pH, acids, based and buffers

Acids and bases, like salts, dissociate in water into electrolytes. Acids and bases can very much change the properties of the solutions in which they are dissolved.

Acids

An acid is a substance that releases hydrogen ions (H⁺) in solution (Figure 3a). Because an atom of hydrogen has just one proton and one electron, a positively charged hydrogen ion is simply a proton. This solitary proton is highly likely to participate in chemical reactions. Strong acids are compounds that release all of their H⁺ in solution; that is, they ionize completely. Hydrochloric acid (HCl), which is released from cells in the lining of the stomach, is a strong acid because it releases all of its H⁺ in the stomach’s watery environment. This strong acid aids in digestion and kills ingested microbes. Weak acids do not ionize completely; that is, some of their hydrogenionsremain bonded within a compound in solution. An example of a weak acid is vinegar, or acetic acid; it is called acetate after it gives up a proton.

Figure 3. Acids and Bases. (a) In aqueous solution, an acid dissociates into hydrogen ions (H⁺) and anions. Nearly every molecule of a strong acid dissociates, producing a high concentration of H⁺. (b) In aqueous solution, a base dissociates into hydroxyl ions (OH^–) and cations. Nearly every molecule of a strong base dissociates, producing a high concentration of OH^–.

Bases

A base is a substance that releases hydroxyl ions (OH^–) in solution, or one that accepts H⁺ already present in solution (see Figure 3b). The hydroxyl ions or other base combine with H⁺ present to form a water molecule, thereby removing H⁺ and reducing the solution’s acidity. Strong bases release most or all of their hydroxyl ions; weak bases release only some hydroxyl ions or absorb only a few H⁺. Food mixed with hydrochloric acid from the stomach would burn the small intestine, the next portion of the digestive tract after the stomach, if it were not for the release of bicarbonate (HCO₃^–), a weak base that attracts H⁺. Bicarbonate accepts some of the H⁺ protons, thereby reducing the acidity of the solution.

The Concept of pH

Hydrogen ions are spontaneously generated in pure water by the dissociation (ionization) of a small percentage of water molecules into equal numbers of hydrogen (H⁺) ions and hydroxide (OH^–) ions. While the hydroxide ions are kept in solution by their hydrogen bonding with other water molecules, the hydrogen ions, consisting of naked protons, are immediately attracted to un-ionized water molecules, forming hydronium ions (H₃0⁺). Still, by convention, scientists refer to hydrogen ions and their concentration as if they were free in this state in liquid water.

The relative acidity or alkalinity of a solution can be indicated by its pH. A solution’s pH is the negative, base-10 logarithm of the hydrogen ion (H⁺) concentration of the solution. As an example, a pH 4 solution has an H⁺ concentration that is ten times greater than that of a pH 5 solution. That is, a solution with a pH of 4 is ten times more acidic than a solution with a pH of 5. The concept of pH will begin to make more sense when you study the pH scale, like that shown in Figure 4. The scale consists of a series of increments ranging from 0 to 14. A solution with a pH of 7 is considered neutral—neither acidic nor basic. Pure water has a pH of 7. The lower the number below 7, the more acidic the solution, or the greater the concentration of H⁺. The concentration of hydrogen ions at each pH value is 10 times different than the next pH. For instance, a pH value of 4 corresponds to a proton concentration of 10^–4 M, or 0.0001M, while a pH value of 5 corresponds to a proton concentration of 10^–5 M, or 0.00001M. The higher the number above 7, the more basic (alkaline) the solution, or the lower the concentration of H⁺. Human urine, for example, is ten times more acidic than pure water, and HCl is 10,000,000 times more acidic than water.

Figure 4. The pH Scale

Buffers

So how can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers readily absorb excess H⁺ or OH^–, keeping the pH of the body carefully maintained in the narrow range required for survival. Maintaining a constant blood pH is critical to a person’s well-being. The buffer maintaining the pH of human blood involves carbonic acid (H₂CO₃), bicarbonate ion (HCO₃^–), and carbon dioxide (CO₂). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, hydrogen ions are removed, moderating pH changes. Similarly, as shown in [Figure 8], excess carbonic acid can be converted to carbon dioxide gas and exhaled through the lungs. This prevents too many free hydrogen ions from building up in the blood and dangerously reducing the blood’s pH. Likewise, if too much OH^– is introduced into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to put survival in jeopardy.

Figure 8: This diagram shows the body’s buffering of blood pH levels. The blue arrows show the process of raising pH as more CO2 is made. The purple arrows indicate the reverse process: the lowering of pH as more bicarbonate is created.

 

Other examples of buffers are antacids used to combat excess stomach acid. Many of these over-the-counter medications work in the same way as blood buffers, usually with at least one ion capable of absorbing hydrogen and moderating pH, bringing relief to those that suffer “heartburn” after eating. The unique properties of water that contribute to this capacity to balance pH—as well as water’s other characteristics—are essential to sustaining life on Earth.

The pH of human blood normally ranges from 7.35 to 7.45, although it is typically identified as pH 7.4. At this slightly basic pH, blood can reduce the acidity resulting from the carbon dioxide (CO₂) constantly being released into the bloodstream by the trillions of cells in the body. Homeostatic mechanisms (along with exhaling CO₂ while breathing) normally keep the pH of blood within this narrow range. This is critical, because fluctuations—either too acidic or too alkaline—can lead to life-threatening disorders.

All cells of the body depend on homeostatic regulation of acid–base balance at a pH of approximately 7.4. The body therefore has several mechanisms for this regulation, involving breathing, the excretion of chemicals in urine, and the internal release of chemicals collectively called buffers into body fluids.

A buffer is a solution of a weak acid and its conjugate base. A buffer can neutralize small amounts of acids or bases in body fluids. For example, if there is even a slight decrease below 7.35 in the pH of a bodily fluid, the buffer in the fluid—in this case, acting as a weak base—will bind the excess hydrogen ions. In contrast, if pH rises above 7.45, the buffer will act as a weak acid and contribute hydrogen ions.