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Bonding 2008


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Bonding 2008

  1. 1. Bonding <ul><li>Occurs when two or more atoms join together to form a stable group </li></ul><ul><li>Affected by the # of valence electrons </li></ul><ul><li>There are several types of forces (BONDS) which hold the atoms together in these groups. </li></ul><ul><ul><li>I. METALLIC (share electrons”communally”) </li></ul></ul><ul><ul><li>II. IONIC (exchange electrons) </li></ul></ul><ul><ul><li>III. COVALENT (share electrons between 2 nuclei) </li></ul></ul>
  2. 2. I. Metallic Bonds <ul><li>A. occur between metal element atoms. </li></ul><ul><li>B. Since metal element atoms tend to lose electrons, the electrons are transferred from atom to atom (this is why metals conduct electricity). </li></ul><ul><li>1. Electron Sea Model - metal atoms are surrounded by a “sea” of electrons that moves around. </li></ul><ul><li>2.The metals share electrons communally </li></ul>
  3. 3. I. Metallic Bonds Cont’d <ul><li>C. Types of Metallic Bonds </li></ul><ul><li>1. Pure - the entire metal sample is composed of atoms of the same element. </li></ul><ul><li>2. Alloys - a substance that contains a mixture of elements. </li></ul><ul><ul><li>Examples: Sterling silver (93% Ag, 7% Cu); Bronze (Copper & Tin; Brass (Copper & Zinc; Steel (Iron & Carbon); Stainless Steel (Fe, Cr, Ni, C) </li></ul></ul>
  4. 4. II. Ionic Compounds
  5. 5. II. Ionic compounds <ul><li>Made of Cations and Anions bound together </li></ul><ul><ul><li>These elements join together because of electrostatic attractions (between oppositely charged IONS) </li></ul></ul><ul><ul><li>Remember ion formation? </li></ul></ul><ul><ul><li>a.cations are usually metallic elements (lose valence e-) </li></ul></ul><ul><ul><ul><li>Na (1 v.e - ) --> Na + </li></ul></ul></ul><ul><ul><ul><li>(e- configuration looks like Ne) </li></ul></ul></ul>
  6. 6. <ul><li>2b. Anions are usually nonmetallic elements (GAIN valence electrons) </li></ul><ul><li> Cl (7 v.e - ) --> Cl - </li></ul><ul><li> (e- configuration looks like Ar) </li></ul><ul><li>3. Named using Criss-Cross Method </li></ul>
  7. 7. Reminder: Valence electrons <ul><li>Valence electrons: electrons in the outermost shell that are involved in bonding </li></ul><ul><li>To determine the number of valence electrons, look at the group number on the periodic table </li></ul>
  8. 8. II. Ionic compounds 4. Salts such as NaCl form crystals a. Crystals = solids w/ repeating patterns (“crystal lattice”)held together by electrostatic attractions
  9. 9. III. CovalentCompounds <ul><li>Aka Molecular Compounds </li></ul><ul><li>These occur when two or more nonmetal elements (no metal elements in molecular compounds) bond together </li></ul><ul><li>These involve sharing electrons in orbitals between two nonmetal atoms </li></ul><ul><li>Named using prefixes </li></ul><ul><ul><li>(mono, di, tri, tetra, penta, hexa, septa, octa, nona, deca) </li></ul></ul>
  10. 10. E. Polarity <ul><li>1. Polarity occurs when a bond or molecule has a negative or positive end </li></ul><ul><li>2. Nonpolar molecules have a symmetrical distribution of e- so there is no + or - end </li></ul><ul><li>3. Polar Covalent bonds occur when the e- are not shared evenly </li></ul><ul><li>A. One end is slightly + and the other is slightly - </li></ul><ul><li>B. This slight charge is indicated using deltas (“∂+” and “∂-”) or an arrow </li></ul>
  11. 11. I. Lewis Structures for Molecular Cmpds <ul><li>A. Both nonmetal atoms “want” to gain e - to become stable </li></ul><ul><ul><ul><li>Electrons are shared in order to allow this to happen </li></ul></ul></ul><ul><ul><ul><li>The number of e - shared depends on the element (valence e-) </li></ul></ul></ul><ul><ul><ul><li>STABLE =NOBLE GAS CONFIGURATION! (Use P.T.) </li></ul></ul></ul>
  12. 12. I. Lewis Structures <ul><li>B. Types of Covalent Bonds </li></ul><ul><li>1. Single - ONE shared pr of e - </li></ul><ul><li>2. Double - TWO shared pr of e - </li></ul><ul><li>3. Triple - THREE shared pairs of e - </li></ul>
  13. 13. <ul><li>Lewis Structures </li></ul>C. Compounds formed with covalent bonds are neutral and follow the OCTET rule ! 1. All atoms are most stable when they have the electron configuration of a noble gases (THIS MEANS HAVING EIGHT* VALENCE ELECTRONS) *Exceptions exist
  14. 14. II. Steps for drawing Lewis Structures <ul><li>A.Sum the valence electrons. </li></ul><ul><li>B. Draw a skeleton structure with a bond between each pair of atoms. </li></ul><ul><li>C. Distribute the remaining electrons as lone pairs (unshared electrons) to fulfill the octet rule </li></ul><ul><li>exceptions </li></ul><ul><li>H =2 </li></ul><ul><li>Be = 4 </li></ul><ul><li>B = 6 </li></ul>
  15. 15. II. Steps for Drawing Lewis Structures D. Some General Guidelines 1. Often the central atom is the first one in the formula 2. Carbon likes to be in the middle with four bonds around it 3. Halogens like to be on ends with only one bond 4. Hydrogen MUST be on an end 5. Oxygen likes two bonds
  16. 16. Try these examples <ul><li>CH 4 </li></ul><ul><li>BeCl 2 </li></ul><ul><li>PCl 3 </li></ul><ul><li>CH 3 OH </li></ul><ul><li>CH 2 O </li></ul><ul><li>C 2 H 2 </li></ul><ul><li>O 2 </li></ul>
  17. 17. III. Special Cases <ul><li>Polyatomic Ions </li></ul><ul><li>1. Have covalent bonds within the ion </li></ul><ul><li>2. Must be bracketed with charge </li></ul><ul><li>Try: </li></ul><ul><ul><li>SO 3 2- </li></ul></ul><ul><ul><li>CN - </li></ul></ul>
  18. 18. III. Special Cases <ul><li>B. Exceptions to the Octet rule </li></ul><ul><li>1. There are some exceptions to the rule of having 8 electrons around each atom </li></ul><ul><li>a. electron deficient (less than 8 e-) </li></ul><ul><li>b. expanded octets (more than 8 e-) </li></ul>
  19. 19. III. B 1a. Electron deficient <ul><li>1. H always has only 2 electrons around it </li></ul><ul><li>2. Beryllium can have just 4 </li></ul><ul><li>3. Boron can have just 6 </li></ul><ul><li>Try BH 3 </li></ul>
  20. 20. III. B 1b Expanded Octets <ul><li>1. Third row/period and larger elements can exceed the octet if necessary </li></ul><ul><li>2. First and second period elements cannot do this!!! </li></ul><ul><li>3. When we must exceed the octet, extra electrons go on central atom. </li></ul><ul><li>Try these </li></ul><ul><li>SF 6 </li></ul><ul><li>PF 5 </li></ul><ul><li>XeF 4 </li></ul><ul><li>I 3 - </li></ul>
  21. 21. III. Special Cases <ul><li>C. Resonance </li></ul><ul><li>1. Sometimes there is more than one valid structure for an molecule or ion. This is called “resonance.” </li></ul><ul><li>2. Use double arrows to indicate it is the “average” of the structures. No one Lewis structure is correct, so you must show all possible. </li></ul><ul><li>Try: </li></ul><ul><li>NO 3 - </li></ul><ul><li>O 3 </li></ul>
  22. 22. Periodic Table Trends <ul><li>Ionization Energy </li></ul><ul><li>Electronegativity </li></ul><ul><li>Atomic Size </li></ul><ul><li>Ionic Size </li></ul><ul><li>Reactivity </li></ul>
  23. 23. I. Ionization Energy <ul><li>A. Amount of energy required to remove an electron from an atom and form an ion </li></ul><ul><li>B. Increases from left to right (across periods) </li></ul><ul><li>C. Increases from bottom to top (up groups) </li></ul>
  24. 24. Ionization Energy & the Periodic Table Elements with the highest ionization energies occur in the upper right corner (most chemically active nonmetals) Elements with the lowest ionization energies are in the lower left corner of the table (most chemically active metals)
  25. 25. II. Electronegativity <ul><li>Ability of atom to attract an electron to itself </li></ul><ul><ul><li>Impacts bond polarity </li></ul></ul><ul><ul><ul><li>Differences in electronegativity between atoms tell us how polar a bond is </li></ul></ul></ul><ul><li>B. Increases from left to right across periods </li></ul><ul><li>C. Increases from bottom to top down groups </li></ul><ul><li>D. Noble gases do not have electronegativity values (Why?) </li></ul>
  26. 26. III. Atomic size <ul><li>A. Atoms get smaller from left to right across periods Why? </li></ul><ul><li>B. Atoms get larger from top to bottom going down groups Why? </li></ul>
  27. 27. III. Atomic Size & the P.T. Size is decreasing Size is decreasing
  28. 28. IV. Ionic size <ul><li>A. Cations are smaller than the atoms they came from Why? </li></ul><ul><li>B. Anions are larger than the atoms they came from Why? </li></ul>