ANCH WEEK 7: Aqueous Solutions and Chemical Equilibria

• Chemical Equilibrium


  Most analytical techniques require this state
• Chemical Equilibrium


  Rate of a forward process or reaction and that of the reverse process are equal
• Chemical Equilibrium


  Rate ( Formula ) = Rate ( Reaction )
• Electrolytes


  forms ions when dissolved in water (or certain solvents) and thus produce solutions that conduct electricity
• Strong Electrolytes


  ionize essentially completely in a solvent
• Weak Electrolytes


  ionize ionize only partially in a solvent
• Classification of Strong Electrolytes


  1. Inorganic acids such as HNO3, HClO4, H2SO4, HCl, HBr, HClO3, HBrO3
  2. Alkali and alkaline-earth hydroxides
  3. Most salts
• Classification of Weak Electrolytes


  1. Many inorganic acids, including H2CO3, H3BO3, H3PO4, H2S, H2SO3
  2. Most organic acids
  3. Ammonia and most organic bases
  4. Halides, cyanides, and thiocyanates of Hg, Zn, and Cd
• According to Brønsted-Lowry Theory


  • ACID : proton donor
  • BASE : proton acceptor
• ACIDS


  • Dissociable substance YIELDS hydrogen ions
• BASE


  • Dissociable substance YIELDS hydroxyl ions
• Conjugate Base


  a potential proton acceptor
• Conjugate Acid/Base pair or Conjugate Pair


  Acid + Base
• Conjugate Acid


  every base accepts proton
• Acid/Base or Neutralization


  The two process (conjugate acid and base) are combined
• Amphiprotic Species


  Species that have both acidic and basic properties with dihydrogen phosphate ion , H2PO4 , as an example.
• Water 


  is a classic examples of an amphiprotic solvent. Other common amphiprotic solvent: methanol , ethanol , anhydrous acetic acid
• Autoprotolysis


  Amphiprotic solvents undergo self-ionization, to form a pair of ionic species. It is also another example of acid/base behavior.

  The extent to which water undergoes autoprotolysis at room temperature ( 22-25 degree Celcius ) is slight
• Strong Acids


  reaction with the solvent is sufficiently complete that no undissociated solute molecules are left in aqueous solution
• Weak Acids


  react incompletely with water to give solutions containing significant quantities of both parent acid and conjugate base. Note that acids can be cationic (negative), anionic (positive), or electrically neutral
• Strongest Base = Weakest Acid


  Strongest Acid = Weakest Base
• Equilibrium-constant expressions


  are algebraic equations that can describe the concentration relationships among reactants and products at equilibrium. Also permits calculation of the error in an analysis resulting from the quantity of unreacted analyte that remains when equilibrium has been reached
• Many reactions in analytical chemistry never result complete conversion and proceeds to a state of chemical equilibrium in which the ratio of the reactants and products are constant
• Concentration relationship at chemical equilibrium can be altered by applying stresses :

  • Changes in temperature
  • Changes in pressure (gas)
  • in total concentration of reactant or product

  These effects can be predicted qualitatively by Le Chatelier's Principle
• Le Chatelier's Principle


  principle state that the position of the chemical equilibrium always shifts in a direction that tends to relieve the effect of an applied stress

  Gases - effects the equilibrium

  Solids and Liquids - has no effects

  
• Catalyst


  speeds up the rate of reaction
• Enzymes

  catalyst of the body

  HOT - degrades

  COLD - slow reaction

  Big deviation in pH - degradation of enzymes
• The addition of arsenic acid (H3AsO4) and hydrogen ions

  causes an increase in color as more triiodide ion and arsenous acid are formed.

  Adding arsenous acid has the reverse effect
• Mass-Action Effect


  is a shift in the position of an equilibrium caused by adding one of the reactants or products to a system
• Equilibrium-Constant Expressions


  • influence of concentration or pressure (if gases) described in quantitative terms
  • derived from thermodynamics
  • important because it allows us to predict the direction and completeness of chemical reactions
  • yields no information concerning the rate of reaction
• Buffer solution


  resist changes in pH when diluted, or acids or bases are added to it
• Buffers


  maintain the pH of a solution at a constant and predetermined level
• Henderson-Hasselbalch Equation


  used to calculate the pH of buffer solutions, frequently encountered in the biological literature and biochemical texts

  pKa = 6.1 (combine hydration and dissociation constants for CO2 in blood)

  conjugate base = bicarbonate, HCO3, kidneys

  weak acid = carbonic acid, H2CO3, lungs

  
• Normal pH range of blood

  pH 7.35-7.45
• Perfect or average pH of blood

  pH 7.4
• Effect of Dilution


  pH of buffer solutions remains essentially independent of dilution
• ACIDS



  • Dissociable substance ACCEPTS hydroxyl ions
• BASE


  • Dissociable substance ACCEPTS hydrogen ions
• If the reactant is high and the product is low it will shift to the:

  right
• If the reactant is low and the product is high it will shift to the:

  left