Chapter 5

1. Chapter 5 Chemical Bonding and Nomenclature

2. 5.1 Ionic Bonding Chemical bonds come from electrons moving between atoms • Metals tend to lose electrons. • Nonmetals tend to gain electrons. • When a sodium atom (group IA) reacts with a chlorine atom (group VIIA), each Na loses an electron and each Cl gains an electron. Positive and negative ions attract. This is an ionic molecule. Its held by an ionic bond.

3. 5.1 Ionic Bonding (Continued) • Ionic compounds • Formed by attraction of ions with opposing charges • The ions stack positive next to negative. • Ions are very strongly stuck together. • This is called electrostatic force.

4. 5.1 Ionic Bonding (Continued) • Electrostatic force • Governed by Coulomb’s law • Strength of the attraction is: The smaller the distance and larger the charges, the stronger the force if charges are opposite.

5. 5.1 Ionic Bonding (Continued) • Electronegativity A scale from 0 to 4 that shows how much attraction a given element has for electrons. Corner rules apply: * F is highest, Fr is lowest * Metals have low electronegativity and easily lose electrons to become positive * Non-metals have high electronegativity and easily take electrons to become negative.

6. 5.1 Ionic Bonding (Continued) • Ionic bond • Chemical bond occurring between ions of a metal and a nob-metal with differences in electronegativity > 1.7 • Forms extended three-dimensional lattices due to the ordering of ions

7. 5.6 Equal Versus Unequal Sharing of Electrons—Electronegativity and the Polar Covalent Bond (Continued) • Electronegativity and bonding types • Two atoms with identical electronegativity results in pure covalent bonding • Differing electronegativity leads to a polarcovalent bond, unequal electron sharing • - shows a partial negative charge

8. 5.6 Equal Versus Unequal Sharing of Electrons—Electronegativity and the Polar Covalent Bond (Continued) • Electronegativity difference (EN) • The larger the EN, the greater the polarity of a bond.

9. 5.1 Ionic Bonding (Continued) • Lattice types • Crystalline—material with a long-range, highly ordered lattice • Amorphous—no overall order in the extended structure

10. 5.1 Ionic Bonding (Continued) • Properties of ionic bonds • Very strong bonding, yielding high melting points

11. 5.1 Ionic Bonding (Continued) • Cleaving • Breaking of a crystal along perfectly straight planes

12. 5.2 Molecules—What Are They? Why Are They? (Continued) • Molecule types • Molecular compound • Comprised of two or more types of elements • Examples would be CH4 and H2O. • Elemental substance • Comprised of only one type of element • Examples include O2, Cl2, and I2.

13. 5.2 Molecules—What Are They? Why Are They? (Continued) • Molecular properties are not an average of the individual atoms properties which makeup the compound • Rather, each molecule’s properties are unique and almost unrelated to the properties of its elements.

14. 5.3 Holding Molecules Together—The Covalent Bond • Covalent bond • Bonding that occurs between electrons in one atom to the nucleus of another atom, mutually sharing electrons. Happens when the difference in electronegativity is < 1.7. • Results in electrons orbiting around both nuclei, holding atoms together .

15. 5.3 Holding Molecules Together—The Covalent Bond (Continued) • Covalent bonding

16. 5.3 Holding Molecules Together—The Covalent Bond (Continued) • Representations of a covalent bond • The line between two atoms denotes a covalent bond.

17. 5.3 Holding Molecules Together—The Covalent Bond (Continued) • Energy in bonds • Bond formation ALWAYS releases energy. • Bond breaking ALWAYS takes energy.

18. 5.3 Holding Molecules Together—The Covalent Bond (Continued) • Counting electrons • Chemists count shared electrons as being in both atoms.

19. 5.4 Molecules, Dot Structures, and the Octet Rule • Covalent compounds follow the octet rule. • Share electrons to reach eight valence electrons.

20. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • Predicting covalent compounds • Decide the number of valence electrons in a neutral isolated atom. • Decide how many electrons are needed to fill the octet. • Each covalent bond adds one electron to the valence count.

21. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • Consider a C atom. • Contains four valence electrons in neutral atom • Must form four covalent bondsto reach octet

22. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued)

23. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) Predict the compounds that form between: (a) Carbon and chlorine (b) Phosphorus and hydrogen

24. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) Predict the compounds that form between: • Carbon and chlorine The carbon needs to form four bonds and each chlorine forms one bond, making the formula CCl4. (b) Phosphorus and hydrogen The phosphorus has five valence electrons, so it will form three more bonds while the hydrogen forms one, each making the formula PH3.

25. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • Lewis dot diagrams (dot diagrams) • Two parts to the diagram: • Contains the element symbol • Number of dots corresponding to valence electrons

26. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • Types of electrons • Electron pairs (paired electrons) • When two electrons are present on one side • Unpaired electrons • When only one electron is on a side

27. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • Octet rule and Lewis dot structures • Elemental fluorine has seven dots around the F. • In diatomic fluorine, each atom shares one electron to become octets.

28. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) Why do two He atoms refuse to form the molecule He2 when brought next to one another?

29. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) Why do two He atoms refuse to form the molecule He2 when brought next to one another? (a) He does not have any unpaired electrons to match up. (b) He already has the correct number of valence electrons according to the octet rule.

30. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • Drawing molecules using electron dot structures

31. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • NF3 redrawn with lines instead of dots for bonding pairs

32. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • How many valence electrons are in NF3? • How many bonding pairs? • How many non-bonding pairs?

33. 5.4 Molecules, Dot Structures, and the Octet Rule (Continued) • How many valence electrons are in NF3? 26 • How many bonding pairs? Three pairs • How many non-bonding pairs? Ten pairs

34. 5.5 Multiple Bonds • Multiple bonds • When two atoms share more than one pair of electrons with each other

36. 5.5 Multiple Bonds A pair of electrons is sometimes represented as a line Oxygen molecules have double bonds O=O

37. 5.5 Multiple Bonds (Continued) • Covalent bond types • Single bond • Sharing of a single bonded pair of electrons • Double bond • Sharing two pairs of electrons • Triple bond • Sharing three pairs of electrons

38. 5.5 Multiple Bonds (Continued) Draw the dot diagram for N2

39. 5.5 Multiple Bonds (Continued) • Steps to draw a dot diagram: • Determine total number of valence electrons. • Connect atoms with a single bond. • Place remaining electrons as lone pairs. • First on terminal atoms, then the central atom • If each atom does not have an octet, change lone pairs to bonding pairs.

40. Carbon dioxide

41. Carbon monoxide

42. 5.5 Multiple Bonds (Continued) • Resonance forms • Valid dot diagrams that differ only by electron placement • Chemists consider each to be correct. • Combination of each possibility

43. 5.5 Multiple Bonds (Continued) • Polyatomic ions • Ions made up of multiple atoms that tend to stay together as a group. • Examples NH4+ would have one less electron due to the positive one charge. NO3- would have one extra electron due to the negative one charge.

44. 5.7 Nomenclature—Naming Chemical Compounds • Common chemical names: H2O—water NH3—ammonia Hg2Cl2–calomel • As the number of compounds increased, a naming system was needed.

45. 5.7 Nomenclature—Naming Chemical Compounds (Continued) • Binary compound naming • Compound made of only two elements • Binary ionic compounds—comprised of a metal and a nonmetal • Binary covalent compounds—comprised of two nonmetals or metalloids

46. 5.7 Nomenclature—Naming Chemical Compounds (Continued) • Naming binary ionic compounds • Metal is considered the cation. • Written as element name • Nonmetal is the anion. • Add suffix –ide to the elemental name.

47. 5.7 Nomenclature—Naming Chemical Compounds (Continued) • Formula from ionic name • Consider the groups of the elements • Charges must balance in final formula • Magnesium chloride • Mg2+ and Cl- • So to balance charges, it must be 1 Mg2+ and 2 Cl-. • Giving final formula of MgCl2

48. 5.7 Nomenclature—Naming Chemical Compounds (Continued) • Transition metals • Can carry multiple charges • Use roman numeral to show charge • Iron (II) chloride is Fe2+ and Cl-. • Iron (III) chloride is Fe3+ and Cl-.

49. 5.7 Nomenclature—Naming Chemical Compounds (Continued) Name the compound TiCl4.

50. 5.7 Nomenclature—Naming Chemical Compounds (Continued) Name the compound TiCl4. Cl is in group VIIA, so it must gain one electron to become Cl- anion. Ti must be +4 to balance the 4 Cl- ions. Therefore, the name is titanium (IV) chloride.