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 (H 3 O + ) * 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 [H 3 O + ] * Hydroxide ions [OH – ] * At 25°C both [H 3 O + ] and [OH – ] is 1.0 × 10 –7 M (in pure water) * pH is related to the exponent of the molarity of [H 3 O + ] Pure water has a pH of 7.
* Concentrations of H 3 O + and of OH – behave like they’re on a teeter totter * Adding acid to water increases [H 3 O + ] and decreases [OH – ] * Example: 0.1M Hydrochloric acid solution * Concentration of H 3 O + 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 H 3 O + * 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 + H 2 O H 3 O + + 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, CH 3 COOH, 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 [H 3 O + ] 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 [H 3 O + ] 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 * HNO 3 - nitric acid * H 2 SO 4 - sulfuric acid * HClO 4 - 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
* CO 2 dissolved in blood forms carbonic acid - a weak acid. [CO 2 + H 2 O H 2 CO 3 ] * Carbonic acid and bicarbonate ions form an important pH buffer [H 2 CO 3 HCO 3 - + H+] * Carbon dioxide is exhaled, shifting the equilibrium: CO 2 + H 2 O H 2 CO 3 HCO 3 - + H+]