pH

Examples of pH in the following topics:

• 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.
  • $-p{ K }_{ a }=-pH+log(\frac { [A^{ - }] }{ [HA] } )$
  • $pH=p{ K }_{ a }+log(\frac { { [A }^{ - }] }{ [HA] } )$
  • ${ 10 }^{ pH-p{ K }_{ a } }=\frac { [base] }{ [acid] }$
  • $pH=p{ K }_{ a }+log(\frac { { [NH_3}] }{ [NH_4^+] } )$

• 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.

• 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.
  • $pH=p{ K }_{ a }+log(\frac { { [A }^{ - }] }{ [HA] } )$

• Acid-Base Titrations

  • strong acid-weak base titration: methyl orange indicator the base is off the scale (e.g., pH > 13.5) and the acid has pH > 5.5: alizarine yellow indicator
  • the base is off the scale (e.g., pH > 13.5) and the acid has pH > 5.5: alizarine yellow indicator
  • the base is off the scale (e.g., pH > 13.5) and the acid has pH > 5.5: alizarine yellow indicator
  • the acid is off the scale (e.g., pH < 0.5) and the base has pH < 8.5: thymol blue indicator
  • You can determine the pH of a weak acid solution being titrated with a strong base solution at various points; these fall into four different categories: (1) initial pH; (2) pH before the equivalence point; (3) pH at the equivalence point; and (4) pH after the equivalence point.

• Relative Amounts of Acid and Base

  • The pH of a buffer depends on the ratio [base]/[acid] rather than on the particular concentration of a specific solution.
  • A buffer's pH changes very little when a small amount of strong acid or base is added to it.
  • It is therefore used to prevent change in the pH of a solution upon addition of another acid or base.
  • Suppose you wish to prepare a buffer solution to keep the pH at 4.30.
  • The more the ratio needs to differ to achieve the desired pH, the less effective the buffer.

• Strong Acid-Weak Base Titrations

  • A strong acid will react with a weak base to form an acidic (pH < 7) solution.
  • A known volume of base with unknown concentration is placed into an Erlenmeyer flask (the analyte), and, if pH measurements can be obtained via electrode, a graph of pH vs. volume of titrant can be made (titration curve).
  • As the equivalence point is approached, the pH will change more gradually, until finally one drop will cause a rapid pH transition through the equivalence point.
  • In strong acid-weak base titrations, the pH at the equivalence point is not 7 but below it.
  • A depiction of the pH change during a titration of HCl solution into an ammonia solution.

• Buffer Range and Capacity

  • A buffer's capacity is the pH range where it works as an effective buffer, preventing large changes in pH upon addition of an acid or base.
  • When H+ is added to a buffer, the weak acid's conjugate base will accept a proton (H+), thereby "absorbing" the H+ before the pH of the solution lowers significantly.
  • In biological systems, buffers prevent the fluctuation of pH via processes that produce acid or base by-products to maintain an optimal pH.
  • Each conjugate acid-base pair has a characteristic pH range where it works as an effective buffer.
  • The buffering region is about 1 pH unit on either side of the pKa of the conjugate acid.