Chemical Bonding

1. Chemical Bonding Chapters 8 and 9

2. Chemical Bonds • What is a bond? • A force that holds atoms together • We will look at it in terms of energy • Bond Energy is the NRG required to break a bond • Why are compounds formed? • Bonds give the system the lowest NRG

3. Bond Energy is the energy required to break a bond. REMEMBER: • It always takes energy to break a chemical bond. AND, • To form a bond, requires a lowering of energy.

4. Types of Bonds • Ionic Bonds – electrostatic forces that exist between ions of opposite charge • Covalent Bonds – sharing of electrons between two atoms • Metallic Bonds – found typically in transition metals; each atom is bonded to several neighboring atoms; bonding electrons are relatively free to move throughout the structure of the metal.

5. Ionic Bonding • An atom with a low ionization energy reacts with an atom with high electron affinity. • Typically a metal with a nonmetal • The electron transfers atoms. • The ions each achieve a Noble Gas electron configuration = low energy state. • Opposite charges hold ions together.

6. Crystal Lattice • A repeating + and  crystal lattice results. NaCl, sodiumchloride

7. Ionic Compounds • Makes a solid crystal. • Ions align themselves to maximize attractions between opposite charges, and to minimize repulsion between like ions. • Chemical formula is actually the empirical formula, called the “formula unit” ClNaClNaClNaClNaClNa NaClNaClNaClNaClNaCl = “NaCl” ClNaClNaClNaClNaClNa

8. Features of Ionic Compounds • Brittle, hard, crystalline solids at room temperature • High melting points • Metals bonded to non-metals • Elements from opposites sides of the Periodic Table • Dissolve in water to form ions • Conduct an electric current in water

9. Coulomb’s Law • Expresses the NRG of interaction between a pair of ions. E= 2.31 x 10-19 J · nm (Q1Q2) / r • E = energy of interaction between a pair of ions (in Joules) • r = distance (in nm) between ion centers • Q1 andQ2 = charges of the ions • Opposite charges means (–E) • Endo or Exo? What does that mean about NRG in the system?

10. Size of Ions • Ion size increases down a group. • Cations are smaller than the atoms they came from. • Anions are larger. • across a row they get smaller, and then suddenly larger.

11. Periodic Trends • Across the period effective nuclear charge increases so they get smaller. • Energy level changes between anions and cations. N-3 O-2 F-1 B+3 Li+1 C+4 Be+2

12. Size of Isoelectronic ions • Positive ions have more protons so they are smaller. • A stronger + charged nucleus pulls the electrons inward. N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2

13. Lattice Energy • The energy release that occurs when separated gaseous ions are packed together to form an ionic solid • X+x(g) + Y-y(g)  XyYx (s) + energy • Lattice NRG (Eel) = k(Q1Q2)/r • k = constant • Q1 andQ2= charges of ions • r = distance between ion centers

14. Example: • Which has the most exothermic lattice energy, NaCl or KCl? • NaCl … why? • Since both have the same charges (+1 and -1), the distance between the charges needs to be considered. Since Na+ is smaller than K+, the distance between the centers of Na and Cl is less and therefore has a greater lattice NRG.

15. Example 2 and 3 • Arrange the following ionic compounds in order of increasing lattice energy: NaF, CsI, and CaO • CsI < NaF < CaO • Which substance would you expect to have the greatest lattice energy, AgCl, CuO or CrN? • CrN

16. Covalent Bonds • Bond is a force which causes a group of atoms to behave as a single unit. • Electrons are shared by atoms. • Electron orbitals must overlap.

17. The Covalent Bond • The electrons in each atom are attracted to the nucleus of the other. • The electrons repel each other, • The nuclei repel each other. • They reach a distance with the lowest possible energy. • The distance between is the bond length.

18. Energy 0 Internuclear Distance

19. Energy 0 Internuclear Distance

20. Energy 0 Internuclear Distance

21. Energy 0 Internuclear Distance

22. Energy 0 Bond Length (They reach a distance apart with the lowest possible energy.) Internuclear Distance

23. Energy Bond Energy 0 Internuclear Distance

24. Features of Covalent Compounds (aka: “Molecular Compounds”) Electrons are shared, (Single, Double or Triple Bonds are possible.) 2. Non-metals bonded to Non-metals. 3. Includes all diatomic molecules. 4. Relatively low melting/boiling points. 5. No repeating formula, particle is a single unit called a “molecule”.

25. How We Represent Covalent Compounds 1. Molecular Formulas CH4 2. Structural Formulas H H - C - H H Show the bonds between the atoms.

26. Covalent vs. Ionic Bonding • Covalent is sharing, ionic is stealing. • Totally different from each other. • Reality Check:There are many compounds which exhibit both traits! • These are called polar covalent bonds. • The electrons are shared, but shared unequally.

27. Polar Covalent Bonds • The electrons are shared, but they are not shared evenly. • One atom has the pair more often than the other. • A “polar” molecule results: One end is slightly positive, while the other is slightly negative. • A partial charge is called a “dipole”.

28. d+ d- H - F Example:

29. d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d-

30. d+ H - F d+ H - F d- d- d- H - F d+ d+ d- H - F H - F d+ d- d- H - F d+ d+ d+ H - F H - F d- d- - +

31. d+ d- d+ d- H - F H - F d+ d- H - F d+ d- d+ d- H - F H - F d+ d- d+ d- d+ d- H - F H - F H - F - +

32. How Do We Know Which Atom Has the Electron Pair More Often? • Answer: Electronegativity Values • Electronegativity - The ability of an atom to attract shared electrons to itself. • The more electronegative an atom, the more often it has the shared pair. • Greater electronegativity =  pole

33. Electronegativity • E.N. values are assigned for almost every element (Figure 8.6, p. 285) • Gives us relative electronegativities of all elements. • Tends to increase left to right and decreases as you go down a group. • Noble gases aren’t discussed. • Difference in electronegativity between atoms tells us how polar.

34. Helpful Number Line: • Determine the Electronegativity Value difference between two atoms Non-polar covalent Polar covalent ionic 0 0.4 2.0 3.4

35. Polar Covalent Ionic Electronegativity difference Bond Type Zero Covalent Covalent Character decreases Ionic Character increases Intermediate Large

36. d+ d- H - F Polar Covalent Bond-How it is drawn

37. Reminder on Writing Formulas and Nomenclature Ionic Compounds: • Name the cation then name the anion • When writing formulas, add subscripts to make sure that the charges balance Covalent Compounds: • When naming, use prefixes for the subscripts, the 2nd atom will end in –ide • Write formulas, assign subscripts based on the prefixes

38. Lewis Dot Structures are Models to represent both ionic and covalent What is a Model? • Explains how nature operates. • Derived from observations. • It simplifies and categorizes the information. • A model must be sensible, but it has limitations.

39. Properties of a Model • A human invention, not a blown up picture of nature. • Models can be wrong, because they are based on speculations and oversimplification. • You must understand the assumptions in the model, and look for weaknesses. • We learn more when the model is wrong than when it is right.

40. Lewis Structures • Show how the valence electrons are arranged. • One dot for each valence electron. • A stable compound has all its atoms with a noble gas configuration. • Hydrogen follows the duet rule. • The rest follow the octet rule. • Bonding pair is the one between the symbols.

41. Electron Dot Placement 6 2 X 3 5 7 1 4 8 The X represents the symbol for the element. The dots are placed around the symbol in the order shown above.

42. Rules • Sum the # of ALL the valence electrons. • Determine the central atom. The least electronegative element is central(H never central, C nearly always central) • Write symbols for the atoms to show which atoms are attached and connect them with a single bond (a dash) • Complete the octets of the atoms attached to the central atom (except for H, follows a duet rule). • Place any leftover electrons on the central atom. Not enough electrons - consider a double or triple bond.

43. A useful equation: ( happy - have )  2 = # bonds (what they want - what they have)  2 = # bonds H2O (12 - 8)  2 = 2 bonds   H  O  H  

44. Practice Structures NH3 C2H4 BrO3-1 O2 ClO2-1 OH-1 PO4-3 • PCl3 • ICl • CH4 • CH2Cl2 • HCN • NO+ • CO2 • H2

45. Partial Ionic Character • There are probably no totally ionic bonds between individual atoms.

46. 75% % Ionic Character 50% 25% Electronegativity difference

47. How do we deal with it? • If bonds can’t be ionic, what are ionic compounds? • An ionic compound will be defined as any substance that conducts electricity when melted. • Also use the generic term salt. • As it turns out, most compounds fall somewhere between ionic and covalent.

48. The bond is a human invention. • It is a method of explaining the energy change associated with forming molecules. • Bonds don’t exist in nature, but are useful. • We have a model of a bond. H - Cl

49. Exceptions to the octet • Less than an Octet BH3 Be and B often do not achieve octet, but form highly reactive compounds • More than an Octet SF6 and I3- Third row and larger elements can exceed the octet. How? Use 3d orbitals • Odd Number of Electrons • NO, NO2, and ClO2

50. Exceptions to the octet • When we must exceed the octet, extra electrons go on central atom. ClF3 XeO3 ICl4- BeCl2