Chemical Bonding

1. Chemical Bonding • Bonding within a molecule is called intramolecular attraction • Ionic bonds • Covalent bonds • Polar covalent bonds

2. Electronegativity • Electronegativity is the tendency of an atom to attract electrons. • Group 1 and 2 lose electron easily, these elements have a low electonegativity • Group 17 and 16 can attain electrons easily, these elements have high electronegativities • Delta represents a partial charge or fractional charge

3. Ionic Bonds • Steals an e-, this causes one atom to have a positive charge and the other a negative charge. This is the reason for the attraction • Opposites attract!! • Group 1 and 2 loose electrons to group 16 and 17 (usually) • When the difference in electonegativities between atoms is greater than 1.7 the molecules is ionic

4. Example NaCl • Na = .9 • Cl = 3.0 • 3.0 - .9 = 2.1 • 2.1 > 1.7 = ionic bond • Cl would be the more negative atom because it has a higher electronegativity (it steals the e- from Na)

5. Covalent Bonds • Atoms share electrons because they have similar electronegativities • Atoms share electron to fill octet rule • Hydrogen form stable molecules were it shares two electrons, this is the duet rule • Single bond formed when atoms share one pair of e- • Double bond formed when atoms share two pair of e- • Triple bond formed when atoms share three pair of e- • The electonegativites differences are 0 - .3

6. Example O2 • O = 3.5 • 3.5 – 3.5 = 0 • 0 falls between 0 and .3 therefore O2 is a covalent bond

7. More examples of covalent bonds Ø

8. Polar covalent bonds • Electrons not always shared equally. • The atom with the higher electronegativity attract the shared electron pair more strongly pulling it away from the other atom. • The shared pair is shifted from the center between the two participating atoms making one end of the molecule positive and the other end negative. The bond is polarized. (Dipole – one side of molecule is slightly negative and one part slightly positive) • The difference in electronegativity among the atoms is .4 to 1.7

9. Now practice with worksheet #55 • Example PO3 • P = 2.1 • O = 3.5 • 3.5 – 2.1 = 1.4 • 1.4 falls between .4 and 1.7 therefore PO3 is a polar covalent bond • Oxygen would be the more negative atom because of the greaterelectonegativity

11. Lewis Structures Represents individual valence electrons Represents a nonbonding pair or lone pair of electrons Represents a pair of e-

12. X 2s 2p B. Lewis Structures • Electron Dot Diagrams • show valence e- as dots • distribute dots like arrows in an orbital diagram • Show a single line for a single bond, double line for a double bond, triple line for a triple bond • 4 sides = 1 s-orbital, 3 p-orbitals • EX: oxygen O

13. Lewis Structures • 1) Count all valence electrons in all the atoms in the molecule; it doesn’t matter which atoms they come from. • 2) If the compound has more than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.) • 3) Make a bond between the central atom and the other atoms using a dash (-) to show a pair of shared electrons. • 4) Place the remaining valence electrons around each of the atoms so that they have an octet.

14. Lewis structure • When more than one Lewis structure can be drawn for a particular molecule this molecule exhibits resonance

15. Example of an electron dot diagram for CCl4 First step determine the number of valence electrons C has 4 (remember 2 from 2s and 2 from 2p) Cl has 7 (4) (remember the 7 comes from 2e- from the 3s and 5ee from 3p) the 4 comes from Cl4 4 + 7 (4) = 32 valence electrons Draw the e- dot diagram CL O O O Now count to see if you have 32 valence e-

16. Your turn, draw the e- dot diagram for CH3I • Step 1 – count valence e- • Step 2 – draw e- dot diagram remember (If the compound has more than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.) • Step 3 count to see if you have the correct number of Valence e-

17. Did you remember to: • Check to see if all of the valence electrons got used? • Check to make sure every atom has an octet of electrons (or 2 for Hydrogen)

18. There are 14 Valence e- (4 +3 + 7) • H only needs two e- to satisfy the duet rule • Iodine and Carbon needs 8 e- to satisfy the octet rule I C H H H How did you do? Now practice with worksheet #56 and #57

19. Metallic Bonding • Metallic bonding is the attraction between metal atoms and a sea of surrounding valence electrons • A result of this bonding is mobile electons which gives rise to the excellent electrical conductivity of most metals.

21. Intermolecular Forces • Intermolecular forces – bond that holds molecules together. • Effect boiling and freezing points • http://www.bcpl.net/~kdrews/interactions/interactions.html

22. Intermolecular Forces • Dipole – Dipole • Molecules are attracted to each other as a result of partial charges of dipole molecules this is called a dipole-dipole intermolecular force • In general, intermolecular forces are about 1% as strong as intramolecular forces.

23. Intermolecular Forces • Hydrogen Bonding • Especially strong dipole-dipole are called hydrogen bonding, example = water • When F,O and N are attached to hydrogen they will form hydrogen bonds with other molecules • Common molecules that form hydrogen bonds are HF, H2O and NH3. • Molecules with an O-Hbond, like alcohols also exhibit hydrogen bonding

24. Hydrogen Bonding – type of attraction that holds two water molecules together. • Cohesion – attractive force between particles of the same kind • Adhesion – attractive force between unlike substances (meniscus)

25. Intermolecular Forces • London Dispersion Forces (LDF) • The force that holds noble gas atoms and nonpolar molecules together • No permanent dipoles however, at any given time, the e- can be mostly on one side of the molecule creating a slightly negative charge on one side compared to the other • Weak and easily broken force call LDF • Larger the molecule the more e- it has therefore the stronger LDF

26. Bonding forces and boiling and freezing points • Strongest bonding force is a hydrogen bond, followed by a dipole-dipole force, and then LDF. • H-Bonds have highest boiling, melting and freezing points • Dipole-dipole force have intermediate boiling, freezing and melting points • LDF have lowest boiling, freezing and melting points

27. Now practice with worksheet #58

28. VSEPR model • Valance Shell Electron Pair Repulsion (VSEPR) is a 3D model of a molecule • Draw the Lewis structure for the molecule • Count the electron pairs and arrange them in the way that minimizes repulsion(that is, put the pairs as far apart as possible) • Determine the positions of the atoms from the way the electron pairs are shared • Determine the name of the molecular structure from the positions of the atoms. • http://wunmr.wustl.edu/EduDev/Vsepr/table1.html

29. Linear • Number of bonding pairs around central atom = 2 • Number of lone pairs around central atom = 0 • 180 bond • Example- BeCl2, CO2

30. Triangle (Trigonal planar structure) • Number of bonding pairs around central atom = 3 • Number of lone pairs around central atom = 0 • Examples- BF3, SO3

31. tetrahedral • Number of bonding pairs around central atom = 4 • Number of lone pairs around central atom = 0 • 109.5 bond angle • Example- CH4

32. Trigonal pyramid • Number of bonding pairs around central atom = 3 • Number of lone pairs around central atom = 1 • 120 bond angle • Example-NH3

33. Bent or v-shaped • Number of bonding pairs around central atom = 2 • Number of lone pairs around central atom = 1 or 2 • 106 bond angle • Example H2O, SO2