# dissociation

(noun)

## Definition of dissociation

Referring to the process by which compounds split into smaller constituent molecules, usually in a reversible manner.

Source: Wiktionary - CC BY-SA 3.0

## Examples of dissociation in the following topics:

• ### Base Dissociation Constant

• The base dissociation constant, Kb, is a measure of basicity -- the general strength of the base.
• 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+.
• $B + H_2O \leftrightarrow HB^+ + OH^−$Using similar reasoning to that used for the acid dissociation constant, we can derive an expression for the base dissociation constant.
• In water, the concentration of the hydroxide ion, [OH−], is related to the concentration of the hydrogen ion by the ionic product of water, Kw = [H+] x [OH−], (Kw is the water dissociation constant).
• The base dissociation constant is the equivalent of the acid dissociation constant for bases and represents the strength of a base.
• ### Acid Dissosociation Constant (Ka)

• Acid Dissociation Constant (Ka)An acid dissociation constant, Ka, is a quantitative measure of the strength of an acid in solution.
• It is the equilibrium constant for a chemical reaction; it is known as dissociation in the context of acid-base reactions.
• $HA \rightleftharpoons H^+ + A^-$The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L):$Ka = [A-][H+]/[HA]$Due to the many orders of magnitude spanned by Ka values, a logarithmic measure of the acid dissociation constant is more commonly used in practice.
• The logarithmic constant, pKa, which is equal to -log10Ka, is sometimes (incorrectly) referred to as an acid dissociation constant.The larger the value of pKa, the smaller the extent of dissociation.
• A strong acid is almost completely dissociated in aqueous solution; it is dissociated to the extent that the concentration of the undissociated acid becomes undetectable. pKa values for strong acids can be estimated by theoretical means or by extrapolating from measurements in non-aqueous solvents with a smaller dissociation constant, such as acetonitrile and dimethylsulfoxide.Acetic acid is a weak acid.
• The acid dissociation constant, Ka, is the measure of the strength of an acid in solution.
• ### Calculating Equilibrium Concentrations of Polyprotic Acids

• The dissociation constant of the first proton may be denoted as Ka1 and the constants for dissociation of successive protons as Ka2, etc.
• The following examples indicate the mathematics and simplifications for a few polyprotic acids under specific conditions.Sulfuric AcidIf water is the solvent, sulfuric acid, H2SO4, loses one proton as a strong acid with an immeasurably large dissociation constant.H2SO4 → H+ + HSO4-It also can lose a second proton as a weak acid with a measurable dissociation constant.
• For the second dissociation of phosphoric acid, for which pKa2 = 7.21:pKa2 = -log(Ka2) = -log([H+]*[HPO42-]/[H2PO4-])pH = -log[H+]Therefore, pH - pKa2 = log([HPO42-]/[H2PO4-])When pH = pKa2, we have the ratio [HPO42-]/[H2PO4-] = 1.00.
• When a weak polyprotic acid, such as carbonic acid, H2CO3, dissociates, most of the protons present come from the first dissociation:$H_2CO_3 \rightleftharpoons H^+ + HCO_3^-$  pKa1 = 6.37Since 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.
• Correspondingly, the second dissociation has a negligible effect on the concentration of the product of the first dissociation, HCO3-.
• Polyprotic acids have complex equilibriums due to the dissociation of hydrogen species at different pHs and the presence of multiple species in solution.
• ### Calculating Percent Dissociation

• Weak acids in water dissociate as: $HA (aq) \rightleftharpoons H^+(aq) + A^- (aq)$ The strength of a weak acid is represented as either an equilibrium constant or as a percent dissociation.
• The percent dissociation is symbolized as α (alpha) and which can range from 0%< α < 100%.
• Instead, the percent dissociation often must be calculated from the given information of pKa and pH.
• Thus the percent dissociation of the acetic acid will decrease and the pH of the solution will increase.
• This, in turn, affects the percent dissociation of the acid.
• The percent dissociation represents the strength of an acid and can be calculated through indirect information, like the pKa and pH.
• ### The Acid Dissociation Constant

• To understand the acid dissociation constant, it is first important to understand the equilibrium equation for acid dissocation.
• This gives rise to a special equilibrium constant known as theacid dissociation constant, Ka.
• 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.
• Therefore, the larger the value of pKa, the smaller the extent of dissociation.
• The acid dissociation constant measures the strength of an acid and is essential for understanding acid-base equilibria in solution.
• ### The Form of K and the Equilibrium Equation

• Binding constants, water ionization constant, association constants, and dissociation constants are all types of equilibrium constant.
• The association, or binding, constant Ka equation is:${K}_{a}=\frac{[RL]}{[R][L]}$The reverse, or dissociation constant, Kd is also frequently used:${K}_{d}=\frac{[R][L]}{[RL]}$These equilibrium constants are critical in the understanding of how proteins interact with each other.
•  The balanced equation is:${H}_{2}O+{H}_{2}O\rightleftharpoons{H}_{3}{O}^{+}+O{H}^{-}$The equilibrium constant, Kw, for this reaction is given by:${K}_{w}=\frac{[{H}_{3}{O}^{+}][O{H}^{-}]}{{[{H}_{2}O]}^{2}}$Because H2O is a pure liquid (that is, itsconcentration is high enough it does not change during the reaction), the equation simplifies to:Sometimes, HO , which is the concentration of the freed hydrogen, is written simply as H, which simplifies the equation to:${K}_{w}=[{H}_{3}{O}^{+}][O{H}^{-}]$Sometimes, H3O+ , which is the concentration of the freed hydrogen, is written simply as H+, which simplifies the equation to:${K}_{w}=[{H}^{+}][O{H}^{-}]$At 25°C, Kw is equal to 1.0 \times 10−14.ACID DISSOCIATION AND BASE IONIZATIONAn acid dissociation constant, Ka, is the equilibrium constant for a chemical reaction known as dissociation in the context of acid-base reactions.
• The formula for the acid dissociation constant is:${K}_{a}=\frac{{[A}^{-}]{[H}^{+}]}{[HA]}$Ka is the measure of the strength of an acid.
• Due to the many orders of magnitude spanned by Ka values, a logarithmic measure of the acid dissociation constant is more commonly used in practice:$p{K}_{a}=-log({K}_{a})$The larger the value of pKa, the smaller the extent of dissociation.
• Common reactions, like binding, water ionization, association, and dissociation, all have specially named equilibrium constants.
• ### Diprotic and Polyprotic Acids

• The dissociation does not happen all at once; the two stages of dissociation have different Ka values.
• ### Weak Acids

• Partial dissociation means that it does not release all of its hydrogens in a solution.
• Dissociation Weak acids ionize in a water solution to only a moderate extent; that is, if an acid is represented by the general formula HA, then in an aqueous solution, a significant amount of undissociated HA remains.
• Weak acids in water dissociate as: $HA \leftrightarrow H^+ (aq) + A^- (aq)$ 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: $K_a = \frac{([H^+]*[A^-])}{[HA]}$ The greater the value of Ka, the more the formation of H+ is favored, which makes the solution more acidic.
• The first Ka refers to the first dissociation: $H_2CO_3 + H_2O \rightarrow HCO_3^{-} + H_3O^+$ This Ka value is 4.46×10−7 (pKa1 = 6.351).
• ### Strong Acids

• For sulfuric acid, which is diprotic, the "strong acid" designation refers only to the dissociation of the first proton:$H_2SO_4 (aq) \rightarrow H^+ (aq) + HSO_4^− (aq)$More precisely, the acid must be stronger in aqueous solution than a hydronium ion (H+), so strong acids are acids with a pKa < -1.74.
• This generally means that in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution.Due to the complete dissociation of strong acids in aqueous solution, the concentration of hydronium ions in the water is equal to the total concentration (ionized and un-ionized) of the acid introduced to solution: [H+] = [A−] = [HA]total and pH = −log[H+].Strong acids, like strong bases, can cause chemical burns when exposed to living tissue.
• Examples of Strong AcidsSome common strong acids (acids with pKa < -1.74) include: Hydroiodic acid (HI): pKa = -9.3Hydrobromic acid (HBr): pKa = -8.7Perchloric acid (HClO4): pKa ≈ -8Hydrochloric acid (HCl): pKa = -6.3Sulfuric acid (H2SO4): pKa1 ≈ -3 (first dissociation only)p-Toluenesulfonic acid: pKa = -2.8p-Toluenesulfonic acid is an example of an organic soluble strong acid Strong Acid CatalysisStrong acids can accelerate the rate of certain reactions.