Titrations Main Idea: Titrations are an application of acid-base neutralization reactions that require the use of an indicator. www.ibchem.com/ppt/shelves/aab/indicpp.ppt

NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l) Stoichiometry The stoichiometry of an acid-base neutralization reaction is the same as that of any other reaction that occurs in solution (they are double displacement reactions, after all). For example, in the reaction of sodium hydroxide and hydrogen chloride, 1 mol of NaOH neutralizes 1 mol of HCl: NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l) Stoichiometry provides the basis for a procedure called titration, which is used to determine the concentrations of acidic and basic solutions.

Titration Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration. If you wish to find the concentration of an acid solution, you would titrate the acid solution with a solution of a base of known concentration. You could also titrate a base of unknown concentration with an acid of known concentration.

In the titration of an acid by a base, the pH meter measures the pH of the acid solution in the beaker as a solution of a base with a known concentration is added from the buret. http://wps.prenhall.com/wps/media/objects/3312/3392202/blb1703.html

How is an acid-base titration performed? The figure on the previous slide illustrates one type of setup for the titration procedure outlined on the next slide. In the procedure pictured on Slide 4, a pH meter is used to monitor the change in pH as the titration progresses.

Titration Procedure A measured volume of an acidic or basic solution of unknown concentration is placed in a beaker. The electrodes of a pH meter are immersed in this solution, and the initial pH of the solution is read and recorded. A buret is filled with the titrating solution of known concentration. This is called the standard solution, or titrant. Measured volumes of the standard solution are added slowly and mixed into the solution in the beaker. The pH is read and recorded after each addition. This process continues until the reaction reaches the equivalence point, which is the point at which moles of H+ ion from the acid equal moles of OH- ion from the base.

In the titration of a strong acid by a strong base, a steep rise in the pH of the acid solution indicates that all of the H+ ions from the acid have been neutralized by the OH- ions of the base. The point at which the curve flexes is the equivalence point of the titration. Bromthymol blue is an indicator that changes color at this equivalence point. Notice that phenolphthalein and methyl red don’t match the exact equivalence point, but the slope is so steep that it doesn’t matter.

COLOUR CHANGES OF SOME COMMON INDICATORS Acid-base indicators Must have an easily observed colour change. Must change immediately in the required pH range over the addition of ‘half’ a drop of reagent. COLOUR CHANGES OF SOME COMMON INDICATORS pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 METHYL ORANGE CHANGE LITMUS CHANGE PHENOLPHTHALEIN CHANGE

Must change immediately in the required pH range Acid-base indicators Must have an easily observed colour change. Must change immediately in the required pH range over the addition of ‘half’ a drop of reagent. To be useful, an indicator must change over the “vertical” section of the curve where there is a large change in pH for the addition of a very small volume of alkali. The indicator used depends on the pH changes around the end point - the indicator must change during the ‘vertical’ portion of the curve. In the example, the only suitable indicator is PHENOLPHTHALEIN. PHENOLPHTHALEIN LITMUS METHYL ORANGE

weak acid (CH3COOH) v. strong alkali (NaOH) pH curves Types There are four types of acid-base titration; each has a characteristic curve. strong acid (HCl) v. strong base (NaOH) weak acid (CH3COOH) v. strong alkali (NaOH) strong acid (HCl) v. weak base (NH3) weak acid (CH3COOH) v. weak base (NH3) In the following examples, alkali (0.1M) is added to 25cm3 of acid (0.1M) End points need not be “neutral‘ due to the phenomenon of salt hydrolysis

strong acid (HCl) v. strong base (NaOH)

strong acid (HCl) v. strong base (NaOH) pH 1 at the start due to 0.1M HCl (strong monoprotic acid)

strong acid (HCl) v. strong base (NaOH) Very little pH change during the initial 20cm3 pH 1 at the start due to 0.1M HCl (strong monoprotic acid)

strong acid (HCl) v. strong base (NaOH) Very sharp change in pH over the addition of less than half a drop of NaOH Very little pH change during the initial 20cm3 pH 1 at the start due to 0.1M HCl (strong monoprotic acid)

strong acid (HCl) v. strong base (NaOH) Curve levels off at pH 13 due to excess 0.1M NaOH (a strong alkali) Very sharp change in pH over the addition of less than half a drop of NaOH Very little pH change during the initial 20cm3 pH 1 at the start due to 0.1M HCl (strong monoprotic acid)

strong acid (HCl) v. strong base (NaOH) PHENOLPHTHALEIN LITMUS METHYL ORANGE Any of the indicators listed will be suitable - they all change in the ‘vertical’ portion

Strong-Strong Titration The previous slide shows how the pH of the solution changes during the titration of 50.0 mL of 0.100 M HCl, a strong acid with 0.100 M NaOH, a strong base. The inital pH of the 0.100 M HCl is 1.00. As NaOH is added, the acid is neutralized and the solution’s pH increases gradually. When nearly all of the H+ ions from the acid have been used up, the pH increases dramatically with the addition of an exceedingly small volume of NaOH. This abrupt change in pH occurs at the equivalence point of the titration. Beyond the equivalence point, the addition of more NaOH again results in the gradual increase in pH.

Acid-base indicators General Many indicators are weak acids and partially dissociate in aqueous solution HIn(aq) H+(aq) + In¯(aq) The un-ionised form (HIn) is a different colour to the anionic form (In¯).

The equivalence point here is not at a pH of 7 The equivalence point here is not at a pH of 7. Phenolphthalein is an indicator that changes color at this equivalence point. Notice that the starting pH is different and the region of change is smaller.

Acid-Base Indicators Chemists often use a chemical dye rather than a pH meter to detect the equivalence point of an acid-base titration. Chemical dyes whose colors are affected by acidic and basic solutions are called acid-base indicators. Many natural substances act as indicators. If you use lemon juice in your tea, you might have noticed that the brown color of tea gets lighter when lemon juice is added. Tea contains compounds called polyphenols that have slightly ionizable hydrogen atoms and therefore are weak acids. Adding acid in the form of lemon juice to a cup of tea lessens the degree of ionization, and the color of the un-ionized polyphenols becomes more apparent. Chemists have several choices in selecting indicators. Bromthymol blue is a good choice for the titration of a strong acid with a strong base, and phenolphthalein changes color at the equivalence point of a titration of a weak acid with a strong base.

Indicators and Titration End Point Many indicators used for titrations are weak acids. Each has its own particular pH or pH ranges over which it changes color. The point at which the indicator used in a titration changes color is called the end point of the titration. It is important to choose an indicator for a titration that will change color at the equivalence point of the titration. Remember that the role of the indicator is to indicate to you, by means of a color change, that just enough of the titrating solution has been added to neutralize the unknown solution. Equivalence point ≠ End point! BUT for strong-strong titrations, the pH change is so steep and so large, that the are approximately equal.

Titration with an Indicator

What’s the Point of a Titration Again? To find the unknown concentration of an acid or a base. So you perform the actual titration noting the volume you started with and how much volume of the titrant you added and then... Math! (Oh no! Not math! Anything but math!)

Titration Calculations: An Example The balanced equation of a titration reaction is the key calculating the unknown molarity. For example, sulfuric acid is titrated with sodium hydroxide according to this equation: H2SO4 (aq) + 2 NaOH (aq)  Na2SO4 (aq) + 2 H2O (l) Calculate the moles of NaOH in the standard from the titration data: molarity of the base (MB) and the volume of the base (VB). In other words, MB VB = (mol/L)(L) = mol NaOH in standard From the equation, you know that the mole ratio of NaOH to H2SO4 is 2:1. Two moles of NaOH are required to neutralize 1 mol of H2SO4. mol H2SO4 titrated = mol NaOH in standard x (1 mol H2SO4 / 2 mol NaOH) MA represents the molarity of the acid and VA represents the volume of the acid in liters. MA = mol H2SO4 titrated/VA

In Short Form... MAVA = MBVB (mol acid/mol base) This is the mole ratio Does this make sense? Let’s find out using the definition of molarity (mol/L) and dimensional analysis... (mol acid/L acid)(L acid) = (mol base/L base)(L base) (mol acid/mol base) mol acid = mol base (mol acid/mol base) mol acid = mol acid

HOMEWORK QUESTIONS What is the purpose of a titration? How is it performed (in general)? What is the difference between the equivalence point and the end point of a titration? Describe the differences in the titration curve of a strong-strong titration vs. a strong-weak titration. When is the equivalence point of a titration not at pH 7? Why does this occur?

MORE HOMEWORK 5) What is the molarity of a nitric acid solution if 43.33 mL of 0.1000 M KOH solution is needed to neutralize 20.00 mL of the acid solution? 6) What is the concentration of a household ammonia cleaning solution if 49.90 mL of 0.5900 M HCl is required to neutralize 25.00 mL of the solution? 7) How many milliliters of 0.500 M NaOH would neutralize 25.00 mL of 0.100 M H2SO4?