Chem. 1B – 9/13 Lecture

Announcements First Mastering Homework due Tonight Starting Lab #7 on Wednesday Today’s Lecture – cont. –Le Châtelier’s Principle (Stresses resulting in equilibrium shifts) –Chapter 15: Definition of Acids/Bases (3 types of definitions) –Acid Strengths and K a –Autoprotolysis and pH –[H + ] and pH in strong acid and weak acid solutions

Chem 1B - Equilibrium Le Châtelier’s Principle Changes in Conditions – Types: –Changes in reactant or product concentrations – covered last time –Effect of a change in volume (compression/expansion or dilution/concentration) –Change in temperature

Chem 1B - Equilibrium Le Châtelier’s Principle – Volume Stress Example: 2NO(g) + O 2 (g) ↔ 2NO 2 (g) –Mathematical explanation: –Initially at equilibrium K C = 10 5 and [NO] = M, [O 2 ] = M and [NO 2 ] = M –Now we reduce the volume from 10.0 to 1.00L –New concentrations: [NO] = M, [O 2 ] = M and [NO 2 ] = 0.10 M (same number of moles in 1/10 th the volume so 10X more concentrated) –Q = (0.10 M) 2 /[(0.010 M) 2 (0.010 M)] = 10 4 < K C, so products favored

Chem 1B - Equilibrium Le Châtelier’s Principle – Volume Stress Note: in aqueous solutions, dilution works in the same way (increase in space due to dilution favors side with more moles) a 1 M HC 2 H 3 O 2 (acetic acid) solution is diluted by adding an equal volume of water. How does this reaction change? HC 2 H 3 O 2 (aq) ↔ H + (aq) + C 2 H 3 O 2 - (aq)

Chem 1B - Equilibrium Le Châtelier’s Principle – Temperature Stress Note: change in T changes K (while initial K becomes Q) If ΔH>0, as T increases, products favored - this also means K increases with T If ΔH<0, as T increases, reactants favored Easiest to remember by considering heat a reactant or product Example: OH - + H + ↔ H 2 O(l) + heat (reaction  H < 0) Increase in T

Chem 1B - Equilibrium Le Châtelier’s Principle Looking at the reaction below, that is initially at equilibrium, AgCl(s) ↔ Ag + (aq) + Cl - (aq) (ΔH°>0) determine the direction (toward products or reactants) each of the following changes will result in a)increasing the temperature b)addition of water (dilution of system) c)addition of AgCl(s) d)addition of NaCl

Chem 1B – Aqueous Chemistry Acid Definitions Arrhenius Definition (most narrow definition, but most familiar to us) –An acid releases H + (essentially the same as H 3 O + ) when dissolved in water –A base releases OH - upon dissolution in water –Acid Example: HClO 4 can write reaction as: HClO 4 (aq) + H 2 O(l) → H 3 O + (aq) + ClO 4 - (aq) or as: HClO 4 (aq) → H + (aq) + ClO 4 - (aq) (simplified view) –Base Example: KOH KOH(aq) → K + (aq) + OH - (aq) makes Arrhenius acid an acid makes KOH a base

Chem 1B – Aqueous Chemistry Acid Definitions Brønsted-Lowry (broader definition because it applies to non-aqueous solutions) –An acid is a proton (H + ) donor –A base is a proton acceptor (normally must have an available electron pair) –Acid-base reactions will have both an acid and a base (different than Arrhenius definition) Example: acetic acid in water HC 2 H 3 O 2 (aq) + H 2 O(l) ↔ H 3 O + (aq) + ClO 4 - (aq) acid base C 2 H 3 O 2 -H + :O-H ↔ H-O + -H + C 2 H 3 O 2 - H H

Chem 1B – Aqueous Chemistry Acid Definitions Examples –In the following examples, indicate Arrhenius and Brønsted-Lowry acids and bases: 1)NH 3 (aq) + H 2 O(l) ↔ OH - (aq) + NH 4 + (aq) 2)HNO 2 (aq) + H 2 O(l) ↔ H 3 O + (aq) + NO 2 - (aq) 3)HCl(aq) → H + (aq) + Cl - (aq) 4)NH 4 + (aq) ↔ H + (aq) + NH 3 (aq) 5)Al(H 2 O) 6 3+ (aq) ↔ Al(H 2 O) 5 OH 2+ (aq) + H + (aq) 6)H 2 SO 4 (sol’n) + HC 2 H 3 O 2 (l) ↔ HSO 4 - (sol’n) + H 2 C 2 H 3 O 2 + (sol’n)

Chem 1B – Aqueous Chemistry Acid Definitions Brønsted-Lowry – Conjugate acids and bases –When a Brønsted-Lowry acid loses its H +, the remainder is called a conjugate base –When a Brønsted-Lowry base extracts a H +, the remainder is called a conjugate acid Examples HNO 2 (aq) + H 2 O(l) ↔ H 3 O + (aq) + NO 2 - (aq) NH 3 (aq) + H 2 O(l) ↔ OH - (aq) + NH 4 + (aq) Conjugate base Conjugate acid

Chem 1B – Aqueous Chemistry Acid Definitions Lewis Acids and Bases (15.10) –A Lewis Acid is an electron pair acceptor –A Lewis Base is an electron pair donor –Brønsted-Lowry acids are also Lewis acids (but not always visa versa) –Examples: :NH 3 (aq) + H 2 O(l) ↔ OH - (aq) + NH 4 + (aq) Electron pair on N makes NH 3 a base Ag + + 2NH 3 (aq) ↔ Ag(NH 3 ) 2 + (aq) (formation of metal- ligand complex) Top example H 2 O(l) is both a Brønsted-Lowry and also a Lewis acid, while Ag + in bottom is only a Lewis acid

Chem 1B – Aqueous Chemistry Acid Strength In water, extent of formation of H + gives strength of acid Generic Acid: HA HA(aq) ↔ H + (aq) + A - (aq) The more strongly the above reaction favors the products, the stronger the acid This is given by the equilibrium constant – called K a in this example where: K a = [H + (aq)][A - (aq)]/[HA(aq)]

Chem 1B – Aqueous Chemistry Acid Strength – Strong Acids These are characterized by a K a >>1 in which the products dominate to an extent that no reactant is expected Also said to fully dissociate Example: HCl HCl(aq) → H + (aq) + Cl - (aq) No HCl(aq) expected Other strong acids: HBr, HI HNO 3 HClO 4 H 2 SO 4 (only first loss of H + )

Chem 1B – Aqueous Chemistry Acid Strength – Weak Acids Partially dissociate in water (so for HA, it will exist to some extent as both HA(aq) and A - (aq) These are characterized by 0.01 > K a > Examples: –HC 2 H 3 O 2 – acetic acid (in vinegar) –HCHO 2 – formic acid (used by some ants) –HClO – hypochlorous acid (bleach) Stronger weak acids have larger K a values Which is stronger: HCN (K a = 4.9 x ) or HC 2 H 3 O 2 (K a = 1.8 x )?

Chem 1B – Aqueous Chemistry Autoprotolysis of Water and pH Water, and some other “protic” solvents, reacts with itself as both an acid and a base 2H 2 O(l) ↔ H 3 O + (aq) + OH - (aq) or H 2 O(l) ↔ H + (aq) + OH - K = K w = [H + ][OH - ] = 1.0 x (at 25ºC) For pure water (using an ICE approach), we can show [H + ] = [OH - ] (= x) And K w = [H + ][OH - ] = [H + ] 2 or [H + ] = (K w ) 0.5 Or [H + ] = [OH - ] = 1.0 x M An acidic solution is where [H + ] > [OH - ] or [H + ] > 1.0 x M

Chem 1B – Aqueous Chemistry Autoprotolysis of Water and pH The pH scale is based on [H + ] but on a log scale pH = -log[H + ] (note: inverse relationship between acid conc. and pH) Very acidic solution (1 M HCl) pH = 0.0 Neutral solution [H + ] = [OH - ] = M: pH = 7.0 Very basic solution (1 M KOH) pH = 14.0

Chem 1B – Aqueous Chemistry Acid Strength and pH A Few Questions: 1)An unknown acid is dissolved in water so that [HA] o = M. The pH is measured and found to be Is this a strong or a weak acid? 2)ClO - is known as the: _________ __________(acid or base) of HClO? 3)At around 45ºC, K w = 4.0 x What is neutral pH at this K w value? 4)At a pH of a solution at equilibrium is (and at 25ºC where K w = 1.0 x ), what is [OH - ]?

Chem 1B – Aqueous Chemistry Equilibrium Problems Involving Acids Strong Acids (large K a values) Example: Determine the pH of a M HCl solution We could set up an ICE table, but with strong acids, we assume 100% dissociation HCl(aq) → H + (aq) + Cl - (aq) So [H + ] = [HCl(aq)] o = M pH = -log[H + ] = 1.85 Note: this approach works as long as [H + ] from HA is greater than [H + ] from H 2 O

Chem 1B – Aqueous Chemistry Equilibrium Problems Involving Acids Weak Acids (K a values < 1) Example: Determine the pH of HCHO 2 (formic acid, K a = 1.8 x ) solution initially made to be 0.20 M.