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Chapter 8

Chapter 8. Chemical Bonding. Chemical Bond: __________________________ ________________________________________________________________________________ The type of bonding is determined by the way the valence e-’s are redistributed. Molecular or Covalent Bonding.

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Chapter 8

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  1. Chapter 8

  2. Chemical Bonding Chemical Bond: __________________________ ________________________________________________________________________________ The type of bonding is determined by the way the valence e-’s are redistributed.

  3. Molecular or Covalent Bonding Molecule:______________ ______________________________________________ e.g. H2O Diatomic Molecule: molecules containing __________ atoms. E.g. O2 , CO , HF, NO

  4. More definitions… Octet Rule: where chemical compounds tend to form, such that ______________________________ _________________________________________ This is done by _______________________________ (becoming an ion or entering a covalent bond)

  5. e.g. Fluorine gas exists as F2 . F –- F Each F atom has achieved a stable octet by _______________of electrons. (do the same with: O2 , PF3 )

  6. How many electrons would it need to fill an octet? Is that possible? e.g. HCl Chlorine has a stable octet, but H does not. That’s because there are _____________________________________________________________________ i.e. __________________

  7. Bonding • When atoms bond they ______________ _________________ • In dot notation this is represented as two dots between symbols, one from each atom

  8. (Do F2 , PF3 , HCl ) The unshared pairs of electrons are also known as ________________. The _________________________________________. (do F2 , & PF3 ) These representations are known as: Lewis Structures:_____________________________ ________________________________________________________________________________________

  9. The dots representing the lone pairs can also be dropped. The new representation is known as ____________ ____________________

  10. A single shared pair is known as a _____________. Let’s consider O2 : (Diagram) The sharing of ____________________ between 2 atoms is known as a ________________.

  11. Let’s consider N2 : (Diagram) The sharing of ____________________ between 2 atoms is known as a ________________ Double & Triple Bonds are also known as: Multiple Bonds.

  12. We still haven’t explained why carbon can form 4 bonds instead of 2…___________________ Let’s look at Carbon (6): (ec, orbital diag., Lewis, & Structural) It makes sense to assume that Carbon forms 2 covalent bonds.

  13. But when Carbon bonds with other atoms, a special thing happens. The 2s & 2p merge together to form an _____________. Now apply Hund’s rule. (Diagram) So now, Carbon has 4 single bonds. Hybridization also applies to Be, B, & Si.

  14. Let’s review with some examples: (Central atom is the least EN atom or C) (e.g’s of NH3 , HCN , C2H6 , C2H4 , C2H2) Show Lewis structure

  15. Two Types of Bonds Ionic: results from 1 atom giving up its valence e-’s (cation) & transferring them to another atom (anion) e.g. NaCl Covalent: _____________________ ______________________________ • Most bonds are between these extremes!

  16. Non-Polar Covalent Non-Polar Covalent: the bonding valence e-’s are ________________by the atoms resulting in _____________________of electrical charge. e.g. N2

  17. Polar covalent bond Polar Covalent: the bonding valence e-’s are more strongly attracted to the more EN atom resulting in an ___________________________________. It is still sharing, not a transfer like in ionic. e.g. CO2

  18. So: Chemical Bonding Ionic Covalent Polar Non-Polar The type of bonding can be determined simply by the _____________________ ______________________________(∆EN) of the 2 atoms. 6.2 p 177.

  19. E.g.’s: A H-F molecule has an EN difference of: 4.0(for F)– 2.1(for H)= For Na-Cl the EN difference is: 3.0(for Cl)– 0.9(for Na)= For H-H (H2) the EN difference is: 2.1(for H)– 2.1(for H)=

  20. The difference tells you what type of bonding that is occurring: > 1.7 = Ionic < 0.3 = Non-Polar Covalent 0.31.7 = Polar Covalent or: EN difference = 0 0.3 1.7 3.3 I--------I--------------------I---------------------------I non- Polar Ionic Polar (see p.162 Fig.6-2 )

  21. Going back to the previous examples: H-F∆EN = 1.9  Na-Cl∆EN = 2.1  H-H∆EN = 0  Other examples: Mg-S CO2

  22. Intermolecular Forces: -forces ______________molecules. -weaker than ionic & covalent bonds. In Polar Covalent (∆EN=0.3-1.7) The EN difference creates a ____________ __________________________________________________ e.g. I-Cl => I---Cl => I Cl (2.5)(3.0) +- 0.5 Difference Dipole

  23. Dipole-Dipole interaction Intermolecular force __________________________________________________________________ Strong intermolecular force.

  24. Hydrogen Bonding The ______________________ force that has H partially bonded to an electronegative atom. E.g’s: Causes higher than normal boiling points water is a liquid instead of a gas @ room temp. __________________________________________________________ e.g. H2O , HCl , HF , H2S … (do diagrams) I love water!!! (why?)

  25. London Dispersion forces Average shape Temporary shift Effects other molecules Non-Polar Molecules (∆EN= 0 - 0.3) There is no dipole because the EN diff is too low. But a ____________________________________________________________________________________________, which effects the next molecule, and so on. e.g. Draw Cl2 gas with London dispersion, aka Van der Waals forces. Other e.g’s? O2 , H2 , F2 …

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