Aqueous Equilibria Part 1

Aqueous Equilibria, Part 1 Name: ______

AP Chemistry Lecture Outline

Acids and Bases

Arrhenius ...acids increase the ______when dissolved in H2O.

...bases increase the ______when dissolved in H2O.

e.g., HCl and NaOH

Bronsted-Lowry acids:

Bronsted-Lowry bases:

-- H+ and H3O+ are used interchangeably

-- concept is not limited to aqueous solutions

Arrhenius and Bronsted-Lowry definitions overlap, in many cases.

NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)

A substance can’t be a Bronsted-Lowry acid unless...

-- the acid must be able to lose H+

-- the base must have a nonbonding pair of e– that can bind with the H+

Amphoteric substances can be acids or bases, depending on the reaction conditions.

HCl(aq) + H2O(l)

NH3(aq) + H2O(l)

-- In acid-base equilibria, protons are donated in forward and reverse reactions.

-- The two substances in a conjugate acid-base pair differ by a H+...

EX. HNO2(aq) + H2O(l) NO2–(aq) + H3O+(aq)

-- Strong acids / bases easily / ______H+.

-- Weak acids / bases do NOT easily / ______H+.

In every acid-base reaction, the position of equilibrium favors the

transfer of protons from the stronger acid to the stronger base.

The Autoionization of Water

In ordinary water, we constantly have... H2O(l) + H2O(l) H3O+(aq) + OH–(aq)

-- reaction is very rapid in each direction

-- at room temp. ~1 out of a billion m’cules are ionized, so pure water is a poor conductor

For the above equation, Kc =

If we exclude the pure liquid...

This equation is taken to be valid for pure water and for dilute aqueous solutions.

[ H+ ] > [ OH– ] [ H+ ] < [ OH– ] [ H+ ] = [ OH– ]

______

EX. At 25oC, calculate the hydrogen ion concentration if the hydroxide ion is 2.7 x 10–4 M.

Is this solution an acid or a base?

The pH Scale

Definition of pH... 0 ACID 7 BASE 14

A change of 1 pH unit means that [ H+ ] (and [ OH– ]) has changed by a factor of _____.

EX. A solution with pH 2 is how many times more acidic than one with pH 5?

Especially in biochemical reactions, the rate law is often dependent on the hydrogen ion concentration. If the rate is 1st order in [ H+ ], then doubling [ H+ ] would double the rate.

-- Human blood pH is between 7.35 and 7.45.

** Significant figures rule for logarithms:

EX. Find the pH if the hydronium ion concentration is 5.2 x 10–9 M.

EX. Find the pH if the hydroxide ion concentration is 3.9 x 10–10 M.

Other equations:

EX. Find [ H+ ] if the pH = 5.22.

Measuring pH

pH meter

acid-base indicators:

e.g.,

pH paper is paper impregnated with mixtures of various indicators

Strong Acids and Bases

-- these are strong electrolytes that exist entirely as ions in aqueous solution

-- memorize the names and formulas of the seven strong acids...

...and the eight strong, hydroxide bases...

Strong electrolytes are often written using a one-sided arrow.

e.g., HNO3(aq) H+(aq) + NO3–(aq)

EX. Find the pH of a 0.012 M perchloric acid solution.

EX. If a nitric acid solution has pH = 2.66, find the solution’s concentration.

** NOTE: Ca(OH)2, Sr(OH)2, and Ba(OH)2 are strong electrolytes, but have limited solubilities. This means that if you put a big chunk of any of these into water, little of the chunk will dissolve, BUT the little that does dissolve will float around as X2+(aq) and

OH–(aq), NOT as X(OH)2(aq).

EX. Find the pH of a 0.0034 M solution of calcium hydroxide.

Metal oxides, hydrides, and nitrides (containing O2–, H–, and N3–, respectively) also react with water to produce strongly basic solutions.

e.g., --

Weak Acids

-- most acids are weak

-- For a weak acid HX... HX(aq) H+(aq) + X–(aq)

-- acid-dissociation constant Ka =

large Ka:

small Ka:

The % of a weak acid that is ionized is given by the equation:

For organic acids (containing only C, H, and O) the “donated” H was connected to...

EX. A 0.020 M niacin solution has pH 3.26.

(a) What % of the acid is ionized?

(b) What is Ka?

EX. If Ka for niacin is 1.6 x 10–5, find the pH of a 0.010 M niacin solution.

% ionization of a weak acid at a given temperature...

Rationale:

Recall that for the weak acid HX(aq) H+(aq) + X–(aq)

If we increase [ HX ], particle [ ] increases. System doesn’t “want” higher [ ] of particles, so it shifts LEFT to reduce the number of particles (one vs. two). [ H+ ] will increase, but not as much as [ HX ]. Therefore, % ionization decreases.

-- The opposite is true if we dilute the solution.

EX. Calculate the % of HF molecules ionized in a 0.10 M HF solution. (Ka = 6.8 x 10–4)

EX. Calculate the % of HF molecules ionized in a 0.010 M HF solution. (Ka = 6.8 x 10–4)

Polyprotic acids – like sulfurous acid, H2SO3 – have more than one ionizable H.

H2SO3(aq) H+(aq) + HSO3–(aq) (Ka1 = 1.7 x 10–2)

HSO3–(aq) H+(aq) + SO32–(aq) (Ka2 = 6.4 x 10–8)

-- Usually, Ka2 is at least 1000X smaller than Ka1. In such cases, one can calculate [ H+ ]

and pH based only on Ka1 (i.e., ignore Ka2 and pretend you have a monoprotic acid).

EX. Find the pH of a 0.0037 M carbonic acid solution. (Ka1 = 4.3 x 10–7, Ka2 = 5.6 x 10–11)

Weak Bases

weak base + H2O conjugate acid + OH–

Kb =

Weak bases are often nitrogen-containing molecules (“amines”) or anions.

Example: H–N–CH3 (aq) + H2O(l) H–N–CH3 (aq) + OH–(aq)

Anions of Weak Acids

NaClO(aq) Na+(aq) + ClO–(aq)

EX. What mass of NaClO is required to make 2.0 L of a pH 10.50 solution?

The Kb for ClO– is 3.3 x 10–7.

EX. What is the [ NH3 ] in a solution having pH 9.35?

Relationship Between Ka and Kb

When two reactions are added to give a third, then K3 = K1K2, where the Ks are the equilibrium constants. For example...

(1) NH4+ NH3 + H+ (Ka = 5.6 x 10–10)

(2) NH3 + H2O NH4+ + OH– (Kb = 1.8 x 10–5)

ADD...

Because they are easily calculated from Ka And just like pH = –log [ H+ ],

values, reference tables often contain no Kb values. pKa = –log Ka

pKb = –log Kb

pKa + pKb = 14.00 (@ 25oC)

EX. For hydrofluoric acid, if Ka = 6.8 x 10–4...

(a) ...find pKa

(b) ...identify the conjugate base

Acid-Base Properties of Salt Solutions

Since ions can exhibit acid or base properties (i.e., donate or accept protons), salt solutions can be acidic or basic.

-- Nearly all salts are strong electrolytes.

-- hydrolysis:

Anions derived from weak acids react with water to form OH–, and are thus basic.

Example: Consider the weak acid HX(aq) H+(aq) + X–(aq)

Now, say we have X–(aq) in solution from a salt.

We’ll get...

i.e.,

Anions of strong acids do NOT influence pH.

e.g.,

Anions that still have ionizable protons (e.g., HSO3–) are amphoteric.

EX. Does Na2HPO4 form an acidic or a basic solution in water?

Ans. Real question is: How does HPO42– behave?

Like an acid...

...or a base?

From a reference table, we find...

(a) H3PO4 H+ + H2PO4–

(b) H2PO4– H+ + HPO42–

All cations except (remember the “b” for strong hydroxide bases?) act as ______in aqueous solutions.

Predicting pH for salts derived from a...

(1) ...strong base and a strong acid

e.g.,

(2) ...strong base and a weak acid

e.g.,

(3) ...weak base and a strong acid

e.g.,

(4) ...weak base and a weak acid

e.g.,

Acid-Base Behavior and Chemical Structure

Three Factors Affecting Acid Strength

1. To transfer H+, the acid must have the polarity...

-- ionic hydrides don’t

-- C–H bonds don’t

As polarity increases, acid strength generally increases. BUT...

2. ...if H–X bond is too strong.. (e.g., )

3. Stable conjugate bases usually indicate a...

For binary acids:

(1) Down a group...bond strength is the determining factor.

...bond strength ___ and acid strength ___

(2) Across a period...polarity governs.

...polarity ___ and acid strength

oxyacids:

Consider H2SO4 (an acid) and Ca(OH)2 (a base).

H–O–S–O–H H–O–Ca–O–H

The electronegativity of S is similar The electronegativity of Ca is low,

to that of O, so the S–O bonds are while for O, it is high. Thus, the Ca–O

covalent. This means that... bonds are ionic. This means that...

Other tips:

(1) For oxyacids with the same number of OH groups and/or the same number of O

atoms, acid strength increases with increasing electronegativity of the central atom.

e.g., HClO

HBrO

HIO

(2) For oxyacids with the same central atom, acid strength increases as the number of

oxygens attached to the central atom increases.

e.g., HClO HClO2 HClO3 HClO4

Carboxylic acids contain the ______.

-- these are the largest category of organic acids

-- acid strength increases with the addition of more electronegative atoms

e.g., CH3COOH vs. CF3COOH

Lewis Acids and Lewis Bases

The “Lewis” definitions greatly broaden

the range of acids because many

species other than H-containing ones

can accept an e– pair.

e.g.,

The simple term “acid” suggests that we are referring to an Arrhenius or a Bronsted-Lowry acid, i.e., an H-containing substance in an aqueous solution. If you are referring to a Lewis acid, then use the term “Lewis acid.”

Substances with an incomplete octet (e.g., BF3) or ones having vacant orbitals (e.g., Fe3+) can function as Lewis acids.

-- The (+) charge attracts (i.e., accepts) the lone pairs of e– on the O of a water molecule.

This process is hydration, and it is responsible for most salts dissolving in H2O.

Cation size and cation charge determine the extent to which the pH is affected.

-- short distance -- larger distance

-- strong interaction -- weaker interaction

-- --