Acids and Bases-Review

1. Acids and Bases-Review • The properties of acids include the following: • Taste sour (but don't taste them!!) • Their water solutions conduct electrical current (electrolytes) • They react with bases to form salts and water • Turns Blue Litmus Paper to Red • The properties of bases include the following: • Have a slippery feel between the fingers • Have a bitter taste (but don't taste them!!) • React with acids to form salts and water • Turns Red Litmus Blue • Their water solutions conduct electrical current (electrolytes)

2. Acids and Bases Arrhenius in 1884 discovered that acids give off H+ ions and allow for a good flow of electricity through a solution. Arrhenius also discovered that bases give off OH- ions and OH- ions also allow for a good flow of electricity through the solution. Traditionally Professor Arrhenius defined: Acid released Hydrogen ion (as Hydronium ions, H3O+) in water solution. Base produced Hydroxide ion in water solution. The limitations on these definitions were: 1. The need for water 2. The need for a protic acid 3. The need for Hydroxide bases

3. Bronsted/Lowry acids and bases Bronsted and Lowry defined these two terms the following: Acid-Proton donor Base-Proton acceptor These definitions are not as restrictive as Arrhenius’ definitions. • No need for water although it can be present, it need not be. • Bases do not have to be Hydroxide compounds. However, one restriction still remaining is the need for a protic acid. Each Bronsted acid is coupled to a conjugate base to constitute a CONJUGATE ACID-BASE PAIR CH3COOH + H2O H3O++CH3COO-

4. Lewis Acids and Bases G.N. Lewis defined these in an even less restrictive manner: Acid- Electron pair acceptor Base- Electron pair donor In this set of definitions there is no longer a need for a protic acid. In other words only electron exchange must occur. These definition sets are NOT contradictory. A Proton donor is the same as an electron acceptor. A Proton acceptor is the same as an electron donor. Also the first set of definitions are less inclusive so that all of the Arrenhius acids are found under the Bronsted definition but not all Bronsted acids will be Arrenhius acids. All Arrenhius and Bronsted acids will be under the Lewis definition but not all Lewis acids will be Bronsted or Arrenhius acids.

5. Acid and Base Strength Strong acids (memorise) dissociate completely in water HClO4, HClO3, HCl, HBr, HI, HNO3 and H2SO4 Strong bases are the metal hydroxides of Group 1 and heavy Group 2 E.g. LiOH, NaOH, KOH, Ba(OH)2 etc Weak acids and bases are not completely ionised in solution CH3COOH + H2O H3O++CH3COO- Ka is an equilibrium constant called the acid dissociation constant

6. Acid and Base Strength :NH3 + H2O NH4++OH- (a molecular base) The magnitude of the Ka or Kb, using water as a common proton donor/acceptor, determines the strength of the acid or base In general (for acids) HA + H2O H3O++A- Water is AMPHOTERIC. It can act as an acid or a base

7. Acid and Base Strength Ka ~1010 1x10-2 6.8x10-4 1.75x10-5 9.5x10-8 5.7x10-10 4.7x10-11 1.8x10-16 Stronger Acid Levelling Effect Each acid will transfer a proton to a base below it in a mixed solution HClO4 ClO4- H2SO4 HSO4- HCl Cl- H3O+ H2O HSO4- SO42- HF F- CH3COOH CH3COO- H2S HS- NH4+ NH3 HCO3- CO32- H2O OH- Stronger Base

8. 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

9. The pH Scale

10. The Common Ion Effect • The solubility of a partially soluble salt is decreased when a common ion is added. • Consider the equilibrium established when ethanoic acid is added to water. • At equilibrium H+ and C2H3O2- are constantly moving into and out of solution, but the concentrations of ions is constant and equal. • If a common ion is added, e.g. C2H3O2- from NaC2H3O2 (which is a strong electrolyte) then [C2H3O2-] increases and the system is no longer at equilibrium. • So, [H+] must decrease, according to Le Chatelier’s Principle.

11. 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

12. 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?

13. 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.

14. 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(mL)/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

15. 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

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

17. 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

18. 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

19. 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)

20. 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

21. Acid-Base Titrations Weak Acid-Strong Base Titrations • The weaker the acid, the smaller the equivalence point inflection. • For very weak acids, it is impossible to detect the equivalence point. • Choose an indicator with a Ka range suited to the weak acid. • Titration of weak bases with strong acids have similar features to weak acid-strong base titrations.

22. 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

23. 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

24. Solubility Product Ksp • Consider • for which • Ksp is the solubility-product constant. (BaSO4 is ignored because it is a pure solid).

25. Factors That Affect Solubility Common-Ion Effect • Solubility is decreased when a common ion is added. • This is an application of Le Châtelier’s principle: • as F- (from NaF, say) is added, the equilibrium shifts away from the increase. • Therefore, CaF2(s) is formed and precipitation occurs. • As NaF is added to the system, the solubility of CaF2 decreases.