Definitions:

Acid: A chemical compound that releases H+ to a solution is called an acid.

–strong acids: include hydrochloric acid (HCL) which is found in your stomach and aids in digestion of food. In solution, HCl breaks apart into the ions H+ and Cl-.

Base: A compound that accepts H+ and removes them from solution is a base. Some bases, such as sodium hydroxide (NaOH), do this by releasing OH-, which combines with H+ to form H2O.

pH: To describe the acidity of a solution, chemists use the pH scale, a measure of the hydrogen ion (H+) concentration in a solution. In fact, pH stands for “power of hydrogen”. The scale ranges from 0 (most acidic) to 14 (most basic). A solution having a pH of 7 is neutral, meaning that its H+ and OH- concentrations are equal. The lower the pH below 7, the more acidic the solution, or the greater its excess of H+ compared with OH-. The higher the pH above 7, the more basic the solution, or the greater the deficiency of H+ relative to OH-. Each pH unit represents a tenfold change in the concentration of H+. For example, lemon juice at pH 2 has 100 times more H+ than an equal amount of tomato juice at pH 4.

Aqueous solutions that are neither acidic nor basic (such as pure water) are said to be neutral; they have a pH of 7. The concentration of H+ and OH- are equal. The pH of the solution of most living cells is close to 7.

Some common examples of acidic solutions from most acidic to less acidic include battery acid, lemon juice/stomach acid, grapefruit jice, soft drinks, tomato juice, black coffee and urine.

Some common examples of basic solutions from most basic to least basic are oven cleaner (pH around 14), household bleach, household ammonia, milk of magnesia and seawater (pH around 9).

Buffers

Even a slight change in pH cn be harmful to an organism becasue the molecules in cells are extremely sensitive to H+ and OH- concentrations. Buffers are substances that minimize changes in pH by accepting H+ when that ion is in excess and donating H+ when it is depleted. For example, buffer in contact lens solution helps protect the surface of the eye from potentially painful changes in pH.

This process is affected by the environment, however. For example, about 25% of the carbon dioxide (CO2) generated by people (primarily by burning fossil fuels) is absorbed by the oceans. When CO2 dissolves in seawater, it reacts with water to form carbonic acid, which lower ocean pH. The resulting ocean acidification can greatly change marine environments. Oceanographers have calculated that the pH of the ocean is lower now than at any time in the past 420k years and it is continuing to drop.

Carbonic acid, which lowers th pH of water, is also a danger to coral reffers. The higher levels of CO2 in the atmosphere lead to more of CDO2 dissolving into seawater, forming carbonic acid, which lowers the pH of the water. This makes calcium carbonate less available. Some species of calcifying marine plants and animals, including corals, have been shown to form less robust skeetons and to form them more slowly as pH levels decreases. Water warmer due to climate change can also cause the symbiosis with zooxanthellae to break down, a phenomenon known as “coral bleeching” becasue the underlying white skeleton of the animal becomes visible through its body in the absence of the symbionts.

Buffers solutions acheive their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A-. When a strong acid is added to an equilibrium mixture of a weak acid and its conjugate base, the equilibrium is shifted to the left in accordance with Le cahtelier’s principle. Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added. Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amoutn expected for the quantity of alkali added.

Many buffers, especially the good buffers, are supplied as crystalline acids or bases. The pH of these buffer materials in solution will not be ear the pKa and the materials will not become buffers until the pH is adjusted. In practice, one selects a buffer material with a pKa near the desired working pH. If the buffer material is a free acid, it is adjusted to the working pH with sodium hydroxide, potassium hydroxide, tetramethylammonium hydroxide or other appropriate base. Buffer materials obtained as free bases must be adjusted by addition of a suitable acid. (Gueffroy “A guide for the preparation and use of buffers in biological systems” 1975).

Importance of pH in Organisms:

In aqueous solutions, most of the water molecules are in tact. However, some of the water molecules break apart into hydrogen ions (H+) and hydroxide ions (OH-). A blance of these two highly reactive ions is critical for the proper functioning of chemical processes within organisms.

Diseases related to pH:

–Blood acidosis: In this condition, human blood, which normally has a pH of about 7.4, drops to a pH of about 7.1. This condition is fatal if not reated immedaitely.

–-Blood alkalosis: The condition is the reversie of blood acidosis above. Here, an increase in blood pH can result in a serious condition called blood alkalosis.