Academic Chemistry Mrs. Teates Newport High School

1. Academic ChemistryMrs. TeatesNewport High School Chapter 6 – Chemical Bonding

2. Lesson 1 – Introduction to Chemical Bonding • Lesson Essential Questions: • Why do atoms form chemical bonds? • How is the type of chemical bond determined? Vocabulary: chemical bond, ionic bonding, covalent bonding, nonpolar-covalent bonding, polar, polar-covalent bonding

3. Vocabulary • Chemical Bond • attractive force between atoms or ions that binds them together as a unit • bonds form in order to… • decrease potential energy (PE) • increase stability

4. Vocabulary ION 2 or more atoms 1 atom Monatomic Ion Polyatomic Ion Na+ NO3-

5. Types of Bonds COVALENT IONIC e- are transferred from metal to nonmetal e- are shared between two nonmetals Bond Formation Type of Structure true molecules crystal lattice Physical State Solid, liquid, or gas solid Melting Point low high Solubility in Water yes usually not yes (solution or liquid) Electrical Conductivity no Other Properties odorous

6. Types of Bonds METALLIC e- are delocalized among metal atoms Bond Formation Type of Structure “electron sea” Physical State solid Melting Point very high Solubility in Water no yes (any form) Electrical Conductivity malleable, ductile, lustrous Other Properties

7. Types of Bonds Ionic Bonding - Crystal Lattice RETURN

8. Types of Bonds Covalent Bonding - True Molecules Diatomic Molecule RETURN

9. Types of Bonds Metallic Bonding - “Electron Sea” RETURN

10. Bond Polarity • Most bonds are a blend of ionic and covalent characteristics. • Difference in electronegativity determines bond type.

11. Bond Polarity • Electronegativity • Attraction an atom has for a shared pair of electrons. • higher e-neg atom  - • lower e-neg atom +

12. Bond Polarity • Electronegativity Trend • Increases up and to the right.

13. Bond Polarity • Nonpolar Covalent Bond • e- are shared equally • symmetrical e- density • usually identical atoms

14. - + Bond Polarity • Polar Covalent Bond • e- are shared unequally • asymmetrical e- density • results in partial charges (dipole)

15. Bond Polarity • Nonpolar • Polar • Ionic View Bonding Animations.

16. Bond Polarity Examples: • Cl2 • HCl • NaCl 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic

17. More Bond Polarity Practice • What type of bonding would be expected between the following atoms? • Li and Cl • Ca and Ga • I and Cl • K and Na

18. Lesson 2 – Covalent Bonding and Molecular Compounds • Lesson Essential Questions: • How is a molecular compound formed? • What are some of the characteristics of a covalent bond? Vocabulary: molecule, chemical formula, molecular formula, bond energy, electron-dot, Lewis structure, structural formula, single bond, multiple bonds, resonance

19. Vocabulary • Covalent bond – bond that is created by the sharing of electrons • Molecule – neutral group of atoms held together by covalent bonds • Molecular compound – chemical compound made of molecules

20. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

21. Energy of Bond Formation • Potential Energy • based on position of an object • low PE = high stability

22. no interaction increased attraction Energy of Bond Formation • Potential Energy Diagram attraction vs. repulsion

23. increased repulsion balanced attraction & repulsion Energy of Bond Formation • Potential Energy Diagram attraction vs. repulsion

24. Bond Energy Bond Length Energy of Bond Formation • Bond Energy • Energy required to break a bond

25. Energy of Bond Formation • Bond Energy • Short bond = high bond energy

26. Lewis Structures • Electron Dot Diagrams • Pick the central atom • Count the valence electrons (they are what electron dot diagrams show) • Place electrons around the atom

27. Ne Lewis Structures • Octet Rule • Most atoms form bonds in order to obtain 8 valence e- • Full energy level stability ~ Noble Gases

28. - + + Lewis Structures • Nonpolar Covalent - no charges • Polar Covalent - partial charges

29. Practice Drawing Lewis Structures • On page 186 in your text book do practice problems #1-4 • Draw the Lewis structure of ammonia, NH3 • Draw the Lewis structure for hydrogen sulfide, H2S • Draw the Lewis structure for silane, SiH4 • Draw the Lewis structure for phosphorus trifluoride, PF3

30. Multiple Covalent Bonds • Some elements can share more than one electron pair. • Double bond (two pairs of electrons are shared) • Triple bond (three pairs of electrons are shared)

31. Practice of Lewis Structures for multiple bonds • Draw Lewis structures for each of the following molecules: • O2 • CO2 • N3 • N2

32. Resonance • Occurs when more than one valid Lewis structure can be written for a particular molecule (due to position of double bond) • These are resonance structures of benzene. • The actual structure is an average (or hybrid) of these structures.

33. Resonance in Ozone • Note the different location of the double bond • Neitherstructure is correct, it is actually a hybrid of the two. To show it, draw all varieties possible, and join them with a double-headed arrow.

34. Polyatomic ions – note the different positions of the double bond. • Resonance in a carbonate ion (CO32-): • Resonance in an acetate ion (C2H3O21-):

35. Molecular Nomenclature • Prefix System (binary compounds) 1. Less e-neg atom comes first. 2. Add prefixes to indicate # of atoms. Omit mono- prefix on first element. 3. Change the ending of the second element to -ide.

36. Molecular Nomenclature PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER 1 2 3 4 5 6 7 8 9 10

37. Molecular Nomenclature • CCl4 • N2O • SF6 • carbon tetrachloride • dinitrogen monoxide • sulfur hexafluoride

38. Molecular Nomenclature • arsenic trichloride • dinitrogen pentoxide • tetraphosphorus decoxide • AsCl3 • N2O5 • P4O10

39. Lesson 3 – Ionic Bonding and Ionic Compounds • Lesson Essential Questions: • How is an ionic bond formed? • What are some of the characteristics of an ionic bond? Vocabulary: ionic compound, formula unit, lattice energy, polyatomic ion

40. Vocabulary: • Ionic compound – composed of positive and negative ions that are combined so that the charges are equal.

41. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

42. Forming Ionic Compounds • Electron dot notation is used to note changes. • Form to create an atmosphere of stability

43. Lewis Structures and Ionic Compounds • Covalent – show sharing of e- • Ionic – show transfer of e-

44. Characteristics of Ionic Bonding • Ions minimize potential energy in crystals by forming a crystal lattice. • Distance between all ions represent a balance of attraction between oppositely charged particles and repulsion between like charged particles

45. Energy of Bond Formation • Lattice Energy • Energy released when one mole of an ionic crystalline compound is formed from gaseous ions

46. Ionic vs. Covalent Ionic • High melting temperature • High boiling point • Hard • Brittle, because slight shift of crystal can cause it to break • Conduct electricity when dissolved in water Covalent • Low melting temperature • Low boiling point • Do not conduct electricity • Not as brittle

47. Ionic Nomenclature Ionic Formulas • Write each ion, cation first. Don’t show charges in the final formula. • Overall charge must equal zero. • If charges cancel, just write symbols. • If not, use subscripts to balance charges. • Use parentheses to show more than one polyatomic ion. • Stock System - Roman numerals indicate the ion’s charge.

48. Ionic Nomenclature Ionic Names • Write the names of both ions, cation first. • Change ending of monatomic ions to -ide. • Polyatomic ions have special names. • Stock System - Use Romannumerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

49. Ionic Nomenclature • Consider the following: • Does it contain a polyatomic ion? • -ide, 2 elements  no • -ate, -ite, 3+ elements  yes • Does it contain a Roman numeral? • Check the table for metals not in Groups 1 or 2. • No prefixes!

50. Ionic Nomenclature Common Ion Charges 1+ 0 2+ 3+ NA 3- 2- 1-