Monday, April 26, 2010

Types Of Chemical Bonds

The forces of attraction that hold atoms together are called chemical bonds. The following is a list of different types of chemical bonds.

Intermolecular bonds refers to the forces of attraction that hold atoms together within a molecule. These types of bonds are considerably stronger than intermolecular bonds. We will study three types of intermolecular bonds - covalent, ionic and metallic bonds.

# Ionic Bonds:

The individual atoms are sodium and chlorine with only their valence electrons shown. Note that chlorine has seven valence
electrons (it wants a full shell of eight), and that sodium has one valence electron (it also wants a full shell of eight). Because chlorine has a high electronegativity (3.5) compared to sodium (0.9), chlorine can easily steal sodium's one lonely valence electron. This makes chlorine very happy, as it now has eight valence electrons. This also makes sodium very happy, as it also now has eight valence electrons. Where did these eight valence electrons come from? They were already a part of the sodium atom, in the 2nd Principle Energy Level (hidden from your view in this animation to simplify matters). The one electron it lost was in the 3rd Principle Energy Level.

The transfer of the electron caused the previously neutral sodium atom to become a positively charged ion (a cation), and the previously neutral chlorine atom to become a negatively charged ion (an anion). The attraction for the cation and the anion is called the ionic bond.

An ionic bond is typically formed between a metal and a non-metal. Metals have low electronegative s (less than 2.0), while non-metals have high electronegative s (above 2.0). Consequently, the non-metal is "stronger" than the metal, and can steal electrons very easily from the metal. This results in the metal becoming a cation, and the non-metal becoming an anion.

*An ionic bond is the resulting attraction for an anion and a cation after an electron is transferred from the metal to the non-metal.*

* Covalent Bonds:

The individual atoms are atoms of chlorine with only their valence electrons shown. Note that each chlorine atom has only seven valence electrons, but really wants eight. When each chlorine atom shares its unpaired electron, both atoms are tricked into thinking each has a full valence of eight electrons. Notice that the individual atoms have full freedom from each other, but once the bond is formed, energy is released, and the new chlorine molecule (Cl2) behaves as a single particle.

A covalent bond is typically formed by two non-metals. Non-metals have similar electronegativities. Consequently, neither atom is "strong" enough to steal electrons from the other.Therefore, the atoms must share the electrons.

# Metallic Bonding:

Metals have low ionization energies, thus they do not have a tight hold on their valence electrons. These outer electrons easily move around, as they do not "belong" to any one atom, but are part of the whole metal crystal. The negatively charged electrons act as a "cement" that hold the positively charged metal ions in their relatively fixed positions.

The fact that the electrons flow easily helps to explain some of the characteristics of metals:

- Metals are good conductors of heat and electricity. This is directly due to the mobility of the electrons.

- The "cement" effect of the electrons determines the hardness of the metal. Some metals are harder than others; the strength of the "cement" varies from metal to metal.

- Metals are lustrous. This is due to the uniform way that the valence electrons of the metal absorb and re-emit light energy.

- Metals are malleable (can be flattened) and ductile (can be drawn into wires) because of the way the metal cations and electrons can "flow" around each other, without breaking the crystal structure.

*Metallic bonds are best characterized by the phrase "a sea of electrons"*

Intermolecular Bonds:

Intermolecular bonds refers to the forces of attraction that hold molecules together. These bonds are considerably weaker than intramolecular bonds. We will study three types of intermolecular bonds - hydrogen bonds, van der Waals forces, and molecule-ion attractions.

* Hydrogen Bonds:
The crystalline structure is due to hydrogen bonding. A hydrogen bond exists between two highly polar molecules containing hydrogen. Hydrogen has a relatively low electronegativity (2.2), and when it is covalently bonded to an atom of either fluorine, oxygen, or nitrogen (electronegativities of 4.0, 3.5, and 3.1, respectively), the resulting bonds are highly polar. So when two such molecules are nearby, they will orient themselves so that the partially negative end of one molecule will face the partially positive end of another molecule.

*The attraction of the partially positive end of one highly polar molecule for the partially negative end of another highly polar molecule is called a hydrogen bond.*

van der Waals Forces:

The hydrogen gas (H2), carbon dioxide (CO2), nitrogen (N2), and in the noble gases (He, Ne, Ar, Kr, etc).

A temporary polarity in the neighboring atom (we'll call it atom #2) is also then induced when its electrons are repelled by atom #1. The temporarily induced polarity allows the two atoms to be attracted to each other very weakly, when the partially negative end of atom #1 is attracted to the partially positive end of atom #2.

At normal conditions of temperature and pressure, van der Waals forces are negligible. As the pressure increases (molecules are forced closer together), and/or the temperature decreases (molecules slow down), van der Waals forces are much more important. Only at conditions of high pressure and/or low temperature are the molecules able to participate in van der Waals forces noticeably---only at these conditions will these gases liquefy. What holds non-polar molecules together in the liquid state? You got it! van der Waals forces!

Another important concept to understand about van der Waals forces is this - the bigger the molecule, the stronger the van der Waals. Why? Bigger molecules have more electrons, which results in bigger electron pile-ups, which results in a bigger induced polarity. Of the noble gases He, Ne, Ar, and Kr, which do you think is the easiest to liquefy? The one that has the strongest van der Waals forces....which will be the biggest on that list...krypton.

*van der Waals forces are weak attractive forces that hold non-polar molecules together.*

# Molecule-Ion Attractions:

The polar water molecule orients itself so that its partially positive end faces the negative ion in the ionic solid. The molecule-ion attraction is stronger than the ions' attraction for each other, and the anion can be removed from the crystal lattice. Once dislodged, the anion is surrounded by more water molecules, all oriented with the positive end facing the anion. The animation continues on, to show the similar process where a cation is dissolved from the ionic solid. This time, however, it is the partially negative end of the water molecule that faces the cation.

*This "molecule-ion attraction" is appropriately named, as a molecule (water) is attracted to an ion.*

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