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FOUR THEORIES

FOUR THEORIES. 4.1 Lewis Dot Diagrams. Electron Distribution in Molecules. Electron distribution is depicted with Lewis dot structures Electrons are distributed as shared or bond pairs and unshared or lone pairs. G. N. Lewis 1875 - 1946. 8A. Valence Electrons. 1A. 7A. 2A. 3A. 4A.

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FOUR THEORIES

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  1. FOUR THEORIES 4.1 Lewis Dot Diagrams

  2. Electron Distribution in Molecules • Electron distribution is depicted with Lewis dot structures • Electrons are distributed as shared or bond pairs and unshared or lone pairs. G. N. Lewis 1875 - 1946

  3. 8A Valence Electrons 1A 7A 2A 3A 4A 5A 6A Number of valence electrons is equal to the Group number.

  4. Lewis Structures • Shows how valence electrons are arranged among atoms in a molecule. • Reflects central idea that stability of a compound relates to noble gas electron configuration.

  5. •• H Cl • • •• Unshared or lone pair (LP) shared or bond pair Bond and Lone Pairs Electrons are distributed as shared or bond pairs and unshared or lone pairs. This is a LEWIS ELECTRON DOT structure.

  6. Ionic Compounds • Use arrows to represent the transfer of electrons. • The final structure will contain ions.

  7. Octet-Rule Covalent Compounds • Simple covalent compounds such as CH4 may be done by pairing electrons from different atoms. The pairing is shown with circles. • More complex covalent compounds and polyatomic ions should be drawn using the Need-Have Method.

  8. Drawing Lewis Dot Structures • Count number of NEEDED electrons. Don’t forget to add an electron for every negative charge and subtract an electron for every positive charge. • each element on the periodic table needs (8 e) except Hydrogen (2 e) • Count number of VALENCE electrons • Number of BONDING e = NEEDED-VALENCE • Determine the central atom. • the unique atom (only one of it) is the central atom • HXO# H is written besides O • When all else fails, put the least electronegative element in the middle • Arrange the rest  a skeleton • Oxygen rarely bonds to itself except in: O2 and O3; peroxides H2O2; superoxides NaO2 • Fluoride never bonds to more than one atom. • Calculate the Formal Charge of each ELEMENT • Distribute the remaining electrons in pairs around the atoms, trying to satisfy the octet rule. Assign them to the most electronegative atom first. • If you have extra electrons and all the atoms have an octet, put the extra electrons on the central atom in pairs. If the central atom has an atomic number greater than fifteen, you are allowed to have more than eight electrons around it.

  9. Completing a Lewis Structure -CH3Cl • Make carbon the central atom • Add up available valence electrons: HAVE • C = 4, H = (3)(1), Cl = 7 Total = 14 • NEED = 3(2) + 8 + 8 = 22 • BONDS = (22-14)/2 = 4 • Join peripheral atoms • to the central atom • with electron pairs. H .. .. .. H .. C .. Cl • Complete octets on • atoms other than • hydrogen with remaining • electrons .. .. H

  10. Resonance Structures(Exception # 1) • For some molecules, there are multiple ways of placing the electrons between the atoms. • Structures that differ only in the arrangement of the electrons are called RESONANCE STRUCTURES. • Resonance structures are indicated using a double headed arrow.

  11. Example:Use the concept of formal charges to determine the most stable Lewis dot structure for CO2. • It may not have occurred to you, but we can draw five different Lewis dot structures for CO2 that satisfy the octet rule. • WHICH ONE? • Calculate Formal Charge of each ELEMENT FC = # of v.e – (# of lone pairs electrons + 1/2 # of bonding electrons) I see that the formal charges are minimized when the carbon atom is in the middle and forms double bonds to each oxygen, structure (a).

  12. N2O • both (a) and (d) do an equally good job of minimizing the formal charges. • (d) is more stable because the negative formal charge is on oxygen rather than nitrogen (recall that oxygen is more electronegative than nitrogen). • Note: the formal charges add up to zero in every case - the formal charges must add up to equal the charge on the molecule (zero for neutral molecules).

  13. Non-Octet Compounds(They do not follow the rule.(Exception # 2) ) • Some compounds will contain central atoms that do not follow the octet rule. • The four possibilities for non-octet compounds are: • Where more than 4 atoms are bonded to the central atom such as PCl5. • A noble gas is participating in bonding such as XeF4. • Where the central atom has less than 8 valence electrons such as BH3. • Where molecules contain an odd number of nonbonding electrons such as NO.

  14. Drawing Non-Octet Diagrams • To draw non-octet compounds count the electrons that you have available in the structure. • Draw the diagram placing the central atom in the middle and the surrounding atoms around it. • Make a single bond between the central atom and each surrounding atom. • Make sure that all surrounding atoms have a full valence shell. • Place any extra electrons on the central atom (Make sure you only use as many electrons as you have)

  15. Example • PCl5 • IF5 • SF6

  16. Coordinate Covalent Molecules (Exception # 3) • A bond in which the 2 electrons shared by a pair of atoms belonged originally to only one of the atoms.

  17. Draw Lewis Dot Diagrams for the following compounds. Determine whether a bond is polar or non-polar covalent. CH4, HCl, NH3, O2, SF2, H2O, CO2, N2, C2H2 ,H2CO Which of these compounds are POLAR (dissolve in water) and which are NON-POLAR (won’t dissolve in water)? PRACTICE !!! http://chemsite.lsrhs.net/d_bonding/flashLewis.html Building Molecules with Lewis Dot structures

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