Chapter 19: Reaction Equilbrium

• A + B <==> C + D
  
• Equilibrium: The condition at which the forward reaction rate equals the reverse reaction rate.
  
  

19-1 Reversibility of Reaction

• In a closed system, most reactions are reversible (ie while forward reaction proceeds, so will the reverse reaction)
  
  

19-2 Dynamics of Equilibrium

• an example: Heat + 2HI <==> H2 + I2
  
• (in a closed system)
  
  • At equilibrium, two opposing reactions occur:
    
    1. the decomposition reaction of HI --> H2 + I2
       
    2. the formation of HI from H2 and I2
       
  • Initially: only the forward reaction occurs
    
    the reverse rate is 0 becase there are no products
    Heat + 2HI --> H2 + I2
    
• As the concentration of products increases, the reverse reaction begins to occur
  
  H2 + I2 --> 2HI + Heat
  
• Eventually, two opposing reaction rates become equal and a state of equilibrium is established.
  There is no more macroscopic change
  

Notes on Equilibrium

• a state of equilibrium can be reached by starting with reactants or with products and can be attained at either end
  
• if there is an imbalance, equilibrium will be re-established
  
• at equilibrium, heat consumed equals the heat produced --> remains at constant temperature
  
  

19-3 Recognizing Equilibrium States

• there is a closed system
  
• the reaction is reversible
  
• there is a consistency in macroscopic properties
  
• the temperature is constant
  
  

19-4 Factors Controlling Equilibrium

1. Enthalpy - heat content
   
2. Entropy - degree of randomness
   
   

• The factors just mentioned above determines whether the reaction will go to:
  #1 completion, #2 not occur at all, or #3 the state of equilibrium
  
• Reactions favor a lower heat content (exothermic reactions) and a higher state of randomness (higher entropy - gases are highest and solids are lowest)
  
• entropy favors free ions and a greater number of particles
  
• reactions go to completion when the products have a lower heat content (delta H = -ve) and higher entropy (free ions or more particles)
  
• reactions do not occur at all when the reaction is endothermic and the reactants have a higher entropy than the products
  
• reactions go to equilibrium when the entropy of the reactants is higher than that of the products and the reaction is exothermic...OR
  
• reactions will also go to completion when the products have a higher entropy than that of the reactants and the reaction is endothermic!
  
  

Factors Affecting Equilibrium

Concentration Changes

• equilibrium is temporarily upset
  
• increase in collision frequency and rate of reaction
  
• adding more reactant will increase the amount of products - excess reactants will react until a new state of equilibrium is reached
  
• adding more products will increase the amount of reactant - excess products will react until a new state of equilibrium is reached
  
  

Volume Changes

• decreasing the volume of the conatiner increases the collision frequency (although this happens, the amount of gas does not change)
  
• to overcome the increase in pressure, a reaction will take place in which the products have fewer molecules than the reactants
  
• the increase in pressure is opposed by a tendency to the side having the fewest molecules (lower entropy)
  
• increase in pressure by adding an inert gas (like Helium - a nonreactant) causes NO shift in equilibrium since the concentration of the gases involved in the reaction is unchanged




Temperature Changes

• Increasing the temperature will increase both the forward and reverse reaction rates, but the ENDOTHERMIC rate will increase more than the exothermic rate.
  
  

Catalysts

• A catalyst only increases the forward and reverse rates
  
• both rates increase equally
  
• no change in equilibrium
  
• if a system has not yet attained equilibrium, then equilibrium will be established faster (with the help of good 'ol cataylst!)
  
  

"...When a stress is applied to a system at eqiulibrium, the system will respond by reducing that stress..."

Equilibrium Constant Expression

• For any reversible reaction at equilibrium:
  xA + yB <===> aC + bD the forward rate = the reverse rate
  
  
• there is a constant ratio between the equilibrium concentrations of the products and reactants:
  
  

Keq = [C]^a[D]^b / [A]^x[B]^y

Equilibrium Solution Rules

• All gases, aqueous ions, and aqueous mixtures (eg Fe + SCN <==> FeSCN) are included in the expression
  
• All solids and pure liquids are never included
  
  

Interpreting Keq Values

• How the hell does temperature affect Keq?
  1.) If it's an exothermic reaction, Keq decreases when temperature rises (reactants are favored) and vice versa.
  2.) If it's an endothermic reaction, Keq increases when temperature rises (products are favored) and vice versa
  
• generally, doubling an equation squares the original constant, and halving the equation produces a constant that is the square root of the original constant
  
• the Keq of a reversed equation is the inverse of the original constant
  eg [B] / [A].....Keq = X ................. [A] / [B].....Keq = X^-1 (OR 1/X)
  
  

• the initial concentration of the reactants and products
  
• the change in concentrations (x, 2x, 3x.....nx) (oh yeah, also specify if it's a POSITIVE change or a NEGATIVE change!)
  
• the final concentration of the reactants and products
  

  GOOD LUCK!

  
  
  

• Solutions:a homogenous (has uniform distribution of particles, unlike colloidal suspensions which has obvious segregation) mixture of two or more substances
  
• Solvent: The component of the solution present in greater quantity
  
• Solute: The component in the solution present in a smaller quantity relative to solvent
  
  

Solubility

• Solubility: the maximum amount of solute that will dissolve in a given amount of solvent
  
• stuff that affect solubility include: the nature of the solvent, temperature, pressure and the presence of other solutes
  
• a saturated solution contains the maximum amount of solute that can be dissolved under the given conditions
  * Excess insoluble material will remain undissolved and sitting at the bottom of the container when the solution is saturated
  
  

What the hell is Solubility Equilibrium?

• didja know there is an equilibrium present between the undissolved solute and its ions in the solution? (refer to fig 16-2 in Heath)
  
• therefore: solid <==> ions
  
• OR rate of recrystalization = rate of dissociation of undissolved solid
  * Concentration remains constant as long as temperature remains constant
  
  

What the hell is Miscibility?

• Miscibility only applies to liquid - liquid mixtures
  
• When two liquids are miscible, they completely mix together in all proportions
  
• When two liquids are partially miscible, they have limited solubility with each other, but some of the particles WILL mix
  
• When two liquids are immiscible, they do not dissolve in each other and form two or more distinct layers
  
  

The Solubility Table

• All alkali compounds (Li, Na, K, Ce, Rb, Fr), Ammonium compounds (NH4), and Nitrate compounds (NO3) are SOLUBLE!!
  
• besides that, only if you're REALLY into learning this stuff, go ahead, memorize the whole darned thing! It is advisable b/c you just might need it in the future...
  

Solubility Product

• The Ksp, or Solubility Product Constant is calculated very similarly to the Equilibrium Constant (refer to Ch 19) - you take the concentrations of each ion and multiply them to get a Constant expression. Most of the time, you'll be given the constant expression, and from there, you're to calculate the concentrations if individual ions
  
• remember, most of the Ksp calculations you do are gonna involve compounds that have low solubility
  
• easily said AND done
  
  • First - write down the equation:
    

    Solid --> xCation(aq) + yAnion(aq)

    
    
    
  • 2nd - write down how Ksp would be calculated:
    Ksp = [Cation]^x * [Anion]^y
    
    

Common Ion Effect

• What is the solubility of AgCl in water and 0.01M NaCl? Ksp AgCl = 1.6 * 10^-10

• AgCl <==> Ag + Cl
• [Ag][Cl] = 1.6 * 10^-10
• [AgCl] = SQRT 1.6 * 10^-10
• [AgCl]= 1.26^-5 M

• AgCl <==> Ag + Cl
• x <==> x + x
  
• NaCl <==> Na + Cl
• 0.01 <==> 0.01 + 0.01
  
• Ksp = (x)(2x + 0.01) = 1.6 * 10^-10
• cross out the 2x above b/c it is NEGLIGIBLE. It is MUCH smaller relative to its 0.01 counterpart.
• 0.01x = 1.6 * 10^-10
• x = 1.6 * 10^-8
• [AgCl] = 1.6 * 10^-8

Chapter 20: Acids and Bases

Characteristics of Acids and Bases

• Some compounds are amphiprotic, that is, they can act either as a proton donator or a proton acceptor. A monoprotic compound can give or take one proton. A diprotic compound can give or take up to two protons. A triprotic compound can give or take up to three protons. Wow...

The Nature of Acids and Bases

• This equation represents a competition for the proton between the two bases: H2O and A^-
• Conjugate base: everything that remains of the acid molecule after a proton is lost
• Conjugate acid: formed when the proton is transferred to the base

What the hell is the diff between concentration and strength?!

• a strong acid dissociates completely if not almost completely (producing very weak conjugate bases). A weak acid on the other hand does not get the pleasure in dissociating as much (but they yield a hell of a strong conjugate base).
• And as for bases, a strong base (alkaline bases) hydrolyzes easily into OH^- when placed in water. Weak bases on the other hand does not dissociate quite as easily.
• Many proton acceptors do not contain OH^-, but they increase the concentration of OH^- because of their reaction with water.
• bases compete with OH^- (which is a very strong base) for the H^+ ion

The Equilibrium Constant for Water

• For those who didn't know, water actually ionizes into H3O^+ and OH^-. H2O + H2O <==> H3O^+ + OH^-
• (H2O molecules are continually dissociating while H3O^+ and OH^- are recombining to form water)
• water autoionizes
• the equilibrium expression is:

• at 25 degrees celsius, the Kw is always 1.0 * 10^-14
• in a neutral solution, [H3O^+] = [OH^-] = 1.0 * 10^-7M
• if [H3O^+] > [OH^-], then the solution is acidic
• if [H3O^+] < [OH^-], then the solution is basic

The (horrifying) effects of temperature on Kw

• as we all know, burning hydrogen (to form water) releases quite an amount of heat, or energy
• to reverse that reaction, energy will be required
• the dissociation of water is therefore endothermic Heat + 2H2O <==> H3O^+ + OH^-
• increasing the temperature favors the products and increases the dissociation of water, forming more ions
• Kw would also increase, but water remains neutral because [H3O^+] = [OH^-]
• decreasing the temperature will decrease the value of Kw, but the water will still remain neutral because [H3O^+] = [OH^-]

Relative Acidity and the pH Scale

• the relative acidity of a solution is best represented by its concentration of H3O^+ ions
• it is inconventient to work with [H3O^+] because of its wide range of values (10M --> 1.0 * 10^-15)
• so we use the exponent of the concentration (or the logarithm)
• since most concentrations are very small, they usually have negative exponents...but if the negative log were stated, the values would always be positive for small conenctrations
  pH = -log[H3O^+]
  
• IMPORTANT!! the number of decimal places in the log is equal the the number of sig figs in the original number
• for example... [H3O^+] = 4.48 * 10^-5 - it has 3 sig figs, so the log should have 3 decimal places...
• pH = 4.349
• pOH is very similar to pH but instead you take...you guessed it! [OH^-] into account instead of [H3O^+]
• you can also get pKw, which is 14 at 25 degrees celsius
• at 25 degrees celsius, the sums of pH and pOH will add up to 14

Relative Acid Strength

• when two acids are combined, the stronger one will act as an acid while the weaker one acts as a base
• to determine which side (products or reactants) are favored, compare the strengths of the two acids and bases from the Relative Strengths of Bronsted-Lowry Acids and Bases Table
• while strong acids dissociate to a greater extent, weak acids do not
• for any weak acid, there is an equilibrium
• HA + H2O <==> H3O^+ + A^-
  

  Ka =	[H3O^+][A^-] ------------- [HA]

  
  
• for weak acids, [HA] is usually the initial concentration... (as long as [H3O^+] is less than 5% [HA] - this is called the 5% rule)
• as acid strength decreases, its dissociation drops and fewer ions ar formed
• Ka is an indicator of acid strength

Molecular Bases

• B^- + H2O <==> HB + OH^-

• Kb =	[HB][OH^-] ---------- [B^-]

• Ka * Kb = Kw
• Kb = Kw / Ka
• Ka = Kw / Kb (depending on what you're given, and what you're looking for)

Indicators

• indicators are weak acids (or bases)) whose conjugate acids and conjugate base have different colours
• a generalized equation:
  HIn + H2O <==> H3O^+ + In^-
  
• the indicator's equilibrium will shift depending on the environment in which it is placed
• HIn is in greater abundance in acids
• In^- is in greater abundance in bases

Combining Acids

• suppose that two different acid equilibra were combined and one of the acid equilibria was coloured (ie indicator)
  Acid #1: HA <==> H^+ + A^- (colourless)
  Acid #2: HIn <==> H^+ + In^- (coloured)
• the colour of acid #2 depends on teh relative amounts of HIn and In^- in the equilibrium
• if HA is a stronger acid than HIn,
  • HA + In^- <==> HIn + A^-
  • then the equilibrium would favor the product side because HA would dissociate to a greater extent than HIn
  • since HIn is coloured differently from In^-, its presence is easily recognized
• if HIn is a stronger acid, then the equilibrium would favor the reactants and the colour of In^- would be in greater abundance

Acids and Bases Part 2

Indicators (cont'd)

• Indicators do not usually change colour when pH is 7 (duh)
• an indicator is a weak acid/base...therefore, it has its own Ka and Kb
• the indicator is at its transition point when [In-] = [HIn]
• different indicators have different Ka's...therefore they have different transition points...or they have different pH's at which they change colour

Hydrolysis of Salts

• salts form when a positive metal ion from the base combines with a negative non-metal from the acid
• if neither ion reacts with water, the solution will remain neutral...but if one of these ions reacts with water, the solution will become either acidic or basic

  Positive Ion Hydrolysis 4 ions that hydrolyse: Al3+, Fe3+, Cr3+, and NH4+ positive ions hydrolize to form acidic solutions

  Negative Hydrolysis There are 6 negative ions that do not hydrolyze (to form basic solutions): ClO4-, I-, Br-, Cl-, NO3-, and HSO4- anions hydrolyze to form basic solutions

  • if the cation hydrolyzes, the salt is acidic, and vice versa for anions (which yield basic solutions)
  • if both ions hydrolyze, the comparison is made of their relative strengths; the cation's Ka value and the anion's Kb value... the bigger value will dominate in the solution...that is, if the Ka is bigger than Kb, the solution will be acidic, and vice versa if Kb is bigger than Ka

  Anhydrides

  • they are compounds without water
  • they attach themselves to water to produce water containing compounds (hydrides)
  • the anhydrides that contain oxygen are a special group that can react with water to produce either acidic or basic solutions
  • alkaline anhydrides (and stuff on the left of the periodic table) produce basic solutions
  • halogen anhydrides (and stuff on the right of the periodic table) produce acidic solutions

  Buffers

  • a solution of a weak acid and its conjugate base [comes from a salt] (or vice versa)
  • buffers maintain a certain pH when small amounts of strong acids or bases are added - if an acid is added, HA incr, and A- decr, all the while keeping the [H+] relatively constant. If a base is added, HA decr, A- incr, all the while, [H+] i still relatively constant
  • a typical buffer can be made by mixing equal molar quantities of a weak acid and the salt of its conjugate base
  • so, you wanna create a buffer at a certain pH? refer to some Ka values. if the concentrations of [HA] and [A-] are close to equal, then the [H] will be approximately equal to the Ka.
  • buffers are important in maintaining homeostasis (eg pH in blood) - even the slightest fluctuations can cause death
  
  
  

  Titration Curve

  • illustrates the change in pH as standard acid is added to a base, or vice versa
  • there are 4 significant points:
    1. before any standard is added
    2. just before the equivalence point
    3. right smack at the equivalence point
    4. and after the equivalence point
  • titrating a strong base w/ a weak acid is kinda complicated, so here's the steps in calculating the 4 points:
    1. before any of the base is added (given)
    2. before the equivalence point is reached
       

       (original # of mol of H) - (current # off mol of OH- added) ------------------------------------------- total volume

    3. at the equivalence point: [OH-] = square root of (Kb * [A-])
    4. after the equivalence point, take the difference between the original molar amount of H+ and the current molar amount of OH-, then divide the difference by the total volume to obtain [OH-]
  
  
  

  Acid Rain

• lakes and rivers are buffered, but to a certain extent
• once this extent is reached, and the water can't be bufferend any more, then the water has become too acidic for its own good
• so how the hell does it get acidic? industrial fumes like sulphur dioxide and nitrous oxide (anhydrides) become hydrated in the atmosphere to produce sulphurous and nitrous acids...these get into the water, and BOOM! acid rain.

Good luck on the upcoming test (march 13, 97)