pH indicator

(noun)

An acid-base indicator.

Related Terms

Examples of pH indicator in the following topics:

  • Acid-Base Indicators

    • For applications requiring precise measurement of pH, a pH meter is frequently used.
    • These commercial indicators (e.g., universal indicator and Hydrion papers) are used when only rough knowledge of pH is necessary.
    • 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.
    • Common indicators for pH indication or titration endpoints is given, with high, low, and transition pH colors.

  • 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.

  • Weak Bases

    • Both actions raise the pH of the solution by decreasing the concentration of H+ ions.
    • The pH of bases in aqueous solution ranges from greater than 7 (the pH of pure water) to 14 (though some bases have pH values greater than 14).
    • The formula for pH is:
    • Sometimes, however, it is more convenient to focus on the pOH of bases, rather than the pH.
    • Smaller pKb values indicate higher values of Kb; this also indicates a stronger base.

  • Weak Acid-Strong Base Titrations

    • A weak acid will react with a strong base to form a basic (pH > 7) solution.
    • The titration curve demonstrating the pH change during the titration of the strong base with a weak acid shows that at the beginning, the pH changes very slowly and gradually.
    • When the NaOH is in excess, the pH change is the same as in any system dominated by NaOH.
    • However, the pH at the equivalence point does not equal 7.
    • This figure depicts the pH changes during a titration of a weak acid with a strong base.

  • 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.

  • 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.

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

  • Buffers Containing a Base and Conjugate Acid

    • Reactions with weak bases result in a relatively low pH compared to strong bases.
    • Bases range from a pH of greater than 7 (7 is neutral like pure water) to 14 (though some bases are greater than 14).
    • The pH of bases is usually calculated using the OH- concentration to find the pOH first.
    • Calculate the pH of a buffer solution consisting of 0.051 M NH3 and 0.037 M NH4+.
    • Calculate the pH of an alkaline buffer system consisting of a weak base and its conjugate acid.

  • 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:

  • 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.