Science Gateway > Resources > Biology Hypertextbook Index

MIT Biology Hypertextbook: Chemistry Review

Chemistry Review


The study of biology requires an understanding of simple organic chemistry and simple biological chemistry. Carbohydrates, lipids, proteins, and nucleic acids, the players in molecular biology, are themselves composed of smaller building blocks. This chapter contains a review of important chemical interactions and concepts you will encounter in this course.

1 Chemical Bonds

This section provides a quick review of chemical bonds. Emphasis is placed on bonds between the six major elements found in biological systems: H, C, N, O, P, and S.

1.1 Covalent Bonds

Covalent Bonds are the strongest chemical bonds, and are formed by the sharing of a pair of electrons. The energy of a typical single covalent bond is ~80 kilocalories per mole (kcal/mol). However, this bond energy can vary from ~50 kcal/mol to ~110 kcal/mol depending on the elements involved. Once formed, covalent bonds rarely break spontaneously. This is due to simple energetic considerations; the thermal energy of a molecule at room temperature (298 K) is only ~0.6 kcal/mol, much lower than the energy required to break a covalent bond.

There are single, double, and triple covalent bonds:


Bond Number			Example		Energy (kcal/mol)

	                    	  H
		                  |
single                         H--C--H		  ~80
                                  |
                                  H

				H  H
				|  |
double			     H--C==C--H		  ~150
                                |  |
				H  H
 
				H
			        |
			        C
triple			       |||                ~200
				C
				|
				H			
Note that carbon-carbon bonds are unusually strong and stable covalent bonds.

The major organic elements have standard bonding capabilities:


ELEMENT		# of Covalent Bonds		     EXAMPLE

Nitrogen (N)		4			     Ammonia
							H
(positive charge					|+
on nitrogen)					     H--N--H
							|
							H

Oxygen (O)		1		         ionized ethanol
							H  H
							|  |   _
(negative charge				     H--C--C--O
on oxygen)						|  |
							H  H

Sulfur (S)		1		     ionized mercaptoethanol
							H  H
							|  |   _
(negative charge	  			     H--C--C--S
on the sulfur)						|  |
							H  H

Covalent bonds can also have partial charges when the atoms involved have different electronegativities. Water is perhaps the most obvious example of a molecule with partial charges. The symbols delta+ and delta- are used to indicate partial charges.

Oxygen, because of its high electronegativity, attracts the electrons away from the hydrogen atoms, resulting in a partial negative charge on the oxygen and a partial positive charge on each of the hydrogens.

The possibility of hydrogen bonds (H-bonds) is a consequence of partial charges.

1.2 Hydrogen Bonds

Hydrogen bonds are formed when a hydrogen atom is shared between two molecules.

Hydrogen bonds have polarity. A hydrogen atom covalently attached to a very electronegative atom (N, O, or P) shares its partial positive charge with a second electronegative atom (N, O, or P). One example, shown above, involves the hydrogen bonding between water molecules.


More examples:
		    H
		    |
	R--O--H ||| N--R	R==N--H ||| O==R
                    |					Note that R stands for
		    H					any side group.

Hydrogen bonds are ~5 kcal/mol in strength. These bonds are frequently found in proteins and nucleic acids, and by reinforcing each other serve to keep the protein (or nucleic acid) structure secure. But, since the hydrogen atoms in the protein could also H-bond to the surrounding water, the relative strength of protein-protein H-bonds vs. protein-H2O bonds is smaller than 5 kcal/mol.

1.3 Ionic Bonds

Ionic bonds are formed when there is a complete transfer of electrons from one atom to another, resulting in two ions, one positively charged and the other negatively charged. For example, when a sodium atom (Na) donates the one electron in its outer valence shell to a chlorine (Cl) atom, which needs one electron to fill its outer valence shell, NaCl (table salt) results. The symbol for sodium chloride is Na+Cl-. Ionic bonds are often 4-7 kcal/mol in strength.

1.4 Van der Waals Bonds

Van der Walls interactions are very weak bonds (generally no greater than 1 kcal/mol) formed between nonpolar molecules or non-polar parts of a molecule. The weak bond is created because a C-H bond can have a transient dipole and induce a transient dipole in another C-H bond.


            H      H
            | ~~~~ |
            CH3    CH3

1.5 Hydrophobic Interactions

Nonpolar molecules cannot form H-bonds with H2O, and are therefore insoluble in H2O. These molecules are known as hydrophobic (water hating), as opposed to water loving hydrophilic molecules which can form H-bonds with H2O. Hydrophobic molecules tend to aggregate together in avoidance of H2O molecules; hydrophobic interactions are clearly demonstrated when you put an oil drop on water. This attraction/repulsion is known as the hydrophobic (fear of water) force. To understand the energetics driving this interaction, visualize the H2O molecules surrounding a "dissolved" molecule attempting to form the greatest number of hydrogen bonds with each other. The best energetic solution involves forcing all of the nonpolar molecules together, thus reducing the total surface area that breaks up the H2O H-bond matrix.

 

2 pH

The pH is a measure of the concentration of hydrogen ions. These are derived, for example, from the dissociation of an acid--HCl--when this is dissolved in water. The pH value is defined as the negative logarithm of the hydrogen ion concentration in mol/L. The equation is:

	pH = -log10[H+]

The [H+] in pure water is 10^-7; therefore the pH of pure water is:

	pH = -log10(10^-7)	pH = - (-7)	pH = 7

pH 7 is often referred to as "neutral pH". Everything below pH 7 has a higher concentration of H+ and is considered acidic. Everything above pH 7 has a lower concentration of H+ and is considered basic; you can also think of this as a higher concentration of OH-.

A lower pH always means a higher concentration of H+. The biochemically useful ends of the scale are 1 M HCl, which is pH 0, and 1 M NaOH, which is pH 14. In general, c