Chemical Equilibria

Chemical Equilibria Hydrolysis, Buffers and Acid Base Titrations

Bronsted Lowry Theory of Acids and Bases • An acid is a proton (hydrogen ion) donor • A base is a proton (hydrogen ion) acceptor • H2O + HCl → H3O+ + Cl- Base 2 Base 1 Acid 1 Acid 2

Bronsted Lowry Acids and Bases • In the preceding slide the conjugate acid is H3O+ since it can donate a proton. • The conjugate base is Cl- ion. It can accept a proton but is a very weak base.

Bronsted Lowry Acids and Bases • A water molecule can act as an acid or base. Proton accepted Proton donated

Bronsted Lowry Acids and Bases • When an acid in solution reacts with a base, the H3O+ (hyroxonium ion) is functioning as the acid – donating H+ • NaOH + HCl → NaCl + H2O • Na+ + OH- + H3O + + Cl-→ Na+ + Cl- + 2H2O • H3O+ + OH-→ 2H2O Water molecule formed by proton donated to OH-

Bronsted Lowry Acids and Bases • Advantage of this theory • NH3(aq) + H3O+→ NH4+(aq) +H2O(l) • NH3(aq) + H3O+ + Cl-→ NH4+(aq) + Cl- +H2O(l) • In solution, the ammonia accepts a proton from the hydroxonium ion. • NH3(gas) + HCl(gas)→ NH4+(s) + Cl-(s) • In the gas state, the ammonia accepts a proton from the HCl gas. • Either way, the ammonia is acting as a base by accepting a proton from the acid.

Lewis Theory of Acids and Bases • An acid is an electron pair acceptor • A base is an electron pair donor • This theory extends to molecules in which protons are not involved in the reaction

The variation of the pH of pure water with temperature According to Le Chatelier, if you increase the temperature of the water, the equilibrium will move to lower the temperature again. It does this by absorbing the extra heat. That means that the forward reaction will be favored, and more hydrogen ions and hydroxide ions will be formed. The effect of that is to increase the value of Kw as temperature increases H2O + heat ↔ H+ + OH-

The variation of the pH of pure water with temperature

Hydrolysis Hydrolysis is the reaction of a material with water Salts are strong electrolytes, this means they dissociate 100% in solution to give free aqueous ions NaCl → Na+ + Cl- When both ions come from strong acid and bases they have no interactions with the ions formed by the dissociation of water (hydrogen and hydroxide ions). However if the ions come from weak acids and bases they can interact with water ions, H+ and OH-, establishing equilibria.

Hydrolysis Sodium acetate is 100% dissociated into ions:- CH3COONa → CH3COO- + Na+ Sodium ions are from a strong base (sodium hydroxide) and do not interact with the water ions. However, the acetate ions do interact with the hydrogen ions from the water equilibrium (H2O ↔H+ + OH-) CH3COO- + H+↔ CH3COOH We know that this equilibrium lies to the side of the acetic acid (to the right), removing the hydrogen ions from the solution.

Hydrolysis As [H+] decreases the pH rises. Hence a solution of sodium acetate has a pH greater than 7. We say that it is basic by hydrolysis.

Hydrolysis Ammonium chloride solution Ammonium chloride dissociates 100% into ions in solution NH4Cl→NH4+ + Cl- The ammonium ions interact with the hydroxide ions from the water removing them from the solution (equilibrium lies to the right) NH4+ + OH-↔NH3 + H2O This increases the concentration of hydrogen ions (as [H+] x [OH-] is constant) increasing the acidity of the solution (pH decreases) We say that a solution of ammonium chloride is acidic by hydrolysis.

Hydrolysis General rules When the negative ion is from a weak acid then the salt is basic by hydrolysis When the positive ion is from a weak base then the salt is acidic by hydrolysis If the salt is formed from a strong acid and strong base then it is neutral If the salt is formed from a weak acid and weak base then its hydrolysis and pH is determined by the relative Ka and Kb values If Ka > Kb it will be acid If Kb > Ka it will be alkaline

Hydrolysis • Ka = ( [H+][A-] / [HA] ) and • Kb = ( [BH+][OH-] / [B] ) • Ka is a constant which describes the ionisation of an acid (ie how strong it is) and Kb does the same for bases. pKa is the log form of Ka, defined as pKa = -log(Ka) and pKb = -log(Kb). • Like previously with the pH scale, a 1 fold change in pKa will signify a ten fold change in Ka and the same for Kb . • Ka x Kb = Kw (ie they equal 1 x 10-14 at 25ºC) • pKa + pKb = pKw (ie 14 at 25ºC)

Hydrolysis More Rules Strong acids have weak conjugate bases. Strong bases have weak conjugate acids. A strong acid has a large Ka value (or a small pKa value). A strong base, likewise, has a large Kb value and a small pKb value. Given the required input data, all the equations in the previous slides need to be applied as appropriate.

Buffers Solutions able to retain constant pH regardless of small amounts of acids or bases added are called buffers. Classical buffer contains solution of weak acid and conjugate base. Small amounts of acids or bases added are absorbed by buffer and the pH changes only slightly. Weak acid + salt of weak acid: Weak base + salt of weak base. In case of high or low pH just solutions of strong acids or bases are used - for example in case of pH=1 acid concentration is relatively high (0.1 M) and small addition of acid or base doesn't change the pH of such solutions significantly.

Buffers The Henderson-Hasselbalch equation (or buffer equation can be used for pH calculation of solution containing pair of acid and conjugate base - like HA/A-, HA-/A2- or B+/BOH.

Buffers For solutions of weak bases sometimes it is more convenient to use the equation in the form Both equations are perfectly equivalent and interchangeable.

Buffers Henderson-Hasselbalch equation is used mostly to calculate pH of solution created mixing known amount of acid and conjugate base (or neutralizing part of acid with strong base). For example, what is pH of solution prepared mixing reagents so that it contains 0.1 M of acetic acid and 0.05 M NaOH? Half of the acid is neutralized, so concentrations of acid and conjugate base are identical, thus quotient under logarithm is 1, logarithm is 0 and pH = pKa.

Buffers This approach - while perfectly justifiable in many cases - is dangerous, as it creates false conviction that the equation can be used this way always. That's not true. Henderson-Hasselbalch equation is valid when it contains equilibrium concentrations of acid and conjugate base. In case of solutions containing not-so-weak acids (or not-so-weak bases) the equilibrium concentrations can be far from concentrations of substances added to solution. Let's replace acetic acid from our example with something stronger - e.g. dichloroacetic acid, with pKa=1.5.

Buffers The same reasoning leads to a result of pH=1.5 - which is wrong. The proper pH value can be calculated using a more rigorous approach and it is 1.78. The reason is simple. Dichloroacetic acid is strong enough to dissociate on its own and the equilibrium concentration of conjugate base is not 0.05 M (as we expected from the neutralization reaction stoichiometry) but 0.0334 M. As a rule of thumb acids with pKa below 2.5 the use of Henderson-Hasselbalch equation for pH prediction can give wrong results, especially in case of diluted solutions.

Buffers For solutions above 10 mM and acids weaker than pKa>2.5, Henderson-Hasselbalch equation gives results with acceptable error. The same holds for bases with pKb>2.5. However, the same equation will work perfectly regardless of the pKa value if you are asked to calculate ratio of acid to conjugate base in the solution of known pH.

Buffer Capacity Calculation of the Buffer Capacity The buffer capacity refers to the maximum amount of either strong acid or strong base that can be added before a significant change in the pH will occur. This is simply a matter of stoichiometry. The maximum amount of strong acid that can be added is equal to the amount of conjugate base present in the buffer. The maximum amount of base that can be added is equal to the amount of weak acid present in the buffer.

Buffer Capacity • Example: What is the maximum amount of acid that can be added to a buffer made by the mixing of 0.35 moles of sodium hydrogen carbonate with 0.50 moles of sodium carbonate? How much base can be added before the pH will begin to show a significant change? • First, write the equation for the ionization of the weak acid, in this case of hydrogen carbonate. Although this step is not truly necessary to solve the problem, it is helpful in identifying the weak acid and its conjugate base. • HCO3-(aq) + H2O(l) ↔ H3O+(aq) + CO32-(aq)

Buffer Capacity • Second, added strong acid will react with the conjugate base, CO32-. • Therefore, the maximum amount of acid that can be added will be equal to the amount of CO32-, 0.50 moles. • Third, added strong base will react with the weak acid, HCO3-. Therefore, the maximum amount of base that can be added will be equal to the amount of HCO3-, 0.35 moles.

Hydrolysis and Buffer Capacity • PowerPoint slides, giving examples of hydrolysis and buffer capacity calculations with pH will be added to shared folders and my Web Site.

Acid Base Titrations

Strong acid (HCl) titrated into a strong base (NaOH)1 M HCl added to 1 M NaOH The pH does not change appreciably until near the equivalence point (EP). Around the EP the pH falls from 11.3 to 2.7 from 24.9 to 25.1 mL of HCl added Equal No moles of HCl has reacted with NaOH at the EP pH4

Strong acid (HCl) weak base (NH3 soln) Buffer solution is being setup due to excess ammonia and the ammonium chloride being formed NH4Cl pH ≈ 5 Acid side of neutrality pH 7

Weak acid (HAc) strong base (NaOH) Equivalence point NaAc pH ≈ 9 Base side of neutrality pH 7 Buffer solution formed from sodium acetate and acetic acid

Weak acid (HAc) weak base (NH3) The lack of a steep change in pH at the end point makes it difficult to titrate a weak acid against a weak base

Summary

Hydrochloric acid to sodium carbonateNa2CO3 + 2HCl → 2NaCl + CO2 + H2O Titration to HCO3-

Addition of NaOH to oxalic acid (COOH)2 COOH COO-Na+ COOH COO-Na+ Volume of base added

0.1 M NaOH added 0.1 M HCl

0.1 M NaOH added 0.1 M HAc pH 8.5

0.1 M HCl added 0.1 M NH3 pH 5

Acid base indicators An acid base indicator is a dye which changes colour at a certain pH. The indicator must be chosen so that the colour change takes place at the required pH A strong acid strong base; B weak acid strong base C weak base strong acid

Titration of carbonic acid with sodium hydroxide 2NaOH + H2CO3 → Na2CO3+ H2O

Sodium hydroxide added to hydrochloric acid

NH3 added to acetic acid and NaOH added to a weak acid Ka= 1 x 10-7

HCl added to NH3 acid and a weak base Kb = 1 x 10-7

Useful Web Sites http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/stoichiometry/acid_base.html http://www.avogadro.co.uk/chemeqm/acidbase/titration/phcurves.htm http://www.chemguide.co.uk/physical/acidbaseeqia/kw.html#top Use the Web to search suitable sites on acid base equilibria, buffers and pH titrations.

Chemical Equilibria