Acids Lesson 16 Titrations

Calculations to find concentration of sample solution. Objectives: Calculations to find concentration of sample solution. Calculations to find molar ratio of acid and base in neutralization equation. Practical Aspects of Titration Primary Standards and secondary standards Indicators Titration curves

Titration

Titration Purpose: To find concentration of known ion Method: Acid Base neutralization

Titration Purpose: To find concentration of know ion Method: acid base neutralization Material: Burette, Erlenmeyer flask, indicator, measuring cylinder to measure volume, known Concentration base solution, acid solution of unknown concentration

Titration Procedure: Wash burette and flask with distilled water Remove water completely from burette and flask Measure 15 mL of acid solution using measuring cylinder and transfer to flask Close tap of burette and fill it to zero

Titration Procedure: Put 5 drops of indicator solution in flask Hang the burette into the ring stand. Open the tap of the burette and carefully allow Base solution to drop into the flask drop-wise. Keep stirring the flask to mix acid and base solution.

Titration Procedure: When colour change is observed then stop and measure volume of base solution that was used to titrate the acid solution

Titration Calculation Find moles of the base solution by volume and concentration of standard solution Find moles of sample solution Find concentration of the sample solution by moles and volume of the sample solution

Titration Terms Standard solution Solution in burette, known concentration Sample solution Solution in burette, known volume, unknown concentration

Titration Terms Indicator A solution that change colour when reaction is complete. This is added into the flask.

Titration Terms Transition point, end point, stoichiometric point A point in time during titration process when moles of acid become equal to moles of base

Setup for titrating an acid with a base

Titration Calculation 6.50 mL of 0.100 M H2C2O4 is required to neutralize 10.0 mL of NaOH solution in a titration. Calculate the base concentration. 1H2C2O4 + 2NaOH  Na2C2O4 + 2H2O 0.00650 L 0.0100 L 0.100 M ? M 0.00650 L H2C2O4 x 0.100 mole 1 L x 2 mole NaOH 1 mole H2C2O4 [NaOH] = 0.0100 L = 0.130 M

Calculate the volume of 0. 100 M H2CO3 required to. neutralize 25 Calculate the volume of 0.100 M H2CO3 required to neutralize 25.0 mL of 0.200 M KOH. 1H2CO3 + 2KOH → K2CO3 + 2HOH ? L 0.0250 L 0.100 M 0.200 M

Simple Titration Calculations Definition of end point, transition point, stoichiometric point Calculations to find % purity of reactants using titration # of protons donated Practical aspects(choice of indicators, unstable standard solution, curves)

Calculate the volume of 0. 100 M H2CO3 required to. neutralize 25 Calculate the volume of 0.100 M H2CO3 required to neutralize 25.0 mL of 0.200 M KOH. 1H2CO3 + 2KOH → K2CO3 + 2HOH ? L 0.0250 L 0.100 M 0.200 M 0.0250 L KOH x 0.200 mol x 1 mole H2CO3 x 1 L = 0.0250 L L 2 mole KOH 0.100 mol

LAB PROBLEM #1: Standardize a solution of NaOH — i. e LAB PROBLEM #1: Standardize a solution of NaOH — i.e., accurately determine its concentration. 35.62 mL of NaOH is neutralized with 25.2 mL of 0.0998 M HCl by titration to an equivalence point. What is the concentration of the NaOH?

2. Calculations to find % purity of reactants using titration

# of protons donated by acid 3. Calculations to find molar ratio of acid and base in neutralization equation. Or # of protons donated by acid

2. Calculations to find molar ratio of acid and base in neutralization equation. One proton donated H3PO4 + KOH → KH2PO4 + H2O Two protons donated H3PO4 + 2KOH → K2HPO4 + H2O Three protons donated H3PO4 + 3KOH → K3PO4 + H2O

# of protons donated by acid 2. Calculations to find molar ratio of acid and base in neutralization equation. Or # of protons donated by acid How can we calculate Which is correct equation(correct molar ratio)? Which salt is being produced?

0.25M, 20mL H3PO4 reacts with 50mL or 0.20M KOH. 1. How many protons does H3PO4 donates to KOH? 2. What salts is produced? (write formula) One proton donated H3PO4 + KOH → KH2PO4 + H2O Two protons donated H3PO4 + 2KOH → K2HPO4 + H2O Three protons donated H3PO4 + 3KOH → K3PO4 + H2O

0.25M, 20mL H3PO4 reacts with 50mL or 0.20M KOH. 1. How many protons does H3PO4 donates to KOH? 2. What salts is produced? (write formula) H3PO4 + KOH → Moles of H3PO4 : moles of KOH =0.25M X 0.0020 M : =0.20M X 0.050L =0.005 : =0.010 =1 : =2 H3PO4 + 2KOH → K2HPO4 + H2O Two protons donated

Two protons donated Salt Produced = K2HPO4 0.25M, 20mL H3PO4 reacts with 50mL or 0.20M KOH. 1. How many protons does H3PO4 donates to KOH? 2. What salts is produced? (write formula) H3PO4 + KOH → Moles of H3PO4 : moles of KOH =0.25M X 0.0020 M : =0.20M X 0.050L =0.005 : =0.010 =1 : =2 H3PO4 + 2KOH → K2HPO4 + H2O Two protons donated Salt Produced = K2HPO4

Question 94 On Page 158 Home Work: 95-99

4. Titration standards

 A standard solution has known molarity. A primary standard is made by weighing a pure solid and diluting in a volumetric flask. A secondary standard requires a titration to calibrate its concentration. It was made using an acid or base with suspect purity. Not suitable Example Reason Arrhenius bases NaOH Hygroscopic Strong acids HCl volatile liquids or gases Suitable solid bronsted bases Na2CO3 pure solid solid weak acids H2C2O4 pure solid

Primary Standard Secondary Standard   NaOH(s) Na2CO3(s) H2SO4(l) H2C2O4(s) HCl(g) KC6H5COO(s) C6H5COOH(s) NH3(g)

Primary Standard Secondary Standard   NaOH(s) Na2CO3(s) H2SO4(l) H2C2O4(s) HCl(g) KC6H5COO(s) C6H5COOH(s) NH3(g)

Primary Standard Secondary Standard   H2SO4(l) NaOH(s) Na2CO3(s) H2C2O4(s) HCl(g) KC6H5COO(s) C6H5COOH(s) NH3(g)

Primary Standard Secondary Standard   H2SO4(l) NaOH(s) Na2CO3(s) H2C2O4(s) HCl(g) KC6H5COO(s) C6H5COOH(s) NH3(g)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NaOH(s) Na2CO3(s) HCl(g) KC6H5COO(s) C6H5COOH(s) NH3(g)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NaOH(s) Na2CO3(s) HCl(g) KC6H5COO(s) C6H5COOH(s) NH3(g)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) NaOH(s) Na2CO3(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) NaOH(s) Na2CO3(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) NaOH(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) NaOH(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) NaOH(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) NaOH(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) NaOH(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) NaOH(s) HCl(g) KC6H5COO(s) C6H5COOH(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) C6H5COOH(s) NaOH(s) HCl(g) KC6H5COO(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) C6H5COOH(s) NaOH(s) HCl(g) KC6H5COO(s)

Primary Standard Secondary Standard   H2SO4(l) H2C2O4(s) NH3(g) Na2CO3(s) C6H5COOH(s) NaOH(s) KC6H5COO(s) HCl(g)

Preparing a Primary Standard 1. Calculate the mass of H2C2O4.2H2O require to prepare 100.0 mL of 0.1000 M standard solution. 0.1000 L

Preparing a Primary Standard 1. Calculate the mass of H2C2O4.2H2O require to prepare 100.0 mL of 0.100 M standard solution. 0.1000 L x 0.1000 mole 1 L

Preparing a Primary Standard 1. Calculate the mass of H2C2O4.2H2O require to prepare 100.0 mL of 0.100 M standard solution. 0.1000 L x 0.1000 mole x 126.06 g 1 L 1 mole

Preparing a Primary Standard 1. Calculate the mass of H2C2O4.2H2O require to prepare 100.0 mL of 0.100 M standard solution. 0.1000 L x 0.1000 mole x 126.06 g = 1.261 g 1 L 1 mole

Titration indicators

  The equivalence point is the end of a titration where the stoichiometry of the reaction is exactly satisfied, or moles H+ = moles OH-. The transition point refers to when an indicator changes color and [HInd] = [Ind-].

Choosing an Indicator When you choose an indicator, you must pick one so that the transition point of the indicator matches the equivalence point of the titration.   Rule of thumb Salt Equivalence Point Neutral 7 Basic 9 Acidic 5

Titration Curves

Titration Curves  A titration curve is a graph of the pH changes that occur during an acid-base titration versus the volume of acid or base added.

1. Titration Curve: Strong Acid and Strong Base HCl + KOH → KCl + HOH neutral salt Indicator pH = 7 Bromothymol Blue- see page 7 50 mL of 0.10 M KOH is added to 25 mL of 0.10 M HCl We start here The pH of 0.10 M HCl is 1.0 The pH of 0.10 M KOH is 13.0 pH Volume .10 M KOH added 14 7 25 50 0.10 M KOH 0.10 M HCl Neutral Salt pH = 7.0

2. Titration Curve: Weak Acid and Strong Base HCN + KOH → KCN + H2O basic salt Indicator pH = 9 Phenolphthalein- see page 7 20 mL of 1.0 M HCN is added to 10 mL of 1.0 M KOH is added to We end here pH = 3 We start here pH = 14 pH Volume 1.0 M HCN added 14 7 1.0 M KOH Basic Salt pH = 9 1.0 M HCN 10 20

3. Titration Curve: Strong Acid and Weak Base HCl + NH3 → NH4+ + Cl- acid salt Indicator pH = 5 Methyl Red- see page 7 60 mL of 1.0 M NH3 is added to 30 mL of 1.0 M HCl 1.0 M NH3 pH  10 pH Volume 1.0 M NH3 added 14 7 Acid Salt pH  5 30 60 1.0 M HCl pH = 0

 4. Match the Curve with the Reaction pH Volume .10 M KOH added 14 7 25 50 A. 3HCl + Al(OH)3 → AlCl3 + 3HOH B. HCl + KOH → KCl + HOH C. HCN + KOH → KCN + HOH 

 5. Match the Curve with the Reaction pH Volume 1.0 M NH3 added 14 7 30 60  A. 3HCl + Al(OH)3 → AlCl3 + 3HOH B. HCl + KOH → KCl + HOH C. HCN + KOH → KCN + HOH

 6. Match the Curve with the Reaction pH Volume 1.0 M NH3 added 14 7 30 60 A. 3HCl + Al(OH)3 → AlCl3 + 3HOH B. HCl + KOH → KCl + HOH C. HCN + KOH → KCN + HOH 

Features of the Strong Acid-Strong Base Titration Curve The pH starts out low, reflecting the high [H3O+] of the strong acid and increases gradually as acid is neutralized by the added base. Suddenly the pH rises steeply. This occurs in the immediate vicinity of the equivalence point. For this type of titration the pH is 7.0 at the equivalence point. Beyond this steep portion, the pH increases slowly as more base is added.

HPr = Propionic Acid

The four Major Differences Between a Strong Acid-Strong Base Titration Curve and a Weak Acid-Strong Base Titration Curve The initial pH is higher. A gradually rising portion of the curve, called the buffer region, appears before the steep rise to the equivalence point. The pH at the equivalence point is greater than 7.00. The steep rise interval is less pronounced.