General Chemistry II 2302102

General Chemistry II2302102 Acid and Base Equilibria Lecture 1 i.fraser@rmit.edu.au Ian.Fraser@sci.monash.edu.au

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

Acids and Bases Objectives - Lecture 1 • By the end of this lecture AND completion of the set problems, you should be able to: • Understand strong and weak electrolytes, Kw and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H3O+] 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.

Acids & Bases Would you evacuate? 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?

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

Acids & Bases - Lecture 1 Objectives • By the end of this lecture AND completion of the set • problems, you should be able to: • Define BrØnsted 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

Large Range of Industrial Processes: Acids - solvents (dissolve other materials) - ores, food (stomach contains HCl) Bases - solvents - cleaners like “Draino”, bleach Foods: Acids - wine, beer, citrus fruits, vinegar, coffee Physiology: Blood - pH 7.3-7.5 falls below 6.8 will be fatal (acidosis) Most bodily functions under ‘circumneutral’ conditions Environment: ‘Natural’ Erosion of Limestone Caves “acid rain” - dissolved H2SO4 & HNO3 in upper atmosphere - rainfall runoff into acid lakes - devoid of life Acids & Bases - Why are we interested?

Acids & Bases BrØnsted-Lowry Definitions: 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. HCO3- (CO32-, H2CO3), HSO4-

Hydrogen Ion Atomic weight = 1 1 proton What is the Hydrogen Ion? • Hydrogen Atom Atomic weight = 1 1 proton + 1 electron • Hydrogen Ion in Water Represented as: H+(aq) or H3O+ Acid-Base Reaction In fact a proton-transfer reaction in which the proton is transferred from the acid to the base.

Acidity - pH Acidity measure of the H3O+ concentration pH convenient measure of H3O+ concentration e.g. pH = 8.5

2H2O(l) H3O+(aq) + OH-(aq) AQUEOUS SYSTEMS Equilibrium between H3O+ and OH-: at equilibrium [H3O+][OH-] = Kw at 25 °C Kw = 1.00 x 10-14 (c.f. Ksp =[Ag+][Cl-])

2H2O(l) H3O+(aq) + OH-(aq) AQUEOUS SYSTEMS Kw = [H3O+][OH-] at 25 °C Kw = 1.00 x 10-14 Pure water at 25 °C. If 2z mol/L of water react: [H3O+] = z and [OH-] = z z2 = 1.00 x 10-14 \z = 1.00 x 10-7 and pH = 7.00

DEFINITIONS “acid” pH < 7.00 [H3O+(aq)] > [OH-(aq)] “basic” pH > 7.00 [H3O+(aq)] < [OH-(aq)] “neutral” pH = 7.00 [H3O+(aq)] = [OH-(aq)]

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: H2O (l) + HCl (g) H3O+ (aq) + Cl- (aq) + : : : : - H O : + H Cl : H O H + :Cl: : : H H Acid (Proton Donor) Base (Proton Acceptor) Hydronium Ion A hydrated proton = H+ (aq) = H3O+ Acid-Base Reaction: Proton-Transfer between an Acid and a Base (Water)

ACIDS

Reaction between acids and water H2O(l) + HA(aq) H3O+(aq) + A-(aq) Equilibrium constant Ka At equilibrium Note HA may be a molecule anion or cation ACIDS

H2O(l) + HA(aq) H3O+(aq) + A-(aq) Usually tabulated as pKa: ACIDS Reaction between acids and water Equilibrium constant Ka Note - this is exactly the same relationship as between [H+] and pH (pH = - log10 [H+])

H2O(l) + HA(aq) H3O+(aq) + A-(aq) 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.

Reaction between acids and water H2O(l) + HA(aq) H3O+(aq) + A-(aq) STRONG ACID Position of equilibrium almost completely to the right (acid is almost totally ionized) e.g. 0.1 M HCl(aq) [H3O+] = 0.1 [Cl-] = 0.1 [HCl] not known ( < 10-10) pKa < -10

Strong Weak Non Few ionsin solution - Many ionsin solution - No ionsin solution Acids are Electrolytes Strong Acids Weak Acids

Reaction between acids and water H2O(l) + HA(aq) H3O+(aq) + A-(aq) 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 Kasmall < 10-3 pKalarge > 3

2. Dissociation in water H2O(l) + HA(aq) H3O+(aq) + A-(aq) H2O(l) + HF(aq) H3O+(aq) + F-(aq) eg. 0.5 M HF(aq) Ka = 6.8 x 10-4 If x mol/L react then we have WEAK ACIDS initial 0.5 M 0 0 equilibrium (0.5 - x) M x M x M

H2O(l) + HF(aq) H3O+(aq) + F-(aq) At equilibrium: WEAK ACIDS initial 0.5 M 0 0 equilibrium (0.5 - x) M x M x M So HF is ca. 4% ionized

H2O(l) + H2A(aq) H3O+(aq) + HA-(aq) H2O(l) + HA-(aq) H3O+(aq) + A2-(aq) Diprotic & Triprotic Acids • Acids that we’ve considered thus far have been monoprotic (donate 1 proton) e.g. HCl, HNO3 • Other acids can donate 2 or 3 protons - diprotic or triprotic acids e.g. H2SO4, H3PO4

Diprotic & Triprotic Acids Successive pKas increase in magnitude - removal of successive protons results in weaker acids

Consider the proton-transfer reaction: • Forward reaction H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq) Acid (Proton Donor) Base (Proton Acceptor) • - Reverse reaction • H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq) Acid (Proton Donor) Base (Proton Acceptor) Conjugate Acid-Base Pairs

H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq) Conjugate Acid-Base Pair Acid 1 Base 1 Acid 2 Base 2 Conjugate Acid-Base Pair Conjugate Acid-Base Pairs

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

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 HCO3- is the acid, CO32- is the conjugate base • If HCO3- is the base, H2CO3 is the conjugate acid • (HCO3- is amphoteric)

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

Acid-Base Reactions proceed spontaneously with the strongest acid and strongest base forming the weakest acid and the weakest base • Returning to: H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq) Direction of Acid-Base Reactions NH4+ is a stronger acid than H2O, and OH- is a stronger base than NH3 So reaction proceeds spontaneously to the left

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 PbF2 reduced the Pb2+ 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)

initial 0.5 M 0 0 equilibrium (0.5 - x) M x M x M H2O(l) + HF(aq) H3O+(aq) + F-(aq) H2O(l) + HF(aq) H3O+(aq) + F-(aq) At equilibrium, [F-] = [H3O+] = 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 The Common Ion Effect

initial 0.5 M 0 0.05 equilibrium (0.5 - x) M x M 0.05 + x M H2O(l) + HF(aq) H3O+(aq) + F-(aq) The Common Ion Effect So (0.05 + x) x = (0.5 - x) 6.8 x 10-4 x = 6.00 x 10-3 M So [H3O+] = 6.00 x 10-3 M Thus pH = 2.22c.f. 1.74 without the NaF

Acids and Bases - End of Lecture 1 Objectives Covered in Lecture 1 • After studying this lecture and the set problems, • you should be able to: • Understand strong and weak electrolytes, Kw and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H3O+] 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.

General Chemistry II 2302102