ACID-BASE TITRATIONS

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1 ACID-BASE TITRATIONS 1 A. TITRATION The progressive addition of one reagent (chemical reactant) to another. An acid-base titration can involve the addition of an acid to a base or vice-versa. Usually a solution of known concentration (the titrant) is placed in a buret and is progressively added to a measured volume of another solution of unknown concentration (the sample). The titrant is added until the reaction mixture reaches the equivalence point equivalence point point where chemically equivalent amounts of acid and base have reacted (& neither is in excess). The acid-base reaction is complete at this point. equivalent = the amount of acid or base that gives 1 mol of H + or OH - ions, respectively Ideally, this is the same point at which an indicator changes colour = endpoint = ph (the endpoint is the experimental estimate of the equivalence point) The volume of titrant added to achieve the equivalence point can be used to calculate the concentration of the sample via stoichiometry calculations Primary Standards HCl and NaOH are the most commonly used titrants in acid-base titrations. However, neither is a primary standard : a solution that can be prepared from a pure form of the solute to a precise molar concentration Solutions of HCl and NaOH have to be standardized. A standardized solution is one whose concentration is determined from data collected by titration with a primary standard solution HCl has to be standardized because it is a gas at room temperature. It is difficult to dissolve precise amounts of any gas in water. Moreover, once dissolved, tiny amounts of HCl have the tendency to escape from solution thereby reducing the concentration of dissolved HCl. NaOH is a deliquescent or hygroscopic solid - it absorbs moisture from the air. When you weigh out a sample of NaOH, you are obtaining the mass of NaOH plus the mass of the water it has absorbed from the air. If you prepare a solution using the measured mass of NaOH, its molar concentration will be slightly lower than expected because of the absorbed moisture. Acids like HCl are standardized using a primary standard such as sodium carbonate (Na 2 CO 3 ), while bases like NaOH are standardized using potassium hydrogen sulfate (KHSO 4 ) or potassium hydrogen phthalate (KHC 8 H 4 O 4 ).

2 2 B. ACID-BASE STOICHIOMETRY Can be used to predict the concentration of an unknown in a titration (most common application) or to find the volume of acid or base needed for a titration. Steps for solving problems: 1. Write a balanced chemical equation (double replacement & nonionic), then identify and record the required and given substances. 2. If not given moles, convert the given quantity (mass, or solution volume and concentration) to number of moles using one of the formulas below: n = m/m n = Cv 3. Use the mole ratio to calculate the moles of the required substance from the moles of the given substance. (Every problem will involve this calculation). n R = n G x R/G where n R = moles of required substance n G = moles of given substance R/G = mole ratio using the coefficients from the balanced equation ie. moles of required/ moles of given 4. If necessary, convert the moles of the required substance to the required quantity (mass, solution volume or concentration) using one of the formulas m = nm v = n/c C = n/v Example 1: (to find unknown concentration) Determine the concentration of NaOH (aq) if 25 ml of NaOH is neutralized by 50 ml of 2.0 mol/l HCl (aq). Example 2: (to find unknown volume) Calculate the volume of 0.50 mol/l H 2 SO 4(aq) required to completely neutralize 40 ml of 0.90 mol/l LiOH (aq).

3 3 Exercises: 1. Predict the volume of titrant (0.10 mol/l NaOH) needed for complete neutralization of 50 ml of 0.10 mol/l HCl. 2. Analysis shows that 9.44 ml of mol/l potassium hydroxide solution are needed for the titration of 10.0 ml of acid from an acid lake. Determine the molar concentration of acid in the lake water, assuming the acid is sulfuric acid. 3. Find the volume of a mol/l hydrochloric acid solution needed to completely neutralize 45.0 ml of a mol/l solution of sodium carbonate.

4 C. TITRATION CURVE 4 A graph of ph (y axis) vs volume of reagent (titrant) progressively added to the sample. Useful for: 1. Determining the equivalence point = midpoint of sharp change in ph (steep vertical portion) 2. Showing the stepwise progress of a reaction involving a polyprotic acid (PP) or polybasic (PB) species 3. Determining the volume of titrant required to achieve the equivalence point 4. Selecting a proper indicator = one that changes color over a ph range that includes the endpoint ph value (within the steep vertical portion of the curve) 5. Identifying type of reaction: SA + SB, SA + WB, SB + WA, PP + B, or PB + A D. INDICATORS Indicators are dyes that exhibit different colours in acidic and basic solutions. All acid-base indicators are either weak acids or weak bases. In aqueous solution, the acid form of the indicator is in equilibrium with its conjugate base. These two forms will exhibit different colours. HIn (aq) + H 2 O (l) H 3 O + (aq) + In - (aq) HIn = acid form of the indicator In - = base form of the indicator Exercises: * In a strong acid solution, the higher [H 3 O + ] will cause the equilibrium to shift to the left, and the colour of the solution will be that of the acid form, HIn. * In a strong basic solution, the extremely low [H 3 O + ] will cause the equilibrium to shift to the right and In - will determine the colour of the solution. 1. For each of the following indicators, write the equation for its addition to water and indicate the color of solution if added to (i) a strong base or (ii) a strong acid a. Methyl red, HMr b. Bromocresol green, HBg 2. According to the table of acid-base indicators what is the colour of each of the following indicators in the solutions of given ph? a. Phenolphthalein in a solution with a ph of b. Bromothymol blue in a solution with a ph of 11.7 c. Litmus in a solution with a ph of 8.2 d. Methyl orange in a solution with a ph of 3.9

5 3. Complete the analysis for each of the following diagnostic tests. If [the specified indicator] is added to a solution, and the colour of the solution turns [the given colour], then the solution ph is _?_. a. methyl red (red) b. alizarin yellow (red) c. bromocresol green (blue) d. bromothymol blue (green) 4. Separate samples of an unknown solution turned both methyl orange and bromothymol blue to yellow, and turned bromocresol green to blue. a. Estimate the ph of the unknown solution. b. Calculate the approximate hydronium ion concentration. 5. Separate samples of an unknown solution were tested with indicators. Congo red was red and chlorophenol red was yellow in the solution. Note: For Congo red, ph range = 3.0 to 5.0 (blue to red) a. Estimate the ph of the unknown solution. b. Calculate the approximate hydronium ion concentration. E. IDENTIFYING TYPE OF TITRATION 1. Vertical Portion Each steep vertical portion represents a quantitative acid-base reaction A titration involving a weak acid or base typically has a relatively shorter vertical portion than that of a strong acid and strong base A titration of a weak acid and a weak base is not generally performed as these reactions are difficult to describe quantitatively ph at beginning or end of titration: ( = less than; = greater than) Acid: strong (ph 2) ; weak (ph 2) Base: strong (ph 12) ; weak (ph 12) 2. Endpoints of titrations: Strong Acid + Strong Base ph = 7 Strong Acid + Weak Base ph 7 Weak Acid = Strong Base ph 7 EXERCISES: 1. The diagram on page 7 shows titration curves for 4 different titrations in which the acid is the sample and the base is the titrant. a) Identify each: Strong acid + strong base: a b c d Strong acid + weak base: a b c d Weak acid + strong base: a b c d Weak acid + weak base: a b c d b) Sketch 4 curves on the graph paper provided for 4 different titrations in which the base is the sample and the acid is the titrant. 5

6 2. For each of the titration curves shown on page 7, determine the volume of titrant needed to achieve the equivalence point, the ph at the equivalence point and choose an appropriate indicator. 6 Reactants Volume (ml) ph Indicator a) NaOH + HCl b) NH 3 + HCl c) HCl + NaOH d) CH 3 COOH + NaOH 3. For each of the titrations in question 2 verify the volume of titrant required to reach the equivalence point using stoichiometric calculations. (Use the data given at the top of each titration curve). (use graphs distributed in class)

7 F. ACID/BASE EXCESS PROBLEMS 7 Excess problems can be recognized when the concentration and volume of both substances are given. Summary of Steps: 1. Write a balanced, non-ionic, double replacement reaction. All reactions will be in a 1:1 mole ratio. 2. Find the number of moles of each substance given. n = Cv. 3. Identify the excess reagent as the one with the greater number of moles. 4. Determine how much (moles) is in excess: n xs = n G - n R Where n G is the moles given; n R is the moles required to completely react. 5. Find the concentration of the excess substance from the excess amount and the total volume of the mixture. C xs = n/v. 6. Convert the concentration to the required quantity. (Concentration of hydroxide or hydronium ions, ph, poh). Problems: 1. Predict the ph of a solution produced by mixing 100 ml of 0.25 mol/l nitric acid with 50 ml of 0.40 mol/l potassium hydroxide solution during an acidbase titration. 2. Calculate the hydronium and hydroxide ion concentrations of a solution after 30 ml of 0.10 mol/l HCl (aq) is used to neutralize 25 ml of 0.10 mol/l NaOH (aq). 3. Calculate the hydronium and hydroxide ion concentrations of a solution after 100 ml of 0.20 mol/l NaOH (aq) is added to 40 ml of 0.20 mol/l HClO 4(aq). Challenge: 4. A 10 ml sample of a hydrochloric acid solution of unknown concentration was titrated with a standardized mol/l sodium hydroxide solution. The endpoint of 11.5 ml was overshot by 1.5 ml of NaOH (aq). According to the evidence, what is the ph of the final solution?

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