Chapter 6 Problems 6-29, 6-31, 6-39, 6.41, 6-42, 6-48,

Outline Equilibrium of Acids and Bases Bronsted-Lowry Acids/Bases Define strong Define weak pH of pure water at 25 o C Define K a and K b Relationship b/w K a and K b Chapter 8 – Activity Relationship with K

Equilibrium Acids and Bases & Equilibrium Section 6-7

Strong Bronsted-Lowry Acid A strong Bronsted-Lowry Acid is one that donates all of its acidic protons to water molecules in aqueous solution. (Water is base – electron donor or the proton acceptor). HCl as example

Strong Bronsted-Lowry Base Accepts protons from water molecules to form an amount of hydroxide ion, OH -, equivalent to the amount of base added. Example: NH 2 - (the amide ion)

Question Can you think of a salt that when dissolved in water is not an acid nor a base? Can you think of a salt that when dissolved in water IS an acid or base?

Weak Bronsted-Lowry acid One that DOES not donate all of its acidic protons to water molecules in aqueous solution. Example? Equilibrium

Weak Bronsted-Lowry base Does NOT accept an amount of protons equivalent to the amount of base added, so the hydroxide ion in a weak base solution is not equivalent to the concentration of base added. example: NH 3

Common Classes of Weak Acids and Bases Weak Acids carboxylic acids ammonium ions Weak Bases amines carboxylate anion

Question: Question: Calculate the Concentration of H + and OH - in Pure water at 25 0 C. Equilibrium and Water

EXAMPLE: Calculate the Concentration of H + and OH - in Pure water at 25 0 C. H 2 O H + + OH - Initialliquid-- Change-x+x Equilibrium Liquid-x +x K W = (X)(X) = ? K w = [H + ][OH - ] = ?

EXAMPLE: Calculate the Concentration of H + and OH - in Pure water at 25 0 C. H 2 O H + + OH - Initialliquid-- Change-x+x Equilibrium Liquid-x +x K W = (X)(X) = 1.01 X K w = [H + ][OH - ] = 1.01 X (X) = 1.00 X 10 -7

Example What is the concentration of OH - in a solution of water that is 1.0 x M in [H + ] 25 o C)? K w = [H + ][OH - ] 1.0 x = [1 x ][OH - ] 1.0 x = [OH - ] “From now on, assume the temperature to be 25 o C unless otherwise stated.”

pH ~ > ~ +16 pH + pOH = - log K w = pK w = 14.00

Is there such a thing as Pure Water? In most labs the answer is NO Why? A century ago, Kohlrausch and his students found it required to 42 consecutive distillations to reduce the conductivity to a limiting value. CO 2 + H 2 O HCO H +

Weak Acids and Bases HA H + + A - HA + H 2 O (l) H 3 O + + A - KaKa K a ’ s ARE THE SAME

Weak Acids and Bases B + H 2 O BH + + OH - KbKb

Relation Between K a and K b

Relation between Ka and Kb Consider Ammonia and its conjugate acid. NH 3 + H 2 O NH OH - KbKb NH H 2 O NH 3 + H 3 O + KaKa H 2 O + H 2 O OH - + H 3 O +

Example The K a for acetic acid is 1.75 x Find K b for its conjugate base. K w = K a x K b

Example Calculate the hydroxide ion concentration in a M sodium hypochlorite solution. OCl - + H 2 O  HOCl + OH - The acid dissociation constant = 3.0 x 10 -8

1 st Insurance Problem Challenge on page 120

Chapter 8 Activity

Write out the equilibrium constant for the following expression Fe 3+ + SCN -  Fe(SCN) 2+ Q: What happens to K when we add, say KNO 3 ? A: Nothing should happen based on our K, our K is independent of K + & NO 3 -

K decreases when an inert salt is added!!! Why? K eq

8-1 Effect of Ionic Strength on Solubility of Salts Consider a saturated solution of Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg IO 3 - K sp =1.3x seemingly strange effect A seemingly strange effect is observed when a salt such as KNO 3 is added. As more KNO 3 is added to the solution, more solid dissolves until [Hg 2 2+ ] increases to 1.0 x M. Why? ICE some- - -x+x+2x some-x+x+2x

Increased solubility Why? LeChatelier’s Principle? NO – not a product nor reactant Complex Ion? No Hg 2 2+ and IO 3 - do not form complexes with K + or NO 3 -. How else?

The Explanation Consider Hg 2 2+ and the IO Electrostatic attraction

The Explanation Consider Hg 2 2+ and the IO Electrostatic attraction - The Precipitate!! Hg 2 (IO 3 ) 2 (s) The Precipitate!!

The Explanation Consider Hg 2 2+ and the IO Electrostatic attraction Add KNO 3 NO 3 - K+

The Explanation Consider Hg 2 2+ and the IO Add KNO 3 K+ NO 3 - K+

The Explanation Consider Hg 2 2+ and the IO NO 3 - Hg 2 2+ and IO 3 - can’t get CLOSE ENOUGH to form Crystal lattice Or at least it is a lot “Harder” to form crystal lattice - K+

The potassium hydrogen tartrate example

Alright, what do we mean by Ionic strength? Ionic strength is dependent on the number of ions in solution and their charge. Ionic strength (  ) = ½ (c 1 z c 2 z …) Or Ionic strength (m) = ½  c i z i 2

Examples Calculate the ionic strength of (a) 0.1 M solution of KNO 3 and (b) a 0.1 M solution of Na 2 SO 4 (c) a mixture containing 0.1 M KNO 3 and 0.1 M Na 2 SO 4. (  ) = ½ (c 1 z c 2 z …)

Alright, that’s great but how does it affect the equilibrium constant? Activity = A c = [C]  c AND

Relationship between activity and ionic strength Debye-Huckel Equation 2 comments  = ionic strength of solution  = activity coefficient Z = Charge on the species x  = effective diameter of ion (nm) (1)What happens to  when  approaches zero? (2)Most singly charged ions have an effective radius of about 0.3 nm Anyway … we generally don’t need to calculate  – can get it from a table

Activity coefficients are related to the hydrated radius of atoms in molecules

Relationship between  and 

Back to our original problem Consider a saturated solution of Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg IO 3 - K sp =1.3x At low ionic strengths  -> 1

Back to our original problem Consider a saturated solution of Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg IO 3 - K sp =1.3x In 0.1 M KNO 3 - how much Hg 2 2+ will be dissolved?

Back to our original problem Consider a saturated solution of Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg IO 3 - K sp =1.3x10 -18