Pharmaceutical Analytical Chemistry-1 Lecture 6 PDF

Pharmaceutical Analytical Chemistry-1 Pharmaceutical Analytical Chemistry-1 PC-101 Dr. Marwa Rifat El-Zahry Lecture 6 Reactions in aqueous solutions & expressions of concentration  Reactions in aqueous solutions solute + solvent...

Pharmaceutical Analytical Chemistry-1 Pharmaceutical Analytical Chemistry-1 PC-101 Dr. Marwa Rifat El-Zahry Lecture 6 Reactions in aqueous solutions & expressions of concentration  Reactions in aqueous solutions solute + solvent solution § The solvent: is the medium in which the substance is dissolved. The solute: is the substance dispersed in the solvent. Solution: is a molecular or ionic dispersion of one substance in some medium. Electrolyte: Substance able to conduct electricity as it ionizes greatly in water. Two types of electrolytes are well known and electrolytes. Non-electrolyte: Substance not ionized in water and no ability to conduct electricity e.g. sugars and glycerol.  Item Strong electrolyte Weak electrolyte Ionization Higher Lower Electric Higher Lower conduction Acids HCl, H2SO4, HNO3, HCN, H3BO3, HClO4 CH3COOH, H3PO4, H2CO3 Bases NaOH, KOH Al(OH)3 q All salts are strong electrolyte such as NaCl, KCl, CH3COONa. q HClO4 is perchloric acid CH3COOH is acetic acid q H3PO4 is phosphoric acid HCN is hydrocyanic acid q H3BO3 is boric acid H2CO3 is carbonic acid  Water solvent ability q Dissociation: ability the molecule to dissociate into smaller ions (cations and anions) and usually reversible. q Water can dissolve salts by separation of cations and anions and formation of new interactions between water and ions. q Water can dissolve many biomolecules because they are hydrophilic (able to attract water). q Hydration shell (hydration sphere): surrounding solute molecules with water molecules. q In summary, water can dissolve many substances (universal solvent) by two mechanism dipole-dipole interaction and hydrogen bond formation. Figure 1 : Dissociation of NaCl in water aided by dipole- dipole interaction.  Concentration q The concentration: the amount of solute present in a given volume of solution. Amount of solute Concentration= Amount of solvent or solution v Units Commonly Used to Express Concentrations: Physical units Chemical units 1- Physical units: Name Units Symbol Weight% Gram solute/100 g solution %, w/w Volume% Milliliter solute/100 ml solution %, v/v Weight to volume % Gram solute/100 ml solution %, w/v Parts per million Gram solute/106 g solution ppm Parts per billion Gram solute/109 g solution ppb 2- Chemical units: Name Units Symbol Molarity Gram molecular weight (Moles) / liter solution M Formality Gram formula weight solute/liter solution F Normality Gram equivalent weight solute/liter solution N Molality No. of moles solute/ 1000 g solution m v Molar standard solution: Gram molecular weight of substance (analyte) in one liter of solution. e.g. the weights required for preparing 1 liter of a molar (1 M) solution of: 1. Sodium hydroxide (NaOH), = 23 + 16 + 1 = 40 g in one liter of solution. 2. Sulphuric acid (H 2 SO 4 ), = (2x1) + 32 + (4x16) = 98 g in one liter of solution. 3. Copper sulphate pentahydrate (CuSO4.5H2O) = 63.6 + 32 + 64 + (5x18) = 249.6 g in one liter of solution. v Normal standard solution (N): Gram equivalent weight of solute (substance) in one liter of solution. Molecular weight of acid or base 1- Equivalent weight of acids or bases= Number of reacted H+ or OH-  Substance No. of H+ or OH- Equivalent weight HCl 1 M.wt/1 NaOH 1 M.wt/1 H2SO4 2 M.wt/2 Ca(OH)2 2 M.wt/2 H3PO4 3 M.wt/3 CH3COOH 1 M.wt/1 Molecular weight 2- Equivalent weight of salts = No. of single atoms × valence Substance No. of single Equivalent atom × valence weight NaCl 1×1 M.wt/1 CaCl2 1×2 M.wt/2 AlCl3 3×1 M.wt/3  Molecular weight 3- Equivalent weight of oxidant and reductant= No. of electrons gained or lost e.g 1. Fe3+ + 1e- Fe2+ Oxidant Reductant Oxidant : any substance gains electrons Equivalent weight of Fe3+= M.wt of Fe3+/1 e.g. 2. Mn7+ + 5e- Mn2+ Reductant : any substance Oxidant Reductant loses electrons Equivalent weight of Mn7+ = M.wt of Mn7+ /5 ü Molarity (M): Number of moles in one liter of solution. Molarity = No. of moles / volume (L) No. of moles= weight (g) / molecular weight So, molarity (M) = weight (g) /[molecular weight ×volume (L)] ü Normailty (N): Number of equivalents in one liter of solution. Normality= No. of equivalents / Volume (L) No. of equivalents= weight (g) / equivalent weight So, normality (N)= weight (g) / [equivalent weight × volume (L)] 1. Calculate the number of moles represented by each of the following: (a) 20 g NaOH (b) 74 g K2CrO4 (c) 148.2 g Ca(OH)2 (d) 50 g CuSO4.5 H2O (e) 57 g Na3PO4. I2 H2O Na= 23, O= 16, H=1 Answer a) Number of moles = weight (g) / molecular weight (M.wt) = 20 / ( 23 + 16 + 1) =20/40 = 0.5 mole 2. How many grams of SnCl2.2H2O are needed to prepare 75 ml of 0.25 M solution? Sn= 118.7, O = 16, H=1, Cl=35.5 M.wt. of SnCl2. 2 H2O = 118.7 + 2 (35.5) + 2 (2+16) =225.7 Molarity (M)= weight (g) / [M.wt x volume (L)] 0.25 = weight (g) / 225.7 x 75/1000 weight= 4.3 g 3. What is the normality of a solution containing 35 g of MnCl2.4 H2O in 300 ml of solution? Mn= 55, Cl=35.5, O= 16, H=1 Molecular weight of MnCl2.4 H2O = 55 + 71 + 72 = = 198 Equivalent weight = 198 / 2= 99 Normality= weight (g) / equivalent weight×volume (L) Normality= 35/99×0.3= 1.18 N. 4. Ferrous salts react with oxidizing agents according to the equation: Fe= 55.8, O= 16, S=32, H=1 Fe2+ → Fe3+ + e How many grams of FeSO4.7H2O are required to prepare 500 ml of a 0.2 N solution of the reducing agent? Answer M.wt of FeSO4.7 H2O = 55.8 + 32 + 64 + 126 = 277.8 Normality= weight (g) / equivalent weight×volume (L) Weight (g)= 0.2×277.8×0.5= 27.78 g. 5. In acid solution the permanganate ion reacts with reducing agents according to the equation :- MnO4- + 8 H+ + 5 e → Mn2+ + 4 H2O Calculate the weight of KMnO4 necessary to prepare 1 liter of a 1 normal solution. K=39, Mn= 55, O= 16 Weight = Normality× equivalent weight× volume (L)= =1× (M.wt of KMnO4/5) ×1 6. What volume in ml of 2N barium chloride solution can be prepared from 8 g of BaCl2.2 H2O ? Ba= 137.3, Cl=35.5, H=1, O=16 Answer To be answered by students  Chemical reactions between ions Ø Combination of ions occurs through the formation of any of the following: (i) Formation of Water: when metallic hydroxide, including sodium or ammonium hydroxide, is mixed with an acid. e.g1. NaOH+ HCl NaCl+H2O e.g2. NH4OH+HCl NH4Cl+ H2O (ii) Formation of Weak Electrolyte: when a solution of a strong acid is mixed with a solution of a salt containing the anion of a weak acid, the weak acid is formed: HCl + CH3COONa NaCl + CH3COOH (weak acid) q When solutions of strong bases are mixed with solutions of ammonium salts, the weak base ammonium hydroxide is formed: NH4Cl + NaOH NH4OH (weak base)+ NaCl (iii) Formation of a Precipitate: If solutions of ferric chloride and sodium hydroxide are mixed, the following ions are present: Fe3+, Cl-, Na+ and OH -. Since ferric hydroxide is insoluble in water, the following reaction occurs: FeCl3+ 3 NaOH Fe(OH)3 + 3 NaCl (iv) Formation of a Gas: The combination of ions may result in the evolution of a gas for two reasons because either the product is gaseous, or the product is unstable and decomposes to form a gas. Examples of the former are: 2H+ + S2- H2S ↑ hydrogen sulphide gas e.g2. H+ + CN- HCN ↑ Hydrogen cyanide Unstable acids formed by the combination of ions are 2 H+ + CO32- H2CO3 H2O + CO2 ↑ carbonic acid 2 H+ + SO32- H2SO3 H2O + SO2 ↑ sulfurous acid 2 H+ + S2O32- H2S2O3 H2O + S  + SO2 ↑ hydrogen thiosulphate 2 H+ + 2NO2- 2 HNO2 H2O + NO ↑ + NO2 ↑ nitrous acid v) Formation of Complex Ions: The following brief rules will aid in predicting the formation of complex ions: - Anions that frequently form complexes are chloride (Cl-), bromide (Br - ), iodide (I - ), fluoride (F - ), cyanide (CN - ), thiocyanate (SCN - ), thiosulphate (S2O32-) and oxalate (C2O42-). - Molecules that frequently form complexes are ammonia (NH 3 ) and water. - The only common cations that do not usually form complexes are sodium, potassium, ammonium. - Complex: Reaction between Lewis acid (electron acceptor) and Lewis base (electron donor). Cu2++ 4NH3 [Cu(NH3)4]2+  Reactions involving oxidation and reduction (Redox-reactions) § Definition. Reactions which occur due to a transfer of electrons from one substance to another are called oxidation-reduction reactions, or simply redox reactions. In the reaction between sodium and chlorine, an electron is transferred from each sodium atom to a chlorine atom: 2 Na + Cl2 2 Na+ + 2 Cl- reductant oxidant Oxidation: loss of electrons. Reduction: Gain of electrons. Oxidant: substance gains electrons. Reductant: substance loses electrons. The common oxidation states of manganese in different species are shown below: Species Mn Mn2+ Mn3+ MnO2 MnO42- MnO4- Oxidation state 0 +2 +3 +4 +6 +7  Chemical reactions and their balancing 1. Types of Chemical Equations. Two types of equations are used to represent chemical reactions: and As an example of the molecular equation: NaOH + HCI NaCl + H2O Ø This equation shows that 1 mole of sodium hydroxide will neutralize exactly 1 mole of hydrogen chloride to form exactly 1 mole of sodium chloride and 1 mole of water; in other words, it represents the exact of the reaction. Ø However, this equation gives no information about the mechanism of the reaction. NaOH, HCl, and NaCl are all strong electrolytes and in solution are 100 % ionized. Therefore, a solution of NaOH contains only Na+ and OH- ions and no NaOH molecules. Similarly, HCl, and NaCl.  Water, on the other hand, is an extremely weak electrolyte and exists almost entirely in the form of H2O molecules. Actually, when a solution of NaOH is mixed with a solution of HCl the expression showing the ionic and molecular species involved would be: Na+ + OH- + H+ + Cl- Na+ + Cl- + H2O It is clear that both Na+ and Cl- do not altered in the reaction. Therefore, the actual change which takes place is given by: H+ + OH- H2O q In the reaction occurring when hydrochloric acid and a solution of silver nitrate are mixed, insoluble silver chloride is formed: Molecular equation: AgNO3 + HCI AgCl  + HNO3 Ionic equation: Ag+ + Cl- AgCl  2. Balancing of Oxidation-Reduction Equations. An important difference in procedure used in balancing molecular and ionic equations should be emphasized: molecular equations are balanced by atoms; and ionic equations are balanced both by atoms and also electrical charge. For example, if we write the for the reduction of the stannic ion by aluminium as follows: Sn4+ + Al Sn2+ + Al3+ The atoms are balanced, but there are 4 positive charges on the left and 5 on the right. Properly balanced the equation becomes: 3 Sn4+ + 2 Al → 3 Sn2+ + 2 Al3+ ü In this equation both atoms and charges are balanced, with 12 positive charges in both sides of the equation. § There are two methods for balancing equations: the oxidation number method for the balancing of molecular equations, and the ion-electron method is best applied to the balancing of ionic equations. (i) The Oxidation Number Method: HNO3 + H2S NO + S + H2O The oxidation number of nitrogen in HNO3 is +5 and in NO it is +2. The oxidation number of sulphur in H2S is -2 and in the free state (S) it is 0. Then: N (+5) N (+2) i.e. loss of 3 units S (-2) S (0) i.e. gain of 2 units Rewriting these expressions so that the total gain and total loss in oxidation numbers are equal (multiplying the first equation by 2, and the second equation by 3, we get the following: 2 N (+5) N (+2) i.e gain 6 units 3 S (-2) S (0) i.e. loss of 6 units Therefore, we find that the N/S ratio is 2:3, and consequently, the reaction equation is written as: 2 HNO3 + 3 H2S 2 NO + 3 S (incomplete) So, balancing hydrogen and oxygen atoms we obtain: 2 HNO3 + 3 H2S 2 NO + 3 S + 4 H2O (complete). (ii) The Ion-Electron Method: Example: K2Cr2O7 + HCl CrCl3 + Cl2 + KCl + H2O For balancing this unbalanced equation, the following steps are followed: (i) K2Cr2O7 is the oxidizing agent, since the oxidation number of chromium decreases from +6 to +3. HCl is the reducing agent, since the oxidation number of chlorine increases from –1 to 0. ii) Since K2Cr2O7, HCl, and CrCl3 are strong electrolytes, the ions involved in the reaction are Cr2O72-, Cl-, and Cr3+. Writing partial skeleton equations for both oxidizing and reducing agent, and using only ions or molecules actually involved in the reaction, we obtain: Cr2O72- Cr3+ (unbalanced) Cl- Cl2 (unbalanced)  iii) Balance chemically each partial equation, using H+ ions in the reaction of the dichromate, since the change takes place in an acid medium: Cr2O72- + 14 H+ 2 Cr3+ + 7 H2O 2 Cl- Cl2 ( iv) Balance the number of charges in each partial equation by adding electrons in the equation for the oxidizing agent and subtracting electrons in the equation for the reducing agent: Cr2O72- + 14 H+ + 6 e 2 Cr3+ + 7 H2O 2 Cl- - 2 e Cl2 (v) Multiply each partial equation by a number so selected that the total number of electrons gained by the oxidizing agent equals the number lost by the reducing agent. 1 Cr2O72- + 14 H+ + 6 e 2 Cr3+ + 7 H2O 3 2Cl- - 2 e Cl2 The results are: Cr2O72- + 14 H+ + 6 e 2 Cr3+ + 7 H2O 6 Cl- - 6 e 3 Cl2 The algebraic sum of the two equation is: Cr2O72- + 6 Cl‑ + 14 H+ 2 Cr3+ + 3 Cl2 + 7 H2O In this form, the ionic equation is balanced. It should be noted that the total ionic charges (-2 - 6 +14) = (2 × +3) are balanced, as well as the atoms. You can convert to molecular equation to calculate the weights of reactants:- K2Cr2O7 + 14 HCl 2 CrCl3 + 3 Cl2 + 2 KCl + 7 H2O Q1 List two examples of combination of ions leads to formation of a gas due to unstable acid formation. Q2 Choose the best answer:- 1-Which of these cations do not commonly form a complex? A- Cobalt B- Copper C- Sodium D- Zinc 2-The O.N. of oxygen in H2O2 is ………… A- zero B- (-1) C- (-2) D- (+2) 3-The equivalent weight of Al2(SO4)3 is…………….. A- M.wt /2 B- M.wt /3 C- M.wt /6 D- M.wt /12 4- This reaction: 2 S2O32-+ 2Fe3+ 2Fe2++ S4O62- is……….. A- unbalanced molecular reaction B- ionic reaction C- balanced molecular reaction D- None of the above Q3 Give reasons: Acetic acid is a weak electrolyte, while hydrochloric acid is a strong one.