The concentration of hydrogen ions in a solution is very important for living things. This is because, since the hydrogen ions are positively charged they alter the charge environment of other molecules in solution. By putting different forces on the molecules, the molecules change shape from their normal shape. This is particularly important for proteins in solution because the shape of a protein is related to its function. The concentration of hydrogen ions is commonly expressed in terms of the pH scale.

– Dr. Paul's Virtually Biology Show (VBS)

• Make your own pH Indicator - Bottle Biology


• The pH Factor - Miami Museum of Science


• Cabbage Juice - pH indicator - MadSci Network


• Make a Cabbage Juice pH Indicator - Howstuffworks


• Measuring pH - Vanderbilt Student Volunteers for Science

• Measuring pH
  
• Determining the pH of Common Substances
  
• Making a Natural pH Cabbage Juice Indicator
  
• Measuring the pH of Natural Water
  
• Measuring Soil pH
  
• Soil Buffering
  
• Observing the Influence of Acid Rain on Plant Growth
  
• Observing Buffers in Lakes, Ponds, and Streams
  
• Looking at Acid's Effects on Metals
  
• What does pH mean? (Simplified Background)
  
• Testing the pH of various substances (Grades K-4)
  
• Understanding how acid rain affects plant growth (Grades 5-8)
  
• How acidic is rain in your area? (Grades 9-12)
  

pH (abbr. power of hydrogen [2]) is a measure of the activity of hydrogen ions (H⁺) in a solution and, therefore, its acidity or alkalinity. In aqueous systems, the hydrogen ion activity is dictated by the dissociation constant of water (K_w = 1.011 × 10^-14 at 25 °C) and interactions with other ions in solution. Due to this dissociation constant a neutral solution (hydrogen ion activity equals hydroxide ion activity) has a pH of approximately 7. Aqueous solutions with pH values lower than 7 are considered acidic, while pH values higher than 7 are considered alkaline.

The concept was introduced by S.P.L. Sørensen in 1909.

Definition

Though a pH value has no unit, it is not an arbitrary scale; the number arises from a definition based on the activity of hydrogen ions in the solution.

The formula for calculating pH is:

pH = - log₁₀[H ⁺ ]

[H⁺] denotes the activity of H⁺ ions (or more accurately written, [H₃O⁺], the equivalent hydronium ions), measured in moles per litre (also known as molarity). In dilute solutions (like river or tap water) the activity is approximately equal to the concentration of the H⁺ ion.

Log₁₀ denotes the base-10 logarithm, and pH therefore defines a logarithmic scale of acidity. For example, a solution with pH=8.2 will have an [H⁺] activity (concentration) of 10^-8.2 mol/L, or about 6.31 × 10^-9 mol/L; a solution with an [H⁺] activity of 4.5 × 10^-4 mol/L will have a pH value of -log₁₀(4.5 × 10^-4), or about 3.35.

In aqueous solution at standard ambient temperature and pressure (SATP), a pH of 7 indicates neutrality (i.e. pure water) because water naturally dissociates into H⁺ and OH^- ions with equal concentrations of 1×10^-7 mol/L. A lower pH value (for example pH 3) indicates increasing strength of acidity, and a higher pH value (for example pH 11) indicates increasing strength of alkalinity.

Neutral pH is not exactly 7; this would imply that the H⁺ ion concentration is exactly 1×10^-7 mol/L, which is not the case. The value is close enough, however, for neutral pH to be 7.00 to three significant figures, which is near enough for most people to assume it is exactly 7. In nonaqueous solutions or non-SATP conditions, the pH of neutrality may not even be close to 7. Instead it is related to the dissociation constant for the specific solvent used. (Note also that pure water, when exposed to the atmosphere, will take in carbon dioxide, some of which reacts with water to form carbonic acid and H⁺, thereby lowering the pH to about 5.7.)

Most substances have a pH in the range 0 to 14, although extremely acidic or basic substances may have pH < 0, or pH > 14.

Measuring

pH can be measured:

• by addition of a pH indicator into the studying solution. The indicator color varies depending on the pH of the solution. Using indicators, qualitative determinations can be made with universal indicators that have broad color variablity over a wide pH range and quantitative determinations can be made using indicators that have strong color variablitiy over a small pH range. Extremely precise measurements can be made over a wide pH range using indicators that have multiple equilibriums (ie H₂I) in conjunction with spectrophotometric methods to determine the relative abundance of each ph dependant component that make up the color of solution.
• by using a pH meter together with pH-selective electrodes (pH glass electrode, hydrogen electrode, quinhydrone electrode and other).

pOH

There is also pOH, in a sense the opposite of pH, which measures the concentration of OH^- ions. Since water self ionizes, and notating [OH^-] as the concentration of hydroxide ions, we have

(*)

where K_w is the ionization constant of water.

Now, since

by logarithmic identities, we then have the relationship:

and thus

(*)

(*) Valid exactly for temperature = 298.15 K (25 °C) only, acceptable for most lab calculations.

Calculation of pH for weak and strong acids

Values of pH for weak and strong acids can be approximated using certain assumptions.

Under the Brønsted-Lowry theory, stronger or weaker acids are a relative concept. But here we define a strong acid as a species which is a much stronger acid than the hydronium (H₃O⁺) ion. In that case the dissociation reaction (strictly HX+H₂O↔H₃O⁺+X^- but simplified as HX↔H⁺+X^-) goes to completion, i.e. no unreacted acid remains in solution. Dissolving the strong acid HCl in water can therefore be expressed:

HCl(aq) → H⁺ + Cl^-

This means that in a 0.01 mol/L solution of HCl it is approximated that there is a concentration of 0.01 mol/L dissolved hydrogen ions. From above, the pH is: pH = -log₁₀ [H⁺]:

pH = -log (0.01)

which equals 2.

For weak acids, the dissociation reaction does not go to completion. An equilibrium is reached between the hydrogen ions and the conjugate base. The following shows the equilibrium reaction between methanoic acid and its ions:

HCOOH(aq) ↔ H⁺ + HCOO^-

It is necessary to know the value of the equilibrium constant of the reaction for each acid in order to calculate its pH. In the context of pH, this is termed the acidity constant of the acid but is worked out in the same way (see chemical equilibrium):

K_a = [hydrogen ions][acid ions] / [acid]

For HCOOH, K_a = 1.6 × 10^-4 (some other Ka values)

When calculating the pH of a weak acid, it is usually assumed that the water does not provide any hydrogen ions. This simplifies the calculation, and the concentration provided by water, 1×10^-7 mol, is usually insignificant.

With a 0.1 mol/L solution of methanoic acid (HCOOH), the acidity constant is equal to:

K_a = [H⁺][HCOO^-] / [HCOOH]

Given that an unknown amount of the acid has dissociated, [HCOOH] will be reduced by this amount, while [H⁺] and [HCOO^-] will each be increased by this amount. Therefore, [HCOOH] may be replaced by 0.1 - x, and [H⁺] and [HCOO^-] may each be replaced by x, giving us the following equation:

Solving this for x yields 3.9×10^-3, which is the concentration of hydrogen ions after dissociation. Therefore the pH is -log(3.9×10^-3), or about 2.4.

Indicators

The Hydrangea macrophylla blossoms in pink or blue, depending on soil pH. In acid soils the flowers will be blue, in alkaline soils the flowers will be pink [1]

An indicator is used to measure the pH of a substance. Common indicators are litmus paper, phenolphthalein, methyl orange, and bromothymol blue.

See also

• Acid-base reaction theories
• Acid
• Base
• Alkali
• Soil pH
• Titration

References

• D. K. Nordstrom, C. N. Alpers, C. J. Ptacek, D. W. Blowes (2000). "Negative pH and Extremely Acidic Mine Waters from Iron Mountain, California." Environmental Science & Technology 34 (2), 254–258. (Available online: DOI | Abstract | Full text (HTML) | Full text (PDF))