Chapter 21

Neutralization

21.1 NeutralizationLesson ObjectivesThe student will:

• explain what is meant by a neutralization reaction and give an example of one.• write a balanced equation for the reaction that occurs when an acid reacts with a base.• describe the formation of a salt in terms of the Arrhenius definitions of acids and bases.• predict the salt that will be produced from the neutralization reaction between a given acid and base.• identify acidic, basic, and neutral salts from a neutralization reaction.

Vocabulary• acidic salt• basic salt• neutral salt• neutralization

IntroductionNeutralization is a reaction between an acid and a base that produces water and a salt. The generalreaction is shown below:

acid + base → salt + water

After reviewing the concept of ionic compounds, we will examine neutralization reactions and look at thedifferent types of salts that can be formed from acids and bases as they react with one another.

Neutralization ReactionsAcids are a combination of a hydrogen cation and a nonmetal anion. Examples include HCl, HNO3, andHC2H3O2. Many bases are a combination of metal cations and nonmetal anions. Examples include NaOH,

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KOH, and Mg(OH)2. According to the Arrhenius definitions of acids and bases, the acid will contributean H+ ion that will react to neutralize the OH− ion contributed by the base, producing neutral watermolecules.All acid-base reactions produce salts. The anion from the acid will combine with the cation from the baseto form the ionic salt. Examples are shown below.

HClO4(aq) +NaOH(aq) → NaClO4(aq) +HOH(l)H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2HOH(l)(Note: HOH = H2O)

No matter what the acid or the base may be, the products of this type of reaction will always be a saltand water. Aside from the fact that the H+ ion will neutralize the OH− ion to form water, we also knowthat these reactions are double displacement reactions, because they consist of cations exchanging anions.We can write this as a total ionic equation:

H+(aq) + ClO−4(aq) +Na+

(aq) +OH−(aq) → Na+(aq) + ClO−4(aq) +H2O(l)

However, since Na+ and ClO−4 are spectator ions, the net ionic equation is:

H+(aq) +OH−(aq) → H2O(l)

This is the net ionic equation for all neutralization reactions.

Salt HydrolysisWhen a salt is dissolved in water, it is possible for the solution to be neutral, acidic, or basic. If a solutionis to be acidic, it must contain more hydrogen ions than hydroxide ions. For the solution to be basic, itmust contain more hydroxide ions than hydrogen ions. Consider the solution produced when the salt KBrdissolves in water. There will be four ions present in the solution.

K+ + Br− +H+ +OH−

When potassium ions in the solution come into contact with hydroxide ions, the KOH ion pair couldpotentially form. However, since KOH is a strong base, it would immediately dissociate back into ions.Similarly, if the bromide ions come into contact with hydrogen ions, the molecule formed would be HBr,a strong acid, so they would immediately dissociate back into ions. Thus, having potassium and bromideions in a water solution would not cause a reaction.Now consider the solution produced when the salt NH4Cl is dissolved in water. There will be four ionspresent in the solution.

NH+4 + Cl− +H+ +OH−

When hydrogen ions come into contact with chloride ions, if they join together, the resultant moleculewould be HCl, which is a strong acid. Therefore, the HCl would immediately dissociate back into the ions.When NH+

4 ions come into contact with OH− ions, however, the resultant molecule would be NH4OH,

which is a weak base and does not dissociate very much. Therefore, when ammonium chloride is dissolvedin water, a reaction occurs.

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NH+4 + Cl− +H+ +OH− → NH4OH(aq) + Cl− +H+

The ammonium hydroxide dissociates very little, so we would have mostly un-dissociated ammoniumhydroxide molecules in solution with hydrogen and chloride ions. The hydrogen ions in this final solutionwould cause the solution to be acidic. Thus, dissolving ammonium chloride in water produces an acidicsolution.By a similar process, dissolving sodium acetate, NaC2H3O2, in water will produce a basic solution. Whenthe sodium acetate is dissolved in water, four ions will be present in the solution.

Na+ + C2H3O−2 +H+ +OH−

If sodium ions contact hydroxide ions, the substance formed would be a strong base which would immedi-ately dissociate. If hydrogen ions contact acetate ions, however, the molecule formed would be acetic acid,which is a weak acid and the ions would NOT dissociate. Therefore, when sodium acetate is dissolved inwater, a reaction will occur as shown below.

Na+ + C2H3O−2 +H+ +OH− → HC2H3O2(aq) +Na+ +OH−

The resultant solution has excess of hydroxide ions, so it is basic. Dissolving sodium acetate in waterproduces a basic solution.

Neutral, Acidic, and Basic SaltsTable 21.1 shows all of the strong acids and bases that we have encountered so far. You can assume thatany other acids and bases we will look at are weak.

Table 21.1: Strong Acids and Bases

Strong Acid Formula Strong Base FormulaHydrochloric Acid HCl Lithium Hydroxide LiOHHydrobromic Acid HBr Sodium Hydroxide NaOHHydroiodic Acid HI Potassium Hydroxide KOHNitric Acid HNO3 Rubidium Hydroxide RbOHPerchloric Acid HClO4 Cesium Hydroxide CsOHSulfuric Acid H2SO4 Calcium Hydroxide Ca(OH)2

Strontium Hydroxide Sr(OH)2Barium Hydroxide Ba(OH)2

The information in the table helps us to determine what type of salt is formed in an acid-base reaction.For example, a reaction between a strong acid and a strong base will form a neutral salt. It is like apower struggle between the acid and the base. Since both are strong, neither overpower the other, so thesalt ends up being neutral. If, however, we have a reaction between a weak acid and a strong base, theresult would be a basic salt.

HC2H3O2(aq) +NaOH(aq) → NaC2H3O2(aq) +H2O(l)

Acetic acid is a weak acid and sodium hydroxide is a strong base. Therefore, the salt formed, sodium

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acetate, is a basic salt.When the basic salt is dissolved in water, a reaction takes places in which extra hydroxide ions, OH−, areproduced from the salt and the water molecules. A similar situation will occur when we have a strong acidreacting with a weak base. When a strong acid reacts with a weak base, an acidic salt is formed.

HCl(aq) +NH4OH(aq) → NH4Cl(aq) +H2O(l)

Hydrochloric acid is a strong acid and ammonium hydroxide is a weak base. Therefore, the salt formed,ammonium chloride, is an acidic salt. The salt will react with water molecules to produce hydrogen ions,H+, so it is referred to as an acidic salt.Example:Complete the following neutralization reactions and identify the type of salt produced.

1. H2SO4 + Ba(OH)2 →?2. HCOOH+ Ca(OH)2 →?

Solution:

1. H2SO4 +Ba(OH)2 → BaSO4 +2 H2O; A strong acid reacts with a strong base and produces a neutralsalt.

2. 2 HCOOH+Ca(OH)2 → Ca(HCOO)2 + 2 H2O; A weak acid reacts with a strong base and producesa basic salt.

By determining which acid and base were used to form the salt, you can figure out if the salt is acidic,basic, or neutral. For example, let’s say you are looking at calcium nitrate, Ca(NO3)2. Remember thatthere is a double displacement reaction that forms the salt, so we can write out the reaction:

Therefore, the neutralization reaction would have been:

2 HNO3 + Ca(OH)2 → Ca(NO3)2 + 2 HOH

This salt would have been produced from a strong acid-strong base reaction, so it is neutral.Let’s try another. Consider the salt copper(II) chloride (CuCl2).The copper would have come from the base, Cu(OH)2, which is a weak base. The chloride would havecome from the acid, HCl, which is a strong acid.

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2 HCl+ Cu(OH)2 → CuCl2 + 2 HOH

The reaction is a strong acid weak base reaction, so the salt should be acidic.

This video is an overview of endothermic and exothermic reactions and includes a demonstration of anexothermic reaction (7b): http://www.youtube.com/watch?v=sAyPDTQHEeE (2:31).

Figure 21.1: (Watch Youtube Video)http://www.ck12.org/flexbook/embed/view/413

Lesson Summary• A neutralization reaction between an acid and a base will produce a salt and water.• In a neutralization reaction, the acid will produce H+ ions that react to neutralize the OH- ionsproduced by the base, forming neutral water. The other product will be an ionic salt.

• A strong acid + a strong base in an acid/base neutralization reaction will form a neutral salt.• A strong acid + a weak base in an acid/base neutralization reaction will form an acidic salt.• A weak acid + a strong base in an acid/base neutralization reaction will form a basic salt.

Further Reading / Supplemental Links• The following link shows a video of a neutralization reaction.

– http://www.youtube.com/watch?v=_P5hGzA6Vb0

Review Questions1. How do an acid and a base fit the definition of an ionic compound? Use examples in your answer.2. Explain neutralization reactions in terms of Arrhenius theory. Use an example in your answer.3. Which salt will form a basic solution when dissolved in water?

(a) KNO3

(b) CaCl2(c) NaClO4

(d) NaNO2

4. Which salt will form an acidic solution when dissolved in water?(a) copper(II) sulfate(b) sodium acetate

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(c) potassium chloride(d) sodium cyanide

5. Milk of magnesia is a common over-the-counter antacid that has magnesium hydroxide as its mainingredient. It is used by the public to relieve acid indigestion. Acid indigestion is caused by excessstomach acid being present. Write the balanced chemical equation for the reaction between milk ofmagnesia and hydrochloric acid. What type of reaction is this? What type of salt is formed?

6. Complete the following neutralization reactions and identify the type of salt produced.(a) H2SO4(aq) + NaOH(aq) →(b) HNO3(aq) + NH4OH(aq) →(c) HF(aq) + NH4OH(aq) →(d) CH3COOH(aq) + KOH(aq) →(e) HCl(aq) + KOH(aq) →

21.2 TitrationLesson ObjectivesThe student will:

• explain what an acid/base indicator is.• explain how acid-base indicators work.• explain the difference between natural and synthetic indicators.• explain how indicators are used in the lab.• explain what a titration is.• describe how titrations can be used to determine the concentration of an acid or a base in solution.• explain the difference between the equivalence point and the end point.• define a standard solution in terms of acid-base titrations.• calculate the concentration of an acid or base solution using a standard solution.• calculate the concentration of unknown acid or base when given the concentration of the other andthe volume needed to reach the equivalence point in a titration.

Vocabulary• endpoint• equivalence point• natural indicator• standard solution• synthetic indicator• titrant• titration• titration curve

IntroductionThe typical laboratory procedure for determining the concentration of acid and/or base in a solution is tocomplete a titration. There are three main types of titration experiments. As we go through this lesson,we will take apply the knowledge we have obtained about acids and bases, chemical reactions, and molaritycalculations to the concept of titrations.

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Indicators

Recall from the chapter on “Acids-Bases” that an indicator is a substance that changes color at a specificpH and is used to indicate the pH of the solution. One example of an indicator is litmus paper. Litmuspaper is paper that has been dipped in a substance that will undergo a color change when it is exposed toeither an acid or a base. If red litmus paper turns blue, the solution is basic (pH > 7), and if blue litmusturns red the solution is acidic (pH < 7).A natural indicator is an indicator that is a naturally occurring substance. For example, the juice fromred cabbage can be used to prepare an indicator paper. It contains the chemical anthrocyanin, which isthe active ingredient in the indicator. Red beets, blueberries, and cranberries are other great examplesof naturally occurring indicators. These are all due to the same anthocyanin molecule found in the redcabbage.Some flowers are also natural indicators. Hydrangea is a common garden plant with flowers that come inmany colors, depending on the pH of the soil. A hydrangea plant with blue flowers indicates that the soilis acidic, while creamy white flowers mean the soil is neutral and pink flowers mean the soil is basic.

Figure 21.2: A hydrangea plant with blue flowers. What does the flower color indicate about the pH ofthe soil?

Synthetic indicators are compounds created in a chemistry lab rather than compounds found in nature.Both naturally occurring indicators and synthetic indicators are weak organic acids or bases. For example, acommon synthetic indicator used in most chemistry laboratories is phenolphthalein. The chemical structureof phenolphthalein is shown in the figure below.

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This indicator changes color at a pH of 8.2. Below 8.2 it is colorless, and above 8.2 it is bright pink. Thereare many common synthetic indicators that are useful in the chemistry laboratory. When dealing with amore acidic range, chemistry students may use methyl orange. The structure for methyl orange is shownbelow.

Methyl orange changes color from pH 3.2 to 4.4. Below 3.2, the color of the indicator is red. Above 4.4,the color of the indicator is yellow. In between 3.2 and 4.4, there are various shades of orange, hence thename.There are two requirements for a substance to function as an acid-base indicator: 1) the substance musthave an equilibrium affected by hydrogen ion concentration, and 2) the two forms of the compound onopposite sides of the equilibrium must have different colors. Most indicators function in the same generalmanner and can be presented by a generic indicator equation. In the equation below, we represent in theindicator ion with a hydrogen ion attached as HIn, and we represent the indicator ion without the hydrogenattached as In−.

Since the indicator itself is a weak acid, the equilibrium between the protonated form and the anionic formis controlled by the hydrogen ion concentration. For the example above, the protonated form is colored

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red and the anionic form is colored yellow. If we add hydrogen ion to the solution, the equilibrium will bedriven toward the reactants and the solution will turn red. If we add base to the solution (reduce hydrogenion concentration), the equilibrium will shift toward the products and the solution will turn yellow. It isimportant to note that if this indicator changes color at pH = 5, then at all pH values less than 5, thesolution will be red and at all pH values greater than 5, the solution will be yellow. Therefore, putting thisindicator into a solution and having the solution turn yellow does NOT tell you the pH of the solution, itonly tells you that the pH is greater than 5. At pH values less than 5, the great majority of the indicatormolecules are in the red form and the solution will be red. At pH values greater than 5, the great majorityof the indicator particles will be in the yellow form and the solution will be yellow. The equilibrium betweenthese indicator particles is such that the particles will be 50% red form and 50% yellow form at exactlypH = 5. Therefore, at pH = 5, the actual color of the solution will be a 50-50 mixture of red and yellowparticles and the solution will be orange, as demonstrated in the figure below.

Many indicators are available to help determine the pH of solutions. A list of the most common indicators isfound in Table 21.2, along with their respective color change pH valuess and corresponding color changes.

Table 21.2: Colors and pH Ranges for Common Indicators

Indicator pH Range Color ChangeMethyl Violet 0.0 − 1.6 Yellow - BlueThymol Blue 1.2 − 2.8 Red - YellowOrange IV 1.3 − 3.0 Red - YellowMethyl Orange 3.2 − 4.4 Red - OrangeBromophenol Blue 3.0 − 4.7 Orange/Yellow - VioletCongo Red 3.0 − 5.0 Blue - RedBromocresol Green 3.8 − 5.4 Yellow - BlueMethyl Red 4.8 − 6.0 Red - YellowLitmus 5.0 − 8.0 Red - BlueChlorophenol Red 4.8 − 6.2 Yellow - RedBromothymol Blue 6.0 − 7.6 Yellow - BluePhenol Red 6.4 − 8.2 Yellow - Red/VioletThymol Blue 8.0 − 9.6 Yellow - BluePhenolphthalein 8.2 − 10.0 Colorless - PinkAlizarin Yellow R 10.1 − 12.0 Yellow - RedMethyl Blue 10.6 − 13.4 Blue - Pale VioletIndigo Carmine 11.4 − 13.0 Blue - Yellow

There are many more indicators than are shown in Table 21.2, but these are ones that you may find incommon chemistry classroom laboratories. One example of an indicator not found in the table is knownas the universal indicator. The universal indicator is a solution that has a different color for each pH from0–14. Universal indicator is produced by creatively mixing many of the individual indicators together sothat a different color is achieved for each different pH. It is used for many types of experiments to determine

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if solutions are acids or bases and where on the pH scale the substance belongs. The chart below indicatesthe colors of universal indicator for different pH values.

Example:If the pH of the solution is 4.8, what would be the color of the solution if the following indicators wereadded?

1. Universal indicator2. Bromocresol Green3. Phenol red

Solution:

1. Universal indicator = Orange to orange-yellow2. Bromocresol Green = green (midway pH = 4.6)3. Phenol red = yellow

Example:A solution found in the laboratory was tested with a number of indicators. These were the results:

• Phenolphthalein was colorless• Bromocresol green was blue• Methyl red was yellow• Phenol red was yellow

What was the pH of the solution?Solution:

• Phenolphthalein was colorless, pH < 8.0• Bromocresol green was blue, pH > 5.4• Methyl red was yellow, pH > 6.0• Phenol red was yellow, pH < 6.4

Therefore, the pH of the solution must be between 6.0 and 6.4.

The Titration ProcessOne of the properties of acids and bases is that they neutralize each other. In the laboratory setting,an experimental procedure where an acid is neutralized by a base (or vice versa) is known as titration.Titration is the addition of a known concentration of base (or acid), also called the titrant, to a solution

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of acid (or base) of unknown concentration. Since both volumes of the acid and base are known, theconcentration of the unknown solution is then mathematically determined.When doing a titration, you need to have a few pieces of equipment. A burette like the one shown below isused to accurately dispense the volume of the solution of known concentration. An Erlenmeyer flask is usedto hold a known volume of the solution whose concentration is unknown. A few drops of the indicator areadded to the flask before you begin the titration. The endpoint is the point where the indicator changescolor, which tells us that the acid is neutralized by the base. The equivalence point is the point wherethe number of moles of acid exactly equals the number of moles of base.

Some laboratories have pH meters that measures this point more accurately than the indicator. Thediagram below shows a simplified version of a pH meter with the probe from the meter immersed in amildly alkaline solution (pH = 8.03). The two knobs on the meter are used to calibrate the instrument.

An example of a typical electronic pH meter with the attached probes is shown below. The main purposeof a pH meter in this experiment is to measure the changes in pH as the titration goes from start to finish.

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A typical titration setup is shown below. The burette is upright and ready to drip the solution into theflask holding the solution of unknown concentration and the few drops of indicator. When the indicatorchanges color, the number of moles of acid equals the number of moles of base and the acid (or base) hasbeen neutralized.

There are three types of titrations that are normally performed in the laboratory in order to determinethe unknown concentration of the acid or base. These three types are:

1. Strong acid vs. Strong base2. Strong acid vs. Weak base3. Weak acid vs. Strong base

In these titrations, a pH meter may be used to measure the changes in the pH as the titration goes tocompletion. If so, a titration curve can be constructed. A titration curve is a graph of the pH versus thevolume of titrant added. Let’s take a look at how each of these types of titrations differs in terms of theirpH curves and their pH at the equivalence point.

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(1) Strong Acid vs. Strong BaseFor a strong acid vs. a strong base titration, let’s assume the strong base is the titrant. Therefore, theErlenmeyer flask contains the strong acid and a few drops of your indicator. The initial pH of the solutionin the flask will likely be low since the solution is a strong acid. As the base is added, the acid is slowlyneutralized. At first the change in pH is minimal. This is due to the fact that the flask has a much greaternumber of H3O+ ions than OH− ions available from the added titrant.As more and more base is added, more OH− ions are added and thus more H3O+ ions get neutralized.Let’s stop here and look at the reaction. The equation below shows the total ionic equation of a reactionbetween a strong acid and a strong base:

H+(aq) + Cl−(aq) +Na+

(aq) +OH−(aq) → Na+(aq) + Cl−(aq) +H2O(l)

The next equation shows the net ionic equation for the reaction between the strong acid and the strongbase:

H+(aq) +OH−(aq) → H2O(l)

As we add more OH− ions, more H3O+ (or H+) ions are being neutralized. Since these two ions react toform water, a neutral solution will eventually be formed. For a strong acid and a strong base, this meansthe pH = 7.0 at the point of neutralization. If we continue to add the titrant (containing OH− ions) afterall of the H3O+ ions have been neutralized, the pH will continue to rise as more base is added and thereare excess OH− ions.Now that we know what happens in a strong acid-strong base titration, what does the titration curve looklike? The main points described above are shown in the titration curve below.

The points A through D sum up the description of the events that take place during the titration. PointA is the start of the titration. Point B is the midpoint, the point where half of the H+ ions have beenneutralized. Point D is the equivalence point.(2) Strong Acid vs. Weak BaseWhat would happen if we were to titrate a strong acid with a weak base or vice versa? The titration curvefor a weak base-strong acid titration is shown below. Try to determine what is happening in the titrationjust by looking at the graph.

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As the acid (the titrant) is added, the pH decreases as the H3O+ ions begin to neutralize the OH− ions.Point D is the equivalence point. Notice that for a weak base and a strong acid titration, the pH atequivalence point is acidic. The equation for the reaction between NH3, a weak base, and HCl, a strongacid, is shown below:

NH3(aq) +HCl(aq) → NH4Cl(aq) +H2O(l)

The ionic equation is:

NH3(aq) +H+(aq) + Cl−(aq) → NH

+4(aq) + Cl−(aq) +H2O(l)

(3) Weak Acid vs. Strong BaseThe third type of titration is that of a weak acid with a strong base. When we follow through with thesame procedure as the previous two titrations, we can determine a great deal of information simply bylooking at the pH curve. For example, let’s consider the titration of a solution of acetic acid, HC2H3O2,with a solution of potassium hydroxide, KOH. We can write the chemical reaction for this acid-baseneutralization and begin to draw a rough sketch of a titration curve:

H+(aq) + C2H3O−2(aq) +K+

(aq) +OH−(aq) → K+(aq) + C2H3O−2(aq) +H2O(l)

The points on the curve represent the same points as with the other two titration curves. Look, however, atthe equivalence point. Notice how the pH for the equivalence point of the weak acid-strong base titrationis above 7.0.

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Example:Draw a rough sketch of the titration curve between nitric acid and ethylamine, CH3NH2. Assume the acidis in the burette. What is the estimated pH at the equivalence point?Solution:

The <H at the equivalence point is approximately 4.6 from this graph.The titrant is the solution of known concentration. For accuracy reasons, this titrant is normally titratedto find its exact concentration before beginning the desired titration. The purpose of this initial titrationis to determine, with as much accuracy as possible, the exact concentration of the solution in the burette.To determine the exact concentration of the titrant, we use a standard solution. A standard solutionis a solution whose concentration is known exactly. Standard solutions have this property because thesechemicals are normally found in pure, stable forms. Examples of chemicals used to prepare standardsolutions are potassium hydrogen phthalate, KHC8H4O4 (sometimes referred to as KHP), and sodiumcarbonate, Na2CO3.When using a standard solution, the standard is first prepared by dissolving the solid in a known volumeof water, adding a few drops of indicator, and titrating with the solution that you want to standardize.Example:What is the concentration of a sodium hydroxide solution if 32.34 mL is required to neutralize a solutionprepared by dissolving 1.12 g of KHC8H4O4(s) in 25.00 mL of H2O?Solution:Step 1: Find the moles of KHC8H4O4.

moles KHC8H4O4 = massmolar mass = 1.12g

204.2 g/mol = 5.48 × 10−3 mol

Step 2: Use mole ratio from the reaction to find the moles of NaOH.

KHC8H4O4(aq) +NaOH(aq) → KNaC8H4O4(aq) +H2O(l)

Since the reaction is 1:1, 1 mole of KHP reacts with every mole of NaOH.

mol NaOH = 5.48 × 10−3 mol

Step 3: Determine the concentration of NaOH.

[NaOH] = 5.48×10−3 mol0.03234 L = 0.170 M

Therefore, the exact concentration of the sodium hydroxide solution used in the titration is 0.170 mol/L.

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Choosing an Appropriate IndicatorTo choose an appropriate indicator for a titration, a titration curve is useful. Knowing the pH at equivalencefor the different types of titrations (see Table 21.3) is also necessary.

Table 21.3: pH at Equivalence for Titrations

Type of Titration pH at EquivalenceStrong Acid – Strong Base pH = 7Strong Acid – Weak Base pH < 7Weak Acid – Strong Base pH > 7

Choosing an indicator close to the equivalence point is essential to see the point where all of the H+ ionsand OH− ions have been neutralized. The color change should occur on or around the equivalence point.So, for example, with a strong acid-strong base titration, the pH at equivalence is 7.0. Indicators such asbromothymol blue (pH range = 6.0 - 7.6) and phenol red (pH range = 6.6 - 8.0) are common. Notice themidpoint color (green) for bromothymol blue would appear at a pH = 6.8, which is close to 7.0. For phenolred, the midpoint color (orange) would appear at pH = 7.3, again close to 7.0.The same process is used for other titration types. For a strong acid-weak base titration where the pHat equivalence is less than 7, the indicators normally chosen are methyl red (pH range = 4.8 - 6.0) andchlorophenol red (pH range = 4.8 - 6.2). For a weak acid-strong base titration, where the pH at equivalenceis greater than 7, the indicators normally chosen are phenolphthalein (pH range = 8.2 - 10) and thymolblue (pH range = 8.0 - 9.6). As with strong acid-strong base titrations, the visual observation of theindicator’s midpoint color should signal close proximity to the equivalence point.Example:Look at the graph below and determine the appropriate indicator.

Solution:We first look at the graph and mark the vertical stretch of the titration curve in order to find the half-waymark on this vertical stretch. Looking at the graph, when we follow this half-way mark over to the y-axis,we can see that the equivalence point occurs at approximately pH = 8.8. The indicator appropriate to usewould be phenolphthalein pH range = 8.2 - 10). As soon as the pink color forms, we are at the equivalencepoint.There is an interesting observation about the endpoint that has yet to be mentioned. The endpoint wasdefined earlier as the point where the indicator changes color. In an acid-base neutralization reaction, this

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point may not be the point where all of the H+ ions have been neutralized by OH− ions, or vice versa. Theexperimenter continues titration until the indicator changes color, that is, the endpoint has been reached.The equivalence point is the point where the moles of hydrogen ion and the moles of hydroxide ion areequal. It requires knowledge by the experimenter to select an indicator that will make the endpoint asclose as possible to the equivalent point.

The Mathematics of TitrationFor the calculations involved here, we will only use our acid and base examples where the stoichiometricratio of H+ and OH− is 1:1. To determine the volume required to neutralize an acid or a base, or in otherwords, to reach the equivalence point, we will use a formula similar to the dilution formula:

Ma × Va = Mb × Vb

where Ma is the molarity of the acid, Va is the volume of the acid, Mb is the molarity of the base, and Vb

is the volume of the base. Note that if the acid and base do not neutralize each other in a 1:1 ratio, thisequation does not hold true.Example:When 10.0 mL of a 0.125 mol/L solution of hydrochloric acid, HCl, is titrated with a 0.100 mol/L solutionof potassium hydroxide, KOH, what is the volume of the hydroxide solution required to neutralize theacid? What type of titration is this?Solution:Step 1: Write the balanced ionic chemical equation.

H+ + Cl− +K+ +OH− → K+ + Cl− +H2O

Step 2: Use the formula and fill in all of the given information.

Ma × Va = Mb × Vb

Ma = 0.125 mol/LVa = 10.0 mLMb = 0.100 mol/LVb =?Ma × Va = Mb × Vb

Vb = Ma×VaMb

= (0.125 mol/L)(10.0 mL0.100 mol/L = 12.5 mL

Therefore, for this strong acid-strong base titration, the volume of base required is 12.5 mL.

This video shows the technique for performing a titration using an indicator: http://www.youtube.com/watch?v=9DkB82xLvNE (5:03).

Lesson Summary• An indicator is a substance that changes color at a specific pH and is used to indicate the pH of thesolution relative to that point.

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Figure 21.3: (Watch Youtube Video)http://www.ck12.org/flexbook/embed/view/414

• A natural indicator is an indicator that is a naturally occurring substance.• Indicators are normally weak organic acids or bases with complicated structures.• Universal indicator is a mixture of indicators that produces a different color for each pH from 0 – 14.• A titration is the addition of a known concentration of base (or acid) to a solution of acid (or base)of unknown concentration.

• The titrant is the solution of known concentration. This solution is normally in the burette.• The endpoint is the point in the titration where the indicator changes color.• The equivalence point is the point in the titration where the number of moles of acid equals thenumber of moles of base.

• The three types of titrations usually performed in the laboratory are: strong acid vs. strong base,strong acid vs. weak base, and weak acid vs. strong base.

• A titration curve is a graph of the pH versus the volume of titrant added.• For a strong acid vs. strong base titration, the pH at equivalence is 7.0. For a strong acid vs. weakbase titration, the pH at equivalence is less than 7.0. For a weak acid vs. strong base titration, thepH at equivalence is greater than 7.0.

• A standard solution is a solution whose concentration is known exactly and is used to find the exactconcentration of the titrant.

• For titrations where the stoichiometric ratio of mol H+ to mol OH− is 1:1, the concentrations orvolumes for the unknown acid or base can be calculated with the formula Ma × Va = Mb × Vb.

Further Reading / Supplemental LinksThe following link is to a video about acid-base neutralization and titration.

• http://link.brightcove.com/services/player/bcpid9113583001?bctid=1405713919

The video at the link below shows the lab techniques needed for titration.

• http://chem-ilp.net/labTechniques/TitrationVideo.htm

This video is a ChemStudy film called “Acid Base Indicators.” The film is somewhat dated but the infor-mation is accurate.

• http://www.youtube.com/watch?v=yi8QrjmV6Sw

Review Questions1. Why do you think there would be more experimental error when using an indicator instead of a pHmeter during a titration?

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2. Which of the following definitions best suits that of an endpoint?

(a) The stoichiometric point where the number of moles of acid equals the number of moles of base.(b) The visual stoichiometric point where the number of moles of acid equals the number of moles

of base.(c) The midpoint of the vertical stretch on the titration curve.(d) None of the above

3. In the following titration curve, what pair of aqueous solutions would best represent what is shownto be happening in the curve?

(a) HCOOH(aq) +NH3(aq)(b) HCOOH(aq) +NaOH(aq)(c) H2SO4(aq) + Ba(OH)2(aq)(d) HClO4(aq) +NH3(aq)

4. What would be the best indicator to choose for the pH curve shown in question 3?

(a) Methyl red(b) Litmus(c) Phenolphthalein(d) Phenol red

5. What is the best indicator to use in the titration of benzoic acid with barium hydroxide?

(a) Methyl violet, range = 0.0 − 1.6(b) Bromothymol blue, range = 3.0 − 4.7(c) Phenolphthalein, range = 8.2 − 10.0(d) Methyl blue, range = 10.6 − 13.4(e) Indigo carmine, range = 11.4 − 13.0

6. If 22.50 mL of a sodium hydroxide solution is necessary to neutralize 18.50 mL of a 0.1430 mol/LHNO3 solution, what is the concentration of NaOH?

(a) 0.1176 mol/L(b) 0.1430 mol/L(c) 0.1740 mol/L(d) 2.64 mol/L

7. Calculate the concentration of hypochlorous acid if 25.00 mL of HClO is neutralized by 32.34 mL ofa 0.1320 mol/L solution of sodium hydroxide.

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21.3 BuffersLesson ObjectivesThe student will:

• define and give an example of a buffer.• explain the effect of a strong acid or base on a buffer system.• explain the mechanism by which a buffer solution resists changes in pH.• given appropriate information, calculate the pH of a buffer.• describe how to make a buffer solution.

Vocabulary• buffer

IntroductionThere are many situations in which it is desirable to keep the pH of a solution close to a particular valueeven though quantities of acids and/or bases are added to the solution. Many organic and biochemicalreactions require acids or bases, but if the pH goes too high or too low, the products will be destroyed.For these reactions, it is necessary to keep the pH within a very small range even while acids or basesare added to the reaction. Chemists use mixtures called buffers to keep the reaction solutions within thenecessary pH range. Buffers are mixtures of chemicals that cause a solution to resist changes in pH.Buffers are very important to many biological reactions. Human blood is a substance whose function isvery dependent on the function of buffers. Human blood must maintain a nearly constant pH between 7.3and 7.5. A change of just 0.2-0.4 pH units outside this range could cause loss of consciousness or evendeath. The pH of human blood can change slightly depending on foods we eat and the rate at which weinhale and exhale CO2. Fortunately, the human blood stream has buffers that are able to resist these pHchanges.

BuffersA buffer is a solution that maintains the pH level when small amounts of acid or base are added to thesystem. Buffer solutions contain either a weak acid and its conjugate base or a weak base and its conjugateacid. A common procedure for producing a buffer in the lab is to make a solution containing both a weakacid and a salt of its conjugate base.For example, you could make a buffer by preparing a solution containing acetic acid and sodium acetate.This common buffer would contain acetic acid, CH3COOH, and the acetate ion, CH3COO−. The solutionwill remain in the pH range of 3.7 – 5.8 even if small amounts of acids or bases are added. Another exampleis a solution of hydrogen phosphate, HPO2−

4 , plus the phosphate, PO3−4 . This will buffer a solution in the

pH range of 11.3 - 13.3.How is it possible that a solution will not change its pH when an acid or base is added? Let’s examine theacetic acid/acetate ion buffer. The ionization equation for acetic acid is shown below.

HC2H3O2(aq) → H+ + C2H3O−2

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If acid (H+ ions) is added to this solution, the equilibrium will shift toward the reactants to use up some ofthe added hydrogen ions. Equilibrium will be re-established in the solution with different concentrationsof the three species in the reaction. Although the H+ concentration will be higher than its initial value,the change will be partially corrected due to the presence of C2H3O−2 ions. Similarly, if a base is addedto this solution, the base will remove hydrogen ions and the equilibrium will shift to the right to partiallycounteract the stress. Again, equilibrium will be re-established with new concentrations. The existence oflarge quantities of both undissociated acid molecules and acetate ions in the solution is what allows thebuffer to consume quite a large amount of added acid or base without the pH changing significantly.Examine what happens to 1.00 liter of pure water to which 0.100 mole of gaseous HCl is added. Theoriginal concentration of hydrogen ion in the pure water is 1.00/times10−7 M and therefore, the pH is 7.After the 0.100 mole of HCl is added, the concentration of hydrogen ions will be 0.100 M (plus the original1.00 × 10−7 M, which can be neglected as insignificant). This new concentration of hydrogen ions willproduce a pH = 1. So, the addition of the 0.100 mole of gaseous HCl caused the pH of the pure water tochange from 7 to 1.Let’s now see what happens if this same amount of gaseous HCl is added to an acetic acid-acetate ionbuffer. Let’s say we made this solution to contain 0.50 M acetic acid and 0.50 M acetate ion (0.50 Msodium acetate which fully dissociated). The acetic acid dissociation equation will reach equilibrium at itsKa is 1.8 × 10−5.

Ka = [H+][C2H3O−2]

[HC2H3O2]= 1.8 × 10−5

We know both the acetic acid and the acetate ion concentrations will be 0.50 M, so we can plug thesevalues into the expression and solve for [H+].

[H+] = (1.8×10−5)[HC2H3O2]

[C2H3O−2]= (1.8×10−5)(0.50)

(0.50)= 1.8 × 10−5 M

Then we can insert the hydrogen ion concentration into the pH formula and determine the original pH ofthe buffer solution.

pH = − log(1.8 × 10−5) = 4.74

Next, we will add the same 0.100 mole of gaseous HCl to this buffer solution and calculate the pH of thesolution after the acid has been added and equilibrium has been re-established.When we add 0.100 mole HCl gas to this solution, the added hydrogen ion will combine with acetate ionto produce more undissociated acid. A small amount of the newly formed acetic acid may dissociate again,but the amount is minimal and can be neglected. The new [HC2H3O2] will equal 0.60 M (the original0.50 M plus the added 0.10 M) and the new [C2H3O−2 ] will equal 0.40 M (the original 0.50 M minus the0.10 M that reacted with the added hydrogen ions). We can now plug these values into the Ka expression,calculate the new [H+], and find the new pH.

Ka = [H+][C2H3O−2]

[HC2H3O2]= 1.8 × 10−5

[H+] = (1.8×10−5)[HC2H3O2]

[C2H3O−2]= (1.8×10−5)(0.60)

(0.40)= 2.7 × 10−5 M

pH = − log(2.7 × 10−5) = 4.57

The same quantity of HCl gas that changed the pH of pure water from 7 to 1 has changed the pH of thisbuffer from 4.74 to 4.57. The change in pH is only 0.17, which is a function of the buffer. Buffers resist

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change to pH. We could do this same calculation by adding a base instead of an acid and show that the pHincreases by this same slight amount. We will leave this as an exercise for you to complete on your own.Example:Which of the following combinations would you expect to make a useful buffer solution:

1. HClO4/ClO−42. CH3NH2/CH3NH+

3

Solution:

1. HClO4/ClO−4 :HClO4 is a strong acid and buffers are made from weak acids and their conjugate basesor weak bases and their conjugate acids. Therefore, this cannot be made into a buffer.

2. CH3NH2/CH3NH+3 :CH3NH2 is a weak base and CH3NH+

3 is the conjugate acid of this base. There-fore, this can be made into a buffer solution.

The first video considers the mathematics of buffer solutions and the second video shows a laboratoryexample of buffering (5g; 1l I&E Stand.): http://www.youtube.com/watch?v=O_QlZe4fv4g (10:18),(5g) http://www.youtube.com/watch?v=g_ZK2ABUjvA (2:39).

Figure 21.4: (Watch Youtube Video)http://www.ck12.org/flexbook/embed/view/415

Figure 21.5: (Watch Youtube Video)http://www.ck12.org/flexbook/embed/view/416

The Buffer in BloodThe primary buffer found in your bloodstream is carbonic acid, H2CO3. The carbonic acid is present dueto carbon dioxide from your respiratory system dissolving in water.

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CO2(g) +H2O(l) H2CO3(aq)

The amount of carbonic acid in your bloodstream is affected by the rate of your respiration. If you breatherapidly, you reduce the amount of CO2 in your bloodstream, and the equilibrium shown above shiftstoward the reactants, thus lowering the amount of H2CO3. If you breathe slowly, the amount of CO2

in your bloodstream increases, and the equilibrium shifts toward the products, increasing the amount ofH2CO3 in your system.Once the H2CO3 is produced by dissolving carbon dioxide, the carbonic acid dissociates in your blood asshown below:

H2CO3(aq) H+ +HCO−3

The buffering system in your blood is composed of the weak acid, H2CO3, and its conjugate base, HCO−3 .Any changes in blood pH that could be caused by food intake will be buffered by this equilibrium system.If acid is added to your blood, the equilibrium will shift toward the reactants, using up the hydrogen ions.If base were added to your blood, thus reacting with hydrogen ions, the equilibrium will shift toward theproducts to generate more hydrogen ions. This buffer in your blood is very efficient at keeping your bloodpH in the necessary range.Some people, when they get nervous, begin breathing very fast or very slow. Breathing too fast is calledhyperventilating, and breathing too slow is called hypoventilating. Your respiratory rate is normallycontrolled by the amount of carbon dioxide in the blood. Your body receives instructions to breathe fasteror slower to adjust the amount of carbon dioxide in your blood in order to properly regulate the buffersystem. When people breathe too fast or too slow because of other reasons, the CO2 content of the bloodbecomes incorrect and the pH of the blood rises or lowers outside the acceptable range of 7.3 - 7.5. Whenthis happens, the person can pass out. People who hyperventilate when excited or nervous are sometimesadvised to carry a lunch sack or something similar to breathe into when they are feeling light-headed.Breathing into a sack returns air with the same concentration of carbon dioxide that was exhaled. Thiskeeps the amount of carbon dioxide in the blood up and prevents a loss of consciousness.

Lesson Summary• A buffer is a solution of a weak acid and its conjugate base or a weak base and its conjugate acidthat resists changes in pH when an acid or base is added to it.

• Adding a strong acid to a buffer only decreases the pH by a small amount.• Adding a strong base to a buffer only increases the pH by a small amount.

Further Reading / Supplemental LinksThe following link is to a video lecture about acid-base buffers.

• http://www.youtube.com/watch?v=O_QlZe4fv4g

To see a short animated video showing concentration changes as strong acid or base is added to a buffer,follow the link below.

• http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf

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Review Questions1. Define a buffer solution.2. One of the following statements of buffers is incorrect. Which one?

(a) A buffer may be prepared from a weak acid and its conjugate base salt.(b) A buffer may be prepared from a weak base and its conjugate acid salt.(c) A buffer is a solution that can resist changes in pH when any amount of acid or base is added

to it.(d) A buffer is a solution that can resist changes in pH when a small amount of acid or base is added

to it.3. Which of the following pairs of aqueous 1.0 mol/L solutions could be chosen to prepare a buffer: i.NH4HSO4(aq) and H2SO4(aq), ii. HNO2(aq) and NaNO2(aq), iii. NH4Cl(aq) and NH3(aq)?(a) i and iii only(b) ii and iii only(c) i, ii, and iii(d) None of these solutions is a buffer.

4. Which of the following would form a buffer solution if combined in appropriate amounts?(a) HCl and NaCl(b) HCN and NaCN(c) H2S and NaSH(d) HNO3 and NaNO3

5. A buffer is made up of a weak acid and a conjugate base. A small amount of acid is added to thebuffer. What happens to the resulting solution?(a) The acid dissociation constant goes up.(b) The concentration of the weak acid in the buffer goes down.(c) The pH of the solution goes up.(d) The pH remains almost the same.

Image Sources(1) David Kauffman. Hydrangeas. Public domain.

All images, unless otherwise stated, are created by the CK-12 Foundation and are under the CreativeCommons license CC-BY-NC-SA.

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