Ionic Bonds  

- metals and nonmetals react chemically by the TRANSFER of electrons (from metals to nonmetals)

- metals form positive ions by losing valence electrons to the nonmetals which then form negative ions

- positive ions are strongly attracted to the negative ions by the electrostatic attraction that exist between unlike charges

- the new substance formed does not resemble either of the original atoms

- this attraction binding unlike ions together is called ionic bonding

 example:  CaF₂

 see classroom drawing

                                                 Covalent Bonds

- two or more atoms both of which tend to gain electrons during reactions (nonmetals) may combine by sharing 1 or 2 or 3 pairs of electrons

- the force holding the atoms together is due to the attraction of each atom for the electrons that are held jointly (a stable condition)

- HYDROGEN, CARBON, NITROGEN, AND OXYGEN are noted for forming covalent bonds

 single covalent bond:       see examples in classroom

 double covalent bond:

 triple covalent bond:

   

                                                      Homework/Test Problems

First determine if the molecule is ionic or covalently bonded.  Then draw the electron dot structures showing an acceptable bonding structure.

 1. H₂S                  2.  F₂                         3. HF                     4. H₂O                   5. AlF₃

 6. MgO                  7. NH₃                      8. PBr₃                   9. CCl₄                   10. CS₂

 11. CO₂                  12. K₂S                   13. CH₄                    14. C₂H₂                 15. MgCl₂

 16. SiO₂                 17. NF₃                  18. HCl                      19. CHCl₃               20. C₂F₂

 21. C₂H₆                 22. C₂H₄                 23. CHN                 24. Si₂F₄                 *25. BF₃

Bonding Information   Extra Bonding Notes  Bonding Review

 The difference in the electronegativities of two elements can be used to predict the nature of the bond.  When this difference is small, the bond is primarily covalent.  As the difference increases, the covalent bonds become increasingly polar. When the difference becomes even greater, the bond becomes ionic. 

Generally the line is drawn at 1.7.  When the differences in electronegativities is greater than 1.7 the bond is ionic (and less than 1.7 is covalent).  Another boundary often is drawn at a difference of 1.0 (sometimes 0.8) to separate polar bonds from nonpolar bonds.

 When a molecule behaves as if one end were negative and the opposite end positive, the molecule is said to be polar. Polar molecules are known as dipoles.  A molecule is polar when there is an uneven distribution of electrons in the molecule.

 When two atoms of the same element form a molecule, the shared electrons are equidistant from the nuclei of the two atoms.  This makes the bond nonpolar.

 HCl is an example of a two-atom polar molecule.  The shared electron pair is attracted toward the highly electronegative chlorine atom and away from the hydrogen atom.  The resulting concentration of negative charge is closer to the chlorine atom and that end of the molecule will be slightly negative.  The other end will be slightly positive but the molecule as a whole will be neutral.

Summary:

1) compounds or bonded atoms in molecules are polar is the center of positive charge does not coincide with the center of negative charge.

2) when a covalent bond is formed between atoms of different electronegativities, the pair of electrons will be more closely associated with the more electronegative atom, and the resulting covalent bond will be somewhat polar.

3) the greater the difference between the electronegativities of the atoms involved in the bond, the greater the polarity of the bond.

4) if the difference in electronegativity is too large, the electrons will be transferred and ionic bonding will result instead.

5. if both atoms in covalent bond have identical ionization potentials and electronegativities, no ions are formed and there is no polarity.

 Hydrogen bonds:  In compounds such as water, ammonia (NH₃), and hydrogen fluoride (HF), the hydrogen atoms are bonded to small atoms of high electronegativity (oxygen, nitrogen, and fluorine, respectively).  The hydrogen atom has only a very small share of the electron pair that forms the bond.  Such molecules are highly polar.  In fact, each hydrogen atom acts largely as exposed proton.  It can be attracted to, and form a weak bond with, the highly electronegative atom of a neighboring molecule.  This is called a hydrogen bond.  It is more than just an electrostatic attraction between opposite charges. It actually has some covalent character.

 Hydrogen bonding is responsible for a number of unusual properties.  Hydrogen bonding occurs between water molecules.  Water must therefore be raised to a much higher temperature before the kinetic energy of its molecules becomes great enough to break the hydrogen bonds between the molecules.  Breaking these hydrogen bonds is necessary in order to boil water.  X ray studies show that the three-dimensional structure caused by hydrogen bonding gives ice crystals a crystalline arrangement with many hexagonal openings.  This open structure accounts for the low density of ice.