs Bonding electrons

25. Bonding Review • Some molecules require more than single bonds to provide each atom with the required octet (formed primarily by C, N, and O). • Double bond – atoms share 2 pair of electrons • Triple bond – atoms share 3 pair of electrons

26. The Lewis Theory ofChemical Bonding: An Overview • Valence electrons play a fundamental role in chemical bonding. • In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases.

27. Lewis Structures ofSimple Molecules • A Lewis structure is a combination of Lewis symbols that represents the formation of covalent bonds between atoms. • In most cases, a Lewis structure shows the bonded atoms with the electron configuration of a noble gas; that is, the atoms obey the octet rule. (H obeys the duet rule.) • The shared electrons can be counted for each atom that shares them, so each atom may have a noble gas configuration.

28. Some Illustrative Compounds • Note that the two-dimensional Lewis structures do not necessarily show the correct shapes of the three-dimensional molecules. Nor are they intended to do so. • The Lewis structure for water may be drawn with all three atoms in a line: H–O–H. • We will learn how to predict shapes of molecules in Chapter 10.

29. Electronegativity • Electronegativity (EN) is a measure of the ability of an atom to attract its bonding electrons to itself. • EN is related to ionization energy and electron affinity. • The greater the EN of an atom in a molecule, the more strongly the atom attracts the electrons in a covalent bond. Electronegativity generally increases from left to right within a period, and it generally increases from the bottom to the top within a group.

30. Pauling’s Electronegativities Electronegativity has no unit because the values are comparative only. It would be a good idea to remember the four elements of highest electronegativity: N, O, F, Cl.

31. Example 9.4 Referring only to the periodic table inside the front cover, arrange the following sets of atoms in the expected order of increasing electronegativity. (a) Cl, Mg, Si (b) As, N, Sb(c) As, Se, Sb

32. Writing Lewis Structures:Skeletal Structures • The skeletal structure shows the arrangement of atoms. • Hydrogen atoms are terminal atoms (bonded to only one other atom). • The central atom of a structure usually has the lowestelectronegativity. • In oxoacids (HClO4, HNO3, etc.) hydrogen atoms are usually bonded to oxygen atoms. • Molecules and polyatomic ions usually have compact, symmetrical structures.

33. Writing Lewis Structures: A Method • Determine the total number of valence electrons. • Write a plausible skeletal structure and connect the atoms by single dashes (covalent bonds). • Place pairs of electrons as lone pairs around the terminal atoms to give each terminal atom (except H) an octet. • Assign any remaining electrons as lone pairs around the central atom. • If necessary (if there are not enough electrons), move one or more lone pairs of electrons from a terminal atom to form a multiple bond to the central atom.

34. Some Handy Rules to Remember • Hydrogen and the halogens bond once. • The family oxygen is in can bond twice. • The family nitrogen is in can bond three times.  So can boron. • The family carbon is in can bond four times.

35. Example Example 9.6 Write the Lewis structure of nitrogen trifluoride, NF3.

36. Example Example 9.7 Write a plausible Lewis structure for phosgene, COCl2.

37. Example • Example 9.6 A – Write the Lewis structure of hydrazine N2H4.

38. Try It Out • Example 9.7 A – Write a plausible Lewis structure for carbonyl sulfide, COS.

39. Try It Out • Example 9.7 B – Write a plausible Lewis structure for nitrosyl chloride, NOCl.

40. What if an element has charge? • Negative = add electrons • Positive = subtract electrons

41. Example • Example 9.8 – Write a plausible Lewis structure for the chlorate ion, ClO3-1.

42. Try it Out! • Example 9.8 B – Write a plausible Lewis structure for the nitronium ion, NO2+1.

43. Resonance • Can draw more than one way because of multiple bonds.

44. Resonance: Delocalized Bonding • When a molecule or ion can be represented by two or more plausible Lewis structures that differ only in the distribution of electrons, the true structure is a composite, or hybrid, of them. • The different plausible structures are called resonance structures. • The actual molecule or ion that is a hybrid of the resonance structures is called a resonance hybrid. • Electrons that are part of the resonance hybrid are spread out over several atoms and are referred to as being delocalized. Three pairs of electrons are distributed among two bonds.

45. Example • Example 9.10 – Write three equivalent Lewis structures for the SO3 molecule that conform to the octet rule.

46. Try it Out! • Example 9.10 A – Write three Lewis structures for the nitrate ion, NO3-1.

47. Molecules that Don’t Followthe Octet Rule • Molecules with an odd number of valence electrons have at least one of them unpaired and are called free radicals. • Some molecules have incomplete octets. These are usually compounds of Be, B, or Al; they generally have some unusual bonding characteristics, and are often quite reactive. • Some compounds have expanded valence shells, which means that the central atom has more than eight electrons around it. • A central atom can have expanded valence if it is in the third period or lower (i.e., S, Cl, P).

48. Example 9.11 Write the Lewis structure for bromine pentafluoride, BrF5.

49. Example • Example 9.11 A – Write the Lewis structure of phosphorus trichloride.

50. Try It Out! • Example 9.11 B – Write the Lewis structure of chlorine trifluoride.