Chemistry Unit 2

1. Chemistry Unit 2 Chemical Bonding

3. e- e- e- e- e- e- 11 p+ 12 n e- e- e- e- Atomic StructureREVIEW • An atom of sodium has 11 protons • Atomic # 11 • Found in nucleus • It has 11 electrons • Same as # of protons • Atoms are neutral • Found outside nucleus • It has 12 neutrons • At. mass – At. # • Only protons and neutrons have mass • Found in nucleus 11 Na Sodium 22.98977 e-

4. Lewis Dot Structures REVIEW • Lewis dot structures – atomic models that show only the valence electrons for an atom. • Only the valence electrons are involved in bonding, therefore Lewis Structures are good models to use when showing how atoms bond to form compounds.

5. e- e- e- e- e- e- 11 p+ 12 n e- e- e- e- Lewis Dot Structures REVIEW • Sodium e- Na

6. 1 2 8 3 Ne 5 6 7 4 Lewis Dot Structures REVIEW • Dots (electrons) should be added to the element’s symbol in this order.

7. Lewis Dot Structures REVIEW The “A” groups tell you the number of electrons in the valence shell. Li Be B C N O F Ne

8. Why Do Atoms Bond? • To become more stable • like the noble gases. • Octet Rule – atoms tend to gain, lose or share electrons in order to acquire a full shell of valence electrons. (usually 8 – except H and He who only need 2!!)

9. What Does This Picture Tell You?

10. Three Main Types of Bonds • Ionic Bond – Atoms transfer electrons to fill their valence shells, oppositely charged ions are formed, opposites attract. • Occurs between a metal and a nonmetal • Covalent Bond – Atoms share electrons to fill their valence shells. • Occurs between nonmetals • Metallic Bonds – Atoms share a “sea of electrons.” • Occurs between atoms of a metal

12. K F IONIC In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds.

13. K F

14. K F

15. K F

16. K F

17. K F

18. K F

19. _ + K F

20. _ + ANION CATION K F • An atom that loses electrons gets a positive charge and is a cation • (Think of the t in cation as a plus sign) • An atom that gains electrons gets a negative charge and is an anion

21. Ionic Bonding • Ion – a charged particle • A neutral atom becomes an ion when it loses or gains an electron. • If an atom loses an electron, it becomes a (+) ion called a cation. • If an atom gains an electron, it becomes a (-) ion called an anion.

22. Anions & Cations • Based on the fact that most of our atoms want 8, determine the ions that each of the following atoms will form, then say whether it is a cation or an anion Na S O Mg Li Ne Si F C He N Cl Have me check your answers when you’re finished. If I’m busy, move on and I will check later.

23. Ionic Bonding • Example Na Cl To become more stable, sodium must loose one electron To become more stable, chlorine must gain one electron

24. Cl Na Ionic Bonding • Example Sodium loses an electron and becomes an Na+1 ion. Chlorine gains an electron and becomes a Cl-1 ion. Opposites attract, and an ionic compound is formed… NaCl

25. Try Another Example Br Al Aluminum will become more stable if it gets rid of three electrons. Bromine will become more stable if it receives one electron. Are both atoms more stable as a result of this transfer? No, Al must donate two more… where?

26. Br Br Br Al Aluminum & Bromine Now, each atom has a full valence shell… all are more stable.

27. Br Br Br Al Aluminum and Bromine Aluminum donated 3 e-, so it becomes Al+3 Each bromine accepted 1 e-, so they each become Br-1 The compound that forms is AlBr3

28. Let’s Wrap it Up • Ionic bonds are held together by electrostatic forces. (Opposite charges attract.) • The result of an ionic bond is called an ionic compound. • Ionic bonds form between a metal and a nonmetal atom due to large differences in electronegativity. (1.7 or greater) • Electronegativity is the ability of an atom to pull shared electrons to itself. (Think of it like the strength of an atom in a tug-of-war for bonded electrons.) • The nonmetal’s EN is so much greater than the metal’s, that it removes the electrons, forming oppositely charged ions.

29. For Example: Na and O EN of Na = 0.9 EN of O = 3.5

30. Why does Sodium and Oxygen form an ionic bond? 3.5 EN of O - 0.9 EN of Na 2.4 Difference in EN • Difference in electronegativity is 2.4(>1.7) • An ionic bond will form. • Oxygen has a greater electronegativity, and is able to yank electrons away from sodium.

31. How will two chlorine atoms react? Cl Cl COVALENT

32. Cl Cl Each chlorine atom wants to gain one electron to achieve an octet

33. Cl Cl Neither atom will give up an electron What’s the solution – what can they do to achieve an octet?

34. Cl Cl

35. Cl Cl

36. Cl Cl

37. Cl Cl

38. Cl Cl octet

39. Cl Cl octet

40. Cl Cl The octet is achieved by each atom sharing the electron pair in the middle

41. Cl Cl

42. Cl Cl This is the bonding pair

43. Cl Cl It is a single bonding pair

44. Cl Cl It is called aSINGLE BOND

45. Cl Cl Single bonds are abbreviated with a dash

46. Covalent Bonding O O Each atom of Oxygen needs two more electrons to become more stable, so they will share two pairs of electrons. A diatomic molecule of oxygen is formed. O2

47. H H O Try another example Each atom of hydrogen needs one more electron to become more stable. Oxygen needs two electrons to become more stable. All atoms become more stable (have full valence shells). A molecule of water is made. H2O

48. Let’s Wrap it Up… Again! • Covalent bonds are held together by a mutual attraction for the shared electrons (electronegativity). • Covalent bonding occurs when a sharing of electrons results in an overlap of valence orbitals. Each electron is attracted to the positive charge of the opposite nucleus. • The result of a covalent bond is called a molecule.

49. Bonding in Metals • Properties of metals • malleable -ductile • conduct heat -conduct electricity Why? e- e- 2+ 2+ 2+ e- 2+ e- e- e- e- e- e- 2+ 2+ 2+ 2+ e- e- e- e- e- e- e-

50. Bonding in Metals • Why? • The valence electrons of metals are held loosely. • In metallic bonding metals don’t lose electrons. • Metal atoms release valence electrons in a sea of electrons shared by all metal atoms. • Valence electrons are free to move.