4TH QUARTERLY

• Spontaneous process


  Any process which occurs without outside intervention
• Spontaneous processes


  • Ice turning to water at T>0°C
  • Water turning to ice at T<0°C
  • A ball rolling down an incline
  • Iron rusting
  • Water flowing downhill
• Non-spontaneous process


  Any process which occurs with outside intervention, needs energy to carry out
• Reversible process


  One which can go back and forth between states along the same path
• Reversible processes


  • Rolling up a ball on inclined plane
  • Diffusion of gas from low pressure to high pressure
  • Flow of water uphill
• Entropy (S)


  Measure of the disorder of a system
• State function


  A property whose value does not depend on the path taken to reach that specific value. Entropy is a state function.
• Trends in entropy


  • Physical state (Gas > Liquid > Solid)
  • Temperature (Increases as temperature increases)
  • Molecular size (Larger molecules have higher entropy)
  • Forces between particles in a solid (Higher entropy if weak forces)
  • Dissolution and mixing (Entropy of solid increases when dissolved in solution)
• Second law of thermodynamics


  Entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process. Explains why spontaneous processes have a direction.
• Spontaneous processes


  • Heat flows in the direction of decreasing temperature
• Molecular interpretation of entropy


  A gas is less ordered than a liquid which is less ordered than a solid. Any process which increases the number of gas molecules leads to an increase in entropy, and vice versa.
• Three atomic models of motion


  • Translation (Moving of a molecule from one point in space to another)
  • Vibration (Shortening and lengthening of bonds, including the change in bond angles)
  • Rotation (Spinning of a molecule about some axis)
• Energy is required to get a molecule to translate, vibrate, or rotate. The more energy stored in translation, vibration, and rotation, the greater the degrees of freedom and the higher the entropy.
• Gibbs free energy


  For a spontaneous reaction, the entropy of the universe must increase. Reactions with large negative ΔH values are spontaneous.
• Conditions for free energy

  • ΔG < 0 = Forward reaction is spontaneous
  • ΔG = 0 = Reaction is at equilibrium / no reaction
  • ΔG > 0 = Forward reaction is not spontaneous
• Third law of thermochemistry

  The entropy of a perfect crystalline substance is zero at the absolute zero temperature.
• Absolute entropies


  The reference point for entropy of a substance is entropy at 0 K.
• Chemical equilibrium


  Rates of the forward and reverse reactions are equal and the concentration of the reactants and products remain constant.
• Concept of equilibrium


  • As a system APPROACHES equilibrium, both the forward and reverse reactions are occurring
  • At EQUILIBRIUM, the forward and reverse reactions are proceeding at the same rate
  • Once equilibrium is ACHIEVED, the amount of each reactant and product remains constant
• Equilibrium constant (Kc)


  The ratio of the equilibrium concentrations of products over the equilibrium concentrations of reactants each raised to the power of their stoichiometric coefficients.
• Equilibrium constant interpretation


  • K > 1 = product favored
  • K < 1 = reactant favored
• Reaction quotient (Q)


  The ratio of the initial concentration of the products to the initial concentrations of the reactants using the equilibrium expression.
• Reaction quotient interpretation


  • Q>K = Shifts to the left
  • Q<K = Shifts to the right
  • Q=K = At equilibrium
• Le Chatelier's principle


  When a system at equilibrium is stressed, the system works to restore equilibrium.
• Stresses on equilibrium


  • Changes in the concentration of reactants and products
  • Changes to the pressure (for gas systems)
  • Changes to temperature
• Whenever you try to increase something, the system will try to decrease it.
• Arrhenius theory


  Acids are substances that dissociate in water to yield electrically charged atoms or molecules, called ions, one of which is a hydrogen ion (H+), and bases ionize in water to yield hydroxide ions (OH-).
• Arrhenius acid


  Any species that can increase the concentration of H+ in aqueous solution.
• Arrhenius base


  Any species that can increase the concentration of OH- in aqueous solution.
• Bronsted-Lowry theory


  Acids are defined as proton (hydrogen ion) donors, and bases are defined as proton (hydrogen ion) acceptors. A compound that acts as both Bronted-Lowry acid and base together are called Amphoteric.
• Conjugate acid


  Product that results from protonation of a Bronsted-Lowry base. Positive.
• Conjugate base


  Anion that results from deprotonation of a Bronted-Lowry acid. Negative.
• Acid and base strength

  • Strong acids are completely dissociated in water, their conjugate bases are weak
  • Weak acids only dissociate partially in water, their conjugate bases are strong
  • The stronger a base, the weaker its conjugate acid
  • The weaker a base, the stronger its conjugate acid
  • In any acid base reaction, equilibrium will favor the reaction that moves the proton to the stronger base