ka

(noun)

The spiritual part of an individual human being or god that survived after death.

Related Terms

Examples of ka in the following topics:

  • Acid Dissociation Constant (Ka)

    • The Ka expression is as follows:
    • The logarithmic constant (pKa) is equal to -log10(Ka).
    • The larger the value of pKa, the smaller the extent of dissociation.
    • Acids with a pKa value of less than about -2 are said to be strong acids.
    • What is the pKa for acetic acid?

  • Weak Acids

    • Weak acids have very small values for Ka (and therefore higher values for pKa) compared to strong acids, which have very large Ka values (and slightly negative pKa values).
    • The Ka of weak acids varies between 1.8×10−16 and 55.5.
    • The first Ka refers to the first dissociation step:
    • This Ka value is 4.46×10−7 (pKa1 = 6.351).
    • The Ka of acetic acid is $1.8\times 10^{-5}$.

  • The Acid Dissociation Constant

    • Ka is a quantitative measure of the strength of an acid in solution.
    • The logarithmic constant, pKa, which is equal to −log10 (Ka), is sometimes incorrectly referred to as an acid dissociation constant as well.
    • Smaller Ka values yield larger pKa values.
    • An understanding of Ka is also essential for working with buffers; the design of these solutions depends on a knowledge of the pKa values of their components.
    • Another important application of Ka is with pH indicators.

  • Strong Acids

    • More precisely, the acid must be stronger in aqueous solution than a hydronium ion (H+), so strong acids have a pKa < -1.74.
    • An example is hydrochloric acid (HCl), whose pKa is -6.3.
    • p-Toluenesulfonic acid is an example of an organic soluble strong acid, with a pKa of -2.8.

  • The Henderson-Hasselbalch Equation

    • The Henderson–Hasselbalch equation connects the measurable value of the pH of a solution with the theoretical value pKa.
    • The Henderson–Hasselbalch equation mathematically connects the measurable pH of a solution with the pKa (which is equal to -log Ka) of the acid.
    • The equation can be derived from the formula of pKa for a weak acid or buffer.
    • With a given pH and known pKa, the solution of the Henderson-Hasselbalch equation gives the logarithm of a ratio which can be solved by performing the antilogarithm of pH/pK­a:
    • What is the pH of a buffer solution consisting of 0.0350 M NH3 and 0.0500 M NH4+ (Ka for NH4+ is 5.6 x 10-10)?

  • Calculating Percent Dissociation

    • Percent dissociation represents an acid's strength and can be calculated using the Ka value and the solution's pH.
    • We have already discussed quantifying the strength of a weak acid by relating it to its acid equilibrium constant Ka; now we will do so in terms of the acid's percent dissociation.
    • Calculate the percent dissociation of a weak acid in a $0.060\;M$ solution of HA ($K_a=1.5\times 10^{-5}$).
    • Calculate percent dissociation for weak acids from their Ka values and a given concentration.

  • Relative Amounts of Acid and Base

    • The exact ratio of the base to the acid for a desired pH can be determined from the Ka value and the Henderson-Hasselbalch equation.
    • Of the acids listed, the Ka value for acetic acid is closest to the desired hydrogen ion concentration.
    • The pKa of acetic acid is
    • Extrapolating further from this, a buffer is most effective when the concentrations of acid and conjugate base (or base and conjugate acid) are approximately equal—in other words, when the log [base]/[acid] equals 0 and the pH equals the pKa.

  • Oxoacids

    • These acids can be arranged in order of their pKa values and, by extension, their relative strengths:
    • HOCl pKa = 7.5 < HOBr pKa = 8.6 < HOI pKa = 10.6
    • Recall that smaller values of pKa correspond to greater acid strength.
    • Consider the family of chlorooxoacids, which are arranged below in order of pKa values:
    • HOClO3 pKa = -8 < HOClO2 pKa = -1.0 < HOClO pKa = 1.92 < HOCl pKa = 7.53

  • Specialized Equilibrium Constants

    • An acid dissociation constant, Ka, is the equilibrium constant for the dissociation of an acid in aqueous solution.
    • As such, strong acids will have large values of Ka that are greater than one, which indicates that the forward reaction of dissociation is strongly favored.
    • Weak acids, on the other hand, will have small values of Ka that are less than one, indicating that the reverse reaction is strongly favored; weak acids dissociate only to a small extent.
    • As such, Ka acts a relative indicator of acid strength.

  • Overview of the Acid-Base Properties of Salt

    • We can determine the answer by comparing Ka and Kb values for each ion.
    • In this case, the value of Kb for bicarbonate is greater than the value of Ka for ammonium.
    • In summary, when a salt contains two ions that hydrolyze, compare their Ka and Kb values: