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Chemical Bonding

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Chemical Bonding

  1. 1. The Microscopic Structure of Matter <ul><li>Types of Chemical Forces </li></ul><ul><li>Formation of Ionic Compounds </li></ul><ul><li>Properties of Ionic and Covalent Compounds </li></ul><ul><li>Formation of Covalent Compounds </li></ul><ul><li>Lewis Structures - introduction </li></ul><ul><li>Molecular Geometry - introduction </li></ul><ul><li>Molecular Polarity – introduction </li></ul><ul><li>Making Lewis structures & determining their characteristics </li></ul>
  2. 2. Between species from same parent species Between species from different parent species
  3. 3. “ Bonding”: General Ideas and Types <ul><li>General Principle: Atoms interact with atoms in such a way as to achieve the best SET of electronic configurations for all the atoms involved </li></ul><ul><li>Chemical “Bond” - The attractive force that holds the atoms of a substance together </li></ul><ul><li>Ionic “bonding” = Electrostatic attraction = Exchanging electrons as a strategy to get more stable configurations </li></ul><ul><li>Covalent (molecular) bonding = Sharing orbitals (with or without electrons) as a strategy to get more stable configurations </li></ul><ul><li>Metallic bonding: Pooling electrons in orbitals as a strategy to get more stable configurations </li></ul><ul><ul><li>common metals: Fe, Cu, Sn, Pb (elements);bronze (an alloy of two metals) </li></ul></ul><ul><ul><li>Will not cover </li></ul></ul>
  4. 4. Principles of Bond Formation <ul><li>Only the valence electrons of atoms are available for chemical bonding </li></ul><ul><ul><li>Valence electrons = electrons in outermost (valence) shell of an atom </li></ul></ul><ul><li>Certain arrangements of electrons are more stable than others </li></ul><ul><ul><li>He, Ne, Ar, Kr, Xe do NOT form compounds </li></ul></ul><ul><li>Chemical bonds are formed because the energy is lowered upon bond formation </li></ul><ul><ul><li>The number of favorable factors outweighs the number of unfavorable ones! </li></ul></ul>
  5. 5. Rationale for Ionic “Bonding” <ul><li>Atoms of elements exchange electrons such as to produce ions: </li></ul><ul><ul><li>that have electronic configurations of the nearest noble gas if at all possible or that have other stable electronic configurations </li></ul></ul><ul><ul><ul><li>“ duet” for He, “octet” for other noble gases </li></ul></ul></ul><ul><ul><li>that produce salts with strong ion-ion attractions (lattice forces) </li></ul></ul><ul><li>There is a limit to the number of electrons that can be donated or accepted (+3 or - 3) </li></ul>
  6. 6. Electron Dot Structures <ul><li>Electron dot structure help to determine the number of electrons that can be lost or gained to reach a stable octet (duet) </li></ul><ul><li>What would Li, Be do? </li></ul><ul><li>Would B lose 3 of its 5 electrons; would Al be more or less reluctant to lose 3 electrons? </li></ul><ul><li>What would F do readily, O less readily and N even less readily? </li></ul><ul><li>What is C’s dilemma? </li></ul>
  7. 7. Formation of Ions Repeat but with electronic configurations Na + + Na + Cl Cl - Cl - Na + Cl Na e - (1s 2 2s 2 2p 6 3s 2 3p 5 ) (1s 2 2s 2 2p 6 ) (1s 2 2s 2 2p 6 3s 1 ) (1s 2 2s 2 2p 6 3s 2 3p 6 )
  8. 8. Formation of Ions – another example Ca + + Ca 2+ O O 2- O 2- Ca 2+ O Ca 2 e - (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 ) (1s 2 2s 2 2p 6 ) (1s 2 2s 2 2p 4 ) (1s 2 2s 2 2p 6 3s 2 3p 6 )
  9. 9. Structures of Ionic Compounds <ul><li>Crystalline state </li></ul><ul><ul><li>Network of cations and anions with overall zero charge </li></ul></ul><ul><ul><li>Arrangement of cations and anions depends on size and charge of ion </li></ul></ul><ul><ul><ul><li>Different types of macroscopic crystals as a consequence </li></ul></ul></ul><ul><ul><li>Size of network can vary (small and large crystals) </li></ul></ul><ul><li>Dispersed state </li></ul><ul><ul><li>Requires “dispersing agent” </li></ul></ul><ul><ul><ul><li>Solvent (usually water); heat </li></ul></ul></ul><ul><ul><li>Solutions </li></ul></ul><ul><ul><ul><li>Requires that water can pull apart crystal  LATER </li></ul></ul></ul><ul><ul><li>Molten salts </li></ul></ul><ul><ul><ul><li>Have to be above melting point of salt </li></ul></ul></ul>
  10. 10. Crystals of Cesium Halides  CsF CsCl  CsI   CsBr
  11. 11. Crystals of copper salts
  12. 12. Formulas of Binary Ionic Compounds <ul><li>Binary Ionic Compounds </li></ul><ul><ul><li>Two different ions, one metal, one nonmetal </li></ul></ul><ul><ul><li>Metal ion  + (cation) </li></ul></ul><ul><ul><li>Nonmetal ion  - (anion) </li></ul></ul><ul><ul><li>Total + = Total – </li></ul></ul><ul><li>Formulas of Binary Ionics </li></ul><ul><ul><li>Metal ion first, nonmetal follows </li></ul></ul><ul><ul><li>Subscripts indicate number of each ion used to get balance of charge </li></ul></ul><ul><li>Types of Binary Ionic Compounds </li></ul><ul><ul><li>Fixed charged metals </li></ul></ul><ul><ul><ul><li>Metals from columns 1, 2, 3 & 13 </li></ul></ul></ul><ul><ul><li>Variable charged metals </li></ul></ul><ul><ul><ul><li>All other metals </li></ul></ul></ul><ul><ul><li>Charges on monatomic nonmetals ions are CONSTANT </li></ul></ul>* Method for Making Compounds O 2- e - K K + 2 O 2 e - 2 2 = K 2 O
  13. 13. Properties of Ionic Compounds <ul><li>Ionic compounds have very high melting points . They are always solids at room temperature. </li></ul><ul><ul><li>Attractions between cations and anions are very strong </li></ul></ul><ul><ul><li>The larger the charges, the greater the attraction, the higher the melting point (usually) </li></ul></ul><ul><ul><li>Usually hard and brittle </li></ul></ul><ul><li>Ionic compounds dissociate into cations and anions when melted. </li></ul><ul><ul><li>They are electrical conductors when melted. </li></ul></ul><ul><li>Ionic compounds dissociate into cations and anions IF they can dissolve in water. </li></ul><ul><ul><li>Aqueous solutions with dissolved ions ( electrolytes ) are electrical conductors </li></ul></ul><ul><ul><li>If an ionic compound cannot dissolve in water,, the crystal lattice forces must be strong </li></ul></ul><ul><ul><li>Ionic compounds with larger charges tend not to dissolve in water  LATER </li></ul></ul>
  14. 14. Rationale for Covalent Bonding <ul><li>NOTE: treatment covalent bonding, geometry, polarity (4.5, 4.7, 4.8) in your book is lame  please go with this treatment instead </li></ul><ul><li>Atoms of electron rich elements share orbitals (with or without electrons) such as to produce: </li></ul><ul><ul><li>electronic configurations of the nearest noble gas (in most cases) </li></ul></ul><ul><ul><li>molecular compounds with strong bonds (in all cases) </li></ul></ul><ul><li>An electron rich element is one that has half or more of its outer valence shell filled </li></ul><ul><ul><li>all nonmetals & hydrogen </li></ul></ul><ul><li>“ Sharing” of orbitals can be: </li></ul><ul><ul><li>Equal  each atom contributes the same number of electrons </li></ul></ul><ul><ul><li>Unequal  atoms contribute different numbers of electrons </li></ul></ul><ul><ul><li>“ Parasitic” (co-ordinate covalent)  one atom contributes all the electrons in a bond </li></ul></ul>
  15. 15. VSEPR not in your textbook <ul><li>Atomic vs molecular orbitals </li></ul><ul><ul><li>Atoms of each element are born with atomic orbitals </li></ul></ul><ul><ul><ul><li>Just fine if you are born perfect and stay single !!! </li></ul></ul></ul><ul><ul><li>Orbitals need to adjust when orbitals of one atom interact with orbitals of another atom </li></ul></ul><ul><ul><ul><li>Hybridization (making friends, marriage) </li></ul></ul></ul><ul><ul><ul><li>Formation of molecular orbitals </li></ul></ul></ul><ul><ul><ul><ul><li>Still only 2 electrons per orbital </li></ul></ul></ul></ul><ul><li>V alence S hell E lectron P air R epulsion </li></ul><ul><ul><li>Orbitals used in bonding as well as “leftover” orbitals stay as far away from each other as possible in order to avoid electron repulsions </li></ul></ul><ul><ul><li>Need to know geometry and angles in order to know how to spread out orbitals </li></ul></ul>
  16. 16. The Octet and Duet Rules <ul><li>For H the nearest noble gas is helium </li></ul><ul><ul><li>H in covalent compounds will have a single bond sharing 2 electrons = DUET RULE </li></ul></ul><ul><li>For other nonmetals, the nearest noble gas has a completed outer valence shell of 8 electrons </li></ul><ul><ul><li>these elements will form enough bonds such that afterwards the total number of electrons is 8 = OCTET RULE </li></ul></ul><ul><ul><li>total number of electrons can be any combination of shared and unshared electrons </li></ul></ul><ul><ul><li>Sharing does not have to be equal or symmetric </li></ul></ul><ul><li>Breaking the octet rule occurs  we won’t do </li></ul>
  17. 17. Types of Covalent Bonds <ul><li>Single Bonds </li></ul><ul><ul><li>A:B (A-B) - one pair of electrons shared by 2 atoms </li></ul></ul><ul><ul><li>Overlap of two orbitals (one from each atom) </li></ul></ul><ul><li>Double Bonds </li></ul><ul><ul><li>A::B (A=B) - two pairs of electrons shared by 2 atoms </li></ul></ul><ul><ul><li>Overlap of 4 orbitals (two from each atom) </li></ul></ul><ul><li>Triple Bonds </li></ul><ul><ul><li>A:::B ( A B) - three pairs of electrons shared by 2 atoms </li></ul></ul><ul><ul><li>Overlap of 6 orbitals (three from each atom) </li></ul></ul><ul><li>IMPORTANT: NOTICE THAT WHERE THE ELECTRONS COME FROM IS NOT SPECIFIED </li></ul>
  18. 18. Lewis Structures <ul><li>Lewis Structure - A picture indicating the way in which the valence electrons are distributed in a molecule </li></ul><ul><li>Electrons in bonds are represented by lines </li></ul><ul><ul><li>no dots are placed between atoms! </li></ul></ul><ul><li>Nonbonding electrons are represented by pairs of dots </li></ul><ul><li>Each atom is given its preferred number of bonds, if possible </li></ul><ul><ul><li>Preferred number of bonds = Number of unpaired dots in the dot structure of the free atom = valence </li></ul></ul><ul><ul><li>If the preferred number of bonds is less or greater than the preferred number, this difference has to be made up by neighboring atoms </li></ul></ul>
  19. 19. Table 4.11 Preferred Number of Bonds
  20. 20. Arrangements of Groups vs Molecular Geometry <ul><li>ARRANGEMENTS = Name for the number of groups around a central atom </li></ul><ul><ul><li>Groups are: </li></ul></ul><ul><ul><ul><li>connected atoms (by bonds due to overlapping orbitals) </li></ul></ul></ul><ul><ul><ul><li>unshared pair of electrons </li></ul></ul></ul><ul><ul><li>Groups are arranged about a central atom in such a way as to keep the groups as far apart as possible </li></ul></ul><ul><li>MOLECULAR GEOMETRIES = Name that describes how atoms are arrayed around a central atom </li></ul><ul><ul><li>Molecular geometries are subclasses of arrangements </li></ul></ul>
  21. 21. LINEAR TETRAHEDRAL PLANAR ARRANGEMENT Two double or one triple bond One double bond All single bonds
  22. 22. Example 1: Correlation between Lewis structure, arrangement and molecular geometry
  23. 23. Example 2: Correlation between Lewis structure, arrangement and molecular geometry
  24. 24. Example 3: Correlation between Lewis structure, arrangement and molecular geometry
  25. 25. Electronegativity The most important idea in chemistry <ul><li>Electronegativity - A measure of the ability of a nucleus in an atom to attract electrons in a bond toward itself. </li></ul><ul><li>The most electronegative elements would be those that have large effective nuclear charge </li></ul><ul><li>Elements can be given numerical values indicating relative ability </li></ul><ul><li>Electrons in bonds will be “polarized” to the more electronegative atom in the bond </li></ul>
  26. 26. Figure 4.5
  27. 27. Bond Polarity <ul><li>Nonpolar Bond - A covalent bond in which the electrons are shared equally between the bonded atoms. </li></ul><ul><ul><li>Bonded atoms have the same electronegativity </li></ul></ul><ul><li>Polar Bond - A covalent bond in which there is unequal sharing of electrons between the bonded atoms. </li></ul><ul><ul><li>Bonded atoms have unequal electronegativities </li></ul></ul><ul><ul><li>The atom having the higher electronegativity will have a slight negative charge (  -) </li></ul></ul><ul><ul><li>The atom having the lower electronegativity will have a slight positive charge. (  +) </li></ul></ul>
  28. 28. Table 4.14
  29. 29. Table 4.15
  30. 30. Molecular Polarity <ul><li>Nonpolar Molecule - A molecule that does not have a net positive end and a net negative end. </li></ul><ul><ul><li>All bonds are either nonpolar or polarities in bonds cancel out because of symmetry </li></ul></ul><ul><li>Polar Molecule - A molecule that has a slight positive charge on one end and a slight negative charge on the other end. </li></ul><ul><ul><li>Bonds are polar and do not cancel each other out = molecules are asymmetric </li></ul></ul>
  31. 31. A “Formula” for Success for Structure, Geometry and Polarity <ul><li>Formula -------->Lewis structure </li></ul><ul><ul><li>Tool = follow the rules for making a Lewis structure </li></ul></ul><ul><li>Lewis Structure -------> Arrangement </li></ul><ul><ul><li>Tool = Count the number of groups around the central atom and assign the arrangement name </li></ul></ul><ul><ul><ul><li>groups = atoms (however connected, single, double or triple) and electron pairs </li></ul></ul></ul><ul><li>Arrangement -------> Geometry </li></ul><ul><ul><li>Tool = Count the number of atoms attached to the central atom (NOT the number of bonds) </li></ul></ul><ul><ul><ul><li>staying in the same arrangement family, assign the geometry subclass name </li></ul></ul></ul><ul><li>Arrangement -------> Bond Angles </li></ul><ul><ul><li>Tool = assign all bond angles that connect atoms to each other </li></ul></ul><ul><ul><ul><li>do NOT leave the arrangement family </li></ul></ul></ul><ul><li>Arrangement -----> Polarity (applies only to neutral species) </li></ul><ul><ul><li>Tool = inspect the molecule in 3D and determine whether one end of the molecule is different than another </li></ul></ul><ul><ul><ul><li>if there is a difference, the molecule is polar, otherwise it is nonpolar </li></ul></ul></ul>
  32. 32. Systematic approach to making Lewis structures & determining their characteristics <ul><li>Making the Lewis structure </li></ul><ul><ul><li>All species will obey octet rule (except for H) </li></ul></ul><ul><ul><li>No non octet species (such as in your text) </li></ul></ul><ul><li>Determining the correct arrangement </li></ul><ul><ul><li>Bond angles always the same for a given arrangement </li></ul></ul><ul><li>Determining the geometry SUBCLASS </li></ul><ul><li>Determining whether the species has overall polarity </li></ul>
  33. 33. Lewis Structures – Facts to start with <ul><li>Determine the number of “normal&quot; covalent bonds by knowing the valence of element </li></ul><ul><li>F, Cl, Br, I = 1 O, S, Se, Te = 2 </li></ul><ul><li>N, P, As = 3 C, Si = 4 </li></ul><ul><li>Some species are charged: this has to be accounted for </li></ul><ul><ul><li>Reasons for charge  Later </li></ul></ul><ul><li>When you make a structure, do NOT worry about “breaking” the rules for bonding for certain elements, either too few or too many </li></ul><ul><ul><li>Some atoms will have more than the usual number, others less, BUT these compensate each other </li></ul></ul><ul><ul><li>Have faith in the rules  you will get the right answer </li></ul></ul>
  34. 34. <ul><li>Determine the valence for each of the elements present in the formula and multiply by the number of each of these element present </li></ul><ul><li>Divide this number by 2 since the valence is how many bonds you need for each atom and bonds are shared between atoms </li></ul><ul><li>Place the least electronegative atom (but never H) in the center and place the other atoms around this central atom </li></ul><ul><li>Put bonds in such a way as to give each atom its preferred number of bonds (valence), if possible . </li></ul><ul><ul><li>You may have to double and triple up bonds </li></ul></ul><ul><ul><li>If an atom has too many bonds, a neighboring atom will have too few: this is because of unequal sharing but overall the normal bonding number will be correct </li></ul></ul><ul><li>Add lone pairs to complete the octet as needed around each atom (but not H!) </li></ul><ul><li>Check your structure </li></ul>Lewis Structures – Neutral Species
  35. 35. Lewis Structures – Charged species <ul><li>Determine the valence for each of the elements present in the formula and multiply by the number of each element present </li></ul><ul><li>Subtract one for each overall negative charge and add one for each overall positive charge. This takes care of the shortage or excess of bonds due to charge </li></ul><ul><li>Divide number you get by 2 since the valence is how many bonds you need for each atom & bonds are shared between atoms </li></ul><ul><li>Place the least electronegative atom (but never H) in the center and place the other atoms around this central atom </li></ul><ul><li>Put bonds in such a way as to give each atom its preferred number of bonds (valence), if possible . </li></ul><ul><ul><li>You may have to double and triple up bonds </li></ul></ul><ul><ul><li>There will be a “missing” bond from the overall normal number for every negative charge and an extra bond from the normal number for every positive charge </li></ul></ul><ul><li>Add lone pairs to complete the octet as needed around each atom (but not H!) </li></ul><ul><li>Check your structure </li></ul>
  36. 36. Making Lewis Structures <ul><li>F 2 , O 2 , N 2 </li></ul><ul><li>HCl, H 2 O, NH 3 , SBr 2 , PF 3 , SiCl 4 </li></ul><ul><li>CO 2 , SO 2 , HCN </li></ul><ul><li>NH 4 + ,NO 3 - , NO 2 - , CO 3 2- , SO 4 2- </li></ul><ul><li>ClO 4 - , BrO 2 - </li></ul><ul><li>H 2 SO 4 , HCO 3 - , CH 2 O, SCN - , OCS </li></ul>*
  37. 37. Determining Arrangements & Bond Angles around central atom(s) – Step by Step <ul><li>Count the number of groups </li></ul><ul><ul><li>GROUPS = ELECTRONS PAIRS AND BONDING GROUPS </li></ul></ul><ul><ul><li>ONE ELECTRON PAIR = ONE GROUP </li></ul></ul><ul><ul><li>ONE BONDING GROUP = ONE GROUP </li></ul></ul><ul><ul><ul><li>A bonding group = a single or a double or a triple bond </li></ul></ul></ul><ul><li>Arrangement & Bond Angle Chart </li></ul><ul><ul><li>2 groups = linear = 180 </li></ul></ul><ul><ul><li>3 groups = (trigonal) planar = 120 </li></ul></ul><ul><ul><li>4 groups = tetrahedral = 109.5 </li></ul></ul>*
  38. 38. <ul><ul><li>IN DETERMINING GEOMETRY DO NOT LEAVE ARRANGEMENT </li></ul></ul><ul><ul><li>Linear  all geometries linear </li></ul></ul><ul><ul><li>Planar </li></ul></ul><ul><ul><ul><li>3 atoms  planar </li></ul></ul></ul><ul><ul><ul><li>2 atoms and an electron pair  bent </li></ul></ul></ul><ul><ul><li>Tetrahedral </li></ul></ul><ul><ul><ul><li>4 atoms  tetrahedral </li></ul></ul></ul><ul><ul><ul><li>3 atoms and an electron pair  pyramidal </li></ul></ul></ul><ul><ul><ul><li>2 atoms and two electron pairs  bent </li></ul></ul></ul>Determining Geometries around central atom(s) – Step by Step *
  39. 39. Rules for Overall Polarity of Neutral Species <ul><li>Arrangement of Groups </li></ul><ul><li>Linear </li></ul><ul><li>Trigonal Planar </li></ul><ul><li>Tetrahedral </li></ul><ul><li>Polarity </li></ul><ul><li>if atoms on either side of the central atom are different, the species is polar </li></ul><ul><li>if any group around central atom is different, the species is polar </li></ul><ul><ul><li>be careful with resonance </li></ul></ul><ul><li>if any group is different, the species is polar </li></ul>ALWAYS USE THE ARRANGEMENT TO DETERMINE POLARITY *
  40. 40. LINEAR TETRAHEDRAL PLANAR ARRANGEMENT Two double or one triple bond One double bond All single bonds 180 120 109.5 BOND ANGLE SUBCLASS THE ULTIMATE USEFUL SUMMARY SLIDE One of the S-O bonds is double

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