1. Chemical Bonding and Nomenclature Adapted from Paul Surko

2. Bonding, the way atoms are attracted to each other to form compounds, determines nearly all of the chemical properties we see. The number “8” is very important to chemical bonding. What is Bonding????????

3. What are Compounds? Compounds are a combination of atoms bonded together. Bonding determines the chemical properties of the compound.

4. Ionic Bonding-Great with 8 All atoms want 8 valence electrons. Metals give up electrons to form positive ions (cations) and non-metal atoms will receive or take additional electrons to become negative ions (anions). IONS are charged particles. N becomes N -3 Al becomes Al +3 Cl becomes Cl - O becomes O -2 Mg becomes Mg +2 Na becomes Na + The positive and negative ions are attracted to each other electrostatically.

5. Properties of Ionic Compounds Made of cations and anions Exist in crystalline structures (solids) at STP High MPs and BPs Conduct electricity in aqueous and molten states

6. Opposites Attract!

7. Putting Ions Together Na + + Cl - = NaCl Ca +2 + O -2 = CaO Na + + O -2 = Na 2 O Al +3 + S -2 = Al 2 S 3 Ca +2 + N -3 = Ca 3 N 2 Ca +2 + Cl - = CaCl 2 You try these! Mg +2 + F - = NH 4 + + PO 4 -3 = K + + Cl - = Al +3 + I - = Sr +2 + P -3 = Li + + Br - = Sr 3 P 2 AlI 3 MgF 2 (NH 4 ) 3 PO 4 KCl LiBr Not NH 43 PO 4

8. Nomenclature Naming of Ionic Compounds with TMs Binary Compounds have two types of atoms (not diatomic which has only two atoms). Metals (Groups I, II, and III) and Non-Metals Metal _________ + Non-Metal _________ideSodium Chlorine Sodium Chloride NaCl Metals (Transition Metals) and Non-Metals Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III Bromine Iron ( III ) Bromide FeBr 3 Compare with Iron ( II ) Bromide FeBr 2

9. Metals (Transition Metals) and Non-Metals Older System Ferrous Bromine Ferrous Bromide FeBr 2 Compare with Ferric Bromide FeBr 3 Metal (Latin) _______ + ous or ic + Non-Metal ________ide Nomenclature--Naming of Ionic Compounds with TM

11. Let’s Practice! Name the following. CaF 2 K2SK2S CoI 2 SnF 2 SnF 4 OF 2 CuI 2 CuI SO 2 SrS LiBr Strontium Sulfide Lithium Bromide Copper ( I ) Iodide or Cuprous Iodide Sulfur dioxide Copper ( II ) Iodide or Cupric Iodide Oxygen diflouride Tin ( IV ) Flouride or Stannic Flouride Tin ( II ) Flouride or Stannous Flouride Cobalt ( II ) Iodide or Cobaltous Iodide Potassium Sulfide Calcium Flouride

12. Polyatomic Ions (partial list from page 195 (193 2 nd edition)) Ammonium……………... Nitrate…………………… Permanganate………….. Chlorate………………… Hydroxide………………. Cyanide…………………. Sulfate…………………... Carbonate………………. Chromate……………….. Acetate………………….. Phosphate………………. NH 4 + NO 3 - MnO 4 - ClO 3 - OH - CN - SO 4 2 - CO 3 2- CrO 4 2- C 2 H 3 O 2 - PO 4 3-

13. Lets Practice! Na 2 CO 3 KMnO 4 NaOH CuSO 4 PbCrO 4 NH 3 ammonia Copper ( II ) sulfate or Cupric sulfate Lead ( II ) chromate or Plubous chromate Sodium hydroxide Potassium permanganate Sodium carbonate

14. The Covalent Bond Atoms can form molecules by sharing electrons in the covalent bond. This is done only among non-metal atoms.

15. Properties of Covalent Compounds Generally exist as liquids and gases at STP NO crystalline structure Low MPs and BPs Do NOT conduct electricity

16. Dot Structures-Octet Rule (All atoms want 8 electrons around them.) Lewis came up with a way to draw valence electrons so that the bonding could be determined.

17. Rules to Write Dot Structures 1. Write a skeleton molecule with the lone atom in the middle (Hydrogen can never be in the middle) 2. Find the number of electrons needed (N) (8 x number of atoms, 2 x number of H atoms) 3. Find the number of electrons you have (valence e - 's) (H) 4. Subtract to find the number of bonding electrons (N-H=B) 5. Subtract again to find the number of non-bonding electrons (H-B=NB) 6. Insert minimum number of bonding electrons in the skeleton between atoms only. Add more bonding if needed until you have B bonding electrons. 7. Insert needed non-bonding electrons around (not between) atoms so that all atoms have 8 electrons around them. The total should be the same as NB in 5 above.

18. Let's Try it! 1. S 2. N 3. H 4. B 5. NB 6. E.. H:O:H ●● H O H Water H 2 O 2 x 2 = 4 for Hydrogen 1 x 8 = 8 for Oxygen 4+8=12 needed electrons 8 – 4 = 4 non-bonding electrons 2 x 1 = 2 for Hydrogen 1 x 6 = 6 for Oxygen You have 8 available electrons 12 - 8 = 4 bonding electrons 8 H 12 N 4 B 4 NB - - H:O:HH:O:H.. H:O:H ●●

19. Let's Try it! 1. S 2. N 3. H 4. B 5. NB 6. E.. H:N:H ●● H H N H Ammonia NH 3 3 x 2 = 6 for Hydrogen 1 x 8 = 8 for Nitrogen 6+8=14 needed electrons 8 – 6 = 2 non-bonding electrons 3 x 1 = 3 for Hydrogen 1 x 5 = 5 for Nitrogen You have 8 available electrons 14 - 8 = 6 bonding electrons 8 H 14 N 6 B 2 NB - -.. H:N:H.. H:N:H ●● H H H

20. Let's Try it! 1. S 2. N 3. H 4. B 5. NB 6. E.. O::C::O ●● ●● O C O Carbon Dioxide CO 2 1 x 8 = 8 for Carbon 2 x 8 = 16 for Oxygen 8+16=24 needed electrons 16 – 8 = 8 non-bonding electrons 1 x 4 = 4 for Carbon 2 x 6 = 12 for Oxygen You have 16 available electrons 24 - 16 = 8 bonding electrons 16 H 24 N 8 B 8 NB - - O::C::O.... O::C::O ●● ●●

21. Let's Try it! 1.S 2.N 3.H 4.B 5.NB 6.E...... O::C: O: ●● ●● O O C O Carbonate CO 3 -2 3 x 8 = 24 for Oxygen 1 x 8 = 8 for Carbon 24+8=32 needed electrons 24 – 8 = 16 non-bonding electrons 3 x 6 = 18 for Oxygen 1 x 4= 4 for Carbon You have 22 + 2 more available e - 's 24 H 32 N 8 B 16 NB - -.. O::C:O...... O::C: O: ●● ●● O.. :O: 32 - 24 = 8 bonding electrons -2

22. Nomenclature of Covalently Bonded Compounds--Molecules Non-Metals and Non-Metals Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc. CO 2 Carbon dioxide CO Carbon monoxide PCl 3 Phosphorus trichloride CCl 4 Carbon tetrachloride N 2 O 5 Dinitrogen pentoxide CS 2 Carbon disulfide

23. VSEPR Theory Valence Shell Electron Pair Repulsion Theory—Geometric Shapes Linear, Bent Trigonal Planar, Trigonal Pyramidal Tetrahedral Trigonal Bipyramidal Octahedral

24. VSEPR Theory Why is H 2 O bent and CO 2 linear? O in water has lone pairs causing bending whereas the C in carbon dioxide does not

25. VSEPR Theory Lone pairs on the central atom cause crowding (INCREASED REPULSION) and result in bending *Remember only the lone pairs on the central atom matter—the lone pairs on the external atoms do not crowd

26. Polarity of Molecules Polar Molecule: a molecule that has uneven distribution of charge—dipole moments do NOT cancel Nonpolar Molecule: a molecule that has even distribution of charge—all of the dipoles cancel

27. Polarity of Molecules Cont’d Examples on the Board H 2 O CH 4 CO 2 NH 3 *BF Why does water exist as a liquid at STP and carbon dioxide exists as a gas at STP?

28. Intermolecular Forces Intermolecular Forces: forces of attraction that exist between two molecules Hydrogen Bonding: an IMF results from the attraction between hydrogen and a highly electronegative element like F, N, or O Rather strong force Responsible for water’s high surface tension, holding together DNA, and varying BPs and MPs

29. Intermolecular Forces Dipole-Dipole Force: an IMF that exists between two polar molecules (hydrogen bonding is a special type of dipole- dipole force) Rather strong Used to predict MPs and BPs Van der Waals Force: an IMF that exists between two nonpolar molecules Very weak Instanteous Predicts the low MPs and BPs of nonpolar molecues

30. Forces between Ionic Solids Electrostatic Forces: a force of attraction that exists between ionic compounds due to opposite charges VERY strong Responsible for high MPs and high BPs

31. Hybridization Theory A theory that suggests that orbitals from atoms will merge and create bonding orbitals of equivalent energy Sigma: bonding that occurs by overlapping orbitals end to end Pi: bonding that occurs by overlapping orbitals side to side

32. Hybridization Theory Sigma Bonds (  )Pi Bond (  )

33. Sigma and Pi Bonds Single Bond: sigma only Double Bond: 1 sigma and 1 pi Triple: 1 sigma and 2 pi How many sigma and pi bonds are found in the following: N 2 C 2 H 4

34. Types of Hybridization sp sp 2 sp 3 sp 3 d sp 3 d 2

35. Note: this theory uses both orbitals involved in bonding and orbitals holding lone pairs so CH 4, NH 3, and CH 4 all have sp 3 hybridization Hybridization Theory

36. Molecular Orbital Theory Based on quantum mechanics Treats the electron as a moving object Relates to probability of location not exact location

37. Molecular Orbital Theory Bonding Orbital: area of high electron probability that has lower energy than the orbitals of the separate atoms Antibonding Orbital: area of high electron probability that has higher energy than the orbitals of the separate atoms

38. Molecular Orbital Theory Nonbonding Orbital: an orbital that does not contribute stability nor does it destabilize the molecule Open parking space near door Open parking space far from door Occupied parking space near door

39. Paramagnetism of Oxygen MOT explains the paramagnetism of oxygen