pH

Physiology

(noun)

In chemistry, pH is a measure of the activity of the hydrogen ion concentration.

Related Terms

  • Many
  • buffer
  • bicarbonate
Chemistry

(noun)

The negative of the logarithm to base 10 of the concentration of hydrogen ions, measured in moles per liter; a measure of acidity or alkalinity of a substance, which takes numerical values from 0 (maximum acidity) through 7 (neutral) to 14 (maximum alkalinity).

Related Terms

  • titration
  • (HC2H3O2)
  • acid-base titration
  • pH indicator
  • buffer
  • titrant
  • analyte
  • equivalence point
  • indicator
  • stoichiometry
  • acid dissociation constant

Examples of pH in the following topics:

  • Microbial Growth at Low or High pH

    • Pure water has a pH very close to 7 at 25°C.
    • Solutions with a pH less than 7 are said to be acidic, and solutions with a pH greater than 7 are said to be basic or alkaline .
    • The pH scale is traceable to a set of standard solutions whose pH is established by international agreement.
    • Neutrophiles are organisms that thrive in neutral (pH 7) environments; extromophiles are organisms that thrive in extreme pH environments.
    • A pH scale with annotated examples of chemicals at each integer pH value
  • The Henderson-Hasselbalch Equation

    • The equation is also useful for estimating the pH of a buffer solution and finding the equilibrium pH in an acid-base reaction.
    • $-p{ K }_{ a }=-pH+log(\frac { [A^{ - }] }{ [HA] } )$
    • $pH=p{ K }_{ a }+log(\frac { { [A }^{ - }] }{ [HA] } )$
    • ${ 10 }^{ pH-p{ K }_{ a } }=\frac { [base] }{ [acid] }$
    • $pH=p{ K }_{ a }+log(\frac { { [NH_3}] }{ [NH_4^+] } )$
  • pH, Buffers, Acids, and Bases

    • Acids dissociate into H+ and lower pH, while bases dissociate into OH- and raise pH; buffers can absorb these excess ions to maintain pH.
    • Human cells and blood each maintain near-neutral pH.
    • Using the negative logarithm to generate positive integers, high concentrations of hydrogen ions yield a low pH, and low concentrations a high pH.
    • 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, lowering pH as more bicarbonate is created.
    • The pH scale measures the concentration of hydrogen ions (H+) in a solution.
  • pOH and Other p Scales

    • Here we have the reason that neutral water has a pH of 7.0 -; this is the pH at which the concentrations of H+ and OH- are exactly equal.
    • Relation between p[OH] and p[H] (brighter red is more acidic, which is the lower numbers for the pH scale and higher numbers for the pOH scale; brighter blue is more basic, which is the higher numbers for the pH scale and lower numbers for the pOH scale).
    • This lesson introduces the pH scale and discusses the relationship between pH, [H+], [OH-] and pOH.
    • Investigate whether changing the volume or diluting with water affects the pH.
    • Convert between pH and pOH scales to solve acid-base equilibrium problems.
  • Calculating Changes in a Buffer Solution

    • What is the pH of the solution?
    • Solving for the buffer pH after 0.0020 M NaOH has been added:
    • After adding NaOH, solving for $x=[H^+]$ and then calculating the pH = 3.92.
    • The pH went up from 3.74 to 3.92 upon addition of 0.002 M of NaOH.
    • Solving for the pH of a 0.0020 M solution of NaOH:
  • The Effect of pH on Solubility

    • By changing the pH of the solution, you can change the charge state of the solute.
    • The pH of an aqueous solution can affect the solubility of the solute.
    • By changing the pH of the solution, you can change the charge state of the solute.
    • As it migrates through a gradient of increasing pH, however, the protein's overall charge will decrease until the protein reaches the pH region that corresponds to its pI.
    • Describe the effect of pH on the solubility of a particular molecule.
  • Regulation of H+ by the Lungs

    • Acid-base imbalances in blood pH can be altered by changes in breathing to expel more CO2, which will raise pH back to normal.
    • Since maintaining normal pH is vital for life, and since the lungs play a critical role in maintaining normal pH, smokers have yet another reason to quit smoking.
    • Acid–base imbalance occurs when a significant insult causes the blood pH to shift out of the normal range (7.35 to 7.45).
    • When blood pH drops too low (acidemia), the body compensates by increasing breathing thereby expelling CO2, shifting the above reaction to the left such that less hydrogen ions are free; thus the pH will rise back to normal.
    • When blood pH drops too low, the body compensates by increasing breathing to expel more carbon dioxide.
  • Acid-Base Indicators

    • For applications requiring precise measurement of pH, a pH meter is frequently used.
    • For example, phenol red exhibits an orange color between pH 6.8 and pH 8.4.
    • Therefore, you would want an indicator to change in that pH range.
    • Both methyl orange and bromocresol green change color in an acidic pH range, while phenolphtalein changes in a basic pH.
    • Common indicators for pH indication or titration endpoints is given, with high, low, and transition pH colors.
  • Preparing a Buffer Solution with a Specific pH

    • It is used to prevent any change in the pH of a solution, regardless of solute.
    • In biology, they are necessary for keeping the correct pH for proteins to work; if the pH moves outside of a narrow range, the proteins stop working and can fall apart.
    • Then, measure the pH of the solution using a pH probe.
    • Once the pH is correct, dilute the solution to the final desired volume.
    • $pH=p{ K }_{ a }+log(\frac { { [A }^{ - }] }{ [HA] } )$
  • Chemical Buffer Systems

    • Chemical buffers such as bicarbonate and ammonia help keep blood pH in the narrow range compatible with life.
    • The body is very sensitive to its pH level, so strong mechanisms exist to maintain it.
    • Its pH changes very little when a small amount of strong acid or base is added to it.
    • Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH.
    • Several buffering agents that reversibly bind hydrogen ions and impede any change in pH exist.
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