Structural Biochemistry/Acid-Base Chemistry
General Info[edit | edit source]
There are several ways to describe acid-base chemistry. A simple proposed model by Bronsted and Lowry is that an acid is a proton donor, while a base is a proton acceptor. An easy way to understand these definitions it to think about acidity and basicity in regards to water. Whereas a base will remove protons from water to form hydroxide ions, an acid will donate a proton to water, forming hydronium ions. Water is neutral and has a pKa of 15.7. This can be used as a reference point in determining the acidity or basicity of a compound. The pH is defined as the negative log of [H+]. For pure water, the pH is 7, and an aqueous solution below 7 is considered acidic, while a solution above 7 is considered basic. The pKa of a solution can be determined by the equation pKa = -log(Ka) in which Ka is the acid dissociation constant. Generally an acid with a pKa lower than 1 is considered a strong acid, while an acid that has a pKa greater than 4 is considered weak. Some examples of strong acids include HI (pKa = -10), HBr (pKa = -9), and HCl (pKa = -8). Examples of weak acids include acetic acid (pKa = 4.7) and methanol (pKa = 15.5). Molecules can that can act as an acid under certain conditions and act as a base in a different environment are known as amphoteric compounds. Water is a common example of a self-ionizable amphoteric compound.[1]
Another way of depicting acids and bases comes from Gilbert N. Lewis. A Lewis acid is one that is at least two electrons short of a closed outer shell, while a Lewis base contains at least one lone pair of electrons. Moreover, Lewis bases can share one of its lone pairs with a Lewis acid to create a new covalent bond. This concept can help understand the fundamentals of nucleophiles and electrophiles. All Lewis acids are electrophiles. Nucleophiles generally refer to a Lewis base, and always have at least one lone pair of electrons.[1]
The relative strength of an acid HA can be determined by a few general rules:
1) Electronegativity: The more electronegative the atom attached to the acidic proton, the more acidic the proton will be because the bond will be more polar. This trend is very helpful since generally electronegativity increases from left to right in the periodic table. For example, HF is more acidic than H2O, since fluorine is more electronegative than oxygen.[1] Electronegativity is the phenomena of an element to draw electrons toward it. This means that as electronegativity increases, the more likely an element will pull electrons from other elements. Fluorine is the element with the highest known electronegativity levels.
2) Size: With an increasing size of A in the acid HA, the acidity of the compound increases. This is a result of the poor overlap of the outer-shell orbital of A with the 1s hydrogen orbital. Thus, dissociation of HA into H+ and A- is favored for a larger A. Also, the large outer-shell orbital allows electrons to occupy a larger region of space, and thus electron-electron repulsion is diminished in A-. For these reasons, acidity will increase when going down a column in the periodic table. This helps explain why HI > HBr > HCl > HF (decreasing acidity, HI being most acidic of the group).[1]
3) Resonance: Resonance in A- allows for the delocalization of charge over several atoms. An example of this is acetic acid vs. methanol. Acetic acid is more acidic than methanol because its deprotonated form is resonance stabilized, whereas methanol’s is not.[1]
A general rule for determining the acidity of a compound is that the acidity of HA increases to the right and down the periodic table, while the basicity of A- decreases in the same manner.
Real Life Applications[edit | edit source]
An example of our every day interactions with acid-base chemistry includes stomach acid. The pH of the stomach juice generally remains within a range of 1.0 and 2.5, varying depending on the stimuli of senses. This stomach acid alters the natural folded shapes of protein molecules, allowing them to be broken down by digestive enzymes. Though stomach acid is extremely useful in this manner, it can also be harmful if unregulated, since it can destroy the protein molecules in the stomach tissue itself. To prevent this from happening, the interior of the stomach is coated with a layer of cells known as gastric mucosa, which insulates the stomach wall from acidic gastric juices. Cells beneath the gastric mucosa are activated via stimuli of taste, smell and histamine (a type of signaling molecule) that results in parietal cells releasing HCl into the stomach. Conditions such as hyperacidity, where there is excessive amounts of acid secreted into the stomach, and peptic ulcers, which are sores resulting from bacterial infections, can be regulated by medications that block histamine from signaling the parietal cells. Some common ingredients used in these medications include cimetidine, famotidine, and ranitidine.[1]
Another example of acid-base reactions is the affects of pH on DNA. The formation of DNA occurs readily at a pH of 7.0; however, altering the pH level of a solution containing the double-helical DNA can destabilize the DNA double helix. For example, in a solution with double-helical DNA and a concentrated base (such as OH-), the DNA will begin to dissociate into its corresponding single strands when pH approaches 9.0. This is a result of the hydroxide ions and their interaction with DNA base pairs, removing specific protons. Similarly, when the pH of this solution drops too low (below 5.0), the DNA double helix is destabilized. This is because some of the hydrogen bond acceptors become protonated and can no longer participate in hydrogen bonding, so the double helix separates. This shows how altering the pH of DNA can disrupt its double-helical structure. [2]