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Chemical Examples for the Fit Equations
The following are just a few of the applications for nonlinear curve fitting. A nice introduction to biological applications of nonlinear curve fitting by Dr. Harvey Motulsky is available on the Web^1.
a exp(-bx) f = a e-bx
Example: Chemical kinetics first order decay of a reactant.
A → P gives [A] = [A]o e-kt
a(1-exp(-bx)) f = a (1 - e-bx )
0
1
(^0 2 4) x 6 8 10
a(1-exp(-bx))
Example: Chemical kinetics first order increase of a product.
A → P gives [P] = [P]o(1 - e-kt )
a(1-exp(-b(x-c))) f = a (1 - e-b(x-c) )
Example: This form is the same function as above with an x axis offset. For first order kinetics a delay in the time of the start of the reaction of to would give the rate law
[P] = [P]o(1 - e-k(t-to) ) where to = c in the fitting function.
a(1-exp(-bx)) + c f = a (1 - e-bx ) + c
Example: This form introduces a constant offset for first order type growth. For a first order chemical reaction
[P] = [P]o(1 - e-kt ) + c
where the offset, c, is due to miscalibration or an instrumental artifact.
This form can be used interchangeably with f = A e-bx + C as given in the next function with
A = -a and C = a+c.
a exp(-bx) + c f = a e-bx + c
0
1
(^0 2 4) x 6 8 10
a exp(-bx) + c
Example: Chemical kinetics first order decay of a reactant towards equilibrium. A →← P with k 1 the forward rate constant and k-1 the reverse rate constant The linear form for the concentration of the reactant in a reversible first order reaction is ln([A]-[A]eq) = – (k 1 + k-1) t + ln([A]o-[A]eq) which rearranges to give:
[A] = ([A]o-[A]eq) e– (k^1 + k-1)t + [A]eq
Comparing to "a exp(-bx) + c" gives a = ([A]o-[A]eq) and b = k 1 + k-1 and c = [A]eq.
a exp(b/x) + c f = a e-b/x + c
0
20
40
60
80
100
(^200 400) x 600 800
a exp(-b/x)
Example 1: Gibb's Free energy relationships.
Starting from the equation ∆rGo- = -RT ln K then K = e-∆rGo-/RT
or K = e-∆rSo-/R^ e-∆rHo-/RT
Comparing to "a exp(b/x) + c" gives a = e-∆rSo-/R and b = -∆rHo-/R and c should be fixed at 0.
Example 2: Arrhenius temperature dependence in chemical kinetics.
The temperature dependence of a second order rate constant is often given by k 2 = A e-Ea/RT.
ax^2 + bx + c f = a x^2 + b x + c
Example: Power series are a generally useful functional form, especially when you don't know which function to use. For example the enthalpy of a chemical reaction can be expanded in a
power series in the temperature: ∆rHT = ∆rHo + ∆aT +
∆b 2 T
1/(ax+b) f =
a x + b
0
0 2 4 x 6 8 10
y
Example: Chemical kinetics second order decay of a reactant. 2A → P with k 2 the rate constant The linear form for the concentration of the reactant in a second order reaction is 1 [A] -^
[A]o = k^2 t Solving for [A] gives
[A] =
k 2 t + 1/[A]o Comparing with "1/(ax+b)" gives a = k 2 and b = 1/[A]o when x = time. The next function is more direct and specific for second order kinetics, however.
1/(1/a+bx) f =
1/a + b x
0
0 2 4 x 6 8 10
y
Example: Chemical kinetics second order decay of a reactant. 2A → P with k 2 the rate constant The linear form for the concentration of the reactant in a second order reaction is 1 [A] -^
[A]o = k^2 t Solving for [A] gives
[A] =
1/[A]o + k 2 t Comparing with "1/(1/a+bx)" gives a = [A]o and b = k 2 when x = time.
1/(1/a+bx) + c f =
1/a + b x + c
Example: This form introduces a constant offset for second order type kinetics. For a second order chemical reaction
[A] =
1/[A]o + k 2 t + c where the offset, c, is due to miscalibration or an instrumental artifact.
ax/(b+x) f =
a x b + x
0
0 2 4 x 6 8 10
y
Example 1: This general form occurs often. Examples include Michaelis - Menten Enzyme kinetics:
Rate =
d[P] dt =^
Vmax [S] (kM + [S]) with kM^ =
( k 1 + k-1 ) k 2 and Vmax^ = k^1 [E]o
It is often better to fit to the nonlinear form above than the linearized forms such as the Lineweaver-Burk form.
Example 2: Scatchard experiments for 1:1 binding also have the same form. A typical Scatchard analysis gives
K =
ν (1 - ν)[A]free with ν = [A]bound / [M],
where [M] is the concentration of binding sites, [A]bound is the concentration of ligand bound, and [M] is the maximum concentration of binding sites. ([A]bound can also be written [AM] for A + M→← AM.) The fraction of bound sites is ν. Rearranging to solve for ν gives: ν =
K [A]free 1+K[A]free and dividing by K gives^ ν^ =^
[A]free 1/K+[A]free Scatchard type binding experiments are sometimes called saturation binding^1 or specific binding experiments. In this context Y = specific binding and
Y =
Bmax X Kd +X where X is the ligand concentration, Bmax is the maximum binding capacity and Kd is the equilibrium constant.
Example 3: The Langmuir isotherm from surface chemistry is given by
a/(1+10^(b-x)) +c f =
a 1 + 10(b-x)^ + c
0
2
4
6
8
10
(^0 2 4) x 6 8 10
y
Example: Binding experiments in biological or pharmaceutical analysis. This type of curve is called a "sigmoid dose response curve," where the response increases with the dose. The form given above ( see "a/(1+10^(x-b))+c") with (x-b) gives a decreasing response with dose. In the drug binding context:
Y = [A]min +
[A]max-[A]min 1 + 10(logEC50-x) Where Y = [A]bound , or the function value can be a pharmacological result. In dose response experiments x = log[A], where [A] is the free ligand concentration. Again comparing to the form "a/(1+10^(b-x)) + c", a = [A]max-[A]min , b = logEC50, and c = [A]min. The EC50 value is the effective concentration for 50% response. The EC50 value is sometimes called ED50, for effective dose, or IC50, for inhibitory concentration. This form assumes a standard slope, giving a 10% to 90% response for a change in drug concentration of about two orders of magnitude (assuming x = log[A]).
(a 10^(x-b) + c)/(1+10^(x-b))
f =
a 10(x-b)^ + c 1+10(x-b)
0
2
4
6
8
10
(^0 2 4) x 6 8 10
y
This form and the previous form are very similar. The difference is how the extremes are expressed. In this form the a parameter is the maximum value, and not max-min. The b and c parameters remain the same. When x = b, f = (a+c)/2, the average value.
Example: This form works well with binding experiments, as discussed above. This form is also particularly useful in other types of chemical equilibrium experiments. One specific example is the use of pH indicator species in NMR. Assume that a particular species can exist in two forms, HI →← H+^ + I-^ with a given pKa. The limiting chemical shifts of the two forms are δHI and δI-. The observed chemical shift is given by the mole fraction weighted average:
δobs =
[HI]
[HI]+[I-] δHI^ +^
[I-]
[HI]+[I-] δI- Defining the mole fraction of I-^ as
X =
[I-]
[HI]+[I-] =
[I-]
[HI]t and noting that XHI = 1 – X gives: δobs = (1-X) δHI + (X) δI-
The pH of the solution is given by the Henderson-Hasselbach equation
pH = pKa + log
[A
- ]
[HI]
The concentration ratio is the same as the mole fraction ratio:
[A
- ]
[HI] =^
X
1-X
Solving for this ratio from δobs = (1-X) δHI + (X) δI- gives
X =
δobs - δHI δI- - δHI^ 1-X =^
δI- - δobs δI- - δHI^ and
[A
- ]
[HI] =^
X
1-X =^
δobs - δHI δI- - δobs Substituting into the Henderson Hasselbach equation then gives
pH = pKa + log
δobs - δHI δI- - δobs This functional form is not useful for curve fitting. The plot of the chemical shift versus the pH is more amenable to curve fitting. So solving this equation for δobs as a function of pH gives:
δobs =
δI- 10(pH- pKa)^ + δHI 1+10(pH- pKa) As the pH increases, the shift changes from δHI to δI- in a sigmoidal fashion. When pH = pKa, [HI] = [I-], and the shift is the normal average of the two forms. Comparing to the curve fit functional form, a = δI- = δmax , c = δHI = δmin , and b = pKa. Once the curve fit parameters are determined, they can be substitutued into the Henderson-Hasselbach equation to calculate the pH of an unknown. This type of relationship should hold whenever the observed response is mole fraction weighted.
a/(1+10^(b(c-x))) f =
a 1 + 10b(c-x) b= 0.2, 0.5, 1.0, 1.
-0.
0
1
0 2 4 x 6 8 10
y
Example: Sigmoid dose response curve with a variable slope.
then the free guest concentration is given by^3 :
[G] =
-(1+K([H]t-[G]t)) + (1+K([H]t-[G]t))^2 + 4K[G]t 2K
The mole ratio of the added guest to host is n =
[G]t [H]t. Multiplying and dividing the equilibrium constant by [H]t gives:
[G] =
-(1+K[H]t (1-n)) + (1+K[H]t (1-n))^2 + 4K[H]t n 2K[H]t/[H]t Or
[G] = [H]t (^)
-(1+K[H]t (1-n)) + (1+K[H]t (1-n))^2 + 4K[H]t n 2K[H]t The mole balance equations can be used to calculate [HG] and [H], or the equilibrium expressions can be solved directly giving:
[H] = [H]t (^)
-(1+K[H]t (n-1)) + (1+K[H]t (n-1))^2 + 4K[H]t 2K[H]t
[HG] = [H]t (^)
(1+K[H]t (1+n)) - (1+K[H]t (1+n))^2 - 4K^2 [H]t^2 n 2K[H]t The equations are in the form:
[G] = a
-(1+b(1-x)) + (1+b(1-x))^2 + 4bx 2b
[H] = a
-(1+b(x-1)) + (1+b(x-1))^2 + 4b 2b
[HG] = a
(1+b(1+x)) - (1+b(1+x))^2 - 4b^2 x 2b
with a = [H]t and b = K[H]t with x = n. The fractional occupation ν=[H]/[H]t can also be calculated from the equation for [H]:
ν =
-(1+b(x-1)) + (1+b(x-1))^2 + 4b 2b
In actual practice, an experimental response proportional to the concentration may be determined ( e.g. the absorbance of the solution), R = k [G] and then
R =k[G] = a
-(1+b(1-x)) + (1+b(1-x))^2 + 4bx 2b Where a = k[H]t. A curve fit of the experimental response verses the mole ratio then gives the equilibrium constant from b = K[H]t. The designation of the guest and host in these equations is arbitrary, so the G and H labels may be exchanged. Another specific example is the determination of equilibrium constants in NMR in the extreme narrowing limit. For host-guest association, consider G + H →← HG with a given K. The limiting chemical shifts of the three forms are δH, δG, and δHG. Assume that the total host concentration is [H]t and the guest is added in varying amounts, as above. The observed chemical shift of the host is given by the mole fraction weighted average:
δobs =
[H]
[H]+[HG] δH^ +^
[HG]
[H]+[HG] δHG
Defining the mole fraction of host and host guest complex as
XH =
[H]
[H]+[HG] =
[H]
[H]t XHG^ =^
[HG]
[H]+[HG] =
[HG]
[H]t and noting that XH = 1 – XHG gives: δobs = (1-XHG) δH + (XHG) δHG The observed chemical shift can also be directly related to the host guest complex concentration:
δobs = δH + (δHG - δH)
[HG]
[H]t The [HG]/[H]t ratio is given by the equation above. The final fitting function is then
δobs = a
(1+b(1+x)) - (1+b(1+x))^2 - 4b^2 x 2b + c with a = (δHG - δH), b = K[H]t.as before, and c = δH. The same equations can easily be used for the other types of chemical equilibria mentioned above^3. For example for acid-base equilibria the [G] is replaced by [H+]. As presented here, these equations cannot be used for fitting titration curves, because the total volume increases in a standard titration and [H]t or its equivalent will not be constant. However, if the dilution factor during the titration is small or the experiment is done at constant total volume, then these equations will be useful.
Nonlinear least squares curve fitting programs available on the Web:
The "Nonlinear Least Squares Curve Fitter" written by John C. Pezzullo will be very useful if the equation form that you want to use is not available in standard packages. In this applet, you can input any functional form: http://statpages.org/nonlin.html
Another example for several simple equation forms is: http://people.hofstra.edu/faculty/Stefan_Waner/newgraph/regressionframes.html
Literature Cited:
- http://curvefit.com/analyzing_data_book.htm
- K. A. Connors, Binding Constants, The Measurement of Molecular Complex Stability, John Wiley & Sons, New York, NY, 1987; pp. 50-65.
- D. J. Eatough, J. J. Christensen, R. M. Izatt, Experiments in Thermometric Titrimetry and Titration Calorimetry, Brigham Young University Press, Provo, Utah, 1974; p. 25.
