1. Chem. 1B – 9/15 Lecture correction to slide #10 Definition of Lewis Acids and Bases

2. Announcements I Lab –Starting Experiment 7 Wed./Thurs. –Remember to complete the pre-lab before lab Website –I’m posting example exams (note: Exam 2 is closer to what we are covering now), and quiz solutions Homework –Ch. 14 – due tonight –Next Set due 9/24 (covers part of Ch. 15)

3. Announcements II Today’s Lecture – Chapter 15 Topics –Acid Definitions (Arrhenius, Brønsted-Lowry, and Lewis definitions) –Strength of Acids and K a values –Autoionization of Water and the pH Scale –Equilibrium Problems involving Acids (if time)

4. Chapter 14 Solving Equilibrium Problems Additional information toward problem solving: –Problems where we are calculating equilibrium concentrations and given initial concentrations (show where on flow diagram) –ICE table needed, but some modifications sometimes needed too: –Example Problem: 2A(g) ↔ B(g) –If K is small, can often ignore x terms in equations (for above, at equilibrium [A] = [A] o – 2x ≈ [A]) –If K is large (probably not on exam), can first treat as limiting reagent and find full forward concentrations ([B] full to products = [A] o /2), then look at backwards reactions

5. Chem 1B - Equilibrium Equilibrium Problems – Overview Does problem ask to calculate K or an unknown concentration at equilibrium? KUnknown conc. Are concentrations of all species given at equilibrium? Yes No ICE table needed ICE table needed along with given equil. conc. No Are concentrations of all but 1 species given at equilibrium? Yes No ICE table needed ICE table needed No

6. 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

7. 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

8. 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)

9. 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

10. 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

11. 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)]

12. 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 + )

13. 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 > 10 -12 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 10 -10 ) or HC 2 H 3 O 2 (K a = 1.8 x 10 -5 )?

14. 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 10 -14 (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 10 -7 M An acidic solution is where [H + ] > [OH - ] or [H + ] > 1.0 x 10 -7 M

15. 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 - ] = 10 -7 M: pH = 7.0 Very basic solution (1 M KOH) pH = 14.0

16. Chem 1B – Aqueous Chemistry Acid Strength and pH A Few Questions: 1)An unknown acid is dissolved in water so that [HA] o = 0.010 M. The pH is measured and found to be 3.90. 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 10 -14. What is neutral pH at this K w value? 4)At a pH of a solution at equilibrium is 12.11 (and at 25ºC where K w = 1.0 x 10 -14 ), what is [OH - ]?

17. Chem 1B – Aqueous Chemistry Equilibrium Problems Involving Acids Strong Acids (large K a values) Example: Determine the pH of a 0.014 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 = 0.014 M pH = -log[H + ] = 1.85 Note: this approach works as long as [H + ] from HA is greater than [H + ] from H 2 O

18. 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 10 -4 ) solution initially made to be 0.20 M.