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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+]