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Chemical Bonding and Nomenclature

Chemical Bonding and Nomenclature. Adapted from Paul Surko. What is Bonding????????. 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 are Compounds?.

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Chemical Bonding and Nomenclature

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  1. Chemical Bonding and Nomenclature Adapted from Paul Surko

  2. What is Bonding???????? 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.

  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. Na becomes Na+ O becomes O-2 Cl becomes Cl- Al becomes Al+3 N becomes N-3 Mg becomes Mg+2 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 + Cl- = CaCl2 Ca+2 + O-2= CaO Na+ + O-2 = Na2O Al+3 + S-2 = Al2S3 Ca+2 + N-3 = Ca3N2 You try these! LiBr Li+ + Br- = MgF2 Mg+2 + F- = (NH4)3PO4 AlI3 NH4+ + PO4-3 = Al+3 + I- = Sr3P2 Sr+2 + P-3 = Not NH43PO4 KCl K+ + Cl- =

  8. NomenclatureNaming 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 _________ide Sodium ChlorineSodium Chloride NaCl Metals (Transition Metals) and Non-Metals Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III BromineIron (III) Bromide FeBr3Compare with Iron (II) Bromide FeBr2

  9. Nomenclature--Naming of Ionic Compounds with TM Metals (Transition Metals) and Non-MetalsOlder System Ferrous BromineFerrous Bromide FeBr2Compare with Ferric Bromide FeBr3 Metal (Latin) _______ + ous or ic + Non-Metal ________ide

  10. Let’s Practice! Name the following. CaF2 Calcium Flouride K2S Potassium Sulfide CoI2 Cobalt (II) Iodide or Cobaltous Iodide SnF2 Tin (II) Flouride or Stannous Flouride SnF4 Tin (IV) Flouride or Stannic Flouride OF2 Oxygen diflouride CuI2 Copper (II) Iodide or Cupric Iodide CuI Copper (I) Iodide or Cuprous Iodide SO2 Sulfur dioxide SrS Strontium Sulfide Lithium Bromide LiBr

  11. Ammonium……………... Nitrate…………………… Permanganate…………. . Chlorate………………… Hydroxide………………. Cyanide…………………. Sulfate…………………... Carbonate………………. Chromate……………….. Acetate………………….. Phosphate………………. Polyatomic Ions(partial list from page 195 (193 2nd edition)) • NH4+ • NO3- • MnO4- • ClO3- • OH- • CN- • SO4 2 - • CO32- • CrO42- • C2H3O2- • PO43-

  12. Lets Practice! Na2CO3 Sodium carbonate Potassium permanganate KMnO4 Sodium hydroxide NaOH Copper (II) sulfate or Cupric sulfate CuSO4 Lead (II) chromate or Plubous chromate PbCrO4 ammonia NH3

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

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

  15. 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.

  16. Rules to Write Dot Structures Write a skeleton molecule with the lone atom in the middle (Hydrogen can never be in the middle) Find the number of electrons needed (N) (8 x number of atoms, 2 x number of H atoms) Find the number of electrons you have (valence e-'s) (H) Subtract to find the number of bonding electrons (N-H=B) Subtract again to find the number of non-bonding electrons (H-B=NB) Insert minimum number of bonding electrons in the skeleton between atoms only. Add more bonding if needed until you have B bonding electrons. 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.

  17. Let's Try it! S N H B NB E H O H Water H2O 2 x 2 = 4 for Hydrogen1 x 8 = 8 for Oxygen4+8=12 needed electrons 12 N - 8 H 2 x 1 = 2 for Hydrogen1 x 6 = 6 for Oxygen You have 8 available electrons 4 B - 4 NB 12 - 8 = 4 bonding electrons H:O:H 8 – 4 = 4 non-bonding electrons .. H:O:H●● .. H:O:H●●

  18. Let's Try it! S N H B NB E HH N H Ammonia NH3 3 x 2 = 6 for Hydrogen1 x 8 = 8 for Nitrogen6+8=14 needed electrons 14 N - 8 H 3 x 1 = 3 for Hydrogen1 x 5 = 5 for Nitrogen You have 8 available electrons 6 B - 2 NB 14 - 8 = 6 bonding electrons H ..H:N:H 8 – 6 = 2 non-bonding electrons H .. H:N:H●● H .. H:N:H●●

  19. Let's Try it! S N H B NB E Carbon Dioxide CO2 O C O 1 x 8 = 8 for Carbon2 x 8 = 16 for Oxygen8+16=24 needed electrons 24 N - 16 H 1 x 4 = 4 for Carbon2 x 6 = 12 for Oxygen You have 16 available electrons 8 B - 8 NB 24 - 16 = 8 bonding electrons O::C::O 16 – 8 = 8 non-bonding electrons .. .. O::C::O●● ●● .. .. O::C::O●● ●●

  20. Let's Try it! OO C O Carbonate CO3-2 S N H B NB E 3 x 8 = 24 for Oxygen1 x 8 = 8 for Carbon24+8=32 needed electrons 32 N 24 H - 3 x 6 = 18 for Oxygen1 x 4= 4 for CarbonYou have 22 + 2 more available e-'s 8 B - 16 NB 32 - 24 = 8 bonding electrons O ..O::C:O 24 – 8 = 16 non-bonding electrons ..:O: ..:O: .. .. .. O::C: O:●● ●● -2 .. .. .. O::C: O:●● ●●

  21. Nomenclature of Covalently Bonded Compounds--Molecules Non-Metals and Non-Metals Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc. CO2 Carbon dioxide CO Carbon monoxide PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride N2O5Dinitrogen pentoxide CS2 Carbon disulfide

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

  23. VSEPR Theory • Why is H2O bent and CO2 linear? • O in water has lone pairs causing bending whereas the C in carbon dioxide does not

  24. 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

  25. 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

  26. Polarity of Molecules Cont’d • Examples on the Board • H2O • CH4 • CO2 • NH3 • *BF • Why does water exist as a liquid at STP and carbon dioxide exists as a gas at STP?

  27. 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

  28. 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

  29. 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

  30. 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

  31. Hybridization Theory Sigma Bonds (s) Pi Bond (p)

  32. 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: • N2 • C2H4

  33. Types of Hybridization • sp • sp2 • sp3 • sp3d • sp3d2

  34. Hybridization Theory • Note: this theory uses both orbitals involved in bonding and orbitals holding lone pairs • so CH4, NH3, and CH4 all have sp3 hybridization

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

  36. 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

  37. 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

  38. Paramagnetism of Oxygen • MOT explains the paramagnetism of oxygen

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