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Bonding Theories – Lewis Theory

Bonding Theories – Lewis Theory. One of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory Lewis Theory emphasizes how the valence electrons arranged among the atoms in a molecule

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Bonding Theories – Lewis Theory

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  1. Bonding Theories – Lewis Theory • One of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory • Lewis Theory emphasizes how the valence electrons arranged among the atoms in a molecule • Using Lewis Theory, we can draw models – called Lewis structures – that allow us to predict many properties of molecules such as molecular shape, size, polarity

  2. Rules to Draw Lewis Dot Structure • Draw skeletal structure with the central atom. The central atom of a molecule is usually the least electronegative atom (other than H). Connect the atoms to the central atom with a straight line representing a bond between the two atoms. • Find the total number of valence electrons from all the atoms in the given molecule or ion. • If the molecule is charged, add an electron for each negative charge and subtract an electron for each positive charge. • Complete the octets around each of the outer (surrounding) atoms except H . Distribute 6 electrons as 3 pairs. • How many electrons are remaining (Count each bond as 2 e-)? ______ • Place the remaining electrons as pairs on the central atom. • Final step: check if all the atoms are happy (H always will have 2 e-, the rest of the atoms will have eight unless there is an exception to octet rule – will discuss later) • If the octet on the central atom is not complete, try sharing lone pairs of outside atoms to form double or triple bonds.

  3. CCl4 HCN XeF4 H3PO4 (oxyacid H is bonded to O) SO32- C2H4 Practice – Drawing Lewis Structures

  4. Exceptions to Octet Rule • There are three types of ions or molecules that do not follow the octet rule: • Ions or molecules with an odd number of electrons (E.g. NO) • Ions or molecules with less than an octet. • Mostly when beryllium (Be), boron (B), and Aluminum (Al) as the central atom (E.g. BCl3) • Ions or molecules with more than eight valence electrons (an expanded octet). • Central atom is an element from third period and below (E.g. SF6)

  5. Practice: Exceptions to Octet Rule 1. The central atom in __________ does not violate the octet rule. a) SF4 b) XeF4 c)CF4 d) ICl4- e)KrF2 2. Which of the following includes an atom with more than eight electrons in its valence shell? a) CO2b) PCl5c) H2O d) NO2-1 e) HF 3. The molecule XeF2is an exception to the octet rule. Draw its Lewis dot structure and then identify the exception

  6. Resonance • Occurs when more than one valid Lewis structure can be drawn for a particular molecule or ion E.g. NO3- • The actual structure is the average of all the resonance structures. • Measurements of bond lengths suggest that all three N-O bond lengths are equal • Electrons are delocalized – they can move around the entire molecule

  7. Practice: Resonance Draw the resonance structures of HCO2- ion.

  8. Formal Charge • Formal charges are charges we assign to each atom in a Lewis dot structure (do not confuse with net charge of an ion) • Why do we need to calculate the formal charges? • It is useful to know which part of the molecule is charged? • In some cases we can draw several different Lewis structures which fulfill the octet rule for a compound. Which one is the most reasonable? • Formal charges will help you to determine which resonance structure is preferred over others • The sum of the formal charges on all the atoms in a molecule will be zero whereas for an ion it will be equal to the charge on the ion. 

  9. Formal Charge • Formal Charge on an atom = number of valence electrons on the free atom – (1/2 bonding electrons + nonbonding electrons on that atom) = Group Number of that atom – (dots + bonds around that atom) Formal charge on Br = Formal Charge on each O =

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