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Chemistry 30 HALLETT 2013

Chapter 16 Notes

Section 16.1 Water Ionization and Acid-Base Strength

-there are numerous chemical reactions and systems that depend on the nature of acids and bases and/or the pH of the solution-examples: permanganate ions oxidizing only with hydronium ions, COOH organic compounds will make solutions with pHs less than 7, electrolysis of potassium iodide produces a basic solution, chromate-dichromate ion system at equilibrium (figure 1)-combining knowledge of equilibrium concepts with acids and bases will develop a more comprehensive understanding

Complete Lab Exercise 16A

-for simplicity in numerous chemical reactions, hydrogen ions H+(aq) is used, however it is more consistent with the evidence to represent it as hydronium ions, H3O+(aq) (chapter 6). (read the learning tip)

The Water Ionization Constant, Kw:-even highly purified water has a slight conductivity (figure 2), which is due to the presence of ions (figure 3)-this means that the conductivity in pure water is due to the ions produced by the ionization of some water molecules into hydronium and hydroxide ions-because the conductivity is slight, the equilibrium at SATP greatly favours the water molecules – evidence shows at 25 degrees, fewer than two of every billion molecules are in ionized form-the equilibrium constant expression for this ionization is an extremely small number, because liquid water (or the water in any dilute aqueous solution) has an essentially constant concentration, which does not appear in the expression, but is incorporated into the equilibrium constant value-the water ionization equilibrium relationship that it is called the ion product or ionization constant for water, Kw

Kw = [H3O+(aq)][OH-(aq)] = 1.00 x 10-14 at SATP-the equilibrium equation shows the ions form in a 1:1 ratio making their concentrations equal, which is also the case for any neutral aqueous solution-using math, the concentrations of these ions can be calculated – see in textbook-the ionization of water is important in studying acidic and basic solutions-recall that acids react with water to form hydronium ions in concentrations higher than 10-7 mol/L-recall that bases have a hydroxide ion concentration greater than 10-7 mol/L – produced either by the complete dissociation of an ionic hydroxide or by the partial reaction of some weak base ion or molecule with water-the most important point about Kw is that it applies to pure water and to any solution that is mostly water


-this ionization equilibrium will be involved in any other reaction in an aqueous solution, if that reaction involves hydronium ions or hydroxide ions in any way-example: Lab 16A – since the reaction contains hydronium ions, the chromate-dichromate equilibrium can be controlled by adjusting either the hydronium or hydroxide ion concentration (shifting the water ionization equilibrium)-other examples are found on the “relative strengths of oxidizing and reducing agents table” (those containing hydroxide or hydronium ions)-the equation may also be switched to calculate either ion concentrations – see textbook-in ordinary dilute aqueous acidic or basic solutions, the presence of substances other than water decreases the certainty of the Kw value at 25 degrees to two significant digits

Read communication examples 1-4 and the learning tips.Complete practice questions #1-6

Communicating Concentrations: pH and pOH-review that: pH = -log[H3O+(aq)] and [H3O+(aq)] = 10-pH

pOH = -log[OH-(aq)] and [OH-(aq)] = 10-pOH

-look at the math used to create the relationship: pH + pOH = 14.00, which allows for a quick conversion between pH and pOH values

Read Learning Tip and complete Practice #7-9 questions

Acid Strength as an Equilibrium Position:-acidic solutions of different substances at the same concentration do not possess acid properties to the same degree – the pH of a 1.00mol/L solution of an acid can vary from 0-7, conductivity and chemical reaction rates may also vary depending on the acid-weak and strong acid concepts were developed to explain these differences-a weak acid is if its characteristic properties are less than those of a common strong acid – weaker electrolytes, react at a slower rate, pH closer to 7-a strong acid is explained as an acid that reacts quantitatively with water to form hydronium ions -example: the reaction of dissolved hydrogen chloride (hydrochloric acid) is virtually complete – the equation could be written with double equilibrium arrows with >99.9%, but it is more common to write it with a single arrow (see textbook)-read learning tip-a weak acid reacts partially with water to form hydronium ions – less than 50%-example: acetic acid reacts only 1.3% in solution 25 degrees and 0.10 mol/L – see textbook-remember that any equilibrium position depends on concentrations as well as temperature – experiments show that the higher the concentration of a weak acid solution, the lower its percent ionization becomes -the hydronium ion concentration of any acid solution can be calculated by multiplying the percent reaction by the initial amount concentration of the acid solute


-example: HCl solution – virtually 100% of the HCl molecules react with water molecules at equilibrium – see calculations in textbook-the six classified strong acids are: hydrochloric, nitric, sulfuric, hydrobromic, hydroiodic, and perchloric acid solutions – the first 3 are common-for any aqueous strong acid, assume that the concentration of hydronium ions in solution is equal to the initial concentration of the acid dissolved -for weak acids (majority) you need to calculate the hydronium ion concentration in solution because only a small proportion of the initial acid concentration will be converted to ions at equilibrium

Read communication example 5

-the stronger the acid, the lower the pH, the higher the hydronium ion concentration, the greater the percent reaction-we can find the percent reaction for ionization of any weak acid solution from the measured pH of a solution of known initial concentration

Read learning tip and communication example 6

Complete bold terms in notebook, add additional notes from textbook and complete P. 721 questions #1-16.


Section 16.2 The Bronsted-Lowry Acid-Base Concept

-even with the more comprehensive understanding of acids and bases there are still problems: no provisions for reactions that do not occur in aqueous solution, some substances that have both acid and base properties- example: sodium hydrogen carbonate (baking soda) – forms a basic solution in water, but will also neutralize a sodium hydroxide solution – see the textbook equations – acid or base?!

The Proton Transfer Concept:-according to Bronsted and Lowry (figure 1), acids in reactions donate a proton to a water molecule (see text for example) and bases in reactions accepts a proton from a water molecule (see text for example)- water does not need to be one of the reactants (see 2 examples in text)-a Bronsted-Lowry neutralization is a competition for protons that results in a proton transfer from the strongest acid present to the strongest base present-a Bronsted-Lowry reaction equation is an equation written to show an acid-base reaction involving the transfer of a proton from one entity (acid) to another (base)-only an entity that is involved in a proton transfer in a reaction can be defined as an acid or a base – and only for that particular reaction – example: reaction of HCl with water – water acts as a base. Reaction of water with NH3 – water acts as an acid-substances that appear to act as a Bronsted-Lowry acid in some reactions and as a Bronsted-Lowry base in other reactions is called amphoteric-amphoteric is an empirical term referring to a chemical substance with the ability to react as either an acid or a base-amphiprotic is a theoretical term describing an entity (ion or molecule) having the ability to either accept or donate a proton-the hydrogen carbonate ion in baking soda (figure 3), like every other hydrogen polyatomic ion, is amphiprotic (see the reactions in text) – note that the Kc value for the ions reacting as a base is 2000 times more than the number reacting as an acid making the resulting solution basic-the amphoteric nature of baking soda (caused by the amphiprotic nature of the hydrogen carbonate ion) is expressed through the evidence of neutralizing strong acids and bases (see reactions in text)-this theory does not explain why a proton is accepted or donated and cannot predict which will occur -the advantage of the Bronsted-Lowry definitions is that they enable us to define acids and bases in terms of chemical reactions rather than simply as substances that form acidic and basic aqueous solutions

Complete Practice questions #1-6

Conjugate Acids and Bases:-the acid-base reactions in the Bronsted-Lowry concept are universally reversible and always result in an acid-base equilibrium


-in a proton transfer reaction at equilibrium, both forward and reverse reactions involve Bronsted-Lowry acids and bases-example: acetic acid solution – forward reaction: proton transfer from acetic acid to water molecules, reverse reaction – proton transfer from hydronium to acetate ions (see text)-this equilibrium is typical of all acid-base reactions – always two acids and two bases-the base on the right is formed by removal of a proton from the acid on the left, the acid on the right is formed by the addition of a proton to the base on the left – a pair of substances with formulas that differ only by a proton is called a conjugate acid-base pair-example: an acetic acid molecule and an acetate ion are a conjugate acid-base pair – acetic acid: conjugate acid, acetate ion: conjugate base-there must always be two conjugate acid-base pairs in any proton transfer reaction-conjugate acid-base pairs appear opposite each other in a table of acids and bases in Appendix I-at equilibrium, only 1.3% of the CH3COOH molecules have reacted with water – this means that the ability of the CH3COO- part of the molecule to keep its proton (H+) is greater than the ability of the H2O to attract the proton away (figure 5), which means the CH3COO- is a stronger base than H2O-a quantitative example is when HCl reacts with water (figure 6): the Cl- of each HCl molecule has a much weaker attraction for its proton (H+) than any colliding water molecule. The water molecules gain the protons easily and at equilibrium essentially all of the HCl molecules have lost protons to water molecules – making the reaction quantitative and the HCl is called a strong acid-Remember: the stronger the base, the more it attracts another proton, the stronger an acid, the less it attracts its own proton

-read learning tip on p. 726-strong and weak acids can be explained by the Bronsted-Lowry concepts and by comparing the reactions of different acids with the same base-HA is the general symbol for any acid, and A- is its conjugate base in the Relative Strengths of Aqueous Acids and Bases table (Appendix I), which lists the position of equilibrium of aqueous solutions of different acids – ordered according to how much they ionize in (react with) the water solvent (See text)-the extent of the proton transfer between HA and H2O determines the relative strength of HA(aq)-when a strong acid reacts with water, an almost complete transfer of protons occurs in the forward reaction and almost no transfer of protons occurs in the reverse reaction – making a nearly 100% reaction -the equilibrium constant value is found to be very large-theoretically, a strong acid holds its proton weakly and easily loses it to any base, this means that the conjugate base of a strong acid must have a weak attraction for protons-the stronger an acid, the weaker its conjugate base; and the weaker an acid, the stronger its conjugate base

Complete Practice questions #7-8Complete Investigation 16.1


Predicting Acid-Base Reaction Equilibria:-to predict the outcome of any acid-base combination, we must rely on empirical evidence as the Bronsted-Lowry theory does not include any theoretical explanation for why entities will attract a proton -this evidence is gained by measuring and recording the relative strengths of acid and base entities

Complete Lab Exercise 16.B

-when making complex predictions, a combination of empirical and theoretical concepts are used to predict both the products and the extent of acid-base reactions-according to the collision-reaction theory, a proton transfer may result from a collision between an acid and a base-in a system of a mixture of several acids and bases, there are countless random collisions, which means that there are many different possible acid-base reactions and that they all occur (to some extent) all the time-evidence shows that one reaction will occur to a greater extent than the other reactions so that it is the only observable reaction-an explanation of which reaction will predominate by using a combination of the collision-reaction theory and the Bronsted-Lowry concept-a proton could theoretically transfer several times if there were several different acids and bases in a system, transferring to a stronger proton attractor each time-once gained by the strongest base, a proton will remain there -once the strongest acid has lost its proton, its conjugate base cannot gain one back from any other entity because it is a weaker proton attractor than anything else in the system-overall, the predominant acid-base reaction should be the one that involves proton transfer from the strongest acid to the strongest base present in the system-for an aqueous solution system, first list all the entities present as they exist in aqueous solution-Table 1 summarizes how to represent the entities that are present

Look at Sample problem 16.1 for the 5 steps for determining the prominent acid-base proton transferRead the Learning Tip for predicting the position of equilibrium

There are specific restrictions that apply when predicting proton transfer reactions in aqueous solution, because they are much more common than acid-base reactions in any other environments:

-hydronium ions are the strongest acid entity that can exist in an aqueous solution. Any stronger entity would dissolve and react with the water to form hydronium ions. This means that the six strongest acids are all written as H3O+(aq) when in aqueous solution

-hydroxide ions are the strongest base entity that can exist in aqueous solution. Any stronger entity would dissolve and react with the water to form hydroxide ions. The only common example is the dissolving of soluble ionic compounds such as Na2O(s). In such cases, the oxide ion is written as OH-(aq) when in aqueous solution.

-no entity in aqueous solution can react as a base if it is a weaker base than water. For this reason, the conjugate bases of the six strong acids are not considered as bases, in aqueous solution.

Read Learning Tip

Even though the products are favoured for the predominant reaction, we can’t accurately predict the actual equilibrium position of all cases of any specific acid-base reaction by this method


-remember that the initial concentration of entities plays a large part in determining the percent reaction at equilibrium-to predict the acid-base equilibrium, the reaction conditions must be restricted-unless otherwise told, assume that for the predominant reaction in an acid-base system:

-the equation represents a single proton transfer between two entities, neither of which is water, and has a stoichiometric ratio of 1:1:1:1

-the strongest acid and the strongest base present in the reaction system are both present in significant chemical amounts, in approximately equal concentrations-under these circumstances, you may assume that a Bronsted-Lowry acid-base reaction where products are favoured has a percent reaction greater than 50%, or that a reaction where reactants are favoured may be labeled as <50%-with these restrictions, you may also assume that a Kc value for any such reaction is >1 if products are favoured, and <1 if reactants are favoured-note that the above restrictions exclude any reaction equation written to describe the ionization (reaction with water) of any single acid or base entity dissolved to make an aqueous solution – in these cases, water is acting as one of the reactants, and the amount concentration of the water is very much greater than that of the other entity-recall that the reaction of hydronium ions with hydroxide ions is always quantitative, therefore it should always be written with a single arrow-there are many other acid-base reactions that are quantitative, but it is not possible to predict them from the table of Relative Strengths of Aqueous Acids and Bases -other quantitative reactions will always be identified for you

Read communication example

A Five-Step Method for Predicting the Predominant Acid-Base Reaction:

1. List all entities (ions, atoms, or molecules including water) initially present as they exist in aqueous solution (table 1, page 728)

2. Identify and label all possible aqueous acids and bases, using the Bronsted-Lowry definitions

3. Identify the strongest acid and the strongest base present, using the table of Relative Strengths of Aqueous Acids and Bases

4. Write an equation showing a transfer of one proton from the strongest acid to the strongest base, and predict the conjugate base and the conjugate acid to be the products

5. Predict the approximate position of the equilibrium, using the generalization developed in Lab Exercise 16B, Learning Tip on p. 729, and the table of Relative Strengths of Aqueous Acids and Bases

Complete: Practice Questions #9-17, Lab Exercise 16CInvestigation 16.2Lab Exercise 16DCase Study



Section 16.3: Acid-Base Strength and the Equilibrium Law

The Acid Ionization Constant, Ka:-the way that chemists communicate the strength of any weak acid is by using the equilibrium constant expression for a Bronsted-Lowry reaction equation showing the acid ionizing in water-the constant is called the acid ionization constant Ka – see the Ka values listed in the Relative Strengths of Aqueous Acids and Bases data table-note that the strong acids are labeled as “very large” because they all react quantitatively with water to form hydronium ions - because there is no acid stronger than H3O+(aq) they are all considered equivalent-all the other acids on the table are weak acids, that vary greatly with their reactivity with water at equilibrium-to empirically rank any weak acid as stronger or weaker than another, we compare the ionization of solutions of the same concentration to see which has a lower pH (this is how appendix I was created)-example: look at a solution of hydrofluoric acid, HF(aq) and its reaction in water on p. 737-note that liquid water is omitted from the Ka equilibrium law expression – this is because we assume that this value will remain essentially constant for aqueous solutions that are not highly concentrated-this isn’t exactly true because the water is not a separate liquid phase in an aqueous solution, but as long as the amount of water solvent is much greater than the amount of acid solute, making this assumption will cause negligible change -all Ka values are calculated with this assumption, making them less accurate as the acid concentration increases – this is why the values are only given to two significant digits because they are usually only known to a certainty of about ±5%-read learning tip-two calculations involving the Ka constant are common for weak acid solutions:1. calculating a Ka value from measured (empirical) amount concentration data2. using a Ka value to predict a concentration of hydronium ions for an aqueous solution where the initial weak acid amount concentration is known

Calculating Ka from Amount Concentrations:-read Sample problems 16.2, and communication examples 1 and 2 and learning tips

Calculating [H3O + (aq)] from K a

-the second type of calculation involving a Ka value allows us to predict the acidity of any weak acid solution-we can calculate the concentration of hydronium ions from the initial acid concentration and the Ka value-read Sample problem 16.3-when the negligible ionization (percent reaction) assumption holds true, predictions of [H3O+(aq)] using Ka values become simple and straightforward-note that you can perform the 1000:1 ratio test for the simplifying equation before starting the problem-read communication example 3-you will not necessarily need to use the quadratic equation to solve these problems, but you may be asked to state whether a problem would require use of the quadratic equation – whether the ionization assumption test fails



Base Strength and the Ionization Constant, Kb:-the Bronsted-Lowry concept specifies that the strongest base possible in an aqueous solution is the hydroxide ion, so the hydroxide ion is “the” strong base – parallel to the strong acids-with strong bases, many of them contain hydroxide ions initially meaning that hydroxide is formed through the dissociation in water, not the reaction with water-therefore, we find the hydroxide ion concentration directly from a dissociation equation

Read Communication example 4

-other entities that react as bases are collectively weak bases, because empirical evidence indicates that they attract protons less than hydroxide ions do-weak bases react only partially (much less than 50%) with water in aqueous solution to produce hydroxide ions-we can communicate the strength of any weak base using the equilibrium constant expression for its reaction with water in dilute aqueous solution – called the base ionization constant, Kb -example: sodium citrate – assume that in an aqueous solution it dissociates into sodium and citrate ions. Sodium ions are spectators, but the citrate ions are found in the conjugate bases in the Relative Strength of Acids and Bases table – this means it is a weak Bronsted-Lowry base and reacts with water to form a basic solution at equilibrium-look in the text for the reaction equilibrium equation for aqueous sodium citrate and the equilibrium law expression-just as with Ka values, there are two kinds of weak base solution calculations that involve the Kb constant:

-calculating a Kb value from empirical (measured) amount concentration data-using a Kb value to predict an amount concentration of hydroxide ions for an aqueous solution

where the initial weak base concentration is known

Calculating Kb from Amount Concentrations:-essentially the same method as calculating Ka for a weak acid-often an extra calculation step must be included because measured pH values may need to be converted to find the equilibrium hydroxide ion concentration

Read communication example 5Complete Practice questions #10-13

Calculating [OH - (aq)] from K b:-the second type of calculation is the prediction of the basicity of any weak base solution-we can calculate the concentration of hydroxide ions from the initial weak base concentration and the Kb value-the problem is that relative acid-base tables do not normally list Kb values so we need to use the relationship between conjugate acid-base pairs to solve the problem

The Ka – Kb Relationship for Conjugate Acid-Base Pairs:-to demonstrate this concept, the weak acid acetic acid CH3COOH(aq) and its conjugate base the acetate ion, CH3COO-(aq) are used


-look in the text for the equilibrium reaction and law expressions and calculations-for any conjugate acid-base pair, Kw = KaKb and Kb = Kw/Ka

-to find a Kb value, you need to identify its conjugate acid and then divide Kw, 1.0 x 10-14, by the Ka value given for that conjugate acid

Read Sample Problem 16.4 and communication example 6

-predicting basicity always involves calculating the hydroxide ion concentration and may also involve calculating a value for pOH, pH, and/or percent reaction-the first step may be determining the Kb value by using a table of Ka values-the form of the calculation then follows the same format as the equivalent type of question for weal acids, with a similar assumption made, and the same restrictions

-assume that the initial weak base aqueous concentration decreases so little that it is numerically unchanged at equilibrium, but only if the initial amount concentration of the weak base is at least 1000 times greater than its Kb value

Read sample problem 16.5 – shows the format for calculations of this type, and an example of each conversion. Recall that in aqueous solution, the amphiprotic ion HCO3

-(aq) reacts with water as a base to a very much greater extent than it does as an acid – assume that the solution is basic because the ion reacts as a base, and ignore its negligible reaction extent as an acid

The Effect of Amphoteric Entities:-if an entity can react as either a Bronsted-Lowry acid or base, how do you know which will be the predominant reaction in its aqueous solution?-we calculate it out by taking the Ka value for the entity from the table of acids and bases, calculate the Kb value for the entity, and compare the numbers-the higher of the K values tells you which reaction predominates and whether the entity reacts as an acid or as a base in aqueous solution

Read Communication example 7



Section 16.4: Interpreting pH Curves

-as you learned last year, a graph showing the continuous change of pH during an acid-base titration, continued until the titrant is in great excess is a pH curve or titration curve-the pH values and changes provide important information about the nature of acids and bases, the properties of conjugate acid-base pairs and indicators, and the stoichiometric relationships in acid-base reactions

Buffering Regions, Endpoints, and Indicators:-you learned last year that a common reason for plotting a pH curve is to determine the value of the solution’s pH when an acid-base reaction reaches its equivalence point -we can use this information to select indicators with an easily seen endpoint and use the acid-base reaction as a titration to that endpoint for titration analysis-in titration analysis, the critical measured quantity is the titrant volume required to reach the titration endpoint – the value needs to be as close as possible to the theoretically exact value needed to reach the reaction equivalence point-any variation between endpoint values and equivalence point values are the titration error-remember:

-endpoint refers to that point in a titration analysis where the addition of titrant is stopped. It is defined empirically by the observed colour change of an indicator

-equivalence point refers to that point in any chemical reaction where chemically equivalent amounts of the reactants have been combined. The equivalence point is defined theoretically by the stoichiometric ratio from the reaction equation-endpoints are easily detectable because pH changes a great deal and very abruptly as the reaction solution changes from an excess of acid to an excess of base (or vice versa)-read the learning tip

-figure 2 shows a titration of NaOH with HCl, which is used as an exemplar for explaining pH curves:-there are three regions of interest: the pH first changes very slowly, then very rapidly, and then very slowly again as the titrant is added-notice that the addition of acid has very little effect on the pH – the high pH hasn’t changed much even after enough acid has been added to react with 90% of the base-this nearly level region on a pH curve is the buffering region. Buffering is the property of some solutions of resisting any significant change in pH when an acid or base is added-in this case, buffering occurs because any strong acid added immediately reacts with excess hydroxide ions – before any acid is added, the sample solution is primarily water molecules and hydroxide ions -as long as a large amount of hydroxide ions is present, any added acid is immediately converted to water, producing a solution that still consists primarily of water molecules and hydroxide ions-the hydroxide ion concentration decreases a little, which only affects the pH slightly


-this leveling effect finally fails near the equivalence point, when the hydroxide ions in the solution become almost completely consumed – until then the solution maintains a nearly constant pH-this buffering property is critically important for reactions in applied chemistry and biology-once excess acid has been added, the solution consists primarily of water molecules and hydronium ions -now the pH remains stable because adding more hydronium ions does not change the nature of the solution, it just slightly increases the hydronium ion concentration slightly-another buffering region is established, now at a low pH-following the first buffering region, there is a rapid change in pH for a very small additional volume of the titrant-the inflection point on the pH curve represents the equivalence point of the reaction: in this case, theoretically we know it is 7-this reaction can be used for titration analysis if an indicator is selected that changes colour in solution at a pH very close to 7-figure 3 shows that bromothymol blue has a colour change pH range with a middle value very close to pH 7 making it an appropriate indicator-the values below show that the theoretical titrant volume required for complete reaction as predicted by calculation, should be 24mL – see calculations on p. 752-figure 3 shows that if the colour change of bromothymol blue to a green intermediate colour is the endpoint, then the titrant volume actually measured at the endpoint of 24mL is negligibly different from the theoretical predicted value – giving very accurate results

Acid-Base Indicator Equilibrium:-any acid-base indicator is really two entities for which we use the same name: a Bronsted-Lowry conjugate acid-base pair-at least one or both of the entities is visibly coloured, so you can tell simply when it forms or is consumed in a reaction-example: phenolphthalein (figure 4) – conjugate base form is bright red and the conjugate acid form is colourless-common indicators such as methyl red (figure 5) have complex molecular formulas, so shorthand is used to identify them-acid form of phenolphthalein as HPh and the bright red base form as Ph-

-when an equilibrium is shifted to the point where equal concentrations of both forms exist, an intermediate colour may be observed-example: equal concentrations of the blue and yellow forms of bromothymol blue (pH of 6.8) appear green to the human eye-an indicator is a conjugate weak acid-weak base pair formed when an indicator dye dissolves in water -using HIn(aq) represent the acid form and In-(aq) to represent the base form of any indicator – an equilibrium equation can be written – see in text-according to Le Chatelier’s principle, an increase in the hydronium ion concentration with shift the equilibrium to the left-more indicator will change to the colour of the acid form (HIn(aq))


-this happens when litmus is added to an acidic solution-in basic solutions the hydroxide ions remove hydronium ions by reacting with them to form water causing the equilibrium to shift to the right and the base colour of the indicator (In-) predominates-since different indicators have different acid strengths the acidity or pH of the solution at which an indicator changes colour varies (figures 5 and 6), which are measured and reported in the acid-base indicators table-read learning tip


Polyprotic Entities and Sequential Reactions:-polyprotic acids can lose more than one proton and polyprotic bases can gain more than one proton, in Bronsted-Lowry transfers-if more than one proton transfers occur, chemists believe the process occurs as a series of single-proton transfer reactions-typical pH curves for reactions of diprotic or triprotic acids and bases differ from those of monoprotic acids and bases-example: figure 7 – the pH curve for the addition of HCl(aq) to Na2CO3(aq), shows two equivalence points or two significant changes in pH, which are explained by two different proton transfer equations

-first protons transfer from hydronium ions to carbonate ions – the strongest base present in the initial sample. The first drop in pH is evidence that the first reaction is quantitative – look in text at the Bronsted-Lowry equation with the acid and base entities present in the solution and the H3O+(aq) that is being added

-since the first reaction is quantitative, essentially all the carbonate ions will be consumed at the equivalence point and an equal amount of hydrogen carbonate ions will form

-then in a second reaction, protons transfer from the conditionally added hydronium ions to the hydrogen carbonate ions – see equation in text

-since the second reaction is quantitative all of the hydrogen carbonate ions have reacted at the equivalence point and only newly formed carbonic acid molecules remain in the solution

-adding more hydronium ions will cause no significant further change as the strongest base in the solution is now water-the carbonate ion is a diprotic base because it can accept a total of two protons-other polyprotic bases are sulfide ions and phosphate ions – see conjugate entities in text-other polyprotic acids are oxalic acid and phosphoric acid – see conjugate entities in text-evidence shows that for every proton transferred by a polyprotic entity the strength of the new acid or base entity formed greatly decreases – explained by electric charge and electrostatic attraction: it should be easier to pull a positively charged proton away from a neutral molecule than it is to pull one away from a negatively charged ion

-Figure 9 shows the pH curve for phosphoric acid titrated with sodium hydroxide


-phosphoric acid is triprotic – three reactions are possible, but note that only two equivalence points are evident

-at the first equivalence point, equal amounts of H3PO4(aq) and OH-(aq) are in solution. As the last of the H3PO4(aq) reacts, the pH rises because as the H3PO4(aq) is used up only newly formed H2PO4

-(aq) is formed and is a much weaker acid – look in the text to see how to write this reaction

-the second equivalence point corresponds to the completion of the reaction of H2PO4-

(aq) with additional OH-(aq) solution. The pH raises because as the H2PO4-(aq) reacted, HPO4

2-

(aq) is left, which is a weaker acid – look in the text for the equation-no pH endpoint is apparent for the possible reaction of HPO4

2-(aq) with additional OH-

(aq) because the table of acids and bases show that the hydrogen phosphate ion is an extremely weak acid that doesn’t quantitatively lose protons to OH-(aq). Instead an equilibrium is established that shifts right as more base is added and there is no pH increase at a third equivalence point because the reaction never goes to completion – look in the text for the equation-as a general rule, only quantitative reactions produce detectable equivalence points in an acid-base titration-sulfuric acid is a unique polyprotic acid because it is the only common one where the first proton loss is already quantitative in aqueous solution – meaning it is the only strong acid that is polyprotic – as its ionization forms hydrogen sulfate ions as well as hydronium ions

-the hydrogen sulfate ion is a weaker base than water, so it cannot act as a base in aqueous solution – as an acid, it is one of the stronger weak acidsand usually reacts quantitatively with bases

-the sulfuric acid will usually react with a base (assuming the base is in excess) in two complete proton transfer reactions - originally you assumed that both protons were transferred, but couldn’t explain how – see equation in text

-both neutralization reactions may not be quantitative when a weaker base is involved and stoichiometric calculation can only be done for a reaction that is quantitative-there are almost no overall reactions of common polyprotic acids or bases that have more than two definite endpoints-an indicator can be selected so that the titration can be stopped at either chosen equivalence point-for the titration of sodium carbonate solution with hydrochloric acid this would include titrating through two sequential quantitative reactions to the second equivalence point (figure 7)-for this titration, the second equivalence point involves a greater pH change and is easier to detect – using methyl orange-for stoichiometric calculation, sequential quantitative Bronsted-Lowry reaction equations can be combined into a single reaction equation – see in text-either way the stoichiometric ration is 2:1-since chemists usually use the simplest representation, so Bronsted-Lowry equations are more often used for correctly predicting products and equilibria, whereas standard chemical formula equations are more often used for stoichiometry in titration analyses


Read the Summary: pH Curve Reaction InformationComplete Practice questions #5-10

pH Curve Shape versus Acid and Base Strength:-a reaction with the strongest base (hydroxide) and strongest acid (hydronium) is quantitative and always results in an equivalence point pH of 7 at 25C-some weak acids react quantitatively with hydroxide (figure 12)-some weak bases react quantitatively with hydronium (figure 13)-the strongest of weak acids also react quantitatively with the strongest of the weak bases-example: the reaction of HSO4

-(aq) with CO32- (aq) is quantitative

-the farther apart the reacting acid and base are on the Relative Strength of Aqueous Acids and Bases table (the larger the difference of Ka), the more likely the reaction will be quantitative-after plotting numerous pH curves, the general pH curve reference is shown in figure 14-read learning tip-the stoichiometric equivalence point pH for a strong acid-strong base reaction is 7-the equivalence point pH of any quantitative strong acid-weak base reaction is less than 7-the equivalence point pH of any quantitative weak acid-strong base reaction is greater than 7-any reactions between weak acids and weak bases cannot be predicted using figure 14 because they often do not have a detectable equivalence point as they are usually not quantitative – this is why analysis and stoichiometric calculations are usually not done for these reactions-the entities present at an equivalence point can be predicted by either following the five-step Bronsted-Lowry method or simply by converting a chemical substance formula equation into a net ionic equation-each case of common quantitative acid-base reaction types is examined and the equivalence point pH explained in text: read strong acid-strong base reaction, strong acid-weak base reaction, learning tip, and weak acid-strong base reaction-read learning tip and summary Titration Generalizations


pH Curve Buffering Regions and Buffer Solutions:-a titration pH curve for any quantitative acid-base reaction shows at least one region where buffering action occurs between the beginning of the titration and the equivalence point-if the titration is continued much past the equivalence point, it enters another buffering region because a partial reaction has occurred and contains significant amounts of both entities of a conjugate acid-base pair-the term buffer refers specifically to the combination of any weak acid with its conjugate base, in the same solution-buffer solutions have a specific pH and change pH very little when other acids or bases are added, if the quantity added is much less than the buffer-read the learning tip-example: acetic acid titrated with sodium hydroxide (figure 15)


-solution pH is 4.8 when 10mL of hydroxide solution is added (one half the equivalence point volume) and one half of the acid will have reacted (see the equation in text)

-the mixture contains equal amounts of the unreacted acid and the conjugate base – this is referred to as an acetic acid-acetate ion buffer-this buffering example can be explained using Bronsted-Lowry equations and the five-step method for predicting the predominant acid-base reaction – see equation in text-figure 15 shows this reaction as being quantitative -a small amount of hydroxide would convert a small amount of acetic acid to acetate ions causing a small decrease in the ratio of acid to acetate ions in the buffer and a slight increase in the pH-the small change in concentration of the acid-base conjugate pair and the complete consumption of the hydroxide explains the small pH change-this buffer would work equally well if a small amount of a strong acid, such as HCl(aq) were added – the quantitative equation is shown in the text-added hydronium ion is consumed and the mixture has a slightly higher ratio of acetic acid to acetate ions and a slightly lower pH-figure 16 demonstrates buffer capacity – the limit of the ability of a buffer to maintain a pH level-when the entity of the conjugate acid-base pair that reacts with an added reagent is completely consumed, the buffering fails and the pH changes dramatically-buffering is important to numerous biological processes: aspects of cell functions and metabolism in living organisms are sensitive to pH changes (example – enzymes), cellular fluid, blood plasma (table 1)-buffering is important in many consumer, commercial, and industrial applications: fermentation and manufacturing antibiotics, production of various cheeses, yogurt, and sour cream, and food preservations-an example of how a buffer with a pH of 10, which can be used to calibrate pH meters for accuracy, can be made is: mix 50mL of 0.050 mol/L NaHCO3(aq) with 10.7mL of 0.10 mol/LNaOH(aq) – see the equation on text-this reaction converts 3/5 of the initial hydrogen carbonate ions into carbonate ions-an aqueous mixture of a 3:2 ratio of these two acid-base conjugates as the specified concentrations has been found empirically to have an equilibrium pH of 10.00

Complete practice questions #16-21Read Summary: Bronsted-Lowry is a Unifying Concept


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