Chapter 8.4

Reminder: 1 mole of a substance = the mass of the substance

Example: sodium chloride•Na = 23.0 amu 1 mole = 23.0 grams•Cl = 35.5 amu 1 mole = 35.5 grams

•1 mole of NaCl = 23.0 + 35.5 = 58.5 grams

One liter of 1M NaCl solution contains 58.5 grams of NaCl.

* To compare the number of solute particles in solutions, chemists often use moles to measure concentration.

* Molarity: moles of a solute per liter of solution

or Mole

L

* pH scale: 0 to 14

* Describes concentration of hydronium ions (H3O+)

* pH 7 is neutral

* Acids: pH < 7 (0-6)

* Bases: pH > 7 (8-14)

* The pH scale classifies solutions as acids or bases.

* Pure water ionizes slightly

* Arrow pointing left is longer than pointing right because:

* The reaction favors the reactant

* Water contains many more water molecules than ions.

* Pure water is neutral

* Small but equal concentrations of:

*Hydronium ions [H3O+]

*Hydroxide ions [OH–]

*At 25°C both [H3O+] and [OH–] is 1.0 × 10–7

M

(in pure water)

* pH is related to the exponent of the molarity of [H3O

+]

•Pure water has a pH of 7.

* Concentrations of H3O+ and of OH– behave like

they’re on a teeter totter

* Adding acid to water increases [H3O+] and decreases

[OH–]

* Example: 0.1M Hydrochloric acid solution

* Concentration of H3O+ is 1.0 × 10–1 M

* Concentration of OH– is 1.0 × 10–13 M

* pH is related to the exponent of the molarity of H3O+

* pH = 1

*When acids and bases form ions in solution:

*Sometimes involves complete dissociation

*Strong Acid or Base

*Sometimes only partially ionize

*Weak Acid or Base

* Quick reminder:* When reactions go to completion: show with “”

* When reactions reach equilibrium : show with “ ”

*Strong Acids and Bases* Formation of ions from the solute goes to completion.

* Examples:

* Hydrochloric Acid is a strong acid - total ionization: HCl + H2O H3O

+ + Cl–

* Sodium Hydroxide is a strong base – total dissociation NaOH Na+ + OH–

*Weak Acids and Bases

*Ionize or dissociate only partially in water.

*Most of the hydrogens and hydroxides continue to hang on

*Only a few go off on their own (dissociate)

*A solution of acetic acid, CH3COOH, and water can be described by the following equation:

*Equilibrium favors reactants over products*few ions form in solution.

* Two acids of same molarity (concentration):

Weak acid:

* Forms fewer ions (dissociates less)*Most of the weak acids still hanging on to their protons

* Lower [H3O+] gives higher pH (closer to neutral)

Strong acid:

* Forms lots of ions (dissociates almost completely)*Most of the strong acids given up all of their protons

* Higher [H3O+] gives lower pH (farther from neutral)

* Concentration and strength both affect pH.*Concentration: molarity (amount of solute dissolved in a given amount of solution).

*Strength: solute’s tendency to form ions in water*Strong : total dissociation

*Weak: partial dissociation.

*There are only 6 common strong acids:

*HCl - hydrochloric acid

*HBr - hydrobromic acid

*HI - hydroiodic acid

*HNO3 - nitric acid

*H2SO4 - sulfuric acid

*HClO4 - perchloric acid

*Common Strong bases come from the hydroxides of metals in Group 1A & 2A.

*Most common are:

*LiOH - lithium hydroxide

*NaOH - sodium hydroxide

*KOH - potassium hydroxide

*RbOH - rubidium hydroxide

*CsOH - cesium hydroxide

*Ca(OH)2 - calcium hydroxide

*Sr(OH)2 - strontium hydroxide

*Ba(OH)2 - barium hydroxide

* Buffer : a solution that is resistant to large changes in pH. *Weak acids and bases can be used to make buffers.

*Buffers can be prepared by mixing:

* a weak acid and its salt

or

*a weak base and its salt.

* Critical for human body to maintain stable pH*Many cellular reactions very sensitive to pH

*Many cellular reactions create excess hydronium ions

* CO2 dissolved in blood forms carbonic acid - a weak acid. [CO2 + H2O H2CO3]

* Carbonic acid and bicarbonate ions form an important pH buffer [H2CO3 HCO3- + H+]

* Carbon dioxide is exhaled, shifting the equilibrium:

CO2 + H2O H2CO3 HCO3- + H+]