Buffer Intensity and Buffer Capacity

pH Buffers and Alkalimetric Titration

A pH buffer is a mixture of a weak acid and a base. Let’s consider a diprotic acid H2A to which a monoacidic base BOH (with B+ = Na+, K+, or NH4+) is added:1

(1) H2A + n BOH  =  BnH2-n A + n H2O

Example: For H2CO3 and NaOH we obtain – in the special case of n = 0, 1, and 2 – pure solutions of H2CO3, NaHCO3 and Na2CO3 (see Equivalence Points).

The parameter n acts as a stoichiometric coefficient that embodies the ratio of added amount of strong base, CB, to the total amount of the diprotic acid CT:

(2)   n = CB /CT     or     CB = n CT

buffer system: acid plus base at chemical equilibrium

Both CB and CT are given in molar units (mol/L) while n is dimensionless (unitless).

The entity BnH2-n A in reaction (1) does not survive in water; it dissociates into several aqueous species – as indicated in the right picture.

As we will see below, the amount of strong base is just the total alkalinity, Alk = CB. Thus we have:

(3)   n = Alk /CT     or     Alk = n CT

The variation of n (or CB) in reaction formula (1) by adding a strong base is called alkalimetric titration. Note: The reason for our focus on diprotic acids is that it comprises carbonic acid H2CO3 as the most prominent representative (i.e. A-2 = CO3-2).

Notation. To keep the notation straight we abbreviate the activity of H+ by x. Its relationship with pH is given by:

(4)   x ≡ {H+} = 10-pH pH = – lg {H+} = – lg x

The (molar) concentrations of the acid species [H2 A], [HA-] and [A-2] are abbreviated as

(5)   [j]  =  [H2-j A-j] with   j = 0, 1, 2

In addition, we introduce the “alkalinity of pure water” by2

(6)   w  =  [OH-] – [H+]  ≈  Kw /x – x

where Kw = 10-14 is the equilibrium constant of autoprotolysis (self-ionization of water). As we will see later, w represents the contribution of self-ionization to the alkalinity. For pure water – defined by x = (Kw)1/2 = 10-7, i.e. pH = 7, – this contribution vanishes: w = 0.

Starting-Point: Basic Set of Equations

The set of mathematical equations to describe the alkalimetric titration is exactly the same as we already introduced for diprotic acids and their conjugate bases:3

(1.1a)   K1 =  {H+} {HA-} / {H2A} (1st diss. step)
(1.1b)   K2 =  {H+} {A-2} / {HA-} (2nd diss. step)
(1.1c)   Kw =  {H+} {OH-} (self-ionization)
(1.1d)   CT =  [H2A] + [HA-] + [A-2] (mass balance)
(1.1e)   n CT =  [HA-] + 2 [A-2] + [OH-] – [H+] (charge balance)

The only difference to the previous approach is that n is now a smooth variable of any value (rather than merely a fixed integer n = 0, 1, 2).

For the special case of n = 0, the set reduces to the description of the pure diprotic-acid system, where {H+} (or pH) is determined by the amount of CT, and vice versa. Now, with the new variable n the system gets an additional degree of freedom.

Alkalinity.  The right-hand side of 1.1e is just the definition of the total alkalinity (also known as M alkalinity in carbonate systems),

(1.2)   Alk  ≡  ([OH-] – [H+]) + [HA-] + 2 [A-2]

This justifies our statement, which we originally made in 3:

(1.3)   Alk  =  CB  =  n CT

Ionization Fractions (a0, a1, a2)

Our equation system consists of 5 equations. Everything concerning the acid itself is contained in a subset of three equations: (1.1a), (1.1b), and (1.1d). The information inherent in this subsystem can be expressed by ionization fractions (instead of the three acid species [j]):

(2.1)   aj  =  [j] / CT with   j = 0, 1, 2

These ‘normalized concentrations’, which indicate the degree of dissociation, are completely independent of the amount of acid CT. Dividing 1.1d by CT yields:

(2.2)   1  =  a0 + a1 + a2 (mass balance)

From a mathematical point of view, the ionization fractions are nothing else than a clever way to combine 1.1a and (1.1b) using 2.2:

(2.3a)   a0   =   [ 1 + K1/x + K1K2/x2 ]-1  
(2.3b)   a1   =   [ x/K1 + 1 + K2/x ]-1  
(2.3c)   a2   =   [ x2/(K1K2) + x/K2 + 1 ]-1  

They are solely functions of x (or pH); the only other ingredients are two equilibrium constants, K1 and K2. Their pH dependence is shown here. Each aj is imprisoned between 0 and 1, i.e. it neither become negative nor greater than 1. Strictly speaking, the functions will come very close to their boundaries, but never really reach 0 and 1:

(2.4) 0 < aj < 1 for   j = 0, 1, 2

There are three exceptional points of pH (or x):4

(2.5a)   pH = pK1 x = K1 a1 = a0 ≈ ½ a2 ≈ 0
(2.5b)   pH = ½ (pK1+pK2) x = (K1K2)1/2 a0 = a2 ≈ 0 a1 ≈ 1
(2.5c)   pH = pK2 x = K2 a1 = a2 ≈ ½ a0 ≈ 0

In what follows, the three ionization fractions a0, a1, and a2 will become the building blocks of almost all relevant acid-base quantities. In this respect, don’t confuse the smooth variable n with the index j for the three integers j = 0, 1, 2. The latter are dedicated to the three ionization fractions aj and the three dissolved species, H2A, HA-, and A-2.

Analytical Solution of the System of Equations

The set of 5 equations in (1.1) can be cast into a single analytical formula:

(3.1a) \(C_T(x) \, = \, \dfrac{x - K_w/x} {a_1 + 2a_2 - n}\)

Here, the pH dependence (in form of x = 10-pH) is also in a1 and a2. To make it evident, you must insert (2.3b) and (2.3c) into 3.1a:

(3.1b) \(C_T(x) \, = \, \left(x-\dfrac{K_w}{x}\right) \ \left(\dfrac{1+2K_2/x} {x/K_1 + 1 + K_2/x}-n\right)^{-1}\)

Both formulas calculate CT for a given pH (or x) and n. Often, however, one is interested in the “inverse task”: to calculate the pH for a given CT and/or n. For this purpose, you need to rearrange 2.1b to isolate x. But that’s impossible. The only thing you can achieve is a polynomial of 4th order in x:

(3.2)   x4 + {K1 + nCT} x3 + {K1K2 + (n–1)CT K1 – Kw} x2 + K1 {(n–2)CT K2 – Kw} x – K1K2Kw  =  0

This quartic equation (quartic!, not quadratic)5 predicts the pH for any given pair of CT and n. Replacing nCT by CB in 3.2, you get an alternative form:

(3.3)   x4 + {K1 + CB} x3 + {K1K2 + (CB – CT) K1 – Kw} x2 + K1 {(CB – 2CT) K2 – Kw} x – K1K2Kw  =  0

Once the pH value (or x) is known, the concentrations of the acid species, H2 A, HA-, and A-2, can be calculated via the ionization fractions aj as follows:

(3.4)   [H2-j A-j]  =  CT aj (x) for   j = 0, 1, 2

Exact Relationships between pH, CT, and n (or Alk = nCT)

The set of five mathematical equations in (1.1a) to (1.1e) contains all information about the “diprotic-acid plus strong base” system (alkalimetric titration). In particular, the actual equilibrium state (i.e. the concentrations of the acid species H2A, HA-, and A-2) is controlled by two parameters chosen from the triple (CT, n, pH) or (CT, Alk, pH). Once you know two of them, the third is automatically determined:

(4.1a)   pH (CT,n) = –lg x1  with  x1 = positive root of 3.2   [j] = CT aj
(4.1b)   pH (CT,Alk) = –lg x1  with  x1 = positive root of 3.3   [j] = CT aj
(4.2a)   n (CT,pH) = a1 + 2a2 + w/CT   [j] = CT aj
(4.2b)   Alk (CT,pH) = CT (a1 + 2a2) + w   [j] = CT aj
(4.3a)   CT (n,pH) = w/(n – a1 – 2a2)   [j] = [w/(n – a1 – 2a2)] aj
(4.3b)   CT (Alk,pH) = (Alk – w)/(a1 + 2a2)   [j] = [(Alk – w)/(a1 + 2a2)] aj

with w(x) defined in 6. Here [j] stands for the molar concentration of the three acid species.

Note: All formulas presented here are different encodings of one and the same thing, namely the set of equations (1.1a) to (1.1e).

Cross Plots.  It’s quite instructive to exhibit the non-linearity of these relationships graphically. The diagrams below display all possible combinations of one dependent and two independent variables (taken from the triplet CT, n, and pH) for the carbonate system:

dependent and independent variables taken from the triplet (CT, n, and pH) of the carbonate buffer system

The diagrams can be redrawn for the case when n is replaced by the alkalinity (Alk = n CT):

dependent and independent variables taken from the triplet (CT, Alkalinity, and pH) of carbonate buffer system

Other Examples.  Application of 4.2a to a pure CO2 system provides the (blue) titration curves in the diagram below. On the other hand, for the special case of n = 0, 1, 2, 4.3a was used to plot the pH dependence of three equivalence points.

Buffer Capacity

Alkalinity is known to be a buffer capacity (and therefore it is what we have called CB). Thus, the titration curves based on 4.2b display the buffer capacity as a function of pH. In this sense, n(pH) = CB/CT can be understood as the “normalized” or unitless buffer capacity. From the above formulas we get:

(5.1a)   buffer capacity (in mol/L):   CB = (a1 + 2a2) CT + w
(5.1b)   normalized buffer capacity:   n = (a1 + 2a2) + w/CT

Each individual term in these “titration formulas” is pH-dependent:  a1(x), a2(x), CT(x), and w(x). In particular, 5.1b can be written as:

(5.2)   n (CT,x)  =  \(a_1 + 2a_2 + \dfrac{w}{C_T} \ =\ \dfrac{1+2K_2/x} {x/K_1 + 1 + K_2/x} + \dfrac{K_w/x - x}{C_T}\)

Buffer Intensity

The buffer intensity is the 1st derivative of the buffer capacity with respect to pH. Again, we distinguish between two types of buffer intensities:

(6.1a)   buffer intensity (in mol/L): βc = dCB / d pH  =  β CT
(6.1b)   normalized buffer intensity: β  = dn / d pH

buffered and none-buffered systems

The steeper the slope of a titration curve the higher is the buffer intensity β = dn/d(pH), i.e. the higher is the system’s resistance to pH changes (caused by a strong base). Thus, the pH where β reaches its maximum represents the optimal buffer range (bounded by pHmax ±1) – see example below:

optimal buffer range (carbonate system)

Formulas.  The 1st pH-derivative of 5.2 yields (see Appendix):

(6.2a)   β  =  ln 10 {(a2 + a0) – (a2 – a0)2} + ln 10 (w+2x) /CT

or, more explicitly,

(6.2b)   \(\beta \ = \ ln\,10 \, \left\{\dfrac{x/K_1+4K_2/K_1+K_2/x}{(x/K_1+1+K_2/x)^2} \, +\ \dfrac{x+K_w/x}{C_T} \right\}\)

Equation (6.2b), that contains only positive summands, tells us that β is always positive (it never becomes negative or zero).

Local Extrema. The local maxima (and minima) of the buffer intensity are at the following points:

x = K1 maximum optimal buffer range
x = (K1K2)1/2 minimum    
x = K2 maximum optimal buffer range (small CT excluded)

Please note that this is not strictly exact, but it’s a very good approximation. Mathematically, local extrema of β are found where the first-order derivative of β is zero.

Example: The buffer intensity of a pure CO2 system is displayed as green curve in the diagram below.

Buffer Capacity. The buffer capacity is just the integral of the buffer intensity:

(6.3) buffer capacity = \(\int_a^b \beta_c \, dpH\)

Derivative of the Buffer Intensity

The derivative of the buffer intensity is given by:

(7.1)   \(\dfrac{d\beta}{d\,pH} \ = \ (ln\,10)^2 \, \left\{\dfrac{f(x)}{(x/K_1+1+K_2/x)^3} \, +\ \dfrac{K_w/x - x}{C_T} \right\}\)

with the abbreviation

(7.2)   f(x) = (x/K1 – K2/x) (x/K1 + K2/x – 1 + 8 K2/K1)

Example: The derivative of the buffer intensity of a pure CO2 system is displayed as red curve in the diagram below. The zeros of dβ/d(pH) are marked by blue circles.

Example Calculation for the Carbonate System

Given is a pure CO2 system with CT = 100 mM, 10 mM, and 1 mM H2CO3.6 Three quantities are plotted as function of pH:

•  n(pH) (normalized) amount of strong base added7 (as blue curve)
•  β = dn/dpH (normalized) buffer intensity (as green curve)
•  dβ/dpH derivative of the buffer intensity (as red curve)

All three quantities as per definition are unitless. The small blue dots mark the zeros of dβ/d(pH).

aqion - buffer intensities of carbonate system

Since the (blue) titration curve is an ever-increasing function, its derivative, i.e. the buffer intensity β, is always positive (see green curve). This is in full accord with Le&nsp;Châtelier’s principle that every solution resists pH changes.

Appendix – Derivatives with Respect to pH

The first and second derivatives of x = {H+} with respect to pH are

(A1)   dx/d(pH)  =  (–ln 10) x and   d2x/d(pH)2  =  (–ln 10)2 x

This result can be used to differentiate any given function f(x) with respect to pH (by application of the chain rule):

(A2)   \(\dfrac{d\,f(x)}{d\,pH} \ = \ \left(\dfrac{dx}{d\,pH}\right) \left(\dfrac{d\,f(x)}{dx}\right) \ = \ (-ln\,10)\, x \ \dfrac{d\,f(x)}{dx}\)

Hence, for w(x) = Kw/x – x, introduced in 6, we get the amazing result:

(A3a)   dw/d(pH) =   ln 10 (Kw/x + x) =   ln 10 (w + 2x)
(A3b)   d2w/d(pH)2 =   (ln 10)2 (w + 2x – 2x) =   (ln 10)2 w

All derivatives of higher degree repeat this pattern:

(A3c)   \(\dfrac{d^{k}w}{d\,(pH)^k} \ = \ (ln 10)^k \, \left\{\begin{matrix}\ w& &for & k\ even & \\ \ w+2x && for & k\ odd & \end{matrix}\right.\)

Ionization Fractions. The three ionization fractions, introduced in 2.3a to (2.3c), obey the following relations (where the sum runs from i = 0 to 2):

(A4a)   Y0  ≡  Σi ai =  a0 + a1 + a2 =  1  
(A4b)   Y1  ≡  Σi i ai =  0⋅a0 + 1⋅a1 + 2⋅a2 =  a1 + 2 a2  
(A4c)   Yk  ≡  Σi ik ai =  0⋅a0 + 1⋅a1 + 2k⋅a2 =  a1 + 2k a2 (for all k > 0)

Quantities so defined are also called kth moments. Here in our case, Y0 represents the mass balance, Y1 enters the alkalinity (as its main part), and Y2 together with Y3 will be used for the buffer intensity and its derivative. All Yk (excluding Y0) are positive numbers, living in the range 0 < Yk < 2k.

The first derivative of the ionization fractions is given by

(A5)   dai /d(pH)  =  ln 10 (i – a1 – 2a2) ai  =  ln 10 (i – Y1) ai for i = 0, 1, 2

Applying it to the two sums in A4a and (A4b) yields

(A6a)   (d/d pH) Σi ai =   ln 10 Σi (i – Y1) ai =   ln 10 (Y1 – Y1)  =  0
(A6b)   (d/d pH) Σi i ai =   ln 10 Σi i (i – Y1) ai =   ln 10 (Y2 – Y12)

The first relation, which delivers zero, is obvious because it represents the derivation of a constant, namely d1/d(pH) = 0. In general we have

(A7)   d Yk /d pH   =   ln 10 Σi ik (i – Y1) ai   =   ln 10 (Yk+1 – Y1Yk)

The second derivative of Y1 is given by

(A8a)   \(\dfrac{d^2\,Y_1}{d(pH)^2}\) =   \((ln\,10) \,\dfrac{d}{d\, pH}\, (Y_2-Y_1^2)\)
      =   (ln 10)2 (Y3 – Y1Y2 – 2Y1 (Y2 – Y12))
      =   (ln 10)2 (Y3 – 3Y1Y2 + 2Y13)

and the third derivative of Y1 is

(A8b)   \(\dfrac{d^3\,Y_1}{d(pH)^3}\) =   \((ln\,10)^2 \,\dfrac{d}{d\, pH}\, (Y_3-3Y_1Y_2+2Y_1^3)\)
        =   (ln 10)3 {(Y4–Y1Y3) – 3Y2 (Y2–Y12) – 3Y1 (Y3–Y1Y2) + 6Y12 (Y2–Y12)}

Buffer Intensity & Co. The above equations can now be employed to derive the buffer intensity and its derivative introduced in 6.2 and 7.1. In particular, we have:

(A9a)   n = (Y1 + w/CT)
(A9b)   β =   d /d pH (Y1 + w/CT)
(A9c)   dβ/d pH =   d2/d (pH)2 (Y1 + w/CT)

which yields

(A10a)   n =   (Y1 + w/CT)
(A10b)   β =   (ln 10)  {(Y2 – Y12) + (w+2x) /CT}
(A10c)   dβ/d pH =   (ln 10)2 {(Y3 – 3Y1Y2 + 2Y13) + w/CT}

If desired, the Y’s can be replaced by the a’s. For example, the central piece of A10b is given by:

(A11)   Y2 – Y12   =   (a1 + 4a2) – (a1 + 2a2)2   =   (a2 + a0) – (a2 – a0)2

Here, in the last relation A4a was applied.

Remarks & Footnotes

  1. This topic is also discussed in the PowerPoint presentation. A detailed mathematical description is provided as pdf

  2. The rounding sign on the right-hand side results from the approximation [H+] ≈ {H+}. 

  3. Here and in the following we assume that – except for H+ – all activities (expressed by curly braces) are replaced by molar concentrations (expressed by square brackets). This is legitimate for dilute systems or by switching to conditional equilibrium constants. 

  4. The equilibrium constant K can also be expressed by pK = –log K

  5. There are algebraic formulas for quadratic, cubic and quartic equations (4th order polynomial), but the effort for the last two is enormous. (By the way: 5th order and higher polynomials are algebraically already not solvable any more.) 

  6. The two equilibrium constants of the carbonic acid are: K1 = 10-6.35 and K2 = 10-10.33 (see also here). 

  7. The formula allows also negative values of n. This mimics the withdrawal of the strong base from the solution or the addition of a monoprotic acid (e.g. HCl). 

[last modified: 2019-01-20]