1. Unit 3: Bonding

2. Energy and Chemical Bonds Chemical bonds are forces that hold atoms together in a compound Bonds are formed by either: 1.The transfer of valence electrons 2.The sharing of valence electrons Chemical energy is stored in bonds (as potential energy)

3. Energy and Chemical Bonds When chemical changes occur, energy is either given off or absorbed –Breaking bonds requires an input of energy Endothermic reaction Reactants + ENERGY (heat in cal or J)  Products –Forming bonds releases energy Exothermic reaction Reactants  Products + ENERGY (heat in cal or J)

4. Energy and Chemical Bonds When a chemical bond is formed –The product has less potential energy than the reactants because energy is released –The greater the amount of energy released, the greater the stability

5. Octet Rule Octet 1.The configuration of 8 valence electrons This is the maximum # of electrons that an atom can have in its valence level An octet results in a very stable electron configuration Noble gases already possess an octet –Noble gases are unreactive (don’t make bonds with other elements easliy) All other elements bond to achieve octet and become more stable

6. Octet Rule Atoms generally react by transferring or sharing valence electrons in order to achieve the octet Exceptions –H–H, He, Li, and Be will always have a max of 2 valence electrons

7. Weak Vs. Strong Bonds Weak Bonds –R–Require little energy to break –R–Release little energy when formed –U–Unstable Strong Bonds –R–Require lots of energy to break –R–Release lots of energy upon formation –S–Stable

8. Lewis Dot Diagrams Lewis dot (electron dot) diagrams are used to indicate the number of valence electrons around an atom Valence electrons are found in the outermost energy levels’ s and p sublevels Valence electrons determine the chemical properties of an atom All elements in the same group or column have the same # of valence electrons, and therefore the same chemical properties

9. Lewis Dot Diagrams □ □ X □ X= element’s symbol □ Top box is the s orbital Left, right, and bottom box are the p orbitals

10. Lewis Dot Diagrams Examples What are the lewis dot structures for 1.Na5. P 2.Mg6. S 3.Al7. Cl 4.Si8. Ar

11. Answers to Lewis Dot Structures Lewis Dot

12. Ionic Bonds How they are created 1.Metals tend to lose electrons and become cations 2.Nonmetals tend to gain electrons and become anions 3.Loss from cations and gain by anions represents the transfer of electrons between atoms 4.The resulting ions cations (+) and anions (-) are oppositely charged so they can attract each other forming an ionic bond

13. Ionic Bonds Example:

14. Ionic Bonds: Electronegativity Difference (Table S) Electronegativity (EN) is the ability of an atom to attract an electron so a bond will form between atoms Usually a bond is ionic if the difference between the EN values of each atom in the bond is 1.7 or higher 1.Examples: NaCl (3.2 – 0.9 = 2.3) LiF (4.0 – 1.0 = 3.0) 2. Exceptions: Metal Hydrides –F–Forms when a H atom combines with an active group 1 metal –T–The H takes the -1 charge –T–The EN difference comes out to be less than 1.7, but still ionic –e–example: NaH (2.1 – 0.9 = 1.2) –O–Other exceptions: LiH, KH, RbH, CsH, FrH

15. Properties of Ionic Compounds (Salts) Most are crystalline solids at room temperature Generally have high melting points Can conduct electricity when they are melted (molten) or dissolved in water (aqueous) 1.This conductivity is possible because the ions that are produced are free to move

16. Ionic Character Most Ionic means that the answer choice with the highest EN difference is the answer Least ionic means that the answer with the lowest EN difference is the answer Sample Q 27-29 Page 88 in RB

17. Covalent Bonding Examples

18. Covalent Bonding Types of covalent bonds 1.Nonpolar –2–2 atoms of the same nonmetallic element share electrons equally –A–All diatomic elements are examples of nonpolar molecules –E–EN difference between 2 atoms in a nonpolar bond is 0

19. Covalent Bonds 2. Polar Bonds –2–2 atoms of different nonmetallic elements unequally share electrons –E–EN difference for atoms in a polar bond will be greater than 0, but less than 1.7 –E–Exception HF (4.0-2.1 = 1.9)

20. Practice Using Table S to obtain EN values, determine whether the following bonds are ionic, polar covalent, or nonpolar covalent 1.HCl 3.2 -2.1 = 1.1 2.NaCl 3.Cl 2 4.CH 4 2.6- 2.1=.5

21. Covalent Bonds It is possible for 2 atoms to share more than 2 electrons –Atoms can share 2, 4 and even 6 electrons with each other –Single bond (share 2 electrons) –Double bond (share 4 electrons) –Triple bond (share 6 electrons)

22. Covalent Bonding

23. Properties of Molecules May exist as solids, liquids, or gases depending on the strength of intermolecular attractions Generally soft (when in solid state) Poor conductors of heat and electricity Low melting and boiling points Diatomic elements are molecules 1.HOFBrINCl 2.Consist of 2 of the same atom covalently bonded together (H 2, O 2 )

24. Polyatomic Ions Table E Group of covalently bonded atoms that have an overall charge greater or less than zero. These ions (+ and -) also have the ability to form ionic bonds with either a metal or a nonmetal. Classic regents questions: Which of the following contains both ionic and covalent bonds? –Always pick the answer choice that has a metal with a polyatomic ion or a polyatomic ion with a nonmetal

25. Practice Which of the following contains both ionic and covalent bonds? 1.KCl 2.NaNO 3 3.HBr 4.K 3 PO 4

26. Drawing Lewis Dot Structures for Covalent Bonds (Regents Only) Draw a skeletal structure (arrangement of the atoms) Connect atoms with bonds –This is represented with a dash ─ –Each dash represents 2 electrons Once you have connected the atoms, then place all other unbonded electrons

27. Resonance Structures There are some molecules for which satisfactory Lewis structures cannot be made. An example is sulfur trioxide (SO 3 ) S: 1 x 6 = 6 O: 3 x 6 = 18 (6+18 = 24 total electrons)

28. Resonance Structures From this Lewis dot you can see that there are 3 configurations that are possible, so all structures must be drawn. You should look at them as being superimposed on one another The bond lengths are somewhere between a single and double bond The electrons shown in the double bond spread out so that they are equally shared among all constituents (accounts for the intermediate bond lengths)

29. Resonance Structures Draw the lewis dot structure for NO 3 -

30. Exceptions to the Octet Rule Odd # of electrons (extra electron) –Common examples are NO and NO 2 Molecules with fewer than 8 valence electrons –Common examples are H, Be, B, and Al

31. Molecules with More than 8 Valence Electrons These are the most common exceptions to the octet rule In these molecules, the atoms use d orbitals to hold more than 8 valence electrons These tend to form from large central atoms surrounded by small, highly EN elements (F, Cl, O)