acid dissociation constant

Examples of acid dissociation constant in the following topics:

• Acid Dissociation Constant (Ka)

  • The acid dissociation constant (Ka) is the measure of the strength of an acid in solution.
  • The acid dissociation constant (Ka) is a quantitative measure of the strength of an acid in solution.
  • Ka is the equilibrium constant for the following dissociation reaction of an acid in aqueous solution:
  • Acid dissociation constants are most often associated with weak acids, or acids that do not completely dissociate in solution.
  • Acetic acid is a weak acid with an acid dissociation constant $K_a=1.8\times 10^{-5}$ .

• The Acid Dissociation Constant

  • The acid dissociation constant measures the strength of an acid and is essential for understanding acid-base equilibria in solution.
  • To understand the acid dissociation constant, it is first important to understand the equilibrium equation for acid dissocation.
  • It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid-base reactions.
  • The logarithmic constant, pKa, which is equal to −log10 (Ka), is sometimes incorrectly referred to as an acid dissociation constant as well.
  • Discuss the quantitative and qualitative relationship between acid dissociation constant (Ka) and the equilibrium constant for solutions

• Specialized Equilibrium Constants

  • The value of the dissociation constant of water, KW, is $1.0\times 10^{-14}$.
  • An acid dissociation constant, Ka, is the equilibrium constant for the dissociation of an acid in aqueous solution.
  • The base dissociation constant, Kb, is analogous to the acid dissociation constant.
  • As with the acid dissociation constant, large values of Kb are indicative of a stronger base, while small values of Kb are indicative of a weaker base.
  • A ball-and-stick model of the dissociation of acetic acid to acetate.

• Diprotic and Polyprotic Acids

  • With any polyprotic acid, the first amd most strongly acidic proton dissociates completely before the second-most acidic proton even begins to dissociate.
  • The first dissociation constant is necessarily greater than the second ( i.e.
  • This first dissociation step of sulfuric acid will occur completely, which is why sulfuric acid is considered a strong acid; the second dissociation step is only weakly dissociating, however.
  • A triprotic acid (H3A) can undergo three dissociations and will therefore have three dissociation constants: Ka1 > Ka2 > Ka3.
  • The following formula shows how to find this fractional concentration of HA-, in which pH and the acid dissociation constants for each dissociation step are known:

• Weak Acids

  • A weak acid is one that does not dissociate completely in solution; this means that a weak acid does not donate all of its hydrogen ions (H+) in a solution.
  • On average, only about 1 percent of a weak acid solution dissociates in water in a 0.1 mol/L solution.
  • The generalized dissociation reaction is given by:
  • The strength of a weak acid is represented as either an equilibrium constant or a percent dissociation.
  • The equilibrium concentrations of reactants and products are related by the acid dissociation constant expression, Ka:

• Calculating Equilibrium Concentrations of Polyprotic Acids

  • The first proton's dissociation may be denoted as Ka1 and the constants for successive protons' dissociations as Ka2, etc.
  • As we are already aware, sulfuric acid's first proton is strongly acidic and dissociates completely in solution:
  • When a weak diprotic acid such as carbonic acid, H2CO3, dissociates, most of the protons present come from the first dissociation step:
  • Since the second dissociation constant is smaller by four orders of magnitude (pKa2 = 10.25 is larger by four units), the contribution of hydrogen ions from the second dissociation will be only one ten-thousandth as large.
  • The above concentration can be used if pH is known, as well as the two acid dissociation constants for each dissociation step; oftentimes, calculations can be simplified for polyprotic acids, however.

• Absolute Concentrations of the Acid and Conjugate Base

  • A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
  • Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of situations.
  • The Ka for acetic acid is 1.8 x 10-5.
  • Once again, using the acid dissociation constant, we can solve for x to get [H+] = 2.11 x 10-5 M.
  • 8.1.3 Deduce the formula of the conjugate acid/base of any Brønsted-Lowry base/acid IB Chemistry SL - YouTube

• Base Dissociation Constant

  • The base dissociation constant, Kb, is a measure of basicity—the base's general strength.
  • It is related to the acid dissociation constant, Ka, by the simple relationship pKa + pKb = 14, where pKb and pKa are the negative logarithms of Kb and Ka, respectively.
  • The base dissociation constant can be expressed as follows:
  • Historically, the equilibrium constant Kb for a base has been defined as the association constant for protonation of the base, B, to form the conjugate acid, HB+.
  • The general equation for a base dissociation constant, where B is the base, HB is its conjugate acid, and OH- is hydroxide ions.

• Calculating Percent Dissociation

  • 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.
  • However, because the acid dissociates only to a very slight extent, we can assume x is small.
  • As we would expect for a weak acid, the percent dissociation is quite small.
  • However, for some weak acids, the percent dissociation can be higher—upwards of 10% or more.
  • Calculate percent dissociation for weak acids from their Ka values and a given concentration.

• 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.
  • The concentration of hydrogen ions dissociating from pure water is 1 × 10-7 moles H+ ions per liter of water.
  • An acid is a substance that increases the concentration of hydrogen ions (H+) in a solution, usually by dissociating one of its hydrogen atoms.
  • For example, hydrochloric acid (HCl) is highly acidic and completely dissociates into hydrogen and chloride ions, whereas the acids in tomato juice or vinegar do not completely dissociate and are considered weak acids; conversely, strong bases readily donate OH– and/or react with hydrogen ions.
  • Maintaining a constant blood pH is critical to a person's well-being.