The henderson hasselbalch equation
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The Henderson-Hasselbalch Equation. The Henderson-Hasselbalch Equationthe significance of pHthe predominant solution species The significance of pH in solutions which contain many different acid/base pairs Diprotic Acids

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The Henderson-Hasselbalch Equation

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The Henderson-Hasselbalch Equation

The Henderson-Hasselbalch Equationthe significance of pHthe predominant solution species

The significance of pH in solutions which contain many different acid/base pairs

Diprotic Acids

Triprotic AcidsExample - calculating concentrations of all phosphate containing species in a solution of KH2PO4 at known pH

pH buffer solutions- choosing the conjugate acid/base pair- calculating the pH


[A-]

log10 Ka = log10 [H+] + log10

[HA]

[A-]

- pKa = - pH + log10

[HA]

[A-]

pH = pKa + log10

[HA]

Henderson-Hasselbalch Equation

[H+] [A-]

HA = H+ + A-Ka =

[HA]

An especially convenient form of the equilibrium equation is obtained by re-writing the equilibrium expression using logs -

Henderson-Hasselbalch Equation


[A-]

pH = pKa + log10

[HA]

Use of the Henderson-Hasselbalch Equation

In most practical cases, the pH of the solution is known either from

(1) direct measurement with a pH meter

(2) use of a pH buffer in the solution

When the pH is known, the H.-H. Equation is much more convenient to use than the equilibrium constant expression. It immediately gives the ratio of concentrations of all conjugate acid/base pairs in the solution.

Henderson-Hasselbalch Equation

calculate ratio of (base/acid) concentrations

measured or set by buffer

known (from tabulations)


=

=1

[A-]

[base]

[A-]

[HA]

[acid]

[HA]

pKa and the Dissociation of Weak Acids

For any conjugate acid/base pair,

HA = H+ + A-

pH = pKa + log10

Note that pH = pKa when

because log10 (1) = 0


mostly CH3COOH

Mostly CH3COO-

pKa=4.75

CH3COOH = H+ + CH3COO-pKa = 4.75

0 7 14

When pH < pKa, acetic acid is mostly protonated.

pH

When pH > pKa, acetic acid is mostly deprotonated.

When pH = pKa, concentrations of the protonated and deprotonated forms are equal.

The Use of pH plots


pH, pKa and % Dissociation

When the pH = pKa, half the conjugate acid/base pair is in the protonated form, half is de-protonated.

If the pKa = 4.7 (as for acetic acid):

At pH 4.7[CH3COO-] /[CH3COOH]= 1 (equal)

pH 5.7[CH3COO-] /[CH3COOH]= 10(mostly de-protonated)

pH 3.7[CH3COO-] /[CH3COOH]= 0.1 (mostly protonated)


pH = pKa + log10

[CH3COO-]

[CH3COOH]

An Exercise in % Dissociation

A 0.050 M acetic acid solution is made pH 7.00 with added NaOH.

Find [CH3COOH], [CH3COO-], and [H+] in this solution.

ratio = 180

7.00

4.75

Thus, [CH3COO-] ≈ 0.050 M

[CH3COOH] ≈ 2.8 x 10-4 M

most of the acetic acid is dissociated (% undissociated is 0.56%)


pH = pKa + log10

r =

pH = pKa + log10

(r)

where

[base]

[base]

[acid]

[acid]

How to Use the [base] / [acid] Ratio

[acetic acid] + [acetate] = 0.050 M

[acetate] / [acetic acid] = r

[acetic acid] = r * [acetate]

[acetic acid] (1 + r) = 0.050 M

[acetic acid] = 0.050 M / (1 + r) = 2.8 x 10-4 M

[acetate] = 0.050 M - 2.8 x 10-4 M ≈ 0.050 M


Summary: Significance of the pH

The solution may contain many conjugate acid/base pairs (biological solutions usually do). In order to reproduce a sample, you need to reproduce the pH. This guarantees that all conjugate acid/base pairs will have the same ratio of protonated/deprotonated concentrations as in the original sample.

2.When the pH is known, you can readily calculate the ratio of (protonated/deprotonated) forms of any acid for which you know the pKa.


Ratio of protonated and deprotonated species

Knowledge of the pH completely determines the state of protonation / deprotonation of every Bronsted conjugate acid/base pair in solution.

An example:

A solution at pH 6.9 contains lactic acid (pKa = 3.9). Is lactic acid predominantly in the protonated form or the deprotonated form (lactate ion)?

Ans: Predominantly deprotonated. The ratio

[Lac-]/[HLac] = 103


0 7 14

pH

Diprotic Weak Acid - H2CO3

H2CO3 = H+ + HCO3-pKa1 = 6.4

HCO3-= H+ + CO32-pKa2 = 10.3

pKa1=6.4

pKa2=10.3

mostly H2CO3

mostly CO32-

mostly HCO3-


0 7 14

pH

Triprotic Weak Acid - H3PO4

H3PO4 = H+ + H2PO4-pKa1 = 2.12

H2PO4- = H+ + HPO42-pKa2 = 7.21

HPO42- = H+ + PO43-pKa3 = 12.67

pKa1=2.12

pKa2=7.21

pKa3=12.67

mostly H3PO4

mostly H2PO4-

mostly HPO42-

mostly PO43-


0 7 14

pH

Effect of pH on Solution Composition: H3PO4

pKa1=2.12

pKa2=7.21

pKa3=12.67

mostly H3PO4

mostly H2PO4-

mostly HPO42-

mostly PO43-

Problem:

A solution is prepared by dissolving 0.100 mole of NaH2PO4 in water to produce 1.00 L of solution. The pH is then adjusted to pH 8.50 with NaOH.

What are the concentrations of H3PO4, H2PO4-, HPO42-, PO43-, and H+?


[base]

[acid]

Calculating the Concentrations

pH = pKa + log

(1)

Apply H-H eqn to [H2PO4-] and [HPO42-] using pKa2 :

8.50 = 7.21 + log10 ( [HPO42-] / [H2PO4-] )

This gives [HPO42-] / [H2PO4-] = 20

(2)

Apply H-H eqn to [H3PO4] and [H2PO4-] using pKa1 :

8.50 = 2.12 + log10 ( [H2PO4-] / [H3PO4] )

This gives [H2PO4-] / [H3PO4] = 2.4 x 106

Apply H-H eqn to [PO43-] and [HPO42-] using pKa3 :

(3)

8.50 = 12.67 + log10 ( [PO42-] / [HPO42-] )

This gives [PO43-] / [HPO42-] = 6.8 x 10-5


0 7 14

pH

Answers to the Exercise

pKa1=2.12

pKa2=7.21

pKa3=12.67

mostly H3PO4

mostly H2PO4-

mostly HPO42-

mostly PO43-

solution

Ans:

[HPO42-] = 0.95 x10-1 M

[PO43-] = 1.6 x10-4 * [HPO42-]

[H3PO4] = 2.0 x106 * [H2PO4-]

[H2PO4-] = 0.047 x10-1 M

[H3O+] = 3.2 x10-8 M


(a)

Calculate the pH of a 500 mL solution prepared from:

0.050 mol of acetic acid and 0.020 mol sodium acetate.

HAc = H+ + Ac- pKa= 4.75

?

[Ac-]

pH = pKa + log10

[HAc]

4.75

?

?

Problem


Problem

Suppose that 0.010 mol NaOH is added to the buffer of part (a). What is the pH?

(b)

HAc + OH- = H2O + Ac-pKa = 4.75

(0.020 + 0.010) / 0.50

[Ac-]

pH = pKa + log10

[HAc]

(0.050 - 0.010) / 0.50

4.75

4.63

As long as the buffer capacity is not exceeded, the change of pH is small, in this case, 4.35 to 4.63


pH Buffers

1.pH buffers resist a change in pH upon addition of small amounts of either acid or base.

Buffer solutions should contain roughly equal concentrations of a conjugate acid and its conjugate base.

The conjugate acid/base pair of the buffer should have a pKa that approximately equals the pH.

For example: a buffer will result from mixing 0.1 M acetic acid and 0.1 M sodium acetate.

Added OH- is neutralized by the conjugate acid

Added H+ is neutralized by the conjugate base

H+ + C2H3O2- = HC2H3O2

OH- + HC2H3O2 = H2O + C2H3O2-


pH = pKa + log10

5.00

4.75

[base]

[acid]

pH of Buffer Solutions

Problem:

Prepare a pH 5.00 buffer using sodium acetate and acetic acid

ratio [base]/[acid] = 1.78

Any solution with this composition (i.e., this ratio of base / acid), will form a buffer, but higher concentrations provide higher buffering capacity. For example, one could use

0.178 M sodium acetate + 0.100 M acetic acid


Exercises with Buffers

Use the data in Table 10.2 to design buffers at:

pH 6.9

pH 9.3

pH 3.6

Find the weight of solid compounds you would use to produce 100 mL each buffer.


Choosing a Weak Acid/Base Pair for a Buffer

6.9


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