rates equilibrium and pH part one

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  • equilibrium pH and rate determining step
    • the rate determining step
      • this is the slowest step in a reaction this sets the rate of the reaction
      • to start off it is best to write a balanced stoichiometric equation this will tell you all the reactants you will be working with
      • this equation does not tell you about the reaction mechanism for the rate determining step to work this out we need to carry out some rate experiments
        • the aim of these experiments is to work out the orders of the reactants as this information is needed to determine the rate determining step
      • an example is  NO2+CO===> NO+ CO2
        • from experimentation we determine that NO2 is second order and CO is zero order
    • the equilibrium constant
      • the equilibrium constant is called Kc
        • here is an example of how to do a basic question on this
          • N2042NO2
            • the kc for this would be Kc=[NO2]^2/[N2O4]
              • The units are worked out like so=(mols dm^3)/mols dm^3 so your units for this can be worked out by canceling giving the unit of mols dm^3
      • this constant can only be calculated when a reaction is in dynamic eqaulibrium
        • this can only occur in a closed system when the concentrations of the products and reactants stay the same
      • the equilibrium constant is Kc= products/reactants
        • for example
        • the amount of moles in the equation for each reactant and product is the power of the concentration of each reactant and product
      • Kc can have no units
      • position of equilibrium and the equilibrium constant
        • when the value for equalibrum = 1 this means that there is equilibrium halfway between the reactants and products
        • if the kc is larger than one the reaction favours the products this means that the position of equilibrium is to the right
        • if the kc is less than one this means that the position of equilibrium is favouring the reactants shifting equilibrium to the left
        • temperature if increased will shift equilibrium in the endothermic direction
          • if temperature is decreased it is shifted in the exothermic direction
        • pressure has no effect on kc as the system will shift to maintain equilibrium in the system
        • catalysts also don't effect equilibrium as they speed up the forward and reverse reaction equally
    • the equilibrium constant Kc vs the rate constant k
      • large Kc= equilibrium lies to the right so more products are produced
        • large k= fast rate of reaction
      • pressure and concentration have no effect on Kc
        • pressure and concentration cause the value of k to change
      • high temperature shifts the equilibrium in the endothermic direction
        • low temperature shifts equilibrium in the exothermic direction
          • temperature increases the rate of reaction increasing the value for k
        • this changes Kc
          • low temperature shifts equilibrium in the exothermic direction
            • temperature increases the rate of reaction increasing the value for k
      • Kc can be worked form a balanced equation k can only be used from experimental data
    • acids
      • three definitions for acids and bases
        • arrhenius
          • acid = H^+
          • base = OH^-
        • bronsted lowry (important need to know)
          • Acid = a proton donor
          • base= proton acceptor
        • lewis
          • acid = electron acceptor
          • base=electron donor
      • all acids have hydrogen in them
      • three types of basic acids
        • monobasic acids these have only one available hydrogen
        • dibasic acid these have two available hydrogens
        • tribasic acids have three available hydrogens
      • conjugate acid base pairings
        • a typical acid base reaction
          • HCL+NaOH NaCl + H2O
            • this reaction is reversible in the forward reaction the HCl is an acid ans the NaOH is a base
              • in the reverse reaction the H2O can give away a H^+ ion to the NaCl which will accept it as such the water molecule becomes a acid and the NaCl becomes a base
            • the NaCl is the conjugate base for HCl and the H2O is the conjugate acid for the NaOH
        • another example
          • HNO2+H2O====> H3O^+) +NO2^-
            • in the forward reaction the HNO2 is the acid and the H2O is the base
              • the conjugate base for HNO2 is NO2^-
              • the conjugate acid for H2O is H3O ^+
        • the definition for a conjugate acid base pairing is a pair of two species that transform into each other by the loss or gain of a proton
      • acid reactions
        • reactions with bases and alkalis
          • acid + base/alkali====> salt + water
        • acid and carbonate

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