Chemical Equilibrium

Chemical Equilibrium

1. Static Equilibrium

2. Dynamic Equilibrium

3. Irreversible Reactions

4. Reversible Reactions

5. Homogeneous mixtures

6. Heterogeneous mixtures

7. Position of Equilibrium

   • Reaction Quotient (Qc)
   • Equilibrium Constant (Kc)
     • Excluded from Kc & Kp expressions
       • Pure solids
       • Pure liquids
       • Solvents
     • Equilibrium Constant (Kp)

8. LCP (Le Chatelier's Princple)

   • 4 disturbances to system at equilibrium
     • Concentration changes
     • Partial Pressure changes (in Gases)
       • Effect of adding an inert gas
     • Temperature changes
     • Effect of Catalyst

9. Industrial application of LCP

   • The Haber Process

10. Gibbs Free Energy and the POE

11. Acid-Base Equilibria

    • Arrhenius theory
      • Definition: Acid-base reactions are characterised by acid, which dissociate in aqueous solutions to form hydrogen ions(H+) and bases, which form hydroxide (OH-) ions
      • Base
        • - a substance that releases OH- ions when dissolved in water
      • Acid
        • - a substance that releases H+ ions when dissolved in water
    • Bronsted-Lowry theory
      • Acid
        • - a H+ donor
        • Conjugate Base
          • - the particle left when the acid has donated its H+ - the stronger the acid, the weaker is its conjugate base, and vice versa
      • Base
        • - a H+ acceptor
        • Conjugate Acid
          • - the particle formed when the base has accepted the H+ -the stronger the base, the weaker is its conjugate acid, and vice versa
      • Conjugate acid-base pair
        • - differs by one H+
      • Amphiprotic
        • - a substance which can act as either an acid or a base 
    • Lewis theory
      • Acid
        • - an electron pair acceptor - substance capable of accepting a lone pair of electrons from a base to form a dative covalent bond
      • Base
        • - an electron pair donor
    • PH
      • pH= -lg[H3O+] pH= -lg[H+] so [H+]= 10^-pH - can be determined accurately using the pH meter - dependent on concentration and degree of dissociation - can be measured by using its dissociation constant, Ka or Kb which is constant at all dilution and is influenced only by temperature
      • Acid dissociation constant, Ka
        • - a weak acid is partially dissociated in aqueous solution - equilibrium constant for the dissociation reaction - Ka=[H3O+/H+][A-]/[HA]- units: mol dm^-3- constant at constant temperature- pKa=-lgKa - good measure of the strength of an acid- larger Ka, stronger acid-smaller Ka, weaker acid
    • pOH
      • pOH= -lg[OH-] so [OH-]= 10^-pOH
      • Base dissociation constant, Kb
        • - weak base - Kb= [BH+][OH-]/[B] - units: mol dm^-3 - constant at a constant temperature - pKb= -lgKb - good measure of the strength of a base - larger Kb, stronger base - smaller Kb, weaker base
    • Self-ionisation of water
      • Ionic product of water, Kw
        •  - pure water ionises to a very slight extend - Kw= [H3O+/H+][OH-} - units: mol^2dm^-6 - Kw= 1.0x10^-14 mol^2dm^-6 @25degreescelsius - pKw=-lgKw=14.00 (@25degreescelsius) - pH+pOH= pKw=14.00 (@25degreescelsius)
        • r/s with water
          • - Kw is a temperature-dependent constant - Kw increases with temperature - dissociation of water is an endothermic process - by LCP, as temperature increases, POE shifts to the right to favour the endothermic reaction by absorbing the excess heat
    • r/s between Ka, Kb & Kw
      • Ka x Kb = Kw (for a conjugate acid-base pair)= 1.0x10^-14 (@25degreescelsius) pKa+pKb=pKw (for a conjugate acid-base pair)=14.00 (@25degreescelsius)
    • Strong acids and bases
      • - strong monobasic acids are 100% dissociated in aqueous solutions - strong bases e.g alkali metal hydroxides - When [H+] >> 1.0x10^-7moldm^-3, self-ionisation of water is suppressed - When [H+]~<1.0x10^-7moldm^-3, self-ionisation of water cannot be ignored
      • monobasic acids
        • - single dissociable proton
      • dibasic acids
        • - 2 dissociable protons
    • Weak acids and bases
      • - with weak acids, change is negligible, e.g. c-x, x<<c, ~c
    • Salt Hydrolysis
      • - salt reacting with water to give an acidic/alkaline/neutral water
      • Salts
        • - anions: strong conjugate base of the weak acid - cations: strong conjugate acid of the weak base OR high charge density
    • Buffer solutions
      • - maintains a fairly constant pH when a small amount of acid/base is added to it - a solution of a weak acid and its salt (contains strong conjugate base) - a solution of a weak base and its salt (contains strong conjugate acid)
      • Acidic
        • - solution made up of weak acid and its salt(contains strong conjugate base) - qualitative discussion (refer p.19)-pH remains almost constant after adding acid/base to the buffer-quantitative discussion (refer p.20)- pH of buffer depends on the Ka value of the acid & the ratio of [A-]/[HA]- ratio of [A-]/[HA] calculated by making 2 approximations: [A-]~[M=A-]initial and [HA]~[HA]initial (because Ka of HA very small, thus all A- assumed to come from fully dissociated salt M+A-, weak dissociation of acid neglected - pH=pKa+lg[salt]/[acid] can be used to calculate pH of buffer solution, how much acid and salt needed to make buffer solution of certain pH< effect on the pH of the buffer solution when a small amount of acid/base added to it
      • Basic
        • - solution made up of a weak base and its salt(contains strong conjugate acid) - qualitative discussion (refer p.23) - pH of the solution remains almost constant when acid/base added to the buffer - quantitative discussion (refer p.23) - pOH of a weak base buffer is dependent on the Kb value of the base & the ratio of [NH4+]/[NH3] - value of the ratio can be calculated by making 2 approximations: [NH4+]=[NH4Cl] & [NH3]~[NH3]initial as Kb of NH3 is small, thus all NH4+ assumed to have come from fully dissociated salt, NH4Cl, NH4+ from weakly dissociated base NH3 is negligible.
      • Buffer capacity
        • - a measure of the amount of acid/base that the buffer can absorb without a significant change in pH - for equal volumes of solution, the more concentrated the solution, the greater the buffer capacity - for solutions having equal concentration, the greater the volume, the greater the buffer capacity - when buffer capacity is before the equivalence point, volume required is halved - when buffer capacity is after the equivalence point, volume required is doubled
        • Acidic buffer
          • - buffer region effective when 0.1<[salt]/[acid]<10 - pH range=pKa+-1 - maximum buffer capacity occurs when [salt]=[acid]initial, so that pH=pKa
        • Basic buffer
          • - buffer region effective when 0.1<[salt]/[base]<10 - pOH range=pKb+-1 - maximum buffer capacity occurs when [salt]=[base]initial, so that pOH=pKb
        • Buffer preparation
          • - usually prepared from 0.10moldm^-3 to 1.0moldm^-3 solutions of reagents 1. choose a weak acid whose pKa is near the desired pH 2. exact value of [H+] then achieved by adjusting acid/conjugate base ratio
        • Uses
          • Blood
            • - pH of 7.40 - death as a result of pH below 7.0/above 7.8 - tightly regulated at 7.40+-0.05 by buffer whose principal acid-base pair is H2CO3 & HCO3- - buffering action provided by H2CO3/HCO3-, H2PO4-/HPO4^2-, proteins
            • Bicarbonate Buffers
              • - carbon dioxide from respiration dissolves in water to form H2CO3(aq) then HCO3-(aq) and H3O+(aq) - when blood is too alkaline, OH- reacts with H2CO3 to form HCO3- - when blood is too acidic, H+ combines with HCO3- to form H2CO3
          • Others
            • e.g. intravenous injections
    • Acid-base Titrations
      • pH changes
        • - @end point, pH changes drastically (except for weak acid/weak base titrations)
      • Curves
        • - pH equivalence point - buffer region - point of maximum buffer capacity
      • Indicators
        • - weak acids or weak bases - formula: Hln  - Hln(aq) <=> H+(aq)+ln-(aq) acid form(colour A) <=> conjugate base form(colour B) - in acidic solution, [H+] is high, hence according to LCP, POE shifts left, hence colour A - in alkaline solutions, OH- ions neutralise H+ ions, hence according to LCP, POE shifts right, hence colour B
        • pH of colour change at end point
          • - [Hln]=[in-] - pH=pKln - methyl orange (red-yellow) - thymol blue (yellow-blue) - phenolphthalein (colourless-pink) - colour observed indicates that pH is above/below a certain value, thus doesn't tell us directly if solution is acidic/alkaline
        • Choice
          • 2 main considerations  1. pKln of indicator should be as close to the equivalence point of the titration so that there is a distinct colour change 2. indicator working range should lie within the range of rapid pH change
          • types
            • 1. strong acid-strong base 4-10 (almost any indicator) 2. strong acid-weak base 4-7 (methyl orange) 3. weak acid-strong base 7-10 (phenolphthalein, thymol blue) 4. weak acid-weak base no sharp pH change (no suitable indicator)
      • Polybasic acids
        • - acids that contain more than o ne dissociable proton - dissociate in a stepwise manner, each step characterised by its own dissociation constant, Ka1, Ka2, and so on - values of dissociation constants decrease along each step - due to electrostatic forces, it is more difficult to remove a positively charged proton from a negative ion than from an uncharged molecule
        • Polybasic acid-strong base titrations
          • - curves show more than one end point (because each acidic proton reacts in turn, with the monoacidic strong base)
          • Titration of dilute phosphoric acid, H3PO4(aq), with NaOH(aq)
            • - tribasic acid - 3 acidic protons - 3 end-points - 3 points of inflexion
          • Titration of dilute sulfurous acid, H2SO3(aq), with NaOH(aq)
            • - dibasic acid - 2 acidic protons - 2 stages

12. Solubility Equilibria

    • Saturated solutions
      • - a solution is saturated when no more solute will dissolve in it
      • Sparingly soluble ionic salt
        • - when excess shaken with water, saturated solution formed in which the dissolved aqueous ions are in dynamic equilibrium with the remaining undissolved solid - rate of dissolution=precipitation processes - concentrations remain constant
    • Solubility product, Ksp
      • - p 3 e.g. - no fixed unit - temperature-dependent, constant at constant temperature and only changes with temperature Ksp for a sparingly soluble ionic salt is the product of the concentrations of ions in a saturated solution at a given temperature raised to the powers indicated by the stoichiometric coefficients in the balanced equation for the equilibrium. AxBy(s)<=>xA^y+(aq) + yB^x-(aq) Ksp= [A^y+]^x[B^x-]^y units: (moldm^-3)^x+y
      • Solubility & Solubility product, Ksp
        • Calculation of Ksp from solubility
        • Calculation of solubility from Ksp
      • Applications (5)
        • Ionic product & Ksp
        • Common ion effect
        • Complex ion formation
        • pH & solubility
        • Selective precipitation
    • Solubility criterion
      • Sparingly soluble
        • <0.01moldm^-3
      • Moderately soluble
        • 0.01-0.1moldm^-3
      • Soluble
        • >0.1moldm^-3

Imágenes y archivos del mapa

• Acidbaseshit (application/pdf)