2. Chapter 6 Chemical Bonds Section 1 Introduction to Chemical Bonding

3. Objectives Define chemical bond. Explain why most atoms form chemical bonds. Describe ionic and covalent bonding. Explain why most chemical bonding is neither purely ionic or covalent. Classify bonding type according to electronegativity differences.

4. Atomic Stability Q. Why do atoms form compounds? Search for outer shell stability by losing or gaining electrons! Nobel Gases  8 electrons in os = stable Ex: Na loses 1 e- to Cl Chemical Bond  attractive force that holds atoms together in a compound; when atoms gain, lose, or share e-, an attractive force pulls them together to form a compound

6. Oxidation Number  the number that tells you how many electrons an atom has gained, lost or shared to become stable

7. Ion Def. Charged particle that has either gained or loses electrons RESULT: Now has either more or fewer number of electrons than protons Ex: sodium fluoride NaF active ingredient in toothpaste  Na loses 1e- to F; Na is +1; F is -1 Ex: potassium iodide KI – ingredient in iodized salt  K loses 1 e -to I; K is +1; I is -1 Cation = positive ion #P > #E Anion = negative ion #P < #E

8. Ionic Bond Def. Force of attraction between the opposite charges of the ions in an ionic compound (cations and anions) Transfer of electrons btwn metals & nonmetals Ex: magnesium chloride MgCl 2  Mg loses 2 electrons to each Cl; Mg is +2; each Cl is -1 RESULT  neutral compound = sum of the charges of the ions equals 0 2(Mg) + -1(Cl) + -1(Cl) = 0

9. Covalent Bond Def. Attraction that forms between atoms when they share electrons Atoms will be more stable by sharing e- rather than losing or gaining e-

10. Unequal Sharing Electrons don’t always share equally between atoms in a covalent bond Strength of attraction of atom to electrons due to : 1. size of atom 2. size of the positive charge in the nucleus (a strong magnet will hold a metal better than a weak magnet) 3. total # of electrons 4. how far are the electrons from the nucleus being shared (a magnet has a stronger pull to a metal when it is next to it rather than a couple inches away)

11. Ex: HCl hydrochloric acid used to clean metal and found in your stomach to digest food Cl – atoms have a stronger attraction for electrons than H atoms  electrons shared will spend most time near the chlorine atom RESULT  Cl atom has a partial negative charge (Greek delta)  H atom has a partial positive charge VISUAL: Tug-of-war  stronger team pulls the rope towards them

12. Polar or Nonpolar Polar molecule  molecule that has a slightly positive end and a slightly negative end although the overall molecule is neutral ex: water Nonpolar molecule  molecule in which electrons are shared equal in bonds; doesn’t have oppositely charged ends; found in 2 identical atoms or molecules that are symmetric ex: CCl 4

15. Ionic or Covalent??? Bonding between atoms of different elements is rarely purely ionic or purely covalent. Falls somewhere between 2 extremes depending on electronegativity  measure of atom’s ability to attract e-

17. Electronegativity Values

19. If you still need help understand polarity and electronegativity, watch this video: http://www.bing.com/videos/search?q=Examples+Polar+and+Nonpolar+Covalent+ Bonds+With&Form=VQFRVP#view=detail&mid=295CAC087126B74D1565295CA C087126B74D1565

20. HOMEWORK Section Review pg 177 #1-5

22. Section 2 Covalent Bonding Molecular Compounds molecule  neutral group of atoms that are held together by covalent bonds Red = O, White = H, Black = C Molecular compound  a chemical compound whose simplest units are molecules

23. chemical formula  indicates what elements atoms and numbers of atoms in a chemical compound by using atomic symbols and numerical subscripts molecular formula  shows types and number of atoms combined in a single molecule of a molecular compound

24. Formation of a Covalent Bond Nature favors chemical bonding  most atoms have lower potential energy when bonded to other atoms than as independent atoms. a.separated H atoms do not affect each other b.PE decreases as atoms are drawn together by attractive forces c.PE minimum when attractive forces are balanced by repulsion forces = ideal distance d.PE increases when repulsion btwn like charges outweighs attraction between opposite charges

25. Characteristic of the Covalent Bond Bond length  average distance between 2 bonded atoms H-H 75 pm Form CB= H atoms release energy= amt of energy equals drop in PE Bond energy  energy required to break a chemical bond and form neutral isolated atoms (kJ/mol)

27. Octet Rule Chemical compounds tend to form so that each atom has an octet of e- in outer energy level by gaining, losing or sharing e- Draw Fluorine electron configurations:

28. Lewis Dot Diagrams Electron Dot Structure or Lewis Dot Diagram (Gilbert Lewis) Def. A notation showing the valence electrons (electrons in outer energy level) surrounding the atomic symbol.

29. Lewis Structures 1)Write the element symbol. 2)Carbon is in the 4 th group, so it has 4 valence electrons. 3)Starting at the right, draw 4 electrons, or dots, counter-clockwise around the element symbol. On your sheet, try these elements on your own: a)H b)P c)Ca d)Ar e)Cl f)Al

30. unshared pair- (lone pair) e- not involved in bonding and belong to one atom Structual formula- shows bonds but not unshared pairs of e- in molecule

31. Single Covalent Bond Made of 2 shared electrons 1 comes from one atom in the bond and 1 comes from the other atom in the bond Ex: water – O now is stable with 8 e in outer shell and H is stable with 2

32. Multiple Bonds Bonds with multiple pairs of shared electrons Ex: N 2 Nitrogen has 5 e in os and needs 3 to be stable  shares 3 e (triple bond)

34. Homework Electron Dot Diagrams and Lewis Structures Worksheet