3. Water - can act as both an acid and base

Can react with itself to form both H₃O⁺  and OH^-

H₂O + H₂O -> H₃O⁺ + OH^-

K_w = [H₃O⁺][OH^-] = 1 x 10^-14  ( ion product constant of water)

This reaction, termed the autoionization of water, maintains the product of the H₃O⁺ and OH^- concentrations constant.  It provides the link between [H₃O⁺] and [OH^-] in aqueous solutions.  Cannot affect one without affecting the other.  Add an acid, the H₃O⁺ concentration goes up so the OH^- concentration must go down.

Neutralization: 

Acid + base --> water + salt

NaOH + HCl --> H₂O  + NaCl (neutral, not acidic or basic)

The acid and base in effect cancel each others influence.  No production of OH^- or H⁺ just H₂O (we will return to this later)

Bronsted-Lowry definition-

Does not have water in the definition so more general.

Acid - proton donor

Base - proton acceptor

Using this definition, we can look at a reaction such as:

NH₃ + HCl --> NH₄⁺ + Cl^-

and see that NH₃ accepted a proton so it is the base and HCl donated a proton so it is the acid.

Acidity, basicity and neutrality:

 Compounds can be acids or bases

Acids – generate H⁺ (H₃O⁺) in aqueous solution

Bases – generate OH^- in aqueous solutions

 When adding an acid (or base) to water, the amount of H⁺ (or OH^-) generated depends on the strength of the acid (or base) and the concentration of acid (or base).  Do not get acid (or base) strength and concentration confused. 

  How do describe the [H⁺] in solution - acidic basic or neutral solutions

Whether a solution is acidic, basic or neutral depends on the [H⁺], which is always related to the [OH^-] by the ion product of water K_w

  Pure water is neutral:

  In such solutions there is some H⁺ and OH^- generated from the autoionization of water

  [H⁺] = [OH^-] in neutral solutions so [H⁺][OH^-] = [H⁺]² = [OH^-]² = 1 x 10^-14 so [H⁺] = [OH^-] = (1x10^-14)^1/2 = 1x10^-7 M.  This is a small concentration (compare it to the molarity of water in pure water, which is 55M)

Acidic solutions – [H⁺] > 1 x 10^-7 M  (result from the addition of an acid to pure water)

Neutral solutions – [H⁺] = 1 x 10^-7 M

Basic solutions – [H⁺] < 1 x 10^-7 M (results from the addition of a base to pure water)

    [H⁺] vs. [OH^-]   

Because [H⁺][OH^-] = constant = 1 x 10^-14, if we increase [H⁺]  the [OH^-] must decrease.  If we increase [OH^-], the [H⁺] must decrease.

  Ex: Calculating the [OH^-] in solution – use K_w

(a)  what is the [OH^-] in a solution with [H⁺] = 1x10^-3 M

  [OH^-] = 1x10^-14 / 1x10^-3 = 1x10^-11 M

  (b) what is the [OH^-] in a solution with [H⁺] = 1 x 10^-9 M

  [OH^-] = 1x10^-14 / 1x10^-9 = 1x10^-5 M

    The pH scale –

  A more convenient way to express the [H⁺] in solution

  pH = - log[H⁺]

  Problem: what is the pH of a solution with [H⁺] = 1x10^-2 M

  pH = -log(1x10^-2) = 2

  Problem: what is the [H⁺] and [OH^-] in a solution with pH = 9

  pH = -log[H⁺]

  [H⁺] = antilog(-pH) = 10^-pH = 10^-9 M

  [OH^-] = K_w / [H⁺] = 10^-14 / 10^-9 = 10^-5 M

  Controlling pH is important in many facets of our lives:

-Agriculture: Plants grow best in proper pH
-Physiology: change body pH by 1 unit, we die.  Many biochemical reactions pH dependent
-Industry – pH can greatly influence reactivity so its control is essential to the manufacture of chemicals
-Municipal services – pH control important in sewage treatment and purification of drinking water.

  Overhead – pH of common substances:

  Measuring pH

 -pH meters
-indicators

indicators are acids/bases themselves.

An indicator has an acid form and a base form:

If we add acid, the base form gets protonated and we see the color of the acidic form

  In^- + HCl à HIn + Cl^-

  If we add base, the acid gets deprotonated and we see the color of the basic form

  HIn + OH^-  à In^- + H­₂O

Neutralization reactions -

 An acid lowers the pH of a solution.  A base raises the pH of a solution.  What happens if you mix them together?

The reaction of a strong acid with a strong base will result in the formation of water and a salt of neutral pH. 

If just enough acid is added to a strong base, a neutral solution will result.

If just enough base is added to a strong acid, a neutral solution will result.

  Stoichiometry of neutralization reactions - What is just enough?

  HCl + NaOH à H₂O + NaCl (one mole of acid required to neutralize one mole of base)

  The stoichiometry is not always one-to-one

  H₂SO₄ + 2NaOH à 2H₂O + Na­­₂SO­­₄

Notice is takes two moles of the base here to neutralize one mole of the sulfuric acid because it can contribute two protons.

  Normality (N) – more convenient for reactions between acids and bases

based on the concept of an equivalent

one equivalent of acid can generate 1 mol of H⁺

one equivalent of base can generate 1 mol of OH^-

Normality = equivalents / L  (Volume*normality = VN = moles of equivalents)

1M HCl = 1N

1M H₂SO₄ = 2N

  The normality of a base is the moles of OH^- it can contribute per liter of solution

  1M NaOH = 1N

1M Ca(OH)₂ = 2N

To neutralize an acidic or basic solution, an equal number of equivalents of OH^-  and H⁺ must react.

Problem: What volume of 1M KOH must be added to neutralize 100mL of 0.25M H₂SO₄.  What will the pH of the resulting solution be?

In neutralization reactions, it is easiest to deal with equivalents:

For neutralization to occur, we need to have the number of equivalents of base = number of equivalents of acid

N_baseV_base = N_acidV_acid

The 1M KOH solution is 1N

The 0.25 M H₂SO₄ solution is 0.5N

(1N)(0.1L) = (0.5N)(V_acid)

V_acid = 5 L

Neutralization reactions results in the formation of salts with little or no acid base character and water. 

2 KOH + H₂SO­₄ à K₂SO₄ + H₂O

  Consequently, the pH of the solutions will be near neutral = 7