1. BONDING Bond = when two nuclei simultaneously attract the same pair of electrons

2. *The electrons experience a force of attraction from both nuclei. *This_______________________ attraction holds the two particles together. *This attraction is called ________________ * positive-negative-positive a chemical bond. One Pair of Electrons = 1 bond

3. Bonding changes = Energy Changes *Breaking Bonds= Endothermic Endothermic=Absorbing Energy The ______ the Energy, The More _______ !!!!!!! *By __________ Energy, Atoms become _____ stable when separated. AB A + B AB + Energy A + B LowerStable absorbing less

4. Remember… _________the Energy, The More _________ !!!!!!! _____________ =____________  Making Bonds= Exothermic =Releasing Energy (Exit, Exhale) When bonds ______, excess energy is ________, resulting in a more ________ compound. A + B AB + Energy A+B AB The Lower Stable High EnergyLess Stable Exothermic Form released stable

5. Chemical Energy * ______________ is associated with the changes in bonds between atoms. ( ____________) * Single, uncombined atoms have a _______________ __________ *When bonds form, some of the _________________ ______________ * The amount of Potential Energy associated with a specific bond depends on: ______, ___________, ____________, etc. Potential Energy Stored Energy high Potential Energy. Potential Energy is released. MassBond Type Atomic Radius

6. Remember the Octet Rule… *Atoms will _____, _____, or ______ enough electrons as to obtain a _____, stable octet. A. Electrons can be ______________ _____ or ______ to form ions with a ______________ = ________________ B. Electrons can be ___________ When _______________, both atoms feel as if they have a __________. = __________________ __________ TO ______________ _____________ TO ________________ lose gainshare full Transferred Lost gained stable octet.IONIC BONDS METALS NON-METALS Shared sharing electrons full octet Covalent Bonds NON-METAL

7. EXCEPTIONS TO THE OCTET RULE *_________ and ______ are Exceptions to the octet rule because ___________ ________________________________ *_________ and ______ are Exceptions to the octet rule because ___________ ________________________________ *They are stable and Happy :) with a full “ ____” of electrons. HydrogenHelium they can only hold a maximum of 2 electrons in their PEL. duet

8. Are all bonds the same strength? NO!! A. Weak bonds Require little energy to break, and release little energy when forming. (unstable) B.Strong Bonds Require lots of energy to break, and release lots of energy when forming (highly stable)

9. Practice: 1.For the following equation: N (g) + N (g) --> N 2(g) + Energy a. Bond broken, energy absorbed b. Bond broken, energy released c. Bonds formed, energy absorbed d. Bond formed, energy is released 2. Which would be true of the reverse reaction? _____________________________ a. Bond broken, energy absorbed

10. More Practice Which amount of energy would be given off in the formation of a highly stable compound? 3. Which amount of energy would be given off in the formation of a highly stable compound? a) -100 kJ/mol b) -800 kJ/mol c) +170 kJ/mol d) +500 kJ/mol 4. Which amount of energy from above would symbolize weak bond formed? ___________________ d) +500 kJ/mol

11. One way to predict bond type is by taking the difference in atom’s electronegativities. What is electronegativity? -Ability to attract electrons Who had the highest electronegativity? Lowest? What did an atom’s electronegativity value mean for its reactivity? Highest = Non- Metals,upper right on table Lowest = Metals, lower left corner of table Higher electronegativity value,more attraction for electrons, Higher reactivity. Lower electro.value, less reactivity. (More on Electronegativity and bonds ahead)

12. Bonds vs. Molecular Attractions Intramolecular Force of Attraction = ________ *When two nuclei simultaneously attract the same pair(s) of electrons to form a covalent bond. (stronger) Intermolecular Force of Attraction *The attraction that exists b/w two separate molecules that is responsible for holding them near one another in space. (weak in comparison) BOND

13. Bond vs Intermolecular forces of attraction

14. Bond Types 1. IONIC 2. METALLIC 3. COVALENT a. nonpolar covalent b. polar covalent c. coordinate covalent 1. IONIC 2. METALLIC 3. COVALENT a. nonpolar covalent b. polar covalent c. coordinate covalent

15. Ionic bonds *Form between a _____ and a _____________ - Nonmetal _____ those electrons and becomes __ *Metal ______(+ ion) attracts Nonmetal _____(a negative ion) = - Metal _____ 1 or more electrons and becomes __ Ionic Bonds *The amount of electrons transferred depends on metal nonmetal loses + gains - cation anion how many each atom needs to lose or gain to complete its octet.

16. Ionic bonds

17. NaCl =most common ionic compound Na Cl Electron Configuration 2-8-1 2-8-7 Valence Electrons 17 New e- config. With full shell [2-8] +1 [2-8-8] -1 Ion Formed Na +1 Cl -1 Lewis Dot Na Cl Lewis Dot of Ion [ Na ] +1 [ Cl ] -1 * * * * * * * *

18. Ionic Bonding

19. Ionic Bonds can also be predicted using electronegativity values *If the difference in electronegativity is greater than or equal to ______________ ___________ *If the difference in electronegativity is greater than or equal to ______________ ___________ NaCl Cl = Na = Remember: 1.7, the bond is ionic 3.2 0.9 2.3= ionic There are exceptions to everything.

20. Using electronegativity values, predict if the following compounds are ionic. 1) LiBr2) H2O3) CaO4) CO2

21. Properties of Ionic Substances *Form when ________ react with __________. *Electronegativity difference of ____________. *Have ______ melting points and _______ boiling points. *All ionic compounds are _________ at room temperature since the bonds are so _________. *Have very low vapor pressure. *Form when ________ react with __________. *Electronegativity difference of ____________. *Have ______ melting points and _______ boiling points. *All ionic compounds are _________ at room temperature since the bonds are so _________. *Have very low vapor pressure. metals Non metals 1.7or higher high solids strong -Atoms do not easily go from liquid to gas because of strong bonds.

22. Properties of Ionic Substances *Are ______________ of electricity in the solid state. *Conducts electricity as a _______ or ______. *Are ______________ of electricity in the solid state. *Conducts electricity as a _______ or ______. Non conductors liquidaqueous (This allows the electrons to be mobile, move freely and conduct electricity)

23. Writing Lewis Dot Ionic Compounds Na 11 Cl 17 *In terms of electronegativity values, how can you tell that NaCl is ionic?

24. Writing Lewis Dot Ionic Compounds Use: Li and O Use: Li and O Use: Ca and F

25. Binary Ionic Compound: One metal and one nonmetal NaCl, CaBr 2, MgO, K 2 S Some Ionic Compounds have more than 2 elements, they usually contain polyatomic ions NaNO 3 CaSO 4 AgNO 2 NH 4 ClLiCr 2 O 7

26. Polyatomic Ions (E) 1)Grouping of 2 or more atoms that act as a unit and carry a charge. 2)If a compound contains a polyatomic ion, it has both ionic and covalents bonds. NaNO 3 Na +1 and NO 3 -1 1)Grouping of 2 or more atoms that act as a unit and carry a charge. 2)If a compound contains a polyatomic ion, it has both ionic and covalents bonds. NaNO 3 Na +1 and NO 3 -1

28. Metallic Bonds *Metals generally have: -Few valence e- -Low ionization energy ( Ionization = How much energy is needed to remove the outermost e-) *Metallic bonds are relatively strong bonds: -Moderately high melting point -Moderately high boiling point Are formed between _________ and ____________ Metal

29. Metallic Bonds *Kernel- Nucleus and non valence electrons *Arranged *Valence electrons *Metals are in a fixed position of a crystal lattice good conductors because of these mobile electrons move freely about the crystal lattice and do not belong to any given atom.

30. Metallic Bonds *Metallic Bonds Result from the force of attraction of the mobile valence electrons for an atoms positively charged kernel. ****SEA OF MOBILE ELECTRONS !!! Metal= crystal lattice Red sphere = (+) kernel of metal cations Static = moving e-

31. Examples of Metallic Bonds ****** ****** The bonds found in copper wire The bonds found in a piece of magnesium strip The bonds found in a piece of gold.

32. Properties of metals * ex. Cu (s) Au (s) Hg (l) * *Within the metal crystal lattice, electrons belong to the whole crystal rather than individual ions. *Mobility of electrons distinguishes the metallic bonds from all others. * ex. Cu (s) Au (s) Hg (l) * *Within the metal crystal lattice, electrons belong to the whole crystal rather than individual ions. *Mobility of electrons distinguishes the metallic bonds from all others. Composed of Monoatomic metal atoms. Moderate Bond Strength Moderate melting and boiling points Conducts electricity in the solid and liquid phase Malleable and ductile

33. Covalent Bonds = *Electrons are -Sharing can be or *Electronegativity difference is usually *Electrons are -Sharing can be or *Electronegativity difference is usually MOLECULES H = * * ** Outer shell Remember… Hydrogen can have only 2 in its outer shell. By sharing 1 pair of e- = 2 e- shared between 2 or more non metals to obtain a stable octet. equalunequal below 1.7 H2H2 is full for BOTH by SHARING.

34. Practice Covalent Bonds Combine : F and F Show the formation of H 2 O

35. Double Bonds *Double Bond- - *Each bond is usually 1 e- from each atom forming the pair. *Double Bond- - *Each bond is usually 1 e- from each atom forming the pair. Example: Show the formation of the following: CO 2 2 Pairs of e- (4 e-) are shared Makes it more stable, Lower Potential Energy

36. Triple Bonds *Triple Bonds- * *Triple Bonds- * Example: N 2 3 Pairs of e- (6 e-) are shared between 2 atoms. Very stable, low Potential Energy

37. How many e- are shared? Single Bond: Cl-Cl, H-H _____pairs of e- _____ e- total Double BondO=O _____ pairs of e- _____ e- total Triple Bond NΞN _____ pairs of e- _____ e- total Single Bond: Cl-Cl, H-H _____pairs of e- _____ e- total Double BondO=O _____ pairs of e- _____ e- total Triple Bond NΞN _____ pairs of e- _____ e- total 1 2 2 4 3 6

38. Polar or Non Polar Covalent?? *Each Covalent Bond is one of 2 types: or *Each Covalent Bond is one of 2 types: or Non Polar Covalent Bonds : -A bond formed between 2 atoms that have the same attraction ( electronegativity) for the shared pair of e-. ***Different non metals can have the same electronegativity values ---- Polar Non Polar Equal pull on bonding pair Symmetrical e- charge distribution

39. Examples of Non Polar Bonds Examples: Cl 2 H 2 CS 2 Examples: Cl 2 H 2 CS 2 Cl - Cl H - H C-S

40. Polar Covalent Bonds *Polar Covalent Bonds: -Bonds formed between 2 non metal atoms which have different attractions for shared e- (different electronegativities) - *Polar Covalent Bonds: -Bonds formed between 2 non metal atoms which have different attractions for shared e- (different electronegativities) - H – Cl Unequal pull on bonding pair of e- Asymmetrical e- charge distribution Leaves partial charges on atoms within molecule

41. Polar Covalent Examples: H 2 OHBrNH 3 Examples: H 2 OHBrNH 3 e- cluster around more electronegative atom

42. Properties of Covalent Molecules *Form when a _________ reacts with a __________ *Have ____ melting points and ____ freezing points *Molecular solids are soft *Covalent Molecules are ______________ of heat and electricity *Molecular substances are generally _____________ *Molecular substances have a ____ vapor pressure *Form when a _________ reacts with a __________ *Have ____ melting points and ____ freezing points *Molecular solids are soft *Covalent Molecules are ______________ of heat and electricity *Molecular substances are generally _____________ *Molecular substances have a ____ vapor pressure non metal Lowlow poor conductors weak bonds high (goes into vapor easily because of weak bonds.)

43. Shapes of Molecules *The shapes of molecules is determined by _____________ ________________ *The shapes of molecules is determined by _____________ ________________ the repelling of electrons There are 5 Basic shapes: *Linear *Angular *Planar *Pyramidal *Tetrahedral

44. Linear When there are 2 atoms in a molecule Any molecule in which the bonds cause the atoms to fall in a straight line. CO 2 C2H2C2H2 H2H2

45. Angular *V-shaped or Bent SO 2 H2OH2O H2SH2S

46. Planar *The atoms of a molecule fall in the same plane. C2H4C2H4

47. Pyramidal *3 shared pairs of electrons will cause a pyramid shape. NH 3 PCl 3

48. Tetrahedral *4 evenly separated bonds CCl 4 CH 4

49. Polarity Of *Molecules can either be polar or non polar. (Do not confuse bond polarity with molecular polarity) Bond Polarity- When checking between 2 atoms if there is difference in electronegativity or not. * Polar bond-difference in electronegativity between 2 atoms *Non Polar Bond- No difference in electronegativity between 2 atoms. *Molecules can either be polar or non polar. (Do not confuse bond polarity with molecular polarity) Bond Polarity- When checking between 2 atoms if there is difference in electronegativity or not. * Polar bond-difference in electronegativity between 2 atoms *Non Polar Bond- No difference in electronegativity between 2 atoms. Molecules

50. NonPolar Molecules *Molecules containing only nonpolar covalent bonds, are always nonpolar molecules! All diatomics: Br 2, I 2,N 2, Cl 2, H 2, O 2, F 2,

51. *Non polar molecules can also contain polar bonds *Bonding e- have to be symmetrically distributed between atoms. CH 4 Tetrahedral CCl 4 CO 2 Linear NonPolar Molecules

52. Polar Molecules * Molecules that contain polar covalent bonds can also be polar molecules depending on overall molecule symmetry (shape). (asymmetrical) H2OH2O HF SO 2 Leads to partial charges at different points in molecule, even though overall molecule is neutral

53. ***Use dipole arrows for each bond to determine the overall symmetry of the molecule*** 1)Draw the dipole arrow facing the more electronegative atom. a) The atom the arrow points towards is  - while where the arrow points away from is  + 2) If dipole arrows are equal in magnitude but opposite in direction, they cancel out and the molecule is nonpolar

54. What is the Polarity of the following Molecules? HClCO 2 CH 4 NH 3 CCl 4 H 2 0 HClCO 2 CH 4 NH 3 CCl 4 H 2 0

55. S.N.A.P S ymmetrical N on polar A symmetrical P olar S ymmetrical N on polar A symmetrical P olar

56. Like Dissolves Like Like Polarities Dissolve Like Polarities Like Polarities Dissolve Like Polarities Polar dissolves Polar Non Polar dissolves Non Polar Alcohol and water Polar NOT Dissolving non Polar Oil and Water Many organic molecules

57. Intermolecular Forces Of Attraction (IMF’s) * The attractions between molecules that holds them near one another in space.

58. Intermolecular Forces Of Attraction (IMF’s) 4 types of Intermolecular forces of attraction: 1. Dipole-Dipole 2. Hydrogen Bonding 3. Van der Waals 4. Molecule-Ion or Dipole-Ion 4 types of Intermolecular forces of attraction: 1. Dipole-Dipole 2. Hydrogen Bonding 3. Van der Waals 4. Molecule-Ion or Dipole-Ion

59. IMF’s explain the change in phase going down the halogen group (17) F 2 Cl 2 Br 2 I 2 Generally IMF’s will INCREASE with: 1.Increasing #e- 2.Increasing mass of atoms 3.Increasing mass and complexity (polarity) of a molecule.

60. 1. Dipole-Dipole 1) When the partially δ + end of one dipole, attracts the partially δ– end of a neighboring dipole. Show H 2 S:

61. Dipole *The temperature at which liquid boils is due to the strength of the dipole attractions *A higher attractive force= increased mp and increased bp *The larger the molecule, the stronger the attractive forces *The temperature at which liquid boils is due to the strength of the dipole attractions *A higher attractive force= increased mp and increased bp *The larger the molecule, the stronger the attractive forces Dipole= Polar

62. 2. Hydrogen Bonds 1) Not an actual bond, just a strong IMF 2) H-bonding is FON! 3) H-bonds are only possible if H is covalently bonded to either F, O, or N. a) those polar molecules then strongly attracted each other which is the “hydrogen bond” b) Common molecules that can H-bond: H 2 O, HF, and NH 3 1) Not an actual bond, just a strong IMF 2) H-bonding is FON! 3) H-bonds are only possible if H is covalently bonded to either F, O, or N. a) those polar molecules then strongly attracted each other which is the “hydrogen bond” b) Common molecules that can H-bond: H 2 O, HF, and NH 3

63. Hydrogen Bonds *Hydrogen has such a small share of the e- pair, it is almost like a bare proton

64. Van Der Waals Forces *Weakest IMF between small, nonpolar molecules *Van der Waals strength ↑ with ↑ mass and ↑ # e- *Weakest IMF between small, nonpolar molecules *Van der Waals strength ↑ with ↑ mass and ↑ # e-

65. Molecule Ion attraction 1. Partially δ– end of polar molecule, dipole, attracts + cation 2. Partially δ + end of dipole attracts - anion

66. Molecule Ion Attraction Example: Saltwater, NaCl (aq) (Dipole is usually water) Example: Saltwater, NaCl (aq) (Dipole is usually water)

67. Order of IMF Strength Molecule Ion Hydrogen Bonding Dipole Van der Waals Molecule Ion Hydrogen Bonding Dipole Van der Waals Decreasing Strength of IMF

68. What do IMF’s account for? 1)Boiling Point 2)Melting Point 3)Vapor Pressure 4)Phase 1)Boiling Point 2)Melting Point 3)Vapor Pressure 4)Phase ↑IMF ↑BP ↑MP ↓IMF ↓BP ↓MP ↑IMF↓VP, ↓IMF ↑VP The ↑IMF, the more likely a substance will be solid ↓IMF, the more likely the substance will be a gas