Analytical Chemistry I Lecture 3 PDF

Analytical Chemistry I PMC 101 Lecture 3 Dr. Dina A. Gawad Assistant Professor of Pharm. Analytical Chemistry  Chemical Composition of Aqueous Solutions Water is the universal solvent Most of the solutes we will discuss are electrolytes Electrolytes Solutes wh...

Analytical Chemistry I PMC 101 Lecture 3 Dr. Dina A. Gawad Assistant Professor of Pharm. Analytical Chemistry  Chemical Composition of Aqueous Solutions Water is the universal solvent Most of the solutes we will discuss are electrolytes Electrolytes Solutes which form ions when dissolved in water, and thus produce solutions that conduct electricity The amount of ionization is dependent on the strength of the electrolyte. Classification of electrolytes 1-Strong Electrolyte: ionize completely in a solvent Most inorganic acids e.g. HCl, HClO4 , HNO3 Alkali and alkaline earth metal hydroxides Most salts 2-Weak Electrolyte: weak electrolytes ionize partially. Some inorganic acids e.g. H3B03 , H3PO4 Most organic acids NH3& most organic bases Some salts e.g. HgCl2 Different definitions of Acids and Bases Three Definitions of Acids and Bases Arrhenius Definition Brønsted-Lowry Definition Lewis Definition  1-Arrhenius theory: An acid ionizes in H2O to give hydronium ion. HA + H2O ↔ H3O+ + A- A base ionizes in H2O to give hydroxyl ion. B + H2O ↔ BH+ + OH- This theory did not discuss the role of solvent in the ionization process. 2-Brönsted-Lowry theory An acid is a proton donor A base is a proton acceptor. In order for a species to behave as an acid, a proton acceptor (or base) must be present. The reverse is also true.  Conjugate acids and bases: When an acid gives up a proton, a conjugate base is formed that is capable of accepting a proton as: Acid1 ↔ Base1 + Proton HCl ↔ Cl- + H+ Here, Acid1 and Base1 represent a conjugate acid- base pair. Similarly, every base produces its conjugate acid as: Base2 + Proton ↔ Acid2 NH3 + H+ ↔ NH4+ When these two half-reactions are combined, the result is an acid-base or neutralization reaction Acid1 + Base2 ↔ Base1+Acid2 HCl + NH3 ↔ Cl- + NH4+ In this theory, solvents molecules are involved either as an acid or as a base. Thus, when ammonia (a base) is dissolved in water an acid-base reaction takes place in which the solvent (H2O) acts as an acid (proton donor). H2O + NH3 ↔ OH- + NH4+ Amphoteric Solvents Species that can act as either an acid or a base Example: Water( H2O) HCl + H2O → Cl-+ H3O+ When HCl dissolves in water, Water acts as a base, accepting H+ 3-Lewis theory: An acid is a substance that can accept an electron pair, and a base is a substance that can donate an electron pair. H+ + :NH3 → H : NH3+ AlCl3 + :OR2 → Cl3Al : OR2 H+ + :OH2 → H : OH2+ H+ + :OH- → H : OH  The chemical equilibrium -The reactions used in analytical chemistry are seldom complete (reversible reaction). -Instead, they proceed to a state of chemical equilibrium in which the ratio of concentration of reactants and products is constant. - As the reaction proceeds, the [reactants] , and so does the rate at which they react. - The [products] and so does the rate at which they react. - A state of equilibrium is reached, wherein the rate of forward reaction between A and B is balanced by the rate of backward reaction between C and D. A + B ↔ C + D - At equilibrium, the rate of formation of each substance equals the rate of consumption and there is no net change in concentration. The law of mass action the rate of a chemical reaction is proportional to the product of the concentrations of the reacting substances, with each concentration raised to a power equal to the number of molecules or ions of the reactants in a balanced chemical equation aA + bB ↔ cC + dD The equilibrium constant is expressed as follows: ❖ The square-bracketed terms have the following meanings: -Molar concentration if the species is a dissolved solute. -Partial pressure in atmospheres if the species is a gas. -Unity if the species is (a) a pure liquid, (b) a pure solid or (c) the solvent in a dilute solution.  What Does the Value of K Mean? H2+ Br2 ↔ 2HBr K = [HBr]2 [H2][Br2] If K >> 1, the reaction is product-favored; product predominates at equilibrium If K 10-7M, pH < 7 BASIC [H+ ] < 10-7M, pH > 7 pH ranges: 0 (very acidic) to 14 (very basic). Change in just one unit of scale = tenfold change in H+ concentration. (Because the pH scale is a log scale based on 10) Methods for measuring the pH of an aqueous solution pH paper pH meter  The pH of a solution pH, pOH, and pKw The ion-product constant for water: Kw = [H3O+ ][OH- ] -log Kw = -log H+ −log OH- = -log 10 -14 pKw = pH + pOH = 14 ❖ Calculate pH of an ammonia solution whose hydroxide ion concentration is 1.9 x 10 -3 M. pOH = – log (1.9 x 10 -3 ) = 2.72 pH = 14.00 – pOH = 14.00 – 2.72 = 11.28