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Analytical Chemistry Analytical chemistry Denise Lowinsohn [email protected] http://www.ufjf.br/nupis 2018 SCHEDULE Date Activities 21/03/18 -Aqueous-solution chemistry 28/03/18 -Neutralization tritations 04/04/18 -Precipitation 11/04/18 -Precipitation tritimetry 18/04/18 -Oxidation/reduction 25/04/18 -Oxidation/reduction titration 02/05/18 -Complex-formation titrations: part 1 -Complex-formation titrations: part 2 09/05/18 -Preparing samples for analysis 16/05/18 1º Test (100 points) 23/05/18 -Preparing samples for analysis -Preparing samples for analysis: activity delivery (100 points) 30/05/18 -Application of Statistics to Analytical Chemistry 06/06/18 -Application of Statistics to Analytical Chemistry 13/06/18 Holiday -Application of Statistics to Analytical Chemistry 20/06/18 -Sampling experiment (100 points) 27/06/18 2º Test (100 points) 04/07/18 Oral presentation (100 points) Aqueous-solution chemistry References Brown, LeMay e Bursten, Química - A ciência central, 9ª edição, Editora Pearson – Prentice Hall, 2005. Daniel C. Harris, Análise Química Quantitativa, Editora LTC, 5a edição, 2001. Skoog, West, Holler, Fundamentals of analytical chemistry, 7th edition, 1995. Acid-base equilibrium The behaviour of acids and bases is very important in all areas of Chemistry and others areas of Science. Industrial processes, Laboratory and Biological Effect of pH - The pH of the medium is an extremely important parameter for many reactions in Analytical Chemistry. Acid and base: a brief review Acid: taste sour and cause colour changes in pigments. Base: bitter taste and slippery feeling. Arrhenius: In aqueous medium, acids are defined as substances that increase [H+] and bases increase [OH-] + + Acids = substances that produce H3O (H ) ions, when dissolved in water Bases = substances that produce OH- ions, when dissolved in water Arrhenius: acid + base salt + water. Problem: the definition applies only to aqueous solutions. Brønsted-Lowry Theory Brønsted-Lowry: acid – donates protons and base - accepts protons Transference of “H+” ion between two substances species that accepts derived from protons acid 1 (base 2) (base 1) A1 + B2 ⇌ A2 + B1 (conjugate acids and bases) species that derived from donates base 2 protons (acid 2) (acid 1) Conjugate acid: is the species formed when a base accepts a proton. Conjugate base: is the species formed when an acid loses a proton. The most used concept in Analytical Chemistry. The ion H+ in water • The H+ ion is a proton without electrons. • In water, the H+(aq) formes clusters. + • The H ion interacts with the nonbonding electron pair of the H2O molecules to form the hydrogen ions hydrates: hydronium ion • The most simple cluster is formed by interaction of one proton with one H2O molecule. + + • We can use both: H (aq) or H3O (aq). Brønsted-Lowry Theory - + Acids: can be uncharged molecules (HCl), anions (HSO4 ), cations (NH4 ) - Bases: can be uncharged molecules (NH3), anions (Cl ) Amphoteric (or Amphiprotic) substances: behave as acids or as bases (H2O) Examples: species that species that derived from accepts derived from accepts acid 1 protons acid 1 protons (base 1) (base 2) (base 1) (base 2) - + - + H2O + NH3 ⇌ OH + NH4 HNO2 + H2O ⇌ NO2 + H3O species that derived from species that derived from + donates H base 2 donates base 2 + (acid 1) (acid 2) H (acid 2) (acid 1) Strength of acids and bases • The stronger the acid, the weaker the conjugate base. • H+ is the strongest acid that can exist in equilibrium in aqueous solution. • OH- is the strongest base that can exist in equilibrium in aqueous solution. Strong and weak acids/bases Acids Strong totally dissociated (ex: HCl, HNO3) Weak partially dissociated (ex: H3PO4, CH3COOH) + - HCl(aq) + H2O(l) ⇾ H3O (aq) + Cl (aq) Bases Strong totally dissociated (ex: NaOH) Weak partially dissociated (ex: NH3) + - NH3(aq) + H2O(l) ⇌ NH4 (aq) + OH (aq) Amphiprotic substances Substances that have both acidic and basic properties. They behave as acids or bases depending on the medium. - - Ex.: H2PO4 , HCO3 , H2O Amphiprotic solvents: solvents, depending of the medium, behave as acid or base. Protic solvent: solvents with H+ reactive. All protic solvent suffers autoprotolysis. Aprotic solvent: solvents without H+ reactive. Autoprotolysis or autoionization: involves the spontaneous reaction of molecules of a substance to give a pair of ions. Ion product constant for water Aqueous solutions contain small amount of hydronium and hydroxide ions as a consequence of the dissociation reaction: + - 2H2O(l) ⇌ H3O (aq) + OH (aq) The concentration of At 25oC water in dilute aqueous solutions is enormous [H3O ].[OH ] when compared with 2 Keq the concentration of hydrogen and [H 2O] hydroxide ions. 2 [H3O ].[OH ] Keq .[H 2O] 14 [H3O ].[OH ] Kw 1x10 CONSTANT Ion product constant for water Lewis Theory Acid = accepts a pair of electrons Base = donates a pair of electrons Lewis acids and bases don’t need contain protons. Example: 3+ - 2+ Fe (aq) + SCN (aq) ⇌ Fe(SCN) (aq) Lewis base: Lewis acid: donates a pair of electrons accepts a pair of electrons The defintion of Lewis is the most general definition of acids and bases. + pH scale pH = -log[H3O ] SÖRENSEN introduced in 1909, the In the most Neutral solution: [H O+] = [OH-] concept of pH, a conveniente way of 3 + - -7 -1 solutions, the [H3O ] = [OH ] = 1.0 x 10 mol L expressing acidity – the negative [H+(aq)] is Acid solution: [H O+] > [OH-] very small. logarithm of hydrogen ion concentration. 3 + -7 -1 [H3O ] > 1.0 x 10 mol L and [OH-] < 1.0 x 10-7 mol L-1 Kw [H3O ].[OH ] + - Alkaline solution: [H3O ] < [OH ] 0 [H O+] < 1.0 x 10-7 mol L-1 and pK w pH pOH 14 (25 C) 3 [OH-] > 1.0 x 10-7 mol L-1 •Most of the pH and pOH values are between 0 and 14. •There are no theoretical limits on pH or pOH values. (for example, pH from 2.0 mol L-1 HCl solution is -0.301.) [H+] (mol L-1) pH pOH Example 1x10-0 0 14 Battery acid 1x10-1 1 13 Gastric acid 1x10-2 2 12 Lemon juice 1x10-3 3 11 Orange juice, soda 1x10-4 4 10 Tomato juice, acid rain 1x10-5 5 9 Black coffee, bananas 1x10-6 6 8 Urine, milk 1x10-7 7 7 Pure water 1x10-8 8 6 Sea water, eggs 1x10-9 9 5 Baking soda 1x10-10 10 4 Milk of magnesia 1x10-11 11 3 Ammonia solution 1x10-12 12 2 Soapy water 1x10-13 13 1 Bleach, oven cleaner 1x10-14 14 0 Liquid drain cleaner Practicing..... What are the molar concentration of H+ and the pH in: a) 0.010 mol L-1 KOH? b) 1.8x10-9 mol L-1 NaOH? A sample of lemon juice with [H+] = 3.8x10-4 mol L-1. What is the pH? A solution for cleaning windows is commonly available with [H+] = 5.3x10-9 mol L-1. What is the pH of this solution? A sample of apple juice freshly squeezed has pH = 3.76. What is [H+]? A solution formed by the dissolution of an antacid tablet has pH = 9.18. What is [H+]? Strong acids •The strongest common acids are HCl, HBr, HI, HNO3, HClO3, HClO4, e H2SO4. •Strong acids are strong electrolytes. •All strong acids are totally dissociated in aqueous solutions. No undissociated solute molecules. The equilibrium of the reaction is totally shifted towards the products: + - HNO3(aq) + H2O(l) H3O (aq) + NO3 (aq) Calculation: pH of 0.010 mol L-1 strong acid solution [ ] The concentration expressed in brackets represents the concentration (mol L-1) at equilibrium. C Analytical concentration, represents the real amount of the substance added in certain solvent to form a solution of known concentration “C”. + - HNO3(aq) H (aq) + NO3 (aq) Initial 0.010 mol L-1 - - Equilibrium - 0.010 mol L-1 0.010 mol L-1 -1 CHNO3 = 0.010 mol L total amount of HNO3 present in solution + - -1 Concentration at equilibrium: [H3O ] [NO3 ] = 0.010 mol L not considering autoionization of H2O Calculation: pH of 0.010 mol L-1 strong acid solution [ ] The concentration expressed in brackets represents the concentration (mol L-1) at equilibrium. C Analytical concentration, represents the real amount of the substance added in certain solvent to form a solution of known concentration “C”. + pH = - log[H3O ] + - HNO3(aq) H (aq) + NO+3 (aq) - [H3O ] = [NO3 ] = CHNO3 Initial 0.010 mol L-1 - - pH = -log(C) = -log 0.010 Equilibrium - 0.010 mol L-1 0.010 mol L-1 pH = 2.0 -1 CHNO3 = 0.010 mol L total amount of HNO3 present in solution + - -1 Concentration at equilibrium: [H3O ] [NO3 ] = 0.010 mol L not considering autoionization of H2O Practicing..... -1 What is the pH of 0.040 mol L HClO4 solution? HNO3, pH = 2.34. What is the molar acid concentration? A solution of HNO3 was prepared from 0.85 mL of the concentrated acid in 250 mL of distilled water. What is the pH of this prepared solution? The concentrated acid has 69.5% m/m and density 1.40 g cm-3. (M.W. = 63 g mol-1) What is the [H+] and pH of each solutions? a) 0.0020 mols of HCl in 500 mL of solution -1 b) 0.15 g de HNO3 (M.W. = 63 g mol ) in 300 mL of solution c) 10.0 mL de HCl 15 mol L-1 in 750 mL of solution Strong bases •Most ionic hydroxides are strong bases (for example, NaOH, KOH, e Ca(OH)2). •Strong bases are strong electrolytes and totally dissociated in solution. •To be a base an hydroxide need to be soluble. •The bases don’t need to contain OH- ion: 2- - O (aq) + H2O(l) 2OH (aq) - - H (aq) + H2O(l) H2(g) + OH (aq) 3- - N (aq) + 3H2O(l) NH3(aq) + 3OH (aq) Calculation: pH of 0.010 mol L-1 strong base solution NaOH(aq) Na+(aq) + OH-(aq) Initial 0.010 mol L-1 - - Equilibrium - 0.010 mol L-1 0.010 mol L-1 -1 CNaOH = 0.010 mol L total amount of NaOH present in solution Concentration at equilibrium: [Na+] [OH-] = 0.010 mol L-1 not considering autoionization of H2O pOH = - log[OH-] pKw = pH + pOH + - [Na ] = [OH ] = CNaOH 14.0 = pH + 2.0 pOH = -log(C) = -log 0.010 pH = 12.0 pOH = 2.0 Practicing....
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