The Acid Dissociation Constant, Ka, is the equilibrium constant for the reaction in which a weak acid is in equilibrium with its conjugate base in aqueous solution. Notice that in the equilibrium expression below the concentration of water is not included. This is because water is vastly in excess and the amount changes negligibly on equilibrium being established. Ka can be thought of as a modified equilibrium constant. For example, CH3COOH(aq) + H2O(l) = CH3COO-(aq) + H3O+(aq) Ka = [CH3COO-(aq)][H3O+(aq)] / [CH3COOH(aq)] Therefore, the larger the value of Ka, the stronger is the acid. The value is sometimes expressed as the logarithm of its reciprocal, called pKa. Therefore, pKa = -log Ka The smaller the value of pKa the stronger the acid. Ka is a better measure of the strength of an acid than pH because adding more water to the acid solution will not change the value of the equilibrium constant Ka, but it will change the H+ ion concentration on which pH depends. In the above reaction, ethanoic acid and ethanoate ions form a conjugate acid-base pair.http://www.mpcfaculty.net/mark_bishop/weak_acid_equilibrium.htm

Acid Dissociation ConstantsWhen an uncharged weak acid is added to water, a heterogeneous equilibrium forms in which aqueous acid molecules, HA(aq), react with liquid water to form aqueous hydronium ions and aqueous anions, A-(aq). The latter are produced when the acid molecules lose H+ ions to water. HA(aq) + H2O(l) H3O+(aq) + A-(aq)

In writing an equilibrium constant expression for this heterogeneous equilibrium, we leave out the concentration of the liquid water. The equilibrium constant for this expression is called the acid dissociation constant, Ka.

= acid dissociation constant When the equilibrium in question occurs in solution, the chemical formulas enclosed in brackets in the equilibrium constant expression represent the molarities of the substances (moles of solute per liter of solution).

Remember that H+ can be used to represent H3O+, thus simplifying our depiction of the reaction between a weak acid and water and its acid dissociation constant expression: HA(aq) H+(aq) + A-(aq)

= acid dissociation constant For example, acetic acid is a weak acid, because when it is added to water, it reacts with the water in a reversible fashion to form hydronium and acetate ions. HC2H3O2(aq) + H2O(l) or HC2H3O2(aq) H3O+(aq) + C2H3O2-(aq) H+(aq) + C2H3O2-(aq)

= 1.8 x 10-5

EXAMPLE 1 - Writing an Acid Dissociation Constant: Write the equation for the reaction between the weak acid nitrous acid and water, and write the expression for its acid dissociation constant. Solution: HNO2(aq) + H2O(l) H3O+(aq) + NO2-(aq)

or HNO2(aq)

H+(aq) + NO2-(aq)

Exercise 1- Writing an Acid Dissociation Constant: Write the equation for the reaction between the weak acid hydrofluoric acid and water, and write the expression for its acid dissociation constant.

The table below lists acid dissociation constants for some common weak acids. These Ka values can be used to describe the relative strength of the acids. A stronger acid will generate more hydronium ions in solution. A larger Ka indicates a greater ratio of ions (including hydronium ions) to uncharged acid. Therefore, a larger Ka indicates a stronger acid. For example, the larger Ka for chlorous acid (1.2 x 10-2) compared to acetic acid (1.8 x 10-5) tells us that chlorous acid is stronger than acetic acid. Acid Dissociation Constants, Ka, for Common Weak Acids Weak Acid acetic acid benzoic acid chlorous acid formic acid hydrocyanic acid hydrofluoric acid hypobromous acid hypochlorous acid hypoiodous acid lactic acid nitrous acid phenol propionic acid HC2H3O2 C6H5CO2H HClO2 HCHO2 HCN HF HOBr HOCl HOI Equation H+ + C2H3O2H+ + C6H5CO2H+ + ClO2H+ + CHO2H+ + CNH+ + FH+ + OBrH+ + OClH+ + OIKa 1.8 x 10-5 6.4 x 10-5 1.2 x 10-2 1.8 x 10-4 6.2 x 10-10 7.2 x 10-4 2 x 10-9 3.5 x 10-8 2 x 10-11 1.38 x 10-4 4.0 x 10-4 1.6 x 10-10 1.3 X 10-5

CH3CH(OH)CO2H H+ + CH3CH(OH)CO2HNO2 HOC6H5 CH3CH2CO2H H+ + NO2H+ + OC6H5H+ + CH3CH2CO2-

Exercise 2 - Comparing the Strengths of Weak Acids: Which is the stronger acid, formic acid (an irritant found in ant bites) or phenol (formerly used as an antiseptic in hospitals)?

The following study sheet describes one procedure for calculating the pH of solutions of weak acids. If you take other chemistry courses, you will find that there are variations on this procedure for some weak acid solutions.

Study Sheet - Calculating pH for Weak Acid SolutionsTip-off - You are given the concentration of a weak acid solution and asked to calculate its pH. General Steps STEP 1 Write the equation for the ionization of the weak acid in water. HA(aq) H+(aq) + A-(aq)

STEP 2 Write the Ka expression for the weak acid.

STEP 3 Describe each equilibrium concentration in terms of x. x = [H+]equilibrium = [A-]equilibrium [HA]equilibrium = [HA]initial - x STEP 4 Assume that the initial concentration of weak acid is approximately equal to the equilibrium concentration. (Weak acids are rarely ionized to a large degree. We can most often assume that the initial concentration added, [HA]initial is much larger than x. Thus, the equilibrium concentration is approximately equal to the concentration added. You may learn how to deal with weak acid solutions for which this approximation is not appropriate in other chemistry courses.) [HA]equilibrium = [HA]initial STEP #5 Plug the concentrations described in terms of x into the Ka expression, and solve for x.

EXAMPLE 2 - pH Calculations for Weak Acid Solutions: Vinegar is a dilute water solution of acetic acid with small amounts of other components. Calculate the pH of bottled vinegar that is 0.667 M HC2H3O2, assuming that none of the other components affect the acidity of the solution. HC2H3O2(aq) H+(aq) + C2H3O2-(aq)

We get the value for the acid dissociation constant for this reaction from the table above.

x2 = 1.2 x 10-5 x = 3.5 x 10-3 [H+] = 3.5 x 10-3 M H+ pH = -log(3.5 x 10-3) = 2.46

Exercise 3 - pH Calculations for Weak Acid Solutions: Hydrofluoric acid is used to make chlorofluorocarbons (CFCs), to etch glass, and in processing uranium for nuclear power plants. Calculate the pH of a 1.5 M HF solution.http://www.mpcfaculty.net/mark_bishop/weak_acid_equilibrium.htm

(equation 1).

HA(aq )

H +(aq) + AG(aq)

The equilibrium constant for this reaction is given by the law of mass action:

Ka =

[H +] [AG] [HA]

where the square brackets signify that the concentrations are those existing at equilibrium. A convenient form of this equation may be derived if one first takes the comm on logarithm of both sides:

log K a = log[H +] + log([AG]/[HA])

Definition: The p-function is defined such that pX = !log10X, where X may be a variable or a constant. The pH of a solution is thus defined to be equal to !log10[H +]. Substituting into equation 3 now yields

!pK a = !pH + log([A G]/[HA])

Rearrrangement of equation 4 gives the Henderson-Hasselbalch equation (equation 5):

pH

=

pK a + log([A G]/[HA])

Th is equ ation s erve s as the basis for the determ ination of the ionization con stan t, K a. If one can determine the conce ntrations of AG and HA in a solution and measure the pH of the solution, the equation can be solved for the

value of pKa. How ever, any single such determ ination will inevitab ly incur experim ental errors, so we will not solve for pKa in this way.

Inspection of equation 5 reveals that it has the form of the equation of a straight line, where the pH depends upo n the value of log([A -]/[HA]). The theoretical slope of the line is exactly one, and the y-intercept gives the pKa value . W e will prep are s olutions with five diffe rent values of th e ratio [A G]/[HA]. The am oun ts of the co njug ate base and unionized acid are adjusted by partial neutralization of the acid with standard base solution:

HA(aq) + OH G(aq)

A G(aq) + H 2O(l)

The pH values of the five solutions will be measured, and a best-fit line will be obtained using linear regression techniques. By using the regression output to obtain the pK a, the random errors in the individual determinations will avera ge o ut and pro vide a better result.

Read about the use of a pH meter (W ikipedia). https://docs.google.com/viewer?a=v&q=cache:ayZoWbsQiJQJ:web.centre.edu/shiba/che132L/pKa.pdf+i onization+constant+of+weak+acid&hl=tl&gl=ph&pid=bl&srcid=ADGEEShiOfmgFseqhhT2W9jFb-hvl8W_I5cFV85Xc7YuUiqcO21mlH8R0iFXnmVYooKLLxWMzQzEWGAvDcYDnfZ8_EeGnz_QBCAzIPrHV8ZcoW1CZjtogN91whgiyiF0pECQNG3mk&sig=AHIEtbTNSsJOxmkdC63r5942a1KspOgMMQ

In buffer soln, the total solubility of a weak acid and its salt (or a weak base and its salt) is the same. Nevertheless, we have seen that when weak acids, bases, or their salts are dissolved in an unbuffered solution, the pH of the solution changes. When a WA is dissolved in water to make saturated solution, the pH of the solution becomes lower than the pKa of the acid, resulting in a low solubility, otherwise resulting to higher solubility. Introduction to pharmaceutical science by nita p. kandit 2007, Lippincott Williams & wilkins: 530 walnut street philidelphia, PA 19106, 1st edition. Solubility and lipophicity page 34.

1) Weak acids are less than 100% ionized in solution. http://www.chemteam.info/AcidBase/Ka-Intro.html

The ionization constant may be expresses as:

The pH measurement is possible because emf of certain chemical cell varies with the hydrogen ion concentration of the solution on the cell. pH electrode/test solution to be measured// reference electrode symbol//signifies the presence of a liquid junction between test solution and reference electrode If other variables in the cells are controlled, emf of the cell can be correlated with pH Pt, H2 (p) \H+ (a) \ reference electrode. left hand electrode is the hydrogen electrode right hand electrode whose potential is not affected by pH

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pH - MEASUREMENT

ELECTROMETRIC DETERMINATION OF pH

Ecell = EH2, H+ + Eref Ecell - Eref = EH2, H+ Ecell - Eref = RT/nF x 2.303 log10 aH+

R = gas constant (8.314 joule k-1mol-1) T = Absolute temperature (Kelvin) F = A electrochemical constant (96487 coulomb mol-1) N = 1 for univalent species such as H+ Ecell - Eref = 0.0591 pH

pH - MEASUREMENT

ELECTROMETRIC DETERMINATION OF pH

Theoretical slope of a pH electrode is such that change in pH of 1 unit results in a change in the potential (E) of the electrode of 59.1 mV at 25C Mathematically Delta E / Delta pH = 59.1 mV The value of ERef include the following The standard potential of the pH electrode (i.e. the potential when a H+ = 1) The potential of the reference electrode The liquid junction potential The asymmetry potential

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pH - MEASUREMENT

ELECTROMETRIC DETERMINATION OF pH

The value of Eref cannot be determined in routine measurements mainly due to the indeterminate Nature of the liquid junction potential Asymmetry potential Value of asymmetry potential tends to drift as the condition of the electrodes changes The equation not used in a basic way to determine pH of a solution The response of the glass electrode is calibrated using standard aqueous buffer solutions with known reference pH values. Reading of pH meter adjusted so that they correspond to the reference values.

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pH - MEASUREMENT

ELECTROMETRIC DETERMINATION OF pH

The pH of a test solution may then be measured pH = -log a H+ pH measurement reflects the activity rather the concentration of hydrogen ion

pH - MEASUREMENT

BUFFERS:-

The resistance of a solution to change in Hydrogen ion concentration upon the addition of small amounts of acid or alkali is termed as buffer action and a solution which posses this property is called buffer solution. The reference pH values for primary pH standards are established by high accuracy potentiometric measurements Using specially designed electrochemical cells and platinumhydrogen gas electrode rather than a glass electrode, as the H+ - sensing electrode Avoids error in potential measurement arising from liquid junction potential and the asymmetry potential

pH - MEASUREMENT

BUFFERS:-

Enables the pH of the standard solution to be determined from the Nernest equation. Such specialized measurements done at NIST NIST supplies a number of high purity salts as standard reference materials for pH Each of which has a certificate Detailed instructions of preparation and use of solution The pH values certified to 3 decimal places with typical uncertainties of 0.005 pH unit at a range temperature Enables pH meters to be calibrated in a manner that is traceable to a services of internationally recognized standards

pH - MEASUREMENT

High purity salts

Concentration/g/L

pH at stated temperature

Potassium tetraoxolate

Potassium hydrogen phthalate

Potassium di-hydrogen phosphate Di sodium hydrogen phosphate

12.61

10.13

3.39

15C

1.67

4.00

6.90

3.53

Sodium tetra borate decahydrate

Sodium hydrogen carbonate Sodium carbonate

Note: The uncertainty of the tabulated pH values are estimated to be 0.01

3.80

9.28

2.09

2.64

10.12

pH - MEASUREMENT

BUFFERS:-

Buffer solutions to be prepared using salts of the highest purity available Certain of the salts should be dried before use Potassium hydrogen phthalate- Dry at 110C for one hour Potassium di-hydrogen phosphate - Dry at 110C for one hour Disodium hydrogen phosphate - Dry at 110C for one hour Sodium carbonate Dry at 270C for one hour Carbon dioxide-free distilled water to be used to prepare the solutions Important for those buffer solution with a pH > 6

pH - MEASUREMENT

BUFFERS:-

Prepared solutions to be stored in a well stoppered Pyrex or polythene bottles Normally to be replaced after 2 to 4 weeks and sooner if mould or sediment is observed

pH - MEASUREMENT

GLASS ELECTRODE

A Glass bulb B Tube filled with 0.1 mol HCl C Silver-Silver chloride D Saturated KCl solution saturated with AgCl E Silver-Silver chloride electrode

pH - MEASUREMENT

GLASS ELECTRODE

Potential is developed in an aqueous solution is proportional to hydrogen activity or pH of the solution Hydrogen ion in the solution forms a dynamic equilibrium with hydrogen ions that are bound to the membrane surface in ion exchange type process, thereby establishing a potential across the membrane Magnitude of potential is proportional to the pH of the solution

pH, electrometric, glass-electrode

Parameters and Codes:

pH lab, I-1586.85 (units): 00403

pH lab, automated, I-2587-85 (units): 00403

Application

This method may be used to determine the pH of any natural or treated water and any industrial or other wastewater.

Summary of method

2.1 See the introduction to electrometry in this chapter for the principles of pH-meter operation. See also Barnes (1964), Bates (1964), and Willard and others (1965).

2.2 This procedure may be automated with commercially available instrumentation.

Interferences

3.1 The determination is not affected by the presence of color or turbidity, or of organic or colloidal material. Oxidizing and reducing substances do not impair the accuracy of method. 3.2 The pH measurement is temperature dependent, and a significant error results if the temperatures of the buffers and samples differ appreciably. However, a variation of less than 5C has no significant effect except in the most exacting work. 3.3 For samples having abnormally high sodium levels, corrections may be necessary. This correction varies with the type of electrodes used; hence, see the manufacturer s instructions for the necessary computations.