1. General Chemistry II 2302102 Acid and Base Equilibria i.fraser@rmit.edu.au Ian.Fraser@sci.monash.edu.au Lecture 1

2. Acids and Bases - 3 Lectures Autoionization of Water and pH Defining Acids and Bases Interaction of Acids and Bases with Water Buffer Solutions Acid-Base Titrations Outline - 5 Subtopics

3. By the end of this lecture AND completion of the set problems, you should be able to: Understand strong and weak electrolytes, K w and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H 3 O + ] and [OH - ]. Know the Arrhenius and Brønsted definitions of acids and bases. Understand and know examples of conjugate acid-base pairs. Be familiar with the ionization reactions for strong and weak acids, know examples of typical monoprotic, diprotic and triprotic acids. Understand the common ion effect in acid ionization. Objectives - Lecture 1 Acids and Bases

4. Acids & Bases Gold mining frequently uses the base, cyanide (CN - ), in the extraction process. Tailings dams sometimes have high concentrations of CN -. The conjugate acid of CN -, HCN, is extremely toxic. “Acid mine drainage” (AMD) is naturally produced by the exposure of sulfide ores to water. This process is greatly exacerbated by mining the sulfide ores (for Cu, Pb, Zn etc). AMD results in streams with elevated H + concentrations and pH values in the range 1-3. A gold mining tailings dam leaks into an AMD-affected stream - would you evacuate? Would you evacuate?

5. Acids & Bases - 2 Lectures Introduction to acids and bases Strong & Weak Acids Conjugate acid-base pairs Common Ion Effect Bases Buffers Indicators Titrations Strong Acids Weak Acids Outline

6. Acids & Bases - Lecture 1 By the end of this lecture AND completion of the set problems, you should be able to: Define Br Ø nsted acids and bases Calculate [H + ], [OH - ] and pH Distinguish between strong & weak acids Calculate equilibrium concentrations of acids & bases using acidity constants Identify conjugate acid-bases pairs Determine the effect of adding a common ion on equilibrium concentrations Objectives

7. Acids & Bases - Why are we interested? Acids - solvents (dissolve other materials) - ores, food (stomach contains HCl) Bases - solvents - cleaners like “Draino”, bleach Large Range of Industrial Processes: Acids - wine, beer, citrus fruits, vinegar, coffee Foods: Blood - pH 7.3-7.5 falls below 6.8 will be fatal (acidosis) Most bodily functions under ‘circumneutral’ conditions Physiology: ‘Natural’ Erosion of Limestone Caves “acid rain” - dissolved H 2 SO 4 & HNO 3 in upper atmosphere - rainfall runoff into acid lakes - devoid of life Environment:

8. Acids & Bases 1. ACID Species which can donate a hydrogen ion (H + ). 2. BASE Species which can accept a hydrogen ion. 3. AMPHOTERIC SPECIES Species which can act as both an acid and a base. e.g. HCO 3 - (CO 3 2-, H 2 CO 3 ), HSO 4 - Br Ø nsted-Lowry Definitions:

9. Hydrogen Ion What is the Hydrogen Ion? Hydrogen AtomAtomic weight = 1 1 proton + 1 electron Hydrogen Ion in Water Represented as: H + (aq) or H 3 O + Atomic weight = 1 1 proton Acid-Base Reaction In fact a proton-transfer reaction in which the proton is transferred from the acid to the base.

10. Acidity - pH pH convenient measure of H 3 O + concentration e.g. pH = 8.5 Acidity measure of the H 3 O + concentration

11. AQUEOUS SYSTEMS 2H 2 O( ) H 3 O + (aq) + OH - (aq) Equilibrium between H 3 O + and OH - : at equilibrium [H 3 O + ][OH - ] = K w at 25 °C K w = 1.00 x 10 -14 (c.f. K sp = [Ag + ][Cl - ])

12. AQUEOUS SYSTEMS 2H 2 O( ) H 3 O + (aq) + OH - (aq) K w = [H 3 O + ][OH - ] at 25 °C K w = 1.00 x 10 -14 Pure water at 25 °C. If 2z mol/L of water react: [H 3 O + ] = z and [OH - ] = z z 2 = 1.00 x 10 -14  z = 1.00 x 10 -7 and pH = 7.00

13. DEFINITIONS “acid” pH [OH - (aq)] “basic” pH > 7.00 [H 3 O + (aq)] < [OH - (aq)] “neutral”pH = 7.00 [H 3 O + (aq)] = [OH - (aq)]

14. Acid-Base Reaction: Proton-Transfer between an Acid and a Base (Water) According to the Brönsted-Lowry concept, when HCl gas is dissolved in water to form the solution of hydrochloric acid, a proton-transfer reaction occurs: H 2 O (l) + HCl (g) H 3 O + (aq) + Cl - (aq) Base(ProtonAcceptor) HydroniumIon A hydrated proton = H + (aq) = H 3 O + Acid(ProtonDonor) Cl OH : H +H : HOH H : Cl : - + + : : : : : :

15. ACIDS

16. ACIDS Note HA may be a molecule anion or cation Reaction between acids and water H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq) Equilibrium constant K a At equilibrium

17. ACIDS Reaction between acids and water Equilibrium constant K a H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq) Usually tabulated as pK a : Note - this is exactly the same relationship as between [H + ] and pH (pH = - log 10 [H + ])

18. Strength of Acids Strong versus Weak acids The strength of an acid is related to the position of the equilibrium above It is NOT related to the corrosive ability (this often causes confusion) As we shall see, HF is a weak acid, yet it is one of the most corrosive acids known. H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq)

19. STRONG ACID Position of equilibrium almost completely to the right (acid is almost totally ionized) e.g. 0.1 M HCl(aq) [H 3 O + ] = 0.1 [Cl - ] = 0.1 [HCl] not known ( < 10 -10 ) pK a < -10 Reaction between acids and water H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq)

20. Acids are Electrolytes Non No ions in solution Strong Many ions in solution - Weak Few ions in solution - Strong AcidsWeak Acids

21. WEAK ACIDS 1. Definition of weak acid Position of equilibrium lies to the left. Only a small fraction of the acid reacts with the water and is ionized. i.e. Acid is weakly ionized K a small < 10 -3 pK a large > 3 Reaction between acids and water H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq)

22. WEAK ACIDS initial 0.5 M 0 0 equilibrium (0.5 - x) M x M x M 2. Dissociation in water H 2 O( ) + HA(aq) H 3 O + (aq) + A - (aq) eg. 0.5 M HF(aq) K a = 6.8 x 10 -4 If x mol/L react then we have H 2 O( ) + HF(aq) H 3 O + (aq) + F - (aq)

23. initial 0.5 M 0 0 equilibrium (0.5 - x) M x M x M At equilibrium: H 2 O( ) + HF(aq) H 3 O + (aq) + F - (aq) So HF is ca. 4% ionized WEAK ACIDS

24. Diprotic & Triprotic Acids Acids that we’ve considered thus far have been monoprotic (donate 1 proton) e.g. HCl, HNO 3 Other acids can donate 2 or 3 protons - diprotic or triprotic acids e.g. H 2 SO 4, H 3 PO 4 H 2 O( ) + H 2 A(aq) H 3 O + (aq) + HA - (aq) H 2 O( ) + HA - (aq) H 3 O + (aq) + A 2- (aq)

25. Diprotic & Triprotic Acids Successive pK a s increase in magnitude - removal of successive protons results in weaker acids

26. Conjugate Acid-Base Pairs - Reverse reaction H 2 O (l) + NH 3 (aq) NH 4 + (aq) + OH - (aq) Consider the proton-transfer reaction: –Forward reaction H 2 O (l) + NH 3 (aq) NH 4 + (aq) + OH - (aq) Acid (Proton Donor) Base (Proton Acceptor) Acid (Proton Donor)Base(ProtonAcceptor)

27. Conjugate Acid-Base Pairs H 2 O (l) + NH 3 (aq) NH 4 + (aq) + OH - (aq) Acid 1Base 1Acid 2Base 2 Conjugate Acid- Base Pair

28. For the general equation B + HA HB + + A - Conjugate Acid-Base Pairs Acid (Proton Donor) Base (Proton Acceptor) Acid(ProtonDonor) Base (Proton Acceptor) The acid HB + results from the base B gaining a proton. HB + -B is a conjugate acid-base pair. The base A - results from the acid HA losing a proton. HA-A - is a conjugate acid-base pair, too.

29. Conjugate Acid-Base Pairs Any two substances that differ by one proton are a conjugate acid-base pair Can write the formula of the conjugate base of any acid simply by removing the proton: If HCO 3 - is the acid, CO 3 2- is the conjugate base If HCO 3 - is the base, H 2 CO 3 is the conjugate acid (HCO 3 - is amphoteric)

30. Conjugate Acid-Base Pairs The conjugate (base) of a strong acid is a weak base The conjugate (base) of a weak acid is a strong base The conjugate (acid) of a strong base is a weak acid The conjugate (acid) of a weak base is a strong acid

31. Direction of Acid-Base Reactions NH 4 + is a stronger acid than H 2 O, and OH - is a stronger base than NH 3 So reaction proceeds spontaneously to the left Acid-Base Reactions proceed spontaneously with the strongest acid and strongest base forming the weakest acid and the weakest base Returning to: H 2 O (l) + NH 3 (aq) NH 4 + (aq) + OH - (aq)

32. The Common Ion Effect Recall that with sparingly soluble salts:The presence of an ion in solution which is common to the electrolyte will decrease the solubility (adding KF to a solution of PbF 2 reduced the Pb 2+ concentration) The presence of an ion in solution which is common to the weak acid will suppress it’s ionization (decrease the concentration of the free ion)

33. The Common Ion Effect initial 0.5 M 0 0 equilibrium (0.5 - x) M x M x M H 2 O( ) + HF(aq) H 3 O + (aq) + F - (aq) At equilibrium, [F - ] = [H 3 O + ] = 1.8 x 10 -2 M so pH = 1.74 What happens if we add 0.05 M NaF? initial 0.5 M 0 0.05 equilibrium (0.5 - x) M x M 0.05 + x M H 2 O( ) + HF(aq) H 3 O + (aq) + F - (aq)

34. The Common Ion Effect initial 0.5 M 0 0.05 equilibrium (0.5 - x) M x M 0.05 + x M H 2 O( ) + HF(aq) H 3 O + (aq) + F - (aq) So (0.05 + x) x = (0.5 - x) 6.8 x 10 -4 x = 6.00 x 10 -3 M So [H 3 O + ] = 6.00 x 10 -3 M Thus pH = 2.22 c.f. 1.74 without the NaF

35. Acids and Bases - End of Lecture 1 After studying this lecture and the set problems, you should be able to: Understand strong and weak electrolytes, K w and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H 3 O + ] and [OH - ]. Know the Arrhenius and Brønsted definitions of acids and bases. Understand and know examples of conjugate acid-base pairs. Be familiar with the ionization reactions for strong and weak acids, know examples of typical monoprotic, diprotic and triprotic acids. Understand the common ion effect in acid ionization. Objectives Covered in Lecture 1