pH

The pH of a solution is a measure of the hydronium ion (H₃O⁺) concentration in that solution. The hydronium ion in a solution results from the self-ionization of water. In other words, because hydronium ions are formed from H⁺ ions and water molecules, pH is also expressed as the concentration of hydrogen ions (or, positively charged hydrogen atoms) of a solution. The pH is measured on a scale from 0 to 14. Therefore, the pH is a measure of the acidity or alkalinity of a solution based upon the dissociation of water. The variability of pH can have a

Approximate pH of Various Substances. (Thomson Gale.)

dramatic effect on biological or physical chemical reactions (e.g., geochemical weathering processes).

The abbreviation pH is derived from the French term pauvoir hydrogéne, which means hydrogen power. The pH scale was developed by Danish chemist Søren Peter Lauritz Sørensen (1868-1939) in 1909. It is generally presented as ranging from 0-14, although there are no theoretical limits on the range of the scale (e.g., there are substances with negative pHs and with pHs greater than 14) the range is generally cited as being from pH=0 to pH=14.

A solution with a pH of less than pH = 7 is acidic and a solution with a pH of greater than pH = 7isbasic (alkaline). The midpoint of the scale, 7, is neutral. The lower the pH of a solution, the more acidic the solution is and the higher the pH, the more basic it is.

Mathematically, the potential hydronium ion concentration (pH) is equal to the negative logarithm of the hydronium ion concentration: pH =–log [H₃0⁺’. The square brackets indicate the concentration of, in moles per liter. [H₃0⁺’ represents the hydro-nium ion—esentially a water molecule with a proton attached. Thus, [H₃O⁺’ indicates the concentration of hydronium ions in moles per liter. Hydronium ions are important participants in chemical reactions that take place in aqueous (water, H₂0) solutions.

Water is a weak electrolyte. Through a process termed self-ionization, a small number of water molecules in pure water dissociate (separate) in a reversible reaction to form a positively charged H⁺ ion and a negatively charged OH^– ion. In aqueous solution, as one water molecule dissociates, another is nearby to pick up the loose, positively charged, hydrogen proton to form a positively charged hydronium ion (H_3O⁺). The water molecule that lost the hydrogen proton— but that kept the hydrogen electron—becomes a negatively charged hydroxide ion (OH^–).

In dilute solutions, the product of the hydronium ion concentration and the hydroxide ion equals the ion product (K_w) or dissociation constant (K_w = 1.0×10^–14 at 25^°C). Calculations of pH using the ion product yield a number between 0 and 14—the standard pH scale.

In a sample of pure water, the concentration of hydronium ions is equal to 1× 10^–7 moles per liter (0.0000001 M). The equilibrium (balance) between hydronium and hydroxide ions that results from self-ionization of water can be disturbed if other substances that can donate protons are put into solution with water.

The pH of solutions may be measured experimentally with an electronic pH meter (highly accurate pH meters can measure to 0.001 pH units) or by using acid base indicators, chemicals that change color in solutions of different pH. A crude but common test for pH involves the use of hydrion paper strips (litmus paper) that undergo changes similar to those found in indicator solutions. For example, red litmus paper turns blue in a basic solution.

Sometimes it is necessary to maintain a solution at a constant pH. This is especially true in bodily fluids such as blood, which needs to be kept neutral (between a pH of 7.35 and 7.45). If the pH of blood is allowed to vary outside of this range, serious illness or death may occur. Buffers are substances that control the pH of a solution. A buffer is usually a mixture of acids and bases. This mix of acids and bases allows the buffer to release or absorb hydronium ions, which keeps the pH of a solution constant. The most common buffers are mixtures of weak acids and their conjugate bases. A buffer cannot keep the pH of a solution under absolute control under all conditions. There is a limit to the ability of a buffer to maintain a constant pH, called the buffer capacity. In general, the greater the concentration of buffer in a solution, the greater the buffer capacity.

See also Acids and bases.

pH

The most common method of indicating the acidity of a solution is by stating its pH. The term pH refers to a mathematical system developed by Danish chemist Søren Sørenson (1868–1939) around 1909. Sørenson originally suggested the term pH as an abbreviation for potential (or power) of hydrogen.

Acids and bases were first defined by Swedish chemist Svante Arrhenius (1859–1927). Arrhenius proposed that acids be defined as chemicals that produce positively charged hydrogen ions, H⁺, in water. By comparison, he suggested that bases are compounds that produce negatively charged hydroxide ions, OH⁻, in water.

The pH of a solution is determined by the concentration of hydrogen ions present—that is, by its acidity. The more hydrogen ions present (the more acidic the solution), the lower the pH. The fewer hydrogen ions present (the less acidic the solution), the higher the pH. The pH scale runs from 0 to 14. A pH value of 7 (in the middle of that range) represents a solution that is neither acidic nor basic.

Strong acids have very low pHs (battery acid has a pH of 0). Strong bases have very high pHs (sodium hydroxide, commonly known as lye, has a pH of 14). Lemon juice has a pH of 2; vinegar of 2.5; coffee of 5; distilled water of 7; borax of 9; and household ammonia of 11.

pH indicators

One way of finding the pH of a solution is with a pH meter, a mechanical device that gives very precise readings. One can easily place the probe of a pH meter into a solution and read the pH of the solution on the meter dial.

Table 1.
Some Common Solutions and Their pH

Table 2.
Some Common Indicators, the pH Range in Which They Change Color, and Their Color Changes

A much older method for estimating the pH of a solution is the use of an indicator. A pH indicator is a material that changes color in solutions of different pH. One of the most common of all indicators is litmus. Litmus is a chemical obtained from lichens. In the presence of a base, litmus is blue; in the presence of an acid, it is red.

Many indicators are extracted from plants. For example, you can make a reasonably good indicator just by boiling red cabbage and extracting the colored material produced. That material, like litmus, changes color in the presence of acids and bases. A number of indicators are synthetic products made just for testing the pH of solutions.

[See also Acids and bases ]

pH

pH is a measure of the acidity or alkalinity of a solution. The variability of pH can have a dramatic effect on geochemical processes (e.g., weathering processes).

The pH scale was developed by Danish chemist Søren Peter Lauritz Sørensen (1868–1939) in 1909 and is generally presented as ranging from 0 to 14, although there are no theoretical limits on the range of the scale (there are substances with negative pH's and with pH's greater than 14, although for most substances the range of 0–14 suffices). A solution with a pH of less than 7 is acidic and a solution with a pH of greater than 7 is basic (alkaline). The midpoint of the scale, 7, is neutral. The lower the pH of a solution, the more acidic the solution is and the higher the pH, the more basic it is. Mathematically, the potential hydronium ion concentration (pH) is equal to the negative logarithm of the hydronium ion concentration: pH = −log [H₃0⁺], where H₃O⁺ represents the hydronium ion.

Essentially, the hydronium ion can be thought of as a water molecule with a proton attached. The square brackets indicate the concentration of, in moles per liter. Thus, [H₃O⁺] indicates the concentration of hydronium ions in moles per liter.

The hydronium ion is an important participant in the chemical reactions that take place in aqueous (water, H₂0) solutions.

Through a process termed self-ionization, a small number of water molecules in pure water dissociate (separate) in a reversible reaction to form a positively charged H⁺ ion and a negatively charged OH⁻ ion. In aqueous solution, as one water molecule dissociates, another is nearby to pick up the loose, positively charged, hydrogen proton to form a positively charged hydronium ion (H₃O⁺).

Water molecules have the ability to attract protons and form hydronium ions because water is a polar molecule. Oxygen is more electronegative than hydrogen. As a result, the electrons in each of water's two oxygen-hydrogen bonds to spend more time near the oxygen atom . Because the electrons are not shared equally—and because the bond angles of the water molecule do not cancel out this imbalance—the oxygen atom carries a partial negative charge that can attract positively charged protons donated by other molecules.

In a sample of pure water, the concentration of hydronium ions is equal to 1 × 10⁻⁷ moles per liter (0.0000001 M). The water molecule that lost the hydrogen proton—but that kept the hydrogen electron—becomes a negatively charged hydroxide ion (OH⁻).

The equilibrium (balance) between hydronium and hydroxide ions that results from self-ionization of water can be disturbed if other substances that can donate protons are put into solution with water.

The pH of solutions may be measured electronically with a pH meter (better pH meters can measure to 0.001 pH units) or by using acid base indicators, chemicals that change color in solutions of different pH.

See also Acid rain; Geochemistry; Weathering and weathering series

pH

The term pH refers to the concentration of hydrogen ions (H+) in a solution. An acidic environment is enriched in hydrogen ions, whereas a basic environment is relatively depleted of hydrogen ions. The pH of biological systems is an important factor that determines which microorganism is able to survive and operate in the particular environment. While most microorganisms prefer pH's that approximate that of distilled water, some bacteria thrive in environments that are extremely acidic.

The hydrogen ion concentration can be determined empirically and expressed as the pH. The pH scale ranges from 0 to 14, with 1 being the most acidic and 14 being the most basic. The pH scale is a logarithmic scale. That is, each division is different from the adjacent divisions by a factor of ten. For example, a solution that has a pH of 5 is 10 times as acidic as a solution with a pH of 6.

The range of the 14-point pH scale is enormous. Distilled water has a pH of 7. A pH of 0 corresponds to 10 million more hydrogen ions per unit volume, and is the pH of battery acid. A pH of 14 corresponds to one ten-millionth as many hydrogen ions per unit volume, compared to distilled water, and is the pH of liquid drain cleaner.

Compounds that contribute hydrogen ions to a solution are called acids. For example, hydrochloric acid (HCl) is a strong acid. This means that the compounds dissociates easily in solution to produce the ions that comprise the compound (H⁺ and Cl^–). The hydrogen ion is also a proton. The more protons there are in a solution, the greater the acidity of the solution, and the lower the pH.

Mathematically, pH is calculated as the negative logarithm of the hydrogen ion concentration. For example, the hydrogen ion concentration of distilled water is 10^–7 and hence pure water has a pH of 7.

The pH of microbiological growth media is important in ensuring that growth of the target microbes occurs. As well, keeping the pH near the starting pH is also important, because if the pH varies too widely the growth of the microorganism can be halted. This growth inhibition is due to a numbers of reasons, such as the change in shape of proteins due to the presence of more hydrogen ions. If the altered protein ceases to perform a vital function, the survival of the microorganism can be threatened. The pH of growth media is kept relatively constant by the inclusion of compounds that can absorb excess hydrogen or hydroxyl ions. Another means of maintaining pH is by the periodic addition of acid or base in the amount needed to bring the pH back to the desired value. This is usually done in conjunction with the monitoring of the solution, and is a feature of large-scale microbial growth processes, such as used in a brewery.

Microorganisms can tolerate a spectrum of pHs. However, an individual microbe usually has an internal pH that is close to that of distilled water. The surrounding cell membranes and external layers such as the glycocalyx contribute to buffering the cell from the different pH of the surrounding environment.

Some microorganisms are capable of modifying the pH of their environment. For example, bacteria that utilize the sugar glucose can produce lactic acid, which can lower the pH of the environment by up to two pH units. Another example is that of yeast . These microorganisms can actively pump hydrogen ions out of the cell into the environment, creating more acidic conditions. Acidic conditions can also result from the microbial utilization of a basic compound such as ammonia. Conversely, some microorganisms can raise the pH by the release of ammonia.

The ability of microbes to acidify the environment has been long exploited in the pickling process. Foods commonly pickled include cucumbers, cabbage (i.e., sauerkraut), milk (i.e., buttermilk), and some meats. As well, the production of vinegar relies upon the pH decrease caused by the bacterial production of acetic acid.

See also Biochemistry; Buffer; Extremophiles

pH

pH is a measure of the acidity or alkalinity of a solution based upon the dissociation of water . The variability of pH can have a dramatic effect on biological or physical chemical reactions (e.g., geochemical weathering processes).

The pH scale was developed by Danish chemist Søren Peter Lauritz Sørensen (1868–1939) in 1909 and is generally presented as ranging from 0–14, although there are no theoretical limits on the range of the scale (e.g., there are substances with negative pH's and with pH's greater than 14) the range is generally cited as being from pH =0 to pH =14.

A solution with a pH of less than pH =7 is acidic and a solution with a pH of greater than pH =7 is basic (alkaline). The midpoint of the scale, 7, is neutral. The lower the pH of a solution, the more acidic the solution is and the higher the pH, the more basic it is.

Mathematically, the potential hydronium ion concentration (pH) is equal to the negative logarithm of the hydronium ion concentration: pH = -log [H30+]. The square brackets indicate the concentration of, in moles per liter. [H 0+3 ] represents the hydronium ion—esentially a water molecule with a proton attached. Thus, [H Op+3 ] indicates the concentration of hydronium ions in moles per liter. Hydronium ions are important participants in chemical reactions that take place in aqueous (water, H20) solutions.

Water is a weak electrolyte . Through a process termed self-ionization, a small number of water molecules in pure water dissociate (separate) in a reversible reaction to form a positively charged H+ ion and a negatively charged OH- ion. In aqueous solution, as one water molecule dissociates, another is nearby to pick up the loose, positively charged, hydrogen proton to form a positively charged hydronium ion (H3o+ ). The water molecule that lost the hydrogen proton—but that kept the hydrogen electron—becomes a negatively charged hydroxide ion (OH-).

In dilute solutions, the product of the hydronium ion concentration and the hydroxide ion equals the ion product (Kw) or dissociation constant (Kw= 1.0×10-14 at 25°C). Calculations of pH using the ion product yield a number between 0 and 14—the standard pH scale.

In a sample of pure water, the concentration of hydronium ions is equal to 1×10-7 moles per liter (0.0000001 M). The equilibrium (balance) between hydronium and hydroxide ions that results from self-ionization of water can be disturbed if other substances that can donate protons are put into solution with water.

The pH of solutions may be measured experimentally with an electronic pH meter (highly accurate pH meters can measure to 0.001 pH units) or by using acid base indicators, chemicals that change color in solutions of different pH. A crude but common test for pH involves the use of Hydrion paper strips (litmus paper) that undergo changes similar to those found in indicator solutions. For example, red litmus paper turns blue in a basic solution.

See also Acids and bases.