topic 18 topic 18 topic 18: acids and bases basic concepts additional concepts table of contents...

Topic 18Topic 18

Topic 18: Acids and Bases

Basic Concepts

Additional Concepts

Table of ContentsTable of ContentsTopic 18Topic 18

• Although taste is not a safe way to classify acids and bases, you probably are familiar with the sour taste of acids.

Properties of Acids and Bases

Acids and Bases: Basic ConceptsAcids and Bases: Basic ConceptsAcids and Bases: Basic ConceptsAcids and Bases: Basic Concepts

• Lemon juice and vinegar, for example, are both aqueous solutions of acids.

• Bases, on the other hand, taste bitter.

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Properties of Acids and Bases

Acids and Bases: Basic ConceptsAcids and Bases: Basic ConceptsAcids and Bases: Basic ConceptsAcids and Bases: Basic ConceptsTopic 18Topic 18

• Bases have a slippery feel. Like taste, feel is not a safe chemical test for bases, but you are familiar with the feel of soap, a base, on the skin.

• Acids and bases cause certain colored dyes to change color. The most common of these dyes is litmus.

Litmus Test and Other Color Changes


• When mixed with an acid, litmus is red.

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Litmus Test and Other Color Changes


• When added to a base, litmus is blue.

• Therefore, litmus is a reliable indicator ofwhether a substance is an acid or a base.

Reactions with Metals and Carbonates


• This property explains why acids corrode most metals.

Topic 18Topic 18

• Another characteristic property of an acid isthat it reacts with metals that are more active than hydrogen.



• Bases do not commonly react with metals.

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• Another simple test that distinguishes acids from bases is the reaction of acids with ionic compounds that contain the carbonate ion, , to form carbon dioxide gas, water, and another compound.

• The submicroscopic behavior of acids when they dissolve in water can be described in several ways.

Submicroscopic Behavior of Acids


• The simplest definition is that an acid is a substance that produces hydronium ions when it dissolves in water.

• A hydronium ion, H3O+, consists of a hydrogen ion attached to a water molecule.

• When HCl dissolves in water, it produces hydronium ions by the reaction shown below.

Submicroscopic Behavior of Acids


• HCl is definitely an acid; it produces H3O+ when dissolved in water.

• At the submicroscopic level, the reaction of an acid with water is a transfer of a hydrogen ion, H+, from an acid to a water molecule.

Acidic Hydrogen Atoms


• This transfer forms the positively charged hydronium ion, H3O+, and a negatively charged ion.

• In an acid, any hydrogen atom that can be transferred to water is called an acidic hydrogen.

• To help distinguish acids from other hydrogen-containing molecules, acidic hydrogens are written first in the formula.

Acidic Hydrogen Atoms


• Any time hydrogen is the first element in a formula of a compound, the substance is an acid.

• Acids such as acetic acid, HC2H3O2, and hydrochloric acid, HCl, are called monoprotic acids.

Monoprotic Acids


• Monoprotic acids contain only one acidic hydrogen.

• All acids that have more than one acidic hydrogen per molecule are called polyprotic acids.

Diprotic and Triprotic Acids


• Polyprotic acids with two acidic hydrogens are diprotic acids.

Diprotic and Triprotic Acids


• Those with three acidic hydrogens are triprotic acids.

• The behavior of bases is also described at the molecular level by the interaction of the base with water.

Submicroscopic Behavior of Bases


• A base is a substance that produces hydroxide ions, OH –, when it dissolves in water.

• There are two mechanisms by which bases produce hydroxide ions when they dissolve in water.

• The simplest kind of base is a water-soluble ionic compound, such as sodium hydroxide, that contains the hydroxide ion as the negative ion.

Simple Bases: Metal Hydroxides


• When NaOH dissolves in water, for example, it dissociates into aqueous sodium ions and hydroxide ions, as shown below.

• Two related classes of compounds do not fit the previous models of acids and bases, but they still act as acids or bases.

Other Acids and Bases: Anhydrides


• These compounds are both oxides, which are compounds containing oxygen bonded to just one other element.

• These oxides are called anhydrides, which means that they contain no water.

• Anhydrides differ, depending upon whether the oxygen is bonded to a metal or a nonmetal.



• Nonmetal oxides form acids when they react with water and are called acidic anhydrides.

• Metal oxides, on the other hand, react with water to form bases and are called basic anhydrides.



• In both of these reactions, water is an active reactant.


• These metal oxides are predominantly those of sodium, potassium, and calcium.

Basic Anhydrides


• These metal oxides are ionic so they are solids, even at the high temperature of a roaring fire.

• They are the major component of the ash that is left when the fire burns out.

• Sodium hydroxide, NaOH, is a strong base because when NaOH dissolves in water, all NaOH formula units dissociate into separate sodium and hydroxide ions.

Strong Bases


• The dissociation of the base is complete.

• The strength of a base is based on the percent of units dissociated, not the number of OH– ions produced. Some bases, such as Mg(OH)2, are not very soluble in water, and they don’t produce a large number of OH– ions.

Strong Bases


• However, they are still considered to be strong bases because all of the base that does dissolve completely dissociates.

• HCl is a strong acid because no HCl molecules are in a water solution of HCl.

Strong Acids


• Because of the strong attraction between the water molecules and HCl molecules, every HCl molecule ionizes.

• Acetic acid, HC2H3O2, is a good example of a weak acid. The molecular structure of a weak acid determines the extent to which the acid ionizes in water. A solution of weak acid contains a mixture of un-ionized acid molecules, hydronium ions, and the corresponding negative ions.

Weak Acids


Weak Acids


• The concentration of the un-ionized acid is always the greatest of the three concentrations.

• Ammonia is a weak base because most of its molecules don’t react with water to form ions.

Weak Bases


• Other examples of bases that produce so few OH– ions that they are considered to be weak bases are Al(OH)3, and Fe(OH)3.

• Although the terms weak and strong are used to compare the strengths of acids and bases, dilute and concentrated are terms used to describe the concentration of solutions.

Strength Is Not Concentration


• The combination of strength and concentration ultimately determines the behavior of the solution.

• For example, it is possible to have a concentrated solution of a weak acid or weak base or a dilute solution of a weak acid or weak base.

Strength Is Not Concentration


• Similarly, you can have a concentrated solution of a strong acid or strong base, as well as a dilute solution of a strong acid or strong base.

• In most applications, the observed range of possible hydronium or hydroxide ion concentrations spans 10–14 M to 1M.

The pH Scale


• This huge range of concentrations presents a problem when comparing different acids and bases.

• To make this range of possible concentrations easier to work with, the pH scale was developed by S.P.L. Sørenson.

• pH is a mathematical scale in which the concentration of hydronium ions in a solution is expressed as a number from 0 to 14.

The pH Scale


• A scale of 0 to 14 is much easier to work with than a range from 1 to 10–14 (100 to 10–14).

• The pH scale is a convenient way to describe the concentration of hydronium ions in acidic solutions, as well as the hydroxide ions in basic solutions.

• Think about the pH numbers 0 to 14 and the hydronium ion concentration range.

The pH Scale


• Notice that the pH value is the negative of the exponent of the hydronium ion concentration.

• For example, a solution with a hydronium ion concentration of 10–11M has a pH of 11.

• A solution with a pH of 4 has a hydronium ion concentration of 10–4M.

• pH is convenient because there are simple methods for measuring it in the lab or in the field.

Measuring pH


• Indicators register different colors at different pHs.

• pH meters are instruments that measure the exact pH of a solution.

• pH of 7 is neutral. A pH less than 7 is acidic, and a pH greater than 7 is basic.

Interpreting the pH Scale


• As the pH drops from 7, the solution becomes more acidic.

• As pH increases from 7, the solution becomes more basic.

• In a neutral solution, the concentration of hydroxide ions and the concentration of hydronium ions are equal.

pH of Common Materials


• The reaction of an acid and a base is called a neutralization reaction.

Types of Acid-Base Reactions





• Consider the following neutralization reaction.



• Hydrochloric acid, HCl, is a common household and laboratory acid.

• Muriatic acid is the common household name of hydrochloric acid.

• A solution of hydrochloric acid, HCl, is added to exactly the amount of a solution of basic sodium hydroxide, NaOH, that will react with it.

Neutralization Reactions



Neutralization Reactions


• Litmus papers show that the resulting salt solution is neither acidic nor basic.


• A typical type of acid-base reaction is one in which both the acid and base are strong.

Strong Acid + Strong Base


• The reaction of aqueous solutions of hydrochloric acid and sodium hydroxide is a good example of this type of reaction.

• HCl is the acid; NaOH is the base. The products are NaCl, which is a salt, and water.

• Now take a closer look at these reactants and their products.

• You also know that sodium hydroxide in water completely dissociates into sodium ions and hydroxide ions because NaOH is a strong base.



• An overall equation for the reaction between NaOH and HCl shows each substance involved in the reaction.



• An overall equation does not indicate whether these substances exist as ions.

• The best way for you to model the submicroscopic behavior of an acid-base reaction is to show reactants and products as they actually exist in solution.

• Instead of an overall equation, an ionic equation, in which substances that primarily exist as ions in solution are shown as ions, can be written.



• How can a lithium battery have an aqueous electrolyte? Two facets of the construction of this new battery keep the lithium metal from reacting with water.



• First, the lithium is in the form of individual atoms embedded in a material such as manganese (IV) oxide, rather than as a solid metal.

• Second, the electrolyte is full of dissolved lithium salts, so the lithium ions that are produced travel to the site of reduction without reacting with water.



• Note that the ionic equation gives more information about how a strong acid-strong base reaction occurs.

Spectator Ions and the Net Ionic Reaction


• When you examine the two sides of the ionic equation, you see that Na+ and Cl– are present both as reactants and as products.

• Although they are important components of an overall equation, they do not directly participate in the chemical reaction.



• They are called spectator ions because they are present in the solution but do not participate in the reaction.

• Just as in a mathematical equation, items common to both sides of the equation can be subtracted.



• This process simplifies the equation so that the reactants and products that actually change can be seen more clearly.

• When ions common to both sides of the equation are removed from the equation, the result is called the net ionic equation for the reaction of HCl with NaOH.



• The net ionic equation describes what is really happening at the submicroscopic level.



• Although solutions of HCl and NaOH are mixed, the net ionic equation is hydrogen ions reacting with hydroxide ions to form water.

Strong Acid + Weak Base


• Consider the reaction of hydrobromic acid and aluminum hydroxide.

• The overall equation shows the reactants and products.



• Hydrobromic acid, a strong acid, completely ionizes in water.

• All of the Al(OH)3 that dissolves dissociates, so it is technically a strong base.

• However, because it is so insoluble, few OH– ions are produced, and Al(OH)3 acts as a weak base.



• Therefore, the ionic equation shows little dissociation of the base.

• The dissociated salt, AlBr3, is also shown as ions.



• The spectator ions in this equation are bromide ions.

• They are removed from both sides of the equation to produce the net ionic equation.



• A solution of ammonia is best represented by NH3(aq).

• A solution of HCl, as in the last case, is best represented as H+ (aq) and Cl– (aq).

• The ionic reaction is written by representing what is actually in the reactant and product solutions.

• The ionic salt NH4Cl is written as dissociated ions.



• A quick look at the ionic reaction shows that the chloride ion is a spectator ion because it appears on both sides of the reaction.

• You get the net ionic equation by subtracting the spectator Cl–

from both sides of the ionic equation.

A Broader Definition of Acids and Bases


• The reaction of a strong acid with a weak base demonstrates the need for a slightly broader definition of acids and bases.



• As you learned in the last chapter, much of the behavior of acids and bases in water can be explained by a model that focuses on the hydrogen ion transfer from the acid to the base.

• This model will also help explain why every acid-base reaction does not result in a neutral solution.



• In the more inclusive Brønsted-Lowry model, an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor.

• When a Brønsted-Lowry acid donates a hydrogen ion, a conjugate base is formed.

• When a Brønsted-Lowry base accepts a hydrogen ion, a conjugate acid is formed.

• Two substances related to each other by the donating and accepting of a single hydrogen ion are a conjugate acid-base pair.

Identifying Conjugate Acid-Base Pairs


• Identify the conjugate acid-base pairs in this reaction.

• A hydrogen ion is donated by HClO2, which is the Brønsted-Lowry acid in the forward reaction.

• The resulting conjugate base is ClO2–.

• The base in the forward reaction is H2O, which accepts a hydrogen ion to form the conjugate acid H3O+.

Identifying Conjugate Acid-Base Pairs


• As in the case of the weak base-strong acid reaction, mixing equal moles of acid and base does not produce a neutral solution.

• Because the final pH in this case is basic, more hydroxide than hydronium ions must be present in the final reaction mixture.

Weak Acid + Strong Base— Equations—Weak Acid, Strong Base


• Write the overall, ionic, and net ionic equations for the reaction of phosphoric acid with lithium hydroxide.

• Phosphoric acid, H3PO4, is a weak acid. Lithium hydroxide, LiOH, is a strong base.



• Write a balanced equation for the overall reaction.

• The salt is lithium phosphate, Li3PO4. • Water is also a product. • You’ll need three moles of LiOH for each

mole of H3PO4. Three moles of water will be produced.



• Now, write the ionic equation. • Because H3PO4 is a weak acid, it is only

partially ionized, and you write it in the ionic equation as H3PO4.

• LiOH is completely dissociated, as is the salt lithium phosphate.

• Be careful to keep the ionic equation balanced.



• Check for spectator ions. Li+ is a spectator ion.

• Subtract Li+ from both sides of the equation to get the net ionic equation.

Weak Acid + Weak Base


• The strong-strong reaction plus the two types of weak-strong reactions are the favorable acid-base reactions.

• Looking at an acid-base reaction as occurring by H+ transfer helps you to understand why the weak-weak reaction is not considered a favorable reaction.

Weak Acid + Weak Base


• Because neither a weak acid nor a weak base has a strong tendency to transfer a hydrogen ion, transfer between the two may occur, but it is uncommon.

• Reactions between a weak acid and a weak base generally do not play an important role in acid-base chemistry.

Basic Assessment QuestionsBasic Assessment Questions

Question 1

Find the pH of the following solution.

Topic 18Topic 18

The hydronium ion concentration equals: 10–2M.


Answer

pH = 2

Topic 18Topic 18


Question 2

Find the pH of the following solution.

Topic 18Topic 18

The hydroxide ion concentration equals: 10–8M.


Answer

pH = 6

Topic 18Topic 18


Question 3

Write the net ionic equation for sulfuric acid, H2SO4, and strontium hydroxide, Sr(OH)2.

Topic 18Topic 18


AnswerTopic 18Topic 18


Question 4

Identify the conjugate acid-base pairs in the following reaction.

Topic 18Topic 18

Acids and Bases: Additional ConceptsAcids and Bases: Additional Concepts Acids and Bases: Additional ConceptsAcids and Bases: Additional Concepts Topic 18Topic 18

Additional Concepts

Acids and Bases: Additional ConceptsAcids and Bases: Additional Concepts Acids and Bases: Additional ConceptsAcids and Bases: Additional Concepts

Buffers Defined

• A buffer is a solution that resists changes in pH when moderate amounts of acids or bases are added.

Topic 18Topic 18



Buffers Defined

• It contains ions or molecules that react with OH– or H+ if one of these ions is introduced into the solution.

Topic 18Topic 18

• Buffer solutions are prepared by using a weak acid with one of its salts or a weak base with one of its salts


Buffers Defined

• For example, a buffer solution can be prepared by using the weak base ammonia, NH3, and an ammonium salt, such as NH4Cl.

Topic 18Topic 18

• If an acid is added, NH3 reacts with the H+.


Buffers Defined

• If a base is added, the NH4+ ion from the salt

reacts with the OH–.

Topic 18Topic 18


Buffers Defined

• Look at another system that contains the weak acid acetic acid, HC2H3O2, and the salt sodium acetate, NaC2H3O2.

Topic 18Topic 18

• If a strong base, OH–, is added to the buffer system, the weak acid reacts to neutralize the addition.


Buffers Defined

• This reaction takes care of the added OH–.

Topic 18Topic 18

• If H+ is added, the acetate ion from the NaC2H3O2 is available to neutralize the added H+.


Buffers Defined

• These pH changes are insignificant when you compare them to the changes that occur in the unbuffered solution.

Topic 18Topic 18

• These two buffer systems are common ones used in many laboratories.

• Blood has many buffer systems to maintain its constant pH of 7.4.


Acid-Base Titrations

• The general process of determining the molarity of an acid or a base through the use of an acid-base reaction is called an acid-base titration.

Topic 18Topic 18



Acid-Base TitrationsTopic 18Topic 18

• The known reactant molarity is used to find the unknown molarity of the other solution.

• Solutions of known molarity that are used in this fashion are called standard solutions.

• In a titration, the molarity of one of the reactants, acid or base, is known, but the other is unknown.



• You know that NaOH and HCl react completely.

• You know the concentration of the NaOH solution, so it is your standard solution.



• You can use the reaction, the volumes of acid and base used, plus the molarity of the base to determine the molarity of the unlabeled HCl.


Determining Concentration: Using Stoichiometry

Topic 18Topic 18

• From the balanced equation for the reaction, you know that one mole of HCl reacts with one mole of NaOH.

• Therefore, the number of moles of HCl in 20.0 mL of the HCl solution equals the number of moles of NaOH in 19.9 mL of 0.100M NaOH solution.



Topic 18Topic 18

• Now, use the factor label method to solve this solution stoichiometry problem, just as you used it to solve other stoichiometry problems. • Because you know the concentration of the NaOH solution, first find the number of moles of NaOH involved in the reaction.



Topic 18Topic 18

• Next, examine the balanced equation for the reaction and determine that, because their coefficients are the same, equal numbers of moles of NaOH and HCl react.



Topic 18Topic 18

• Because 1.99 x 10–3mol NaOH react, 1.99 x 10–3 mol HCl present in solution also react.

• Finally, use the volume to find the molarity of the acid.



Topic 18Topic 18

• Based on your single titration, the molarity of the HCl solution is 0.0995M.

• However, before you put this value on the label, you probably would repeat the titration for several additional trials in order to verify your analysis and be more confident of the value on the label.


Finding MolarityTopic 18Topic 18

• A 15.0-mL sample of a solution of H2SO4 with an unknown molarity is titrated with 32.4 mL of 0.145M NaOH to the bromothymol blue endpoint.

• Based upon this titration, what is the molarity of the sulfuric acid solution?



• Because the molarity of the NaOH solution is known, the number of moles of NaOH involved in the titration can be calculated.

• The corresponding number of moles of H2SO4 can then be determined, and this figure can be used to calculate the molarity of the acid.



• Write the balanced equation for the reaction. Remember that sulfuric acid is a diprotic acid.

• Because the concentration of the NaOH solution is known, find the number of moles of NaOH used in the titration.



• Using the balanced equation, find the number of moles of H2SO4 that react with 4.70 x 10–3 mol NaOH.

• Find the concentration of the H2SO4.


pH and pOHTopic 18Topic 18

• Because the concentrations of H+ ions are often very small numbers, the pH scale was developed as a more convenient way to express H+ ion concentrations.

• The pH of a solution equals the negative logarithm of the hydrogen ion concentration.



• The pH scale has values from 0 to 14.

• Acidic solutions have pH values between 0 and 7, with a value of 0 being the most acidic.

• The pH of a basic solution is between 7 and 14 with 14 representing the most basic solution.

• A neutral solution has a pH of 7.



• Chemists have also defined a pOH scale to express the basicity of a solution.

• The pOH of a solution is the negative logarithm of the hydroxide ion concentration.



• If either pH or pOH is known, the other may be determined by using the following relationship.

• The pH and pOH values for a solution may be determined if either [H+] or [OH–] is known.

• The following example problem shows you how to calculate pH and pOH.


Calculating pH and pOH from [H+]Topic 18Topic 18

• If a certain carbonated soft drink has a hydrogen ion concentration of 7.3 x 10–4M, what are the pH and pOH of the soft drink?

• Because [H+] is given, it is easier to calculate pH first.



• A log table or calculator shows that log 7.3 = 0.86 and log 10–4 = – 4.

• Substitute these values in the equation for pH.

• The pH of the soft drink is 3.14.



• To find pOH, recall that pH + pOH = 14.00.

• Isolate pOH by subtracting pH from both sides of the equation.



• Substitute the value of pH and solve.

• The pOH of the solution is 10.86.

• As you might expect, the carbonated soft drink is acidic.


Calculating ion concentrations from pHTopic 18Topic 18

• When the pH of a solution is known, you can determine the concentrations of H+ and OH–.

• First, recall the equation for pH.

• Multiply both sides of the equation by –1.


Calculating ion concentrations from pHTopic 18Topic 18

• Now take the antilog of both sides of the equation.

• Rearrange the equation.

• A similar relationship exists between [OH–] and pOH.


Calculating [H+] and [OH–] from pHTopic 18Topic 18

• What are [H+] and [OH–] in an antacid solution with a pH of 9.70?

• Use pH to find [H+].

• Use a log table or calculator to find that the antilog of – 9.70 is 2.0 x 10–10.



• To determine [OH–], first use the pH value to calculate pOH.

• Now use the equation relating [OH–] to pOH.



• A log table or calculator shows that the antilog of – 4.30 is 5.0 x 10–5.

• As expected, [OH–] > [H+] in this basic solution.

Additional Assessment QuestionsAdditional Assessment Questions

A 0.100M LiOH solution was used to titrate an HBr solution of unknown concentration. At the endpoint, 21.0 mL of LiOH solution had neutralized 10.0 mL of HBr. What is the molarity of the HBr solution?

Question 1 Topic 18Topic 18


0.210M HBr



Calculate the pH and pOH of aqueous solutions having the following ion concentrations.


pH: 14.00;pOH: 0.00

Answer 2a

Question 2a

Additional Assessment QuestionsAdditional Assessment QuestionsTopic 18Topic 18

pH: 6.75; pOH: 7.25

Answer 2b

Question 2b



The pH is given for two solutions. Calculate [H+] and [OH–] in each solution.


pH = 2.80

Answer 3a

Question 3a


pH = 13.19

Answer 3b

Question 3b


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topic 18 topic 18 topic 18: acids and bases basic concepts additional concepts table of contents...

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