Lewis Theory. Arrhenius Theory

Transcription

1 d/base Theories Lewis Theory Acids are electron pair acceptors ases are electron pair donol Acids are proton-donors Bases are proton-acceptors * Arrhenius Theory Acids ionize in water and produce hydrogen ions Bases dissociate in water and produce hydroxide ions

2 Writing and Naming Acids 1. Complete the following table. Binary Acid Ionization Strong Acid Acid Name or Constant at or Formula Ternary Weak 1. H2SO4 2. Hypof luorous acid 3. CHSCOOH 4. HCI 5. Chloric acid 6. HCIO4 7. Hydrofluoric acid 8. Boric acid (B0 3 3-) 9. HIO3 10. H Carbonic acid 12. Sulfurous acid 13. H3PO4 14. HCIO 15. Chlorous acid 16. Hydroiodic acid 17. HBr 18. H 2 S 2. Write the balanced ionization equations for acids 1 to 5.

3 ph Stoichiometry: Strong Acids & Bases Calculate the [HteO*], ph, poh, and [OH-] for a solution of mol/l of hydroiodic acid. Consult tables for K a. =>

4 ph Stoichiometry: Strong Acids & Bases 26 g of sodium hydroxide in 1 50 ml of solution unclogs a drain. What is the [OH - ], ph, and poh of the solution? Should you wear gloves? Why?

5 Percent Ionization of Acids Percent ionization is the extent to which a weak acid separates into its ions in aqueous solution. [H ]eguilibnum ^ i Q Q % [HA]initial Sample % ionization problem Calculate the percent ionization of mol/l phosphoric acid solution at 25 C. For polyprotic acids, always use K a i. 23%

6 Percent Ionization of Acids A mol/l solution of methanoic acid has a ph of Calculate the percent ionization of the acid, (formula HCOOH) r

7 Percent Ionization of Acids Calculate the acid ionization constant, K a, of mol/l acetic acid with an ionization of 1.30%. (CH3COOH) 1.7 x 10-5 mt

8 Neutralization Reactions When solutions of acids and bases are combined they react according to a neutralization reaction.

9 Write the balanced neutralization equation for the following reactions: 1. Sulfuric acid + lithium hydroxide 2. Phosphoric acid + calcium hydroxide <=> 3. Nitric acid + strontium hydroxide 4. Boric acid + aluminum hydroxide

10

11 Strong Acid with Strong Base Titrations Titration is a technique designed to accurately determine stoichiometric values when a substance with a known concentration and volume reacts with another substance. Changes in ph may be monitored throughout the titration and a graph of ph versus volume of titrant plotted. Calculate the ph values every 1.00 ml when ml of M HCI is titrated up to ml with M sodium hydroxide solution.

12 Assignment: Date: From: To: Page No. Form 4A-BW Mathematics Help Central mathematicshelpcentral. com

13 Equivalence is the point in a neutralization reaction where the moles of acid are stoichiometrically equal to the moles of base. In the reaction of a strong acid with a strong base, ph at equivalence is 7.0. Calculate the ph of the solution at 1, 2, and 3 ml after equivalence.

14 ph Titration Question Calculate the ph at each of the following points in a titration of 0.05 M lithium hydroxide with ml of M hydrochloric acid: a) before any base titrant has been added b) after 5.00 ml of base added c) after ml of base added d) after ml of base added

15 ml of M sodium hydroxide is titrated to an endpoint with M sulfuric acid. a) What is the volume of acid at neutralization? Calculate the ph after: b) 4 ml of acid is added c) ml of acid is added

16 Calculate the concentration of acetic acid in a ml sample of vinegar if ml of 0.10 M sodium hydroxide completely neutralizes the sample.

17 Appendix 4.5.e Titration Curves Teacher Support Material Plotting the ph of the solution during an acid/base titration generates a "titration curve". The general shape of the curves generated in a series of titrations may be grouped into families, according to the solution titrated and the titrating solutions. Typical examples of the general shapes are illustrated for four classes of titration: I: A solution of a strong acid titrated with a solution of a strong base. In this example 25.0 ml of mol I. ' aqueous solution of hydrochloric acid, HC1, (a monoprotic strong acid) is titrated with a mol Lr 1 hydroxide). The equation representing the reaction may be written as: HCl(aq) + NaOH(aq) NaCl(aq) + H 2 0(1) aqueous solution of sodium hydroxide, NaOH(aq) (an ionic or as the net ionic equation: H}O f (aq) + OH (aq) > 2 lh()(l) At equivalence, moles HC1 originally present = moles NaOH added Volume MCI original x Concentration HC1 solution = Volume NaOH added x Concentration NaOH solution At equivalence, [H^O' ] = [OH"']; [Na + J = [Crj; and since Na" (aq) is a weaker acid than water and CI a weaker base than water, the solution is described as "neutral". At 25 C, the ph will be 7.0 (under the usual set of assumptions). The expected titration curve is shown in figure 1 below. is ph «~ equivalence ^ ) ph 10 - equivalence > 4 2 } ^equivalence ^ i i i i i ' i i i i i i Volume NaOH(aq) /ml Figure 1: 0.1 M HCl(aq) vs. 0.1 M NaOH(aq) Volume HCI(aq) /ml Figure 2: 0.1 M NaOH(aq) vs. 0.1 M HCl(aq) 40

18 Appendix 4.5.c Titration Curves (Contd.) Teacher Support Material II. A solution of a strong base with a solution of a strong acid. This is analogous to the strong acid/strong base titration except that the acid is the independent variable (i.e., is added from a burette or equivalent). The example chosen is the titration of 25.0 ml of a mol/l aqueous solution of sodium hydroxide with an mol/l solution of hydrochloric acid. The expected titration curve is shown in figure 2 above. For both cases I and II, the end point (assumed the equivalence point) is found at the steepest part of the curve, the inflection point where the curve changes direction. III. A solution of a weak acid with a solution of a strong base In this example 25.0 ml of mol I. ' aqueous solution of acetic acid, HCH3CO2, (a monoprotic weak acid with A' a = 1.8 x 10^5 mol L" 1 ) is titrated with a mol L~~' aqueous solution of sodium hydroxide, NaOH(aq) (an ionic hydroxide). The equation representing the reaction may be written as: HCH.C0 2 (aq) + NaOH(aq) -> NaCH 3 C0 2 (aq) + FI 2 0(1) or as the net ionic equation: HCH 3 C0 2 (aq) + OH - (aq) -> CH 3 C0 2 ~(aq) + H 2 0(1) (For a weak acid, the acid of highest concentration in an aqueous solution is the undissociated acid, not the hydronium ion.) At equivalence, moles HCH 3 C0 2 originally present = moles NaOH added At equivalence, [HCIl 3 C0 2 j = [OH ]; [Na + ] = [CII 3 C0 2 ); and since Na + (aq) is a weaker acid than water while CT L,C0 2 is a stronger base than water, the solution will be basic. At 25 C, the ph will be greater than 7.0 (under the usual set of assumptions). The expected titration curve is shown in figure 3 below \- excess NaOH ph 10 h 8 h ph equivalence PH ph = pk Volume NaOH(aq) /ml Volume NaOH(aq) /ml 40 Figure 3: 0.1 M HOAc(aq) vs. 0.1 M NaOH(aq) Figure 3(b) The end point (assumed the equivalence point) is found at the steepest part of the curve, the inflection point where the curve changes direction.

19 Titration Curves Student Activity ph J I LJ I L I I I I I I I J I I I Volume NaOH(aq) /ml 40 A student pipetted 25.0 ml of an aqueous solution of an unknown acid into a conical flask, added 25 ml. of water, and then titrated the resulting mixture with a standard mol L ' solution of aqueous sodium hydroxide, measuring the ph of the mixture after each addition. The above graph shows the titration curve obtained. Using this graph, answer the following questions: 1. What was the concentration of the unknown acid solution? 2. What would be a suitable indicator for this solution? 3. Assuming the unknown acid is a weak monoprotic acid, estimate its AT a value.

20 Appendix 4.5.f Titration Curves (Contd.) Student Activity 14 ph 2 I I I I I I I I I I I I I I I I I I I I I I I I I I I I I I I Volume NaOH(aq) /ml A student weighed g of a solid unknown acid into a conical flask, added about 50 ml. of water, and then titrated the resulting mixture with a standard mol L ' solution of aqueous sodium hydroxide, measuring the ph of the mixture after each addition. The above graph shows the titration curve obtained. Using this graph, answer the following questions: 1. What would be a suitable indicator for this solution? 2. Assuming the unknown acid is a weak monoprotic acid, estimate its A" a value. 3. Estimate the molar mass of the unknown acid. 4. Would the results obtained have been different if she had added 100 ml. of water'.' Explain your answer.

21 Appendix 4.5.f Titration Curves (Contd.) Student Activity ph ll l l I I I I I l I I t l I I l I I I I I I I I I 1 I l l I l l l l I Volume HCI(aq) /ml A student weighed g of a solid unknown base into a conical flask, added about 100 ml. of water,, and then titrated the resulting mixture with a standard mol Lr' solution of hydrochloric acid, measuring the ph of the mixture after each addition. The above graph shows the titration curve obtained. Using this graph, answer the following questions: 1. What would be a suitable indicator for this solution? 2. Assuming the unknown base is a weak monoprotic base, estimate its value. Estimate the A'., of its conjugate acid. 3. Estimate the molar mass of the unknown base.

22

23 CExp2: VOLUMETRIC ANALYSE - ACETIC ACID IN VINEGAR Carefully read the text describing this experiment then complete the pre-laboratory exercise (c pages 41 and 42) to prepare for the laboratory experiment This experiment focuses on techniques and must he done individually. For accurate results you must measure carefully to the limits of the equipment used and employ only clean glassware. Your mark will depend on technique and accuracy of results. Be sure to follow till safety regulations. OBJECTIVES; A. To prepare a primary acid standard (potassium hydrogen phthalate) and use it to standardize a solution of sodium hydroxide. INTRODUCTION B. To determine the concentration of acetic acid in a sample of vinegar by titration with the standardized sodium hydroxide solution. Acid-base Stoichiometry Any acid will react with any base. Provided either the acid or the base is strong the reaction goes to completion. Thus, the strong base sodium hydroxide dissociates completely to form hydroxide ions which react completely with both strong and weak acids, according to the equations: OH-(aq) + H+(aq) -> H2CKI) [strong acid] OH "(aq) + HA(aq) <=> H2O (1) + A _ (aq) [weak monoprotic acid] 20H-(aq) + TfeUaq) o> 2 H2O (1) + I>(aq) [weak diprotic acid] From the balanced equations, you can see that 1 mol of hydroxide ion reacts with 1 mol of hydrogen ion; 1 mol of hydroxide ion reacts with 1 mol of weak monoprotic acid; 2 mol of hydroxide ion are required to neutralize 1 mol of weak diprotic acid. Indicators The point of exact neutralization of acid by base is determined by using an indicator that is added to the reaction mixture. An acid-base indicator can usually exist in two forms (acidic, HIn, and basic, In - ), and there is an equilibrium between the two states which depends on the acidity of the solution. For example: HIn o In- + H + (acidic) (basic) For a useful indicator, the acidic and basic forms of the indicator are two distinctly different colors. In an acidic medium, the indicator exists as HIn and has its specific color. In a basic medium, the indicator exists as In - and has a different color. Thus, when an indicator color change occurs in a titration, you know that a change from an acidic to basic medium (or vice versa) had undergone in the solution itself meaning that the analyte had reacted. This is the end-point that you are looking for. While the choice of indicators can be important, only a trace amount of indicator solution is used and it does not appreciably affect the volume of base needed to reach the endpoint in the titration. Primary Standards Volumetric analysis depends upon the ability of the chemist to prepare solutions of an exactly known concentration. chemical substance suitable for use as a primary standard must satisfy the following requirements: A i) It must be readily available in an extremely pure form (99% pure or greater) ii) iii) It must be stable under normal conditions of storage and use. It must be reasonably soluble (usually in water). These requirements allow a primary standard solution to be prepared by weighing out the substance and dissolving it in water. In addition, it is desirable to use a substance with a high molar mass (molecular weight) to minimize weighing Chemistry Laboratory CExp2 page 33

24 errors. In practice, very few chemicals meet these requirements. NaOH CANNOT be used as a primary standard because it cannot be obtained in a pure form, and it is not stable (it quickly absorbs water from the air). One of the most useful primary standards for volumetric analysis of acids and bases is potassium hydrogen phthalate, KHC8H4O4, often abbreviated as KHP. This compound is stable and non-hygroscopic. Its molar mass is g/mol and it is available 99.9% pure in the anhydrous form. This acidic salt is monoprotic and reacts stoichiometrically in a 1:1 ratio with the sodium hydroxide solution used in this analysis. OCOOH COO " K + Use of.glassware Glassware should be clean before undertaking an analysis (see cleaning procedures in the Preface). When using a (see below), it is important to check that the stopcock is free flowing but does not leak. Test your buret with distilled water after it has been cleaned. Wash titration flasks with warm water then rinse with distilled water before repeating trials. Titration When it is necessary to find the exact amount of an acid or base, with at least one of them in solution, the process of titration is often used. In a titration you add a specific volume of one reagent to either a known volume or a known mass of the second reagent. Then, from the balanced equation, you can determine the amount (number of moles) of reagent that must be present for complete reaction, and hence either its concentration or its molar mass. A titration depends upon the proper preparation and use of a buret. After you fill the buret and make sure that it is operating properly (Figure 1) you must record your initial volume (Figure 2). buret Burets Figure 1: Preparing buret for Titration a. Filling Burets are long graduated cylindrical tubes of very uniform bore, with some device to control the flow of liquid (a stopcock). Burets are calibrated to deliver the measured volume of liquid starting with the tip full of liquid. Most common laboratory burets have, a total capacity of 50 ml, with graduations every 0.1 ml. If the buret is dirty so that water does not run cleanly down its sides but leaves droplets behind, scrub it with detergent and water. Rinse the buret thoroughly two or three times with tap water and then twice with distilled water. Do not forget to run water through the buret tip. Then drain the buret tip thoroughly. Use a funnel to fill the buret with about 10 ml of the solution to be dispensed from it. Tilt and rotate the buret tc rinse down the buret walls with the solution. Run some of the solution through the buret tip. Then drain the buret and tip thoroughly, discarding the rinsing solution. Repeat this whole procedure a second time. Be sure to have a freely turning stopcock for controlled delivery. Clamp the buret on its stand, using the special buret clamps. Fill the buret above the zero mark with the solution and run some rapidly through the buret tip so that no air bubbles remain in the tip. Run the solution from the buret until the solution meniscus is slightly below, the zero mark. After allowing 20 seconds for the solution level to stabilize, read and immediately record the initial buret level. To facilitate reading, prepare a white card with a dark area and place it behind the buret. The reflection of the dark portion in the meniscus makes the meniscus stand out more clearly, and a reading to 0.01 ml can be estimated. When reading the buret, make sure that the meniscus is at eye level. b. check for proper delivery be sure there is no air bubble in tip Now place the flask containing the measured amount of material to be titrated and an end-point indicator under the buret tip so that the tip is within the volume of the flask under operating conditions (Figure 3). Add a few drops of an endpoint indicator then begin to titrate by swirling the flask and its contents while adding titrant. Note the hand positions shown in Figure 3. Generally, the color change characteristic of the indicator will disperse rapidly at first. Continue adding solution from the buret fairly quickly until the color begins to disperse more slowly. At this point make a mental Chemistry Laboratory CExp2 P a g e 34

25 note of the approximate volume and begin to add solution from the buret drop by drop. Continue this slow addition while swirling the flask until one drop (or fraction of a drop) causes a permanent color change of the indicator. Wait about 30 seconds to be sure that the indicator color is permanent and the buret level has stabilized. Record the final buret reading. (Note: burets generally deliver about 20 drops/ml. Since titrations are usually uncertain to one drop, your precision on a titration will be of the order of 0.05 ml. Be sure to record all buret readings to two decimal places.) Figure 3: How to control the flow of solution from the buret tip. The procedure illustrated is for a right-handed person. Put your left hand on the stopcock part of the buret as illustrated. For most people it is easiest to place your thumb in front and your fingers behind. The tension on the stopcock shoud be such that it can be turned easily but does not leak. A simple flick movement will place the hole in line with the barrel so that solution can flow. It is possible to control this so drops come individually or in a stream. With your thumb and forefinger around the handle you will always be applying an inward pressure to keep the stopcock seated and avoid leaking At the same time your right hand should control the flask. Swirl it continuously as solution is added from the buret. In this way the solution is always mixed so that there is less chance that you will miss a color change and end-point. For a left-handed person; the right hand is on the buret and the left hand is used to swirl the flask. NOTE: There is information about titration procedures on the Chemistry U of M web site, specifically in the Laboratory Section. Check this and additional sources before coming to the lab. PROCEDURE: A. Standardization of a Sodium Hydroxide Solution Chemistry Laboratory CExp2 page 35

26 The solutions of sodium hydroxide provided for this experiment will be approximately 0.1 mol/l; you will have to standardize the solution to find its exact concentration. Fill about % of a plastic 14-teaspoon with potassium hydrogen phthalate (KHP) and transfer all of the crystals in your plastic weighing bottle. Take this sample and two labeled 125-mL flasks to the balance room with your data sheet. Record (to the nearest g) the combined mass of the sample and weighing bottle. While gently rotating the weighing bottle, tip it over one of the flasks in order to transfer about one-half of the crystals into it (this should correspond to a transfer of about 0.5 g of solid). Reweigh the remaining sample and weighing bottle. The difference in mass is the sample size for the first titration. Transfer the remaining crystals into the second flask. Reweigh the weighing bottle and, by difference, deterrrtine the mass of the second sample. Obtain about 200 ml of the 0.1 M sodium hydroxide solution in a clean 250 ml Erlenmeyer flask and stopper it. Rinse your previously cleaned buret with a small portion (about ml) of the sodium hydroxide solution then discard the rinsing liquid. Now fill the buret with this solution, eliminating any air bubbles from the tip after filling. Be sure the solution flows evenly from the buret. Wait 20 seconds for the liquid level to stabilize slightly below the zero mark, then read and record the initial volume (V,) of NaOH solution on your data sheet (your reading should be to the nearest 0.01 ml, i.e. to 2 decimal places). Add about 25 ml of water and 2 or 3 drops of phenolphthalein* indicator to one of your samples of potassium hydrogen phthalate in the 125 ml Erlenmeyer flask. Although all of the solid may not dissolve immediately it will go into solution as the titration proceeds. Add sodium hydroxide solution from the buret while swirling the flask and contents. Watch for localized pink coloration. At the endpoint, one drop of base will cause the color change from colorless to pink, so proceed carefully. Once the end point is reached, wait 20 seconds for the level of solution in the buret to stabilize, then read and record the final volume (V f ) of NaOH solution. * Phenolphthalein has been found to be the best indicator for most students; the color change is from colorless (in acid) to pink (in basic solution). The endpoint is signaled by the first pale pink coloration that is reasonably permanent. If you cannot see this color change, consult your T.A. for other suggestions of indicators you might use. Repeat the titration with your second sample. If the results are not essentially the same, titrate a third sample, and if necessary, more until constant results are obtained. [While your results should not differ by more than 2% to be acceptable, good work will be even better. A quick check on the precision may be obtained by dividing the volume of base titrated by the sample size; the results should be within 2% of each other.] Chemistry Laboratory CExp2 page 36

27 B. Acid Analysis in Vinegar Obtain between 10 and 20 ml of the vinegar your T.A. has assigned you. Using the mL Class A transfer pipet provided, transfer ml of the vinegar into a clean 100 ml volumetric flask. Record the code of your vinegar sample on your data sheet. Mix the solution as you fill the mL volumetric flask exactly to the calibration line with distilled water. Invert the closed flask many times to mix thoroughly. Using a mL transfer pipet, transfer an aliquot of this diluted vinegar solution into a labeled 125 ml Erlenmeyer flasks. Add 2 or 3 drops of indicator solution and titrate with your standardized sodium hydroxide solution. The end-point is again determined by a persistent color change upon the addition of one drop of the basic solution. Record all volume readings (V, and V t ) to the nearest 0.01 ml. Repeat the titration of this vinegar solution until you have three acceptable trials. MEASUREMENTS, NOTES, and OBSERVATIONS Chemistry Laboratory CExp2 page 37

28 TABLE B. DETERMINATION OF ACETIC ACID IN VINEGAR Standard [NaOH] (from Part A) Brand (or code) of vinegar used: Titration: Indicator used (2) TRIAL 3 (if needed) 4 (if needed) Volume of dilute vinegar titrated (ml) Final volume reading on buret, V. (ml) Initial volume reading on buret, V, (ml) ' Volume NaOH solution used (ml.) CH,CO : H] in dilute sample (mol/l) [CFLCO-.H] in vinegar sample (mol/l) "'<> Acetic Acid bv volume in vinegar (*) (*) Report acetic acid concentration in vinegar as % acetic acid by volume, as it is done in the marketplace. To do so, assume that densities of acetic acid, vinegar solution, and water are all the same at 1.00 g/ml. V(CHXO?H) 3 2 A concentration expressed in percent by volume (% % ) is defined as _ n / ' - V(Vinegar) t x 100% Sample calculation for concentration of acetic acid in vinegar (in mol/l and % % ): (2) Average % acetic acid by volume in vinegar: ± % (2) Observations: (1) DISCUSSION and CONCLUSION (attach another page for further explanations) (2) Chemistry Laboratory CExp2 page 40

29 Problem Set m: Aqueous Equilibrium 1. Name the following acids, write their ionization equations, and write the Ka expression for each. (2 points) a) HCIO b) HF c) HCN d) HBr 2. Order the adds in question 1 from weakest to strongest. (0.5 point) 3. Which acid in question 1 is the strongest electrolyte? The weakest electrolyte? (0.5 point) 4. Calculate the equilibrium concentrations of all species given a 0.5 M hydrocyanic acid solution at 25 C. Use tables for IC This is an ICE table question (2 points) f jj 5. Calculate the ph, [H + ], [H ], poh, [OH"], and percent ionization of a 0.1 M nitrous acid solution. Use tables for Ka. (3 points) 6. Given a ph of 3.13 for a 0.15 M weak acid solution, HA, calculate the acid ionization constant, K a. (2 points) 7. Calculate the ph of a 0.1AA hypochlorous acid and 0.2 M sodium hypochlorite solution. Use tables for Ka. (2 points) ( /12

30

31 T.A. Name: Day /Section: Room / Locker: Date: CExp2: VOLUMETRIC ANALYSIS - ACETIC ACID IN VINEGAR DATA and RESULTS: Record information on Tables A and B or in the space provided. TABLE A. PREPARATION OF STANDARD NaOH SOLUTION (marks) Titration: NaOH Sample # Indicator used Mass of weighing bottle with KHCgH 4 04 Mass of bottle minus sample: Mass of KHC 8 H used: (1) TRIAL (if needed) 4 (if needed) Final volume reading on buret, V, (ml) Initial volume reading on buret, V. (ml.) Volume NaOH solution used (ml) Concentration NaOH solution (mol/l) Average [NaOH] (± average deviation): ± mol/l (2) Sample calculation for concentration of NaOH solution: Observations: Chemistry Laboratory CExp2 page 39

32

33 T.A. Name: Day / Group: Room, / Locker: Date: CExp2: Pre-Lab Exercise for Volumetric Analysis Experiment To be submitted to your T.A. before you begin the experiment. It will be returned with your report which must be submitted in one week for full credit. Late labs will be penalized. 1. Potassium hydrogen phthalate (KHP) is a weak monoprotic acid. a) How many moles of NaOH are required to neutralize g of KHP? (1) b) If the neutralization of g of KHP requires ml of NaOH solution, what is the concentration of the NaOH solution? (1) 2. It is suggested that phenolphthalein is the best indicator for this titration. a) What color would the solution be if you added phenolphthalein to a solution that was: (0.5) - acidic? - basic? b) Why is the first indication of a pink color not necessarily the true end-point? (0.5) Chemistry Laboratory CExp2 page 41

34 3. What is the difference between a primary and a secondary standard solution? In this experiment, is the sodium hydroxide solution provided a primary or a secondary standard solution? Explain. (1) 4. If the acetic acid content of wine vinegar is measured as 1.00 mol/l, what weight of acetic acid would there be in 1 litre of the vinegar solution? Assume that densities of acetic acid, water, and wine vinegar are all the same at 1.00 g/ml. What is the percent by volume (% % ) of acetic acid in wine vinegar (1) 5. A student added NaOH solution until the phenolphthalein indicator was an intense pink color (i.e. they overshot the end-point). Explain what effect this would have on the concentration of acetic acid reported for the vinegar sample. a ) Chemistry Laboratory CExp2 page 42

35 BLA BLB Theory & Titrations Bronsted Lowry Acids and Bases Arrhenius acids and bases are classified as containing H + ions and OH" ions in water, respectively. According to Arrhenius, the identities of acids and bases are static and do t change. H + ions also exist as H \ The Arrhenius definition is impractical as it does not account for: ^ Q + Q n cj c y c n 1. the existence of acids and bases in non-aqueous solutions H O + 2. the fact that H + ions exist as rlo + 3. the observed behaviour of some chemicals that act either as acids or bases called amphoteric or amphiprotic substances depending on the chemicals with which they are combined. A broader definition of acids and bases is required to account for the limitations of Arrhenius theory. A Bronsted Lowry Acids (BLA) is a substance capable of donating protons, also called a proton donor. Note that a hydrogen ion, h +, is simply a proton! Bronsted Lowry Bases (BLB) are described as substances that accept protons, also called a proton accceptor The Amphoteric Nature of Water Example of water behaving as a Bronsted base. HClfej + H 2 0(i) H 3 0 +^; + CI ( aq ) Example of water behaving as a Bronsted acid. NH 3fe; + n 2 o 0 ) <-» NHA,, + OH' Conjugate Acid Base Pairs Conjugates exist for every BLA and BLB pair. H (aq) Any reactant that acts as a BLA forms a conjugate Bronsted base (CB) on the product side of the equation and any BLB reactant molecule forms a conjugate Bronsted acid (CA) conjugate on the product side of the equation. (aq) BLA BLB CA CB HN0 3 + H 2 0 <r> H N0 3 H \AAAW.pembinatrails.ca/shaftesbury/mrdeakin 1_J adeakin@pembinatrails.ca (204) Shaftesbury High School, 2240 Grant Ave, Wpg, MB, R3P 0P7

36 BLA BLB Theory & Titrations Sample Problem Identify all Bronsted Lowry components (BLA, BLB, CA, CB) in the following equation. ^COlfaq) + HlO(l) <r> ( aq ) + HCO3 ( aq ) Solutions of Weak Bases Weak bases do not react completely with water to form hydroxide ions. Just like acid dissociation constant values, K, describe the strength of acids, base dissociation constant values, Kb, are used to quantitatively determine the relative strength of a weak base. Kb expression and value for pyridine, C 5 H 5 N Furthermore, Kb values are related to K a and K K b is determined from K a and Kw Hydrolysis of Anions and Cations Hydrolysis can be defined as a reaction of a catioritir m anion. Wfh wm &ater resulting in a ph change of the solution. Hydrolysis occurs in solutions of dissolved salts containing the anions or cations of weak conjugate acids or bases. hydrolysis of CN" and NH 4 4 H ^ adeakin@pembinatrails.ca (204) l Shaftesbury High School, 2240 Grant Ave, Wpg, MB, R3P 0P7