Acids, Bases, And PH
Acids, Bases, and pHChapter Preview8.1 Explaining theProperties of Acidsand Bases8.2 The Equilibrium ofWeak Acids and Bases8.3 Bases and Buffers8.4 Acid-Base TitrationCurvesPrerequisiteConcepts and SkillsFor many people, the word “acid” evokes the image of a fuming, highlycorrosive, dangerous liquid. This image is fairly accurate for concentratedhydrochloric acid, a strong acid. Most acids, however, are not as corrosiveas hydrochloric acid, although they may still be very hazardous. Forexample, hydrofluoric acid can cause deep, slow-healing tissue burns ifit is handled carelessly. It is used by artists and artisans who etch glass. Itreacts with the silica in glass to form a compound that dissolves, leavingthe glass with a brilliant surface. Hydrofluoric acid is highly corrosive.Even a 1% solution is considered to be hazardous. Yet chemists classifyhydrofluoric acid as a weak acid.You learned about acids and bases in your previous chemistry course.In this chapter, you will extend your knowledge to learn how the structure of a compound determines whether it is an acid or a base. You willuse the equilibrium constant of the reaction of an acid or base with waterto determine whether the acid or base is strong or weak. You will applyyour understanding of dissociation and pH to investigate buffer solutions:solutions that resist changes in pH. Finally, you will examine acid-basetitrations that involve combinations of strong and weak acids and bases.Before you begin this chapter,review the following conceptsand skills: writing net ionic equations(Concepts and SkillsReview) calculating molar concentrations (Concepts andSkills Review) solving equilibriumproblems (Chapter 7,sections 7.3 and 7.4) explaining themathematical propertiesof logarithms (previousstudies) performing acid-basetitrations (previousstudies)376 MHR Unit 4 Chemical Systems and EquilibriumThe hydrohalic acids (HX(aq) ), where X represents ahalogen) include HF, HCl, HBr, and HI. Only HF is aweak acid. The rest are strong acids. What factorsaccount for this difference?
Explaining the Propertiesof Acids and Bases8.1Table 8.1 outlines properties of acids and bases that you have examined inprevious courses. In this section, you will review two theories that help toexplain these and other properties. As well, you will use your understanding of molecular structure to help you understand why acids andbases differ in strength.Section Preview/Specific ExpectationsIn this section, you will compare strong acidsand bases, and weak acidsand bases, in terms ofequilibrium identify conjugate acid-basepairs solve problems that involvestrong acids and strongbases communicate your understanding of the followingterms: hydronium ion(H3O (aq) ), conjugateacid-base pair, monoproticacids, polyprotic acidsTable 8.1 Examples and Common Properties of Acids and BasesExampleAcidsBasessolidacetyl salicylic acidsodium hydroxideliquidacetic acidanilinegashydrogen chlorideammoniaPropertytasteAcidsBasesAcids taste sour.Bases taste bitter.texture of solutionCAUTION Neverdeliberately touchchemicals. Strong,concentrated acidsand bases will burnyour skin.Acids do not have acharacteristic texture.Bases feel slippery.reaction withphenolphthaleinAcidic phenolphthaleinis colourless.Basic phenolphthaleinis pink.reaction with litmuspaperAcids turn blue litmusred.reaction with metalsAcids react withmetals above hydrogenin the activity series todisplace H2(g).Bases react with certainmetals (such as Al) toform H2(g).reaction withcarbonatesCarbon dioxide is formed.No reaction occurs.reaction withammonium chlorideNo reaction occurs.Ammonia, NH3, a gaswith a characteristicodour, is produced.neutralizationreactionAcids neutralize basicsolutions.Bases neutralize acidicsolutions.reaction with fattyacidsNo reaction occurs.Bases react to form soap(a saponification reaction).aqueous propertyof oxidesNon-metal oxides formMetal oxides form basicsolutions: for example,acidic solutions:CaO(s) H2O( ) Ca(OH)2(aq)for example,CO2(g) H2O( ) H2CO3(aq)CAUTION Never tastechemicals in a lab.amount of dissociation Strong acids dissociatein aqueous solutioncompletely. Weak acids(strength)dissociate only partially.ases turn red litmus lue.Strong bases dissociatecompletely. Weak basesdissociate only partially.Chapter 8 Acids, Bases, and pH MHR377
The following ExpressLab highlights concepts that you will examine inthis section, as well as later in the chapter.ExpressLabComparing Acid-Base ReactionsYou will perform three acid-base reactions. Beforeyou begin, read the Procedure and make a prediction about the relative rates of these reactions.Safety PrecautionsThe solutions that are used in this lab are irritantsand should be handled with care. Wash any spillson your skin or clothing with plenty of water. Informyour teacher immediately.Materialspowdered calcium carbonate, CaCO3(s)3 squeeze bottles, each containing one of thefollowing solutions: 2.0 mol/L HCl(aq) ;2.0 mol/L CH3COOH(aq) ; mixture of2.0 mol/L CH3COOH(aq) and 2.0 mol/L NaCH3COO(aq)scoopula3 test tubestest tube racklabels or grease pencilProcedure1. Label each test tube to identify the solutionit will contain. Then fill each test tube with thecorresponding solution, to a depth of about 2 cm.2. Add a small amount of CaCO3(s) (enough to coverthe tip of a scoopula) to each test tube. Try toadd the same amount of CaCO3(s) to all threetest tubes.Analysis1. (a) What rate-related change did you observein each test tube?(b) If you wanted to collect quantitative datafor each reaction, how could you modifythe experiment?2. In all three test tubes, the following reactionoccurred.CaCO3(s) 2H3O (aq) CO2(g) Ca2 (aq) 3H2O( )The only difference between the test tubes wasthe concentration of H (aq) in the acidic solutionsthat reacted with CaCO3(s) . Explain your rankingof the rates of reaction in terms of [H3O ].3. The concentrations of HCl(aq) and CH3COOH(aq)were identical. The following dissociationreactions occurred.HCl(aq) H2O( ) H3O (aq) Cl (aq) H3O (aq) CH3COO (aq)CH3COOH(aq) H2O( ) Explain your ranking of the [H3O ] in thesesolutions in terms of the extent of theequilibrium dissociation.4. The third solution was a mixture of 2.0 mol/LCH3COOH(aq) and 2.0 mol/L NaCH3COO(aq) . Howdid the addition of sodium acetate affect theequilibrium of the dissociation reaction ofacetic acid?5. Explain your ranking of the rate of the reactionbetween calcium carbonate and the solutionthat was a mixture of acetic acid and sodiumacetate.3. Record your observations. Rank the rates of thethree reactions from fastest to slowest.The Arrhenius Theory of Acids and BasesAccording to the Arrhenius theory (1887), acids and bases are definedin terms of their structure and the ions produced when they dissolvein water. An acid is a substance that dissociates in water to form H (aq) . Twoexamples of Arrhenius acids are hydrochloric acid, HCl, and sulfuricacid, H2SO4 . A base is a substance that dissociates in water to form OH (aq) . Twoexamples of Arrhenius bases are sodium hydroxide, NaOH, andpotassium hydroxide, KOH.378 MHR Unit 4 Chemical Systems and Equilibrium
The Arrhenius theory explains acid-base reactions as a combination ofH (aq) and OH (aq) . It provides insight into the heat of neutralization forthe reaction between a strong acid and a strong base. (Strong acids andbases dissociate completely into ions in solution.) For example, considerthe following reaction.HCl(aq) NaOH(aq) NaCl(aq) H2O( )ΔH 56 kJThe total ionic equation for this reaction isH (aq) Cl (aq) Na (aq) OH (aq) Na (aq) Cl (aq) H2O( )ΔH 56 kJSubtracting spectator ions from both sides, the net ionic equation isH (aq) OH (aq) H2O( )ΔH 56 kJDifferent combinations of strong Arrhenius acids and bases react with thesame exothermic result. Measurements always show the release of 56 kJ ofenergy per mole of water formed. This makes sense, because the net ionicequation is the same regardless of the specific neutralization reaction thatoccurs.The Arrhenius theory has limitations, however. For example, H (aq) ,a bare proton, does not exist in water. The positive charge on a proton isattracted to the region of negative charge on the lone pair of electrons ona water molecule’s oxygen atom. The combination is a hydrated protoncalled a hydronium ion, H3O (aq) .H (aq) H2O( ) H3O (aq)The hydronium ion, itself, forms hydrogen bonds with other watermolecules. (See Figure 8.1.) Thus, a better formula for the ion that ispresent in acidic solutions is [H(H2O)n] , where n is usually 4 or 5. Forconvenience, however, chemists usually use a single hydronium ionwhen writing equations. hydrogen bondH3Figure 8.1 In aqueous solution,the hydronium ion, H3O , formshydrogen bonds with other watermolecules.O The Arrhenius theory also has limitations for explaining certain reactions.For example, aqueous solutions of ammonia are basic. They react withacids in neutralization reactions, even though ammonia does not containthe hydroxide ion. Many aqueous solutions of salts with no hydroxideions are basic, too. Some reactions take place without any liquid solvent.For example, ammonium chloride can be formed by the reaction betweenammonia and hydrogen chloride, which are both gases:NH3(g) HCl(g) NH4Cl(s)Chapter 8 Acids, Bases, and pH MHR379
The Brønsted-Lowry TheoryThe limitations of the Arrhenius theory of acids and bases are overcomeby a more general theory, called the Brønsted-Lowry theory. This theorywas proposed independently, in 1923, by Johannes Brønsted, a Danishchemist, and Thomas Lowry, an English chemist. It recognizes anacid-base reaction as a chemical equilibrium, having both a forwardreaction and a reverse reaction that involve the transfer of a proton. TheBrønsted-Lowry theory defines acids and bases as follows: An acid is a substance from which a proton can be removed. (Somechemists describe Brønsted-Lowry acids as “proton-donors.”) A base is a substance that can accept a proton. (Some chemists describeBrønsted-Lowry bases as “proton-acceptors.”)Note that the word “proton” refers to the nucleus of a hydrogen atom — anH ion that has been removed from the acid molecule. It does not referto a proton removed from the nucleus of another atom, such as oxygen orsulfur, that may be present in the acid molecule. As mentioned previously,H ions share electrons with any species (ion or molecule) that has a lonepair of electrons. In aqueous solution, the proton bonds with a watermolecule to form the hydronium ion. Unlike the Arrhenius theory,however, the Brønsted-Lowry theory is not restricted to aqueous solutions.For example, the lone pair of electrons on an ammonia molecule can bondwith H , and liquid ammonia can act as a base.Definition TermArrhenius TheoryBrønsted-Lowry Theoryacida substance that containshydrogen and dissociatesin water to form H (aq)a substance from whicha proton can be removedbasea substance that containsthe hydroxide group anddissociates in water toform OH (aq)a substance that canaccept a proton froman acidConjugate Acid-Base PairsThe dissociation of acetic acid in water is represented in Figure 8.2.This dissociation is an equilibrium reaction because it proceeds in bothdirections. Acetic acid is weak, so only a few ions dissociate. The positionof equilibrium lies to the left, and the reverse reaction is favoured. Inthe reverse reaction, the hydronium ion gives up a proton to the acetateion. Thus, these ions are an acid and a base, respectively, as shown inFigure 8.3. The acid on the left (CH3COOH) and the base on the right(CH3COO ) differ by one proton. They are called a conjugate acid-basepair. Similarly, H2O and H3O are a conjugate acid-base pair.OCH3H COHacid(acetic acid)Figure 8.2OOHbase(water) HHHacid(hydronium ion)The dissociation of acetic acid, a weak acid, in water380 MHR Unit 4 Chemical Systems and EquilibriumO CH3CObase(acetate ion)
OCH3H COOHOHacidbase HO HCH3 COHconjugate acidconjugate baseconjugate pairconjugate pairFigure 8.3Conjugate acid-base pairs in the dissociation of acetic acid in waterUnlike the Arrhenius theory, the Brønsted-Lowry theory of acids andbases can explain the basic properties of ammonia when it dissolves inwater. See Figure 8.4.conjugate pairHN OH H HbaseH NHacid OHH H Hconjugate acidconjugate baseconjugate pairFigure 8.4The dissociation of ammonia, a weak base, in waterAqueous ammonia is a weak base, so relatively few hydroxide ions form.The position of equilibrium lies to the left. In the forward reaction, thewater molecule gives up a proton and acts as an acid. A substance thatcan act as a proton donor (an acid) in one reaction and a proton acceptor(a base) in another reaction is said to be amphoteric. (Water acts as anacid in the presence of a stronger base, and as a base in the presence of astronger acid.Sample ProblemIdentifying Conjugate Acid-Base PairsProblemIdentify the conjugate acid-base pair in each reaction. H2PO4 (aq) H3O (aq)(a) H3PO4(aq) H2O( ) HPO42 (aq) H2O( )(b) H2PO4 (aq) OH (aq) Continued .Chapter 8 Acids, Bases, and pH MHR381
Continued .SolutionOn the left side of the equation, the acid is the molecule or ion thatdonates a proton. The base is the molecule or ion that accepts theproton. On the right side of the equation, you can identify the conjugateacid and base by the difference of a single proton from the base and acidon the left side.(a) The conjugate acid-base pairs are H3PO4 /H2PO4 and H2O /H3O .conjugate pairbaseconjugate acidH2PO4 (aq)H3PO4(aq) H2O( )acid H3O (aq)conjugate baseconjugate pair(b) The conjugate acid-base pairs are H2PO4 /HPO42 and OH /H2O.conjugate pairconjugate acidbaseH2PO4 (aq) OH (aq)acidHPO42 (aq) H2O( )conjugate baseconjugate pairCheck Your SolutionIn each case, the acid has one more proton than its conjugate base.Practice Problems1. Name and write the formula of the conjugate base of each moleculeor ion.(a) HCl(b) HCO3 (c) H2SO4(d) N2H5 2. Name and write the formula of the conjugate acid of each moleculeor ion.(a) NO3 (b) OH (c) H2O(d) HCO3 3. Identify the conjugate acid-base pairs in each reaction. H2S(aq) OH (aq)(a) HS (aq) H2O( ) (b) O2 (aq) H2O( ) 2OH (aq)4. Identify the conjugate acid-base pairs in each reaction. NH4 (aq) HS (aq)(a) H2S(aq) NH3(aq) (b) H2SO4(aq) H2O( ) H3O (aq) HSO4 (aq)382 MHR Unit 4 Chemical Systems and Equilibrium
Molecular Structure and the Strength of Acids and BasesWhen a strong acid or base dissolves in water, almost every acid or basemolecule dissociates. While there are many acids and bases, most areweak. Thus, the number of strong acids and strong bases is fairly small.Strong Acids binary acids that have the general formula HX(aq) , where X Cl,Br, and I (but not F): for example, hydrochloric acid, HCl, andhydrobromic acid, HBr (HCl and HBr are hydrohalic acids: acidsthat have hydrogen bonded to atoms of the halogen elements.) oxoacids (acids containing oxygen atoms) in which the number ofoxygen atoms exceeds, by two or more, the number of protons thatcan be dissociated: for example, nitric acid, HNO3 , sulfuric acid,H2SO4 , perchloric acid, HClO4 , and chloric acid, HClO3The binary acids of non-metals exhibit periodic trends in their acidstrength, as shown in Figure 8.5. Two factors are responsible for thistrend: the electronegativity of the atom that is bonded to hydrogen,and the strength of the bond.increasing acid strengthincreasing H4NH3H2OHFH2SHClH2SeHBrH2TeHIincreasing acid strength15(VA)decreasing bond strength14(IVA)The binary acidsshow periodic trends, which arerelated to electronegativity andbond strength.Figure 8.5Across a period, electronegativity is the most important factor. The acidstrength of hydrides increases as their electronegativity increases. Thishappens because an electronegative atom draws electrons away from thehydrogen atom, making it relatively positive. The negative pole of a watermolecule then strongly attracts the hydrogen atom and pulls it away.Down a group, bond strength is the most important factor. Acidstrength increases as bond strength decreases. A weaker bond means thatthe hydrogen atom is more easily pulled away from the atom to which itis attached. For example, hydrofluoric acid is a stronger acid than water,but HF is the weakest of the hydrohalic acids because the H-F bond isrelatively strong.Chapter 8 Acids, Bases, and pH MHR383
CHEMFA C TYou might wonder how HCl(aq),HBr(aq) , and HI(aq) can bedescribed as increasing instrength, since each aciddissociates completely inwater. This trend becomesapparent if you add equalconcentrations of the acidsto a solvent that is less basicthan water, such as pureacetic acid. You will findthat the acids dissociate todifferent extents.Oxoacids increase in strength with increasing numbers of oxygenatoms, as shown in Figure 8.6. The hydrogen atoms that dissociatein water are always attached to oxygen atoms. Oxygen is moreelectronegative than hydrogen, so oxygen atoms draw electrons away fromhydrogen atoms. The more oxygen atoms there are in a molecule, thegreater is the polarity of the bond between each hydrogen atom and theoxygen atom it is attached to, and the more easily the water molecule cantear the hydrogen atom away.increasing acid strengthOHO δClO δδhypochlorousacidFigure 8.6HδClHOClOδδOchlorousacid OchloricacidH δO δClOOperchloricacidThe relative strength of oxoacids increases with the number of oxygen atoms.Acids such as HCl, CH3COOH, and HF are monoprotic acids. They haveonly a single hydrogen atom that dissociates in water. Some acids havemore than one hydrogen atom that dissociates. These acids are calledpolyprotic acids. For example, sulfuric acid has two hydrogen atoms thatcan dissociate.H2SO4(aq) H2O( ) H3O (aq) HSO4 (aq)HSO4 (aq) H2O( ) H3O (aq) SO42 (aq)Sulfuric acid is a far stronger acid than the hydrogen sulfate ion, becausemuch more energy is required to remove a proton from a negativelycharged ion. The strength of a polyprotic acid decreases as the number ofhydrogen atoms that have dissociated increases.Strong bases are confined to the oxides and hydroxides from Groups 1(IA) and 2 (IIA).Strong Bases all oxides and hydroxides of the alkali metals: for example, sodiumhydroxide, NaOH, and potassium hydroxide, KOH alkaline earth (Group 2 (IIA)) metal oxides and hydroxides belowberyllium: for example, calcium hydroxide, Ca(OH)2 , and bariumhydroxide, Ba(OH)2The strong basic oxides have metal atoms with low electronegativity.Thus, the bond to oxygen is ionic and is relatively easily broken by theattraction of polar water molecules. The oxide ion always reacts withwater molecules to produce hydroxide ions.O2 (aq) H2O( ) 2OH (aq)Magnesium oxide and magnesium hydroxide are not very soluble. Theyare strong bases, however, because the small amount that does dissolvedissociates almost completely into ions. Beryllium oxide is a weak base.(It is the exception in Group 2 (IIA).) It is a relatively small atom, so thebond to oxygen is strong and not easily broken by water molecules.384 MHR Unit 4 Chemical Systems and Equilibrium
Calculations That Involve Strong Acids and BasesWhen a strong acid dissociates completely into ions in water, theconcentration of H3O (aq) is equal to the concentration of the strongacid. Similarly, when a strong base dissociates completely in water, theconcentration of OH (aq) is equal to the concentration of the strong base.Sample ProblemCalculating Ion Concentrations in Acidic and Basic SolutionsProblemDuring an experiment, a student pours 25.0 mL of 1.40 mol/L nitricacid into a beaker that contains 15.0 mL of 2.00 mol/L sodiumhydroxide solution. Is the resulting solution acidic or basic? What is theconcentration of the ion that causes the solution to be acidic or basic?What Is Required?You must determine the ion in excess and its concentration.What Is Given?You have the following data:Volume of nitric acid 25.0 mL[HNO3] 1.40 mol/LVolume of sodium hydroxide 15.0 mL[NaOH] 2.00 mol/LPlan Your StrategyStep 1 Write the chemical equation for the reaction.Step 2 Calculate the amount of each reactant using the following equation.Amount (in mol) Concentration (in mol/L) Volume (in L)Step 3 Determine the limiting reactant.Step 4 The reactant in excess is a strong acid or base. Thus, the excessamount results in the same amount of H3O or OH .Step 5 Calculate the concentration of the excess ion by using the amountin excess and the total volume of the solution.Act on Your StrategyStep 1 HNO3(aq) NaOH(aq) NaNO3(aq) H2O( )Step 2 Amount of HNO3 1.40 mol/L 0.0250 L 0.0350 molAmount of NaOH 2.00 mol/L 0.0150 L 0.0300 molStep 3 The reactants combine in a 1:1 ratio. The amount of NaOH is less,so this reactant must be the limiting reactant.Step 4 Amount of excess HNO3(aq) 0.0350 mol 0.0300 mol 0.005 0 molTherefore, the amount of H3O (aq) is 5.0 10 3 mol.Step 5 Total volume of solution 25.0 mL 15.0 mL 40.0 mL5.0 10 3 mol0.0400 L 0.12 mol/L[H3O ] Continued .Chapter 8 Acids, Bases, and pH MHR385
Continued .The solution is acidic, and [H3O ] is 0.12 mol/L.Check Your SolutionThe chemical equation has a 1:1 ratio between reactants. The amount ofacid is greater than the amount of base. Therefore, the resulting solutionshould be acidic, which it is.Practice Problems5. Calculate the concentration of hydronium ions in each solution.(a) 4.5 mol/L HCl(aq)(b) 30.0 mL of 4.50 mol/L HBr(aq) diluted to 100.0 mL(c) 18.6 mL of 2.60 mol/L HClO4(aq) added to 24.8 mL of 1.92 mol/LNaOH(aq)(d) 17.9 mL of 0.175 mol/L HNO3(aq) added to 35.4 mL of0.0160 mol/L Ca(OH)2(aq)6. Calculate the concentration of hydroxide ions in each solution.(a) 3.1 mol/L KOH(aq)(b) 21.0 mL of 3.1 mol/L KOH diluted to 75.0 mL(c) 23.2 mL of 1.58 mol/L HCl(aq) added to 18.9 mL of 3.50 mol/LNaOH(aq)(d) 16.5 mL of 1.50 mol/L H2SO4(aq) added to 12.7 mL of 5.50 mol/LNaOH(aq)7. Determine whether reacting each pair of solutions results in anacidic solution or a basic solution. Then calculate the concentrationof the ion that causes the solution to be acidic or basic. (Assumethat the volumes in part (a) are additive. Assume that the volumes inpart (b) stay the same.)(a) 31.9 mL of 2.75 mol/L HCl(aq) added to 125 mL of 0.0500 mol/LMg(OH)2(aq)(b) 4.87 g of NaOH(s) added to 80.0 mL of 3.50 mol/L HBr(aq)8. 2.75 g of MgO(s) is added to 70.0 mL of 2.40 mol/L HNO3(aq) . Is thesolution that results from the reaction acidic or basic? What is theconcentration of the ion that is responsible for the character of thesolution?Section SummaryStrong acids and bases (and strong electrolytes) dissociate completely inwater. Therefore, you can use the concentrations of these compounds todetermine the concentrations of the ions they form in aqueous solutions.You cannot, however, use the concentrations of weak acids, bases, andelectrolytes in the same way. Their solutions contain some particlesthat have not dissociated into ions. Nevertheless, important changes in[H3O ] and [OH ] take place because dissolved ions affect the dissociationof water.386 MHR Unit 4 Chemical Systems and Equilibrium
In the next section, you will focus on the equilibrium of water. Youwill discover how the pH scale is related to the concentrations of the ionsthat form when water dissociates. As well, you will learn how to calculatethe pH values of solutions of weak acids and bases.Section Review1K/U Phosphoric acid, H3PO4(aq) is triprotic. It has three hydrogen ionsthat may be dissociated.(a) Write an equation to show the dissociation of each proton.(b) Show that H2PO4 (aq) can act as either an acid or a base.(c) Which is the stronger acid, H3PO4(aq) or H2PO4 (aq) ? Explainyour answer.2Para-aminobenzoic acid (PABA) is a weak monoprotic acid that isused in some sunscreen lotions. Its formula is C6H4NH2COOH. What isthe formula of the conjugate base of PABA?3Boric acid, B(OH)3(aq), is used as a mild antiseptic in eye-washsolutions. The following reaction takes place in aqueous solution.B(OH)3(aq) 2H2O( ) B(OH)4 (aq) H3O (aq)K/UK/U(a) Identify the conjugate acid-base pairs.(b) Is boric acid strong or weak? How do you know?4Classify each compound as a strong acid, weak acid, strongbase, or weak base.K/U(a) butyric acid, CH3CH2CH2COOH (responsible for the odour ofrancid butter)(b) hydroiodic acid, HI(aq) (added to some cough syrups)(c) potassium hydroxide, KOH (used in the manufacture of soft soaps)(d) red iron oxide, Fe2O3 (used as a colouring pigment in paints)5C Distinguish between a concentrated solution of a weak base, anda dilute solution of a strong base. Give an example of each.Chapter 8 Acids, Bases, and pH MHR387
8.2Section Preview/Specific ExpectationsIn this section, you will define and performcalculations that involvethe ion product constantfor water, Kw, and the aciddissociation constant, Ka compare strong acids andbases in terms of equilibrium compare weak acids andbases in terms of equilibrium communicate yourunderstanding of thefollowing terms: ion productconstant for water ( Kw ),pH, pOH, acid dissociationconstant (Ka), percentdissociationThe Equilibrium ofWeak Acids and BasesThe dissociation of an acidic or basic compound in aqueous solutionproduces ions that interact with water. The pH of the aqueous solution isdetermined by the position of equilibrium in reactions between the ionsthat are present in solution and the water molecules. Pure water containsa few ions, produced by the dissociation of water molecules: H3O (aq) OH (aq)2H2O( ) At 25 C, only about two water molecules in one billion dissociate. Thisis why pure water is such a poor conductor of electricity. In neutralwater, at 25 C, the concentration of hydronium ions is the same as theconcentration of hydroxide ions: 1.0 10 7 mol/L. These concentrationsmust be the same because the dissociation of water produces equal numbers of hydronium and hydroxide ions. Because this is an equilibriumreaction, and because the position of equilibrium of all reactions changeswith temperature, [H3O ] is not 1.0 10 7 mol/L at other temperatures.The same is true of [OH ] .The Ion Product Constant for WaterThe equilibrium constant, Kc , for the dissociation of water is given by thefollowing expression.[H3O ][OH ]Kc [H2O]2So few ions form that the concentration of water is essentially constant.The product Kc[H2O]2 is equal to the product of the concentrations ofhydronium ions and hydroxide ions. The equilibrium value of theconcentration ion product [H3O ][OH ] at 25 C is called the ion productconstant for water. It is given the symbol Kw .Kc[H2O]2 [H3O ][OH ]1.0 10 7 mol/L 1.0 10 7 mol/L1.0 10 14KwThe units are commonly dropped, as in other equilibrium expressions youhave encountered.The concentration of H3O in the solution of a strong acid is equalto the concentration of the dissolved acid, unless the solution is verydilute. Consider [H3O ] in a solution of 0.1 mol/L hydrochloric acid.All the molecules of HCl dissociate in water, forming a hydronium ionconcentration that equals 0.1 mol/L. The increased [H3O ] pushes thedissociation reaction between water molecules to the left, in accordancewith Le Châtelier’s principle. Consequently, the concentration of hydronium ions that results from the dissociation of water is even less than1 10 7 mol/L. This [H3O ] is negligible compared with the 0.1 mol/Lconcentration of the hydrochloric acid. Unless the solution is very dilute(about 1 10 7 mol/L), the dissociation of water molecules can be ignoredwhen determining [H3O ] of a strong acid.388 MHR Unit 4 Chemical Systems and Equilibrium
Similarly, the concentration of hydroxide ions can be determined fromthe concentration of the dissolved base. If the solution is a strong base,you can ignore the dissociation of water molecules when determining[OH ], unless the solution is very dilute. When either [H3O ] or [OH ] isknown, you can use the ion product constant for water, Kw , to determinethe concentration of the other ion. Although the value of Kw for water is1.0 10 14 at 25 C only, you can use this value unless another value isgiven for a different temperature.[H3O ] and [OH ] in Aqueous Solutions at 25 CIn an acidic solution, [H3O ] is greater than 1.0 10 7 mol/L and[OH ] is less than 1.0 10 7 mol/L.In a neutral solution, both [H3O ] and [OH ] are equal to1.0 10 7 mol/L.In a basic solution, [H3O ] is less than 1.0 10 7 mol/L and [OH ]is greater than 1.0 10 7 mol/L.Sample ProblemDetermining [H3O ] and [OH–]ProblemFind [H3O ] and [OH ] in each solution.(a) 2.5 mol/L nitric acid(b) 0.16 mol/L barium hydroxideSolutionYou know that nitric acid is a strong acid and barium hydroxide is astrong base. Since both dissociate completely in aqueous solutions,you can use their molar concentrations to determine [H3O ] or [OH ].You can find the concentration of the other ion using Kw :Kw 1.0 10 14 [H3O ][OH ](a) [HNO3] 2.5 mol/L, so [H3O ] 2.5 mol/L1.0 10 14 mol/L2.5 4.0 10 15 mol/LH2OBa2 (aq) 2OH (aq)(b) Ba(OH)2[OH ] Each mole of Ba(OH)2 in solution forms two moles of OH ions. [OH ] 2 0.16 0.32 mol/L1.0 10 14 mol/L0.32 3.1 10 14 mol/L[H3O ] Check Your SolutionFor a solution of a strong acid, as in part (a), [H3O ] should be greaterthan 1.0 10 14 and [OH ] should be less than 1.0 10 14. For a solutionof strong base, [OH ] should be greater than, and [H3O ] should be lessthan, 1.0 10 14.Chapter 8 Acids, Bases, and pH MHR389
Practice ProblemsSulfuric acid is the onlycommon strong diprotic acid.Explain why the concentrationof hydronium ions in a solutionof 1.0 mol/L H2SO4(aq) is1.0 mol/L, not 2.0 mol/L.9. Determine [H3O ] and [OH ] in each solution.(a) 0.45 mol/L hydrochloric acid(b) 1.1 mol/L sodium hydroxide10. Determine [H3O ] and [OH ] in each solution.(a) 0.95 mol/L hydrobromic acid(b) 0.012 mol/L calci
Acids, Bases, and pH The hydrohalic acids (HX (aq)), where X represents a halogen) include HF, HCl, HBr, and HI. Only HF is a weak acid. The rest are strong acids. What factors account for this difference? 8.1 Explaining the Properties of Acids and Bases 8.2 The Equilibrium of Weak Acids and Bases 8.3
2. Describe the common properties of acids and bases 3. Identify acids and bases using indicators, pH papers 4. Name some common lab acids and bases, acids at and bases at home 5. Describe reactions of acids with metals, bases and carbonates 6. Describe the application of acids, bases and p
_7. Which statement describes an alternate theory of acids and bases? (1) Acids and bases are both H acceptors. (2) Acids and bases are both H donors. (3) Acids are H acceptors, and bases are H donors. (4) Acids are H donors, and bases are H acceptors. _8. Which substance is the
Properties of Acids and Bases Return to the Table of contents Slide 5 / 208 What is an Acid? Acids release hydrogen ions into solutions Acids neutralize bases in a neutralization reaction. Acids corrode active metals. Acids turn blue litmus to red. Acids taste sour. Properties of Acids Slide 6 / 208 Properties
properties of acids and bases. 5. Describe the colors that form in acidic and basic solutions with litmus paper and phenolphthalein. 6. Explain the difference between strong acids or bases and weak acids or bases. 7. Memorize the strong acids and bases. 8. Define the terms polyprotic and amphiprotic. 9. Perform calculations using the following .
Unit 12 Acids and Bases- Funsheets Part A: Name and write the formula for the following acids and bases. 1) Carbonic acid _ . Part E: Using your knowledge of acids and bases, answer the following questions. 1) Fill in the following Venn diagram about properties of acids and bases. You must fill in at least 4 facts in each.
Lecture Notes for Chapter 16: Acids and Bases I. Acids and Bases a. There are several ways to define acids and bases. Perhaps the easiest way to start is to list some of the properties of acids and bases. b. The table below summarizes some properties that will be helpful
Unit 11 Objectives: Acids & Bases Acid-Base Nomenclature Content Objectives: I can name ionic compounds containing acids, and bases, using (IUPAC) nomenclature rules. I can write the chemical formulas of acids and bases. Criteria for Success: I can identify an acid as a binary acid or an oxyacid. I can name common binary acids, oxyacids, and bases given their chemical formula.
Young integral Z t 0 y sdx s; x;y 2C ([0;1]) Recall theRiemann-Stieltjes integral: Z 1 0 y sdx s B lim jPj!0 X [s;t]2P y s ( x t{z x s}) Cx s;t () Pa finite partition of [0;1] Th
