module 5 definitions

• Arrhenius Equation


  k = Ae^(-Ea/RT) where k is the rate constant, A is the pre-exponential factor, E_a is the activation energy, R is the gas constant and T is the temperature
• Clock Reaction


  a reaction in which the time taken for an abrupt visible change to occur is measured. The abrupt change typically indicates the formation of the product
• Colorimetry


  a technique used to measure the amount of light absorbed by a solution, used to determine the rate of a reaction. Typically, a calibration curve is plotted and used to relate the absorbance recorded by the colorimeter to the concentration of the solution
• Concentration-time Graph


  a graph in which concentration is plotted on the y axis and time on the x axis. The gradient of the line is equal to the rate of reaction
• Continuous Monitoring


  during a rate experiment, continuous measurements are taken as the reaction progresses. The results can then be plotted on a concentration-time graph
• First Order Reactant


  doubling the concentration of a first order reactant will double the rate (if all other conditions remain the same)
• Gradient


  change in y ÷ change in x
• Half-life (t_1/2)


  the time taken for the concentration of a reactant to half
• Initial Rate


  the rate of a reaction at t=0
• Order


  a number that relates the rate of a reaction to the concentrations of each reactant
• Overall Order


  the sum of the orders with respect to each reactant
• Rate Constant (k)


  a constant value that relates the rate of a reaction at a given temperature to the concentrations of the reactants. For a first order reaction, this can be determined using the relationship k = ln 2/t_1/2
• Rate-concentration Graph


  a graph that has concentration plotted on the x axis and rate on the y axis
• Rate-determining Step


  the slowest step of a reaction
• Rate equation


  relates rate to the concentrations of the reactants multiplied by the rate constant. Each concentration is raised to the power of the order with respect to that reactant
• Rate of Reaction


  a measure of how quickly a reactant is used up / a product is formed
• Reaction Mechanism


  a step-by-step sequence of the individual reactions that make up the overall reaction
• Second Order Reactant


  doubling the concentration of a first order reactant will quadruple the rate (if all other conditions remain the same)
• Zero Order Reactant


  doubling the concentration of a first order reactant will have no impact on the rate (if all other conditions remain the same)
• Catalyst


  a substance which speeds up the rate of a reaction without being used up
• Concentration


  the amount of a substance that is dissolved per unit volume of solution
• Endothermic


  a reaction that takes in energy from the surroundings. The energy of the products is higher than the reactants
• Equilibrium


  the forward and reverse reactions of a process occur at exactly the same rate meaning there is no overall change
• Equilibrium Constant (K)


  a value that relates the amount of products and reactants at equilibrium in a reversible reaction at a specific temperature. K is unaffected by pressure and presence of a catalyst but is affected by temperature
• Exothermic


  a reaction that gives out energy into the surroundings. The energy of the reactants is higher than the products
• Heterogeneous Equilibrium


  an equilibrium reaction that involves substances in different states (solid, liquid gaseous or aqueous)
• Homogeneous Equilibrium


  an equilibrium reaction that involves substances all in the same state (solid, liquid gaseous or aqueous)
• K_c


  the equilibrium constant that is equal to the concentration of products raised to their stoichiometric coefficients divided by the concentration of reactants to the power of their stoichiometric coefficients. Liquids and solids are not included in heterogeneous K_C expressions as their concentrations effectively remain constant
• K_p


  the equilibrium constant that is equal to the partial pressures of the products raised to the power of their stoichiometric coefficients divided by the partial pressures of the reactants raised to the power of their stoichiometric coefficients
• Mole Fraction


  X_A = n_A/n where X_A is the mole fraction of A, n_A is the number of moles of A and n is the total number of moles
• Partial Pressure


  the pressure that would be exerted by one gas in a mixture if it occupied the container alone. P_A = PX_A where P_A is the partial pressure of A, P is the total pressure and X_A is the mole fraction of A
• Acid Dissociation Constant (K_a)


  the extent of acid dissociation. pK_a = -log(K_a) and K_a = 10^(-pKa)
• Bronsted-Lowry Acid


  a proton donor
• Bronsted-Lowry Base


  a proton acceptor
• Buffer Solution


  a system that minimises pH change on addition of small amounts of an acid or base. A buffer solution can be formed from a weak acid and a salt of the weak acid or from excess weak acid and a strong alkali
• Carbonic Acid-Hydrogencarbonate Buffer


  the buffer system present in blood plasma, used to maintain blood pH between 7.35 and 7.45
• Conjugate Acid-Base Pair


  a pair of compounds that transform into each other by the transfer of a proton. Conjugate acid-base pairs are important in the formation of buffers to control pH
• Dibasic Acid


  an acid that can donate 2 hydrogen atoms per molecule of the acid
• End Point


  the point during a titration when the indicator changes colour. A suitable indicator should change colour near the equivalence point (it should have a pH range within the vertical section of the titration curve)
• Equivalence Point


  the point during a titration when the amount of acid is exactly equal to the amount of base (full neutralisation occurs). When titrating an acid with a base, [H+] is equal to [OH-] at this point