Bonding

1. Bonding Elements form bonds in order to reach the lowest potential energy state. Electrons are transferred or shared in order to obey the ‘rule of eight’.

2. Three types of bonds Ionic Elements transfer electrons from a metal to a nonmetal in order to form an octet in the valence level Covalent Two nonmetals share electrons in order to form an octet in the valence level Metallic Delocalized electrons ‘float’ in a sea of electrons among metal cations

3. Ionic bonds metal + nonmetal lose e- to form cation gain e- to form anion Electrons are transferred until the net charge is zero Example: aluminum oxide Al2O3 O Al Total charge O 2 x 3+ = +6 Al 3 x 2- = -6 O 0 •• 2- • •• • 3+ •• • •• 2- •• • • 3+ •• • •• 2- •• • •

4. Properties of Ionic Compounds Crystalline solid Held together by electrostatic forces (positive-negative attraction between cation and anion) High melting point Soluble in water Conduct in molten state Do not conduct in solid state Brittle

5. Metallic Bonding Results from the attraction between metal atoms and a surrounding sea of mobile electrons. Valence electrons are lost to form cations.  These are delocalized, which means they do not belong to any one atom but move freely throughout the entire metal. Metallic bonding is not directional but uniform throughout the solid.  One group of atoms can slide past another group of atoms amid the electron sea without breaking any attractions.  Explains metallic properties like electrical and thermal conductivity, malleability, ductility, and luster.

6. + + − + + − − − − − − − − − − − − − + + + + − − − − − − − − − − − − − − − + + + + − − − − Metal bond ‘sea of electrons’

7. Covalent bonds nonmetal + nonmetal Elements share electrons Bond forms because of an overlap of orbitals

8. Properties of covalent compounds Solid, liquid, gas low melting point < 300oC High to low solubility Poor conductor Exist as true discrete molecules

9. Bond facts • Number of covalent bonds that an atom can form = number of valence electrons • Atoms will position themselves so as to achieve the lowest possible energy • Distance where energy is minimum is the bond length • Bond will form if the energy of the aggregate is lower than that of the separated atoms covalent ionic 0polarity increases as electronegativity difference increases

10. Electronegativity-ability to attract electrons Higher the electronegativity value, stronger the attraction The difference (absolute value; no negative numbers ) between electronegativity values can be used to predict the type of bond formed by elements. ≤ 0.2 nonpolar covalent bond < 0.2 but ≤ 1.7 polar covalent bond > 1.7 ionic bond CO2     C---O   2.5 – 3.5  =  1.0   polar covalent bond  O2 O = O 3.5 – 3.5 = 0 nonpolar covalent bond

11. polar covalent bonds—electrons are shared unequally The element that has the higher electronegativity value will pull electrons closer This will set up a partial positive (d+) and partial negative (d-) charge.  These partial charges are called dipoles and can be indicated using this symbol          The arrow points toward the atom with the higher electronegativity CO2     C O  or C O 2.5 – 3.5  =  1.0    d+ d -

12. nonpolar covalent bonds-electrons are shared equally no dipoles Nonpolar covalent bonds are usually between identical atoms or elements located very close on the periodic table. Analogy: Imagine a Tug-O-War between two defensive linemen on a football team.  Each pulls with a strong force and no one wins.  Tug-O-War between two horseracing jockeys.  Each pulls with a weak force and no one wins.  Now, what if the two defensive linemen pull against the two horse-racing jockeys, the football players pull with a stronger force and win the Tug-O-War.  Same thing can happen in a molecule. If one nucleus has a stronger attraction for the electrons, the shared pair of e- is pulled closer to that nucleus.

13. Polar and Nonpolar Molecules Polar molecules generally result from polar bonds, but the shape of the molecule influences properties.       CO2 C       O 2.5 - 3.5 = 1.0 polar covalent bond Draw the electron dot structure        O C O The linear molecule with dipoles equal and opposite will cause the net pulling force to equal zero, so this is a nonpolar molecule. CH4               C H   2.5 - 2.1  =  .4                          polar covalent bond Draw the electron-dot structure                      H                                                             H         C         H                                                                         H the tetrahedral molecule with dipoles equal and opposite will cause the net pulling force to equal zero, so this is a nonpolar molecule.

14. CHCl3           C H   2.5 - 2.1  =  .4 polar covalent bond                      C Cl  2.5 - 3.0  =  .5 polar covalent bond Draw the electron dot structure                      H                                                             Cl        C         Cl                                                                        Cl The tetrahedral molecule with dipoles that are unequal will cause the net pulling force to not equal zero, so this is a polar molecule. Polar molecules Lone pair will make a molecule polar Central atom bonded to different atoms

15. Solubility and polarity Polar molecules and ionic substances dissolve in polar solvents and nonpolar molecules dissolve in nonpolar solvents.   Remember…. “Like dissolves like.” You can predict the properties of the molecule by trying to dissolve a substance in the polar solvent, water or the nonpolar solvent, hexane.

16. Bond strength and bond length Single bond (sigma bond)  longer weaker end to end overlap Double or triple bond (pi bond)  shorter stronger O C O side to side overlap First bond formed is always a sigma bond, second or third bond would be a pi bond 1  1  1  1 

17. Hybridization(blending of orbitals) 2 bonding sites sp 3 bonding sites sp2 4 bonding sites sp3 5 bonding sites sp3d 6 bonding sites sp3d2 7 bonding sites sp3d3 lone pair on the central atom = bonding site double or triple bond = one bonding site http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/hybrv18.swf

18. Geometry of Molecules Valence Shell Electron Pair RepulsionVSEPR Valence electrons will repel each other strongly and cause the peripheral atoms to move as far from each other as possible This repulsion will determine the shape of the molecule Every central atom will hybridize

19. G D Example: CO2 linear All atoms in a line Bond angle of 180o bent Usually three atoms that could be linear but instead will bend One or two lone pairs (unbonded electrons) Bond angle two lone pairs =104.5º one lone pair = < 120 º Always polar because of lone pair Example: H2O

20. A L trigonal planar Atoms in same plane but in a triangle arrangement Bond angle 120o Usually a double bond Tetrahedral Central atom with 4 other atoms arranged around this atom Bond angle 109o Example: NO3- Example: CH4

21. F K Trigonal pyramidal One central atoms with three other atoms and a lone pair arranged in a pyramid shape Bond angle 107o Always polar because of lone pair Trigonal bipyramidal Central atom with 5 other atoms arranged around it Exception to the octet rule Bond angles 120o (trigonal planar middle) 180o and 90o Example: NH3 Example: PCl5

22. H Octahedral Central atom with 6 other atoms around it Expanded octet possible because of ‘d’ orbitals Exception to the octet rule Example: SF6

23. Intermolecular forcesvan der Waals forcesattractive forces between molecules • London dispersion forces • Dipole-dipole forces • Hydrogen bonding

24. London dispersion forces • Weakest of all IMF’s • Found in both polar and nonpolar molecules • Strength increases with increased number of electrons • Strength increases with more mass • Induced dipole--results from collision with other molecules- ‘squishy cloud’; electron cloud is polarizable; (creates partial positive and partial negative side) • Reason why: • Fluorine is a gas • Bromine is a liquid • Iodine is a solid

25. d+d- d+d- H F H F Dipole-dipole forces Slightly stronger than LDF’s Only polar molecules • strong attraction between the partial positive charge and partial negative charge • The more polar the molecule, the stronger the dipole-dipole force

26. d+d- d+d- d+d- d+d- d+d- d+d- d+d- Dipole Interactions d+d-

27. Hydrogen bonding • ( NOT a bond) • Strongest of the IMF’s • Only in polar molecules with these bonds: • H—F • H—O • H—N • F, O and N are very electronegative so it is a very strong dipole

28. H O O H H O H H H H O H H H H O O O H H H Hydrogen bonding

29. IMFs Predict high boiling point, strong surface tension, high vapor pressure for compounds with strong intermolecular forces Predict low boiling point, splattering, fast evaporation for compounds with weak intermolecular forces LDF’s are weakest type of IMF Dipole-dipole IMFs are strong Hydrogen bonds are the strongest type of IMF LDF's  increase with increasing mass LDF’s increase with increased number of electrons