pH

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.