1. Chemical Bonding

2. What is a chemical bond? “attachment” of an atom to another atom

3. How does this happen? Chemical bonds are formed when valence electrons are: 1.Transferred from one atom to another 2.Shared between atoms 3.Mobile within a metal

4. Remember stability? Everything wants to be more stable so they want to be at the lowest energy possible Sharing or transferring electrons results in a more stable system…THIS IS BONDING

5. Making Bonds Releases energy EXOTHERMIC

6. Breaking bonds Requires Energy ENDOTHERMIC

7. Don’t forget about electronegativity The ability of an atom to attract electrons when the atom is in a compound F: 4.0 Fr: 0.7 Noble gases- no assigned value (don’t actively search for electrons)

8. EN values come in handy They can be used to determine the type of bond that forms between two atoms Bigger difference in EN  ionic bonding (>1.7)

9. Three Types of Chemical Bonds Ionic Covalent metallic

10. IONIC Bonds Metals- lose electrons  + charge Nonmetals- gain electrons  - charge OPPOSITES ATTRACT Force of attraction that holds ions together is called an ionic bond

11. Ionic Compounds (aka salts) Electrically neutral Additional atoms may be necessary to ensure that all the ions formed have a stable electron configuration Charge of 0!!! Mg 2+ Cl - MgCl 2

12. Which species does not have a noble gas configuration? 1. Na + 2. Mg 2+ 3. Ar 4. S

13. Ionic Compounds Most are crystalline solids at room temperature, which leads to many strong attractions that hold compound together Ionic compounds generally have high melting points

14. Ionic Compounds Dissolving or heating destroys the crystalline structures, ions become free to move. They then conduct electricity well (ELECTROLYTES)

15. Which electron configuration is correct for a sodium ion? 1. 2-7 2. 2-8 3. 2-8-1 4. 2-8-2

16. Covalent Bonding < 1.7 EN difference When atoms are held together by a “tug of war” or sharing electrons Form molecules and molecular compounds

17. Molecules A neutral group of atoms joined together by covalent bonds A compound composed of molecules is called a molecular compound

18. This means That there is no such thing as a molecule of ionic compounds-- they exist as a collection of positively charged ions arranged in repeating 3-D patterns FORMULA UNITS

19. Covalent Compounds Molecules All atoms have filled electron shells “diatomic” molecules: elements -gen or -ine Network Solids Carbon forms continuous bonding pattern (graphite and diamond) Melting and Boiling Points Variation due to polarity and bonds between molecules Nonconductors No free electrons

20. Network Solids Solids in which all of the atoms are covalently bonded to each other Melting a network solid would require breaking covalent bonds throughout the solid Diamond- vaporizes to a gas at 3500+ degrees Celsius

21. Multiple bonds Single bonds- a bond formed when two atoms share a pair of electrons Double bonds- a bond in which 2 atoms share 2 pairs of electrons Triple bonds- 3 pairs of electrons shared by two atoms MULTIPLE BONDS ARE SHORTER, STRONGER AND MORE STABLE

22. The Nature of Covalent Bonding Octet Rule (duet) In forming covalent bonds, electron sharing usually occurs so that atoms attain the electron configuration of noble gases H..HH..H

23. Single Covalent Bonds Two atoms held together by sharing a pair of electrons An electron dot structure such as H..H represents the shared pair of electrons of the covalent bond by two dots A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms H-H

24. Examples 2H. + O --> O H or O-H H H

25. Carbon When carbon forms bonds with other atoms, it usually forms four bonds. You would not predict this based on carbon’s electron configuration

26. So…what happens? One electron gets promoted from 2s to 2p (requires a small amount of energy) CH 4 is much more stable than CH 2

27. Practice: Electron Dot Structures CH 4 H 2 O 2 PCl 3

28. Unshared pairs A pair of valence electrons that is not shared between atoms is called an unshared pair, also knows as a lone pair or nonbonding pair

29. Double and Triple Covalent Bonds Oxygen and it’s family (grp 16)- double Nitrogen and it’s family (grp 15)- triple Carbon can form double and triple bonds with itself, grp 16 and grp 15 elements- carbon dioxide

30. Coordinate Covalent Bonds Covalent bonds in which one atom contributes both bonding electrons In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms

31. Coordinate Covalent Bonds C + O --> C O --> C O Oxygen has a stable electron configuration but the carbon does not The oxygen can donate one of its unshared electrons for bonding

32. In a structural formula… You can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them C O Once formed, these bonds are like any other covalent bond

33. Ammonium Ion NH 4 H H H H + + N H H N H or H N H H H H +

34. Which of the following solids has the highest melting point? 1. H 2 O (s) 2. Na 2 O (s) 3. SO 2(s) 4. CO 2(s)

35. Electron sharing and Molecules The bonding pairs of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons When the atoms pull equally  nonpolar covalent bond Diatomic molecules (H 2, N 2, O 2, Cl 2 )

36. Polar Bond A covalent bond between atoms in which the electrons are shared unequally The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge

37. Polar Bond- Example H- 2.1 Cl- 3.0

38. Water- A Polar Molecule!!!!!!!!

39. Polar Molecules One end of the molecule is slightly negative and the other is slightly positive A molecule that has two poles is called a dipolar molecule, or DIPOLE ASYMMETRICAL

40. Polarity of molecules The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds

41. Molecular Shape and Polarity VSEPR- valence shell electron pair repulsion theory Linear- CO 2 Angular/Bent- H 2 O Pyramidal- NH 3 Tetrahedral- CH 2 Cl 2 SYMMETRY

42. Shapes

43. Ionic vs. Covalent Formula unit Transfer of electrons Metal and nonmetal Solid High M.P. High solubility in H 2 O Good conductivity Molecule Sharing electrons Nonmetals Solid, liquid and gas Low M.P. Low to high solubility in H 2 O Poor to nonconductors

44. Exceptions to the Octet Rule The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons

45. Other exceptions PCl 5 SF 6

46. Attractions Between Molecules Intermolecular attractions are weaker than either ionic or covalent bonds

47. Vander Waals Forces The two weakest forces Dipole interactions Dispersion forces

48. Dispersion Forces WEAKEST- caused by the motion of electrons, occur even between non-polar molecules Electrical forces influence neighboring molecules

49. Dipole Interactions Occur when polar molecules are attracted to one another The attraction involved occurs between the oppositely charged regions of polar molecules Similar to, but much weaker than ionic bonds

50. Hydrogen Bonds Attractive forces in which a hydrogen covalently bonded to a very electronegative atom (F, O, N) is also weakly bonded to an unshared pair of another electronegative atom

51. Ion-Molecule Attraction ION- charged particle MOLECULE- covalently bonded compound

53. The stronger the intermolecular forces, the higher the M.P. and B.P. 1. Ionic solid 2. Molecules with H bonds 3. Polar molecules 4. Nonpolar molecules The greater the mass, the greater the force of attraction

54. Testing of an unknown solid shows that it has the following properties: 1. Low M.P. 2. Nearly insoluble in water 3. Nonconductor of electricity 4. Relatively soft solid State the type of bond expected Explain in terms of attractions between particles why it has a low M.P. Explain why particles are nonconductors of electricity

55. The strength of an atom’s attraction for the electrons in a chemical bond is the atom’s 1. Electronegativity 2. Ionization energy 3. Heat of reaction 4. Heat of formation

56. Which of these formulas contains the most polar bond? 1. H-Br 2. H-Cl 3. H-F 4. H-I

57. Each of the molecules listed below is formed from sharing electrons between atoms when the atoms within the molecule are bonded together Molecule A- Cl 2 Molecule B- CCl 4 Molecule C- NH 3 1. Draw a Lewis structure for NH 3 2. Explain why CCl 4 is classified as a nonpolar molecule 3. Explain why NH 3 has stronger intermolecular forces of attraction than Cl 2 4. Explain how the bonding in KCl is different from the bonding in molecules A, B and C

58. METALLIC Bonding Metals  want to get rid of valence electrons! closely packed cations The “orphaned” valence electrons can be modeled as a sea of electrons They are mobile and can drift freely from one part of the metal to another

59. Metallic Bonds attraction of the free floating valence electrons for the positively charged metal ions “Take me back!”

60. The “sea of electrons” Explains many properties of metals Good conductors of heat and electrical current Ductile and malleable- when subjected to pressure, cations easily slide past one another

61. Metal atoms Are arranged in very compact and orderly patterns Simple crystalline patterns Body-centered cubic Face-centered cubic Hexagonal close-packed

62. The high electrical conductivity of metals is primarily due to 1. High ionization energy 2. Filled energy levels 3. Mobile electrons 4. High electronegativities