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Ionisation of water and pH

pH concept pH = -log[H + ] pX = -logX. Ionisation of water and pH. pH scale [H + ] > 10 -7 M, pH < 7 ACIDIC [H + ] < 10 -7 M, pH > 7 BASIC [H + ] = 10 -7 M, pH = 7 NEUTRAL. For any Bronsted conjugate Acid-Base pair. K a . K b = K w. CH 3 COOH H + +CH 3 COO -. Buffers.

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Ionisation of water and pH

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  1. pH concept pH = -log[H+] pX = -logX Ionisation of water and pH pH scale [H+] > 10-7M, pH < 7 ACIDIC [H+] < 10-7M, pH > 7 BASIC [H+] = 10-7M, pH = 7 NEUTRAL For any Bronsted conjugate Acid-Base pair Ka . Kb = Kw

  2. CH3COOH H++CH3COO- Buffers Every life form is extremely sensitive to slight pH changes. Human blood for example needs to remain within the range 7.38-7.42. Buffers: buffer the system against extreme changes in pH Buffer solutions normally consist of two solutes: a weak Bronsted acid and its conjugate base

  3. Buffers In general for: HAA- + H+ Henderson-Hasselbach Equation Buffer capacity Q. If we generate 0.15mol H+ in a reaction vessel of 1L (with no accompanying volume change) containing 1mol each of CH3COOH and CH3COO-, what will the solution pH change be? For the same reaction in water what is the pH change?

  4. Acid-Base Reactions Acid/Base reactions are reactions that involve the neutralisation of an acid through the use of a base. HCl + NaOH NaCl + H2O In this reaction, the Na+ and the Cl- are called spectator ions because they play no role in the overall outcome of the reaction. The only thing that reacts is the H+ (from the HCl) and the OH- (from the NaOH). So the reaction that actually takes place is: H+ + OH-H2O  If in the end, the OH- was the limiting reagent and there are H+'s still left in the solution then the solution is acidic, but if the H+ was the limiting reagent and OH-'s were left in the solution then the solution is basic. Titration Titration is the process of mixing acids and bases to analyse one of the solutions. For example, if you were given an unknown acidic solution and a 1 molar NaOH solution, titration could be used to determine what the concentration of the other solution was.

  5. Acid-Base Titrations The goal of titration is to determine the equivalence point. The equivalence point is the point in which all the H+ and the OH- ions have been used to produce water. Titration also usually involves an indicator. An indicator is a liquid that turns a specific colour at a specific pH. (Different indicators change colours at different pH's). Indicators are chosen to allow a colour change at the equivalence point. Titration of a strong acid with a strong base 50.00mL of 0.020M HCl with 0.100M NaOH H+ + OH-H2O Kc=1/Kw=1014 at equivalence pt.: nb mol HCl = nb mol NaOH 0.02mol/L x 50/1000 L = 0.1mol/L x Ve/1000 L Ve = 0.001mol HCl  (0.1mol/L x 1/1000 L) = 10 mL pH determined by dissociation of H20: Kw = [H+][OH-] = 10-14 [H+] = 10-14 = 10-7 mol/L => pH = 7.00

  6. Acid-Base Titrations Titration of a strong acid with a strong base Initial pH: 0.02mol/L strong acid. pH = 1.70 before equivalence pt.: when 3.00mL of NaOH has been added Initial conc. pH = 1.88 Fraction of H+ remaining Dilution factor after equivalence pt.: 10.1mL NaOH added pOH = 3.78 pH = 10.22 Initial conc. of base Dilution factor

  7. Titration Curves Titration curve of a strong acid with a strong base

  8. CH3COOH H++CH3COO- 0.02-x x x Titration of a weak acid with a strong base Take the example of a titration of 50.0mL 0.020M CH3COOH (Ka = 1.8 x 10-5) with 0.10M NaOH CH3COOH + NaOH  CH3COONa + H2O Reaction is the reverse of Kb for CH3COO- base K = 1/Kb = 1/(Kw /Ka) = 1.8 x 109 Ve = 10mL (as before) Initial pH: a weak acid equilibrium problem x = 6 x 10-4, pH = 3.22

  9. Titration of a weak acid with a strong base Before eq. pt.: buffer system One of the simplest ways to treat these problems is to evaluate the quotient in the log using relative concentration before and after the reaction. Imagine we have added 3.00mLs of base CH3COOH + NaOH  CH3COONa + H2O Relative Initial: 1 3/10 Relative final: 7/10 3/10

  10. Titration of a weak acid with a strong base When volume of base added = 1/2Ve at equivalence pt.: we have a solution of base in water CH3COONa + H2O CH3COOH + OH- F-x x x Kb = (Kw /Ka) = 5.56 x 10-10 = x2/(F-x) x = 3.05 x 10-6, pOH = 5.52, pH=8.48 (BASIC)

  11. Titration of a weak acid with a strong base after equivalence pt.: pH is determined by excess base added For 10.1mL base added in total pOH = 3.78 pH = 10.22

  12. HIn(aq.)H+ (aq.)+In- (aq.) Acid-Base Indicators Usually dyes that are weak acids and display different colours in protonated/deprotonated forms. In general we seek an indicator whose transition range (±1pH unit from the indicator pKa) overlaps the steepest part of the titration curve as closely as possible

  13. Acid-base indicators Indicator pH range pKa Acid Form Base Form methyl violet 0.0- 1.6 0.8 yellowblue thymol blue 1.2- 2.8 1.6 redyellow methyl yellow 2.9- 4.0 3.3 redyellow methyl orange 3.1- 4.4 4.2 redyellow bromocresol green 3.8- 5.4 4.7 yellowblue methyl red 4.2- 6.2 5.0 redyellow bromothymol blue 6.0- 7.6 7.1 yellowblue phenol red 6.4- 8.0 7.4 yellowred thymol blue 8.0- 9.6 8.9 yellowblue phenolphthalein 8.0- 9.8 9.7 colourlessred thymolphthalein 9.3-10.5 9.9 colourlessblue alizarin yellow R 10.1-12.0 11.0 yellowred indigo carmine 11.4-13.0 12.2 blueyellow

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