《分析化学》专业英语阅读材料03

《分析化学》阅读材料03 摘自Analytical Chemistry (FECS)

●In the Brösted-Lowry theory, an acid is a proton donor and a base is a proton acceptor. Each acid is related to its conjugate base and vice versa:

Acid = base + proton

Therefore, the conjugate base of a strong acid must be a weak base and the conjugate base of a weak acid must be a strong base. Together they form a couple, and an acid without its conjugate base is a meaningless concept. In order to release a proton, the acid must find a base to accept it. In an aqueous solution, the proton, H+, having an extremely small ionic radius, cannot exist as such. It is hydrated, forming the hydronium ion H3O+ and higher hydrates. Thus, an acid-base equilibrium is not a simple dissociation equilibrium, but the result of a proton transfer reaction in which there are at least two reagents and two products. Such a process is also called protolysis. The overall reaction is expressed by: HX + H2O = H3O+ + X-. The overall equilibrium constant is

K = [H3O+][X-]/[HX][H2O].

The acid dissociation constant Ka is given by: Ka = K[H2O] = [H3O+][X-]/[HX]. Ka reflects not only the acid strength of HX, but also the base strength of water. This why different acid dissociation constants are observed for the same acid in different solvent.

●Similar proton transfer reactions exist in all solvents possessing proton donor and acceptor properties. Proton transfer reactions are extremely fast. This makes them very suitable for analytical applications and acid-base reactions have found wide use in volumetric methods and other analytical techniques.

●The pH value is a measure for the acidity or basicity of a solution, aqueous or nonaqueous.

●Acid-base indicators are chemical substances with acid-base properties, having different colors in their protonated and deprotonated forms.

● A most important application of acid-base systems is related to the property of such a system to act as a buffer. Many chemical reactions produce protons (in aqueous solutions hydroniums) or hydroxide ions. If these products remain in the system, a corresponding pH change is observed. However, if a buffer is present in the solution it reacts with the liberated hydrogen or hydroxide ions so that only a relative small change of pH occurs. Buffer consist of a mixture of a weak acid and its conjugate base. The most efficient buffer for a given pH consists of a 1:1 ratio of the protonated and deprotonated forms of a weak acid (with pKa = pH). This cannot always be achieved, but if we wish to prepare a solution of a certain pH, we select a weak acid with a pKa value close to the desired pH. Buffer solution resists changes in pH upon adding of strong acids or strong bases. Depending on the relative concentrations of the acid and base forms of the buffer, the system can resist small or large additions of strong acid or base. This buffer capacity is defined as the number of moles of strong acid of base required to change the pH of 1 L of buffer solution by one pH unit. Solutions with high or low pH values, formed as a result of dissolution of large quantities of a strong base or acid, are characterized by a large buffer capacity, although the electrolyte practically consists of only one of the conjugate forms (e.g., HCl or NaOH solutions).

● A general requirement for all volumetric methods is that the titration process is fast and that it proceeds in a definite stoichiometric ratio, the endpoint of the reaction must by easy to detect and the reaction should be specific and not influenced by other constituents of the solution, i.e., there should be no interference. Question:

1. A H3PO4 solution is brought to pH = 7.00 by the addition of NaOH. Calculate the concentration of the various

forms of orthophosphate if the total phosphate concentration in buffer is 0.200 mol /L. pKa1 = 2.16, pKa2 =

7.21, pKa3 = 12.32.

2. A buffer solution is prepared from acid, HA, Ka = 5 10-5, and its salt. The concentration of HA in the buffer

is 0.25 mol / L. To 100 mL of the buffer added 5.0 mmol of NaOH, and the pH of the resulting solution is

5.60. What was the pH of the original buffer?

3.It is desired to change the pH of 100 mL of 0.100 mol / L HCl from 1.00 to

4.40 by the addition of sodium

acetate, CH3COONa3H2O. How much solid sodium acetate salt must be added in grams? Assume no volume change for the solution as the result of the addition.

4.What is the buffer capacity of a solution which is 0.100 mol / L of NH3 and 0.200 of NH4Cl?

5.Calculate the pH of each of the following solution: (a) Water in equilibrium with CO2 of the air; pKa1 =

6.38,

pKa2 = 10.25. (b) Water as in part (a) brought to pH = 7.00 with NaOH and allowed to regain equilibrium with CO2.

6. Derive the following expression for the pH at the first stoichiometric point in the titration of a mixture of two

weak acid: HA, the stronger, Ka1, concentration C1; HB, the weaker, Ka2, concentration C2:

pH = 1/2 (pKa1 + pKa2) – 1/2 lg (C1/C2)