The Chemical Industry

Note

a Rate of reaction:

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Rate of reaction measures the rate of reactant→product conversion. It can easily be calculated from the gradient of a concentration-time graph
Rate equations can be used to calculate the rate of reaction from concentrations and a known constant for a given temperature
rate = k[A][B]... where:
- k is the rate constant, which is different for each reaction and temperature
- [X] represents the concentration of a reactant/catalyst
- ^etc represents the order of that reactant/catalyst (see below)
As temperature increases, the rate constant k increases

Order:

Order is an integer unique to each reactant, which can be determined by experimental data only
Overall order is the sum of all individual orders in a reaction
Orders vary for each reaction and are calculated by measuring rate of reaction increase when the concentration of that reactant doubles
When finding the order of a reactant, ensure that all other reactants are in excess
For example, if doubling concentration for reactant A doubles rate of reaction, it is linear - the order is 1
If doubling concentration for a reactant quadruples rate of reaction, it is a quadratic relationship with order 2
This can be written with the proportionality symbol ∝
- With order 1 for reactant A, rate [A]
- If it was order 2, it would be rate [A]
Alternatively, the ∝ symbol could be removed and the equations written like rate = k[A] or rate = k[A]

b Reaction half-lives:

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Reaction half-life is the time it takes for half of a reactant to be used up
A half-life can be easily calculated from a concentration-time graph by finding the time taken for the concentration to halve
By measuring 2-3 half-lives, the order can be determined:
- If half-life decreases with time, it's zero order
- If half-life remains constant, it's first order
- If half-life increases with time, it's second order
The symbol for a half-life is

Calculating rate constant from half-life (first order only):

The half-life remains constant for a first-order reaction, so the rate constant can be calculated with (units are )
The derivation of this equation is not required

c Initial rate:

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The initial rate of reaction is the rate at the start of the reaction. It is calculated by measuring the time taken for a small amount of a reactant to be used or a small amount of product to form, and using the equation initial rate = amount of reactant used/product formed time
It can also be estimated from the gradient at the start of a concentration-time graph

Clock reactions:

A clock reaction has a clear end point (e.g. a colour change) once a certain amount of product has formed. Therefore, clock reactions are useful for measuring initial rate of reaction
TO DO: Cards from here
The iodine clock reaction is H₂O₂ (hydrogen peroxide) + 2I^- + 2H⁺ → 2H₂O + I₂. Sodium thiosulfate instantly removes iodine, so a small amount is added. Once this is all used up, there will be iodine molecules in the solution, which can be tested for with starch (which turns blue/black in the presence of iodine)

Concentration-time graphs:

On a concentration-time graph, if the reaction is zero-order (with respect to the reactant being investigated), it is a negative-gradient straight line, because the reactant concentration decreases at a constant rate
If first order, the graph has an exponential decay
TO DO: or is it quadratic?
curve, because rate increases with reactant concentration, so rate decreases over time as the reactant is used up
If second order, the graph has a similar shape, but steeper gradient, as the rate decrease per second is increasing over the course of the reaction

Rate-concentration graphs:

For a zero-order reactant, this is a horizontal line - the rate remains constant for any concentration
For a first-order reactant, the rate-concentration graph has a straight line with positive gradient - rate is increasing with concentration
Second order reactants will have a quadratic graph (or a straight line when rate is plotted against concentration)

d The Arrhenius equation:

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The Arrhenius equation is k = Ae (given in data sheet), where:
- k is the rate constant
- is the activation enthalpy (J mol-1)
- T is the temperature (K)
- R is the gas constant (8.314, given in data sheet)
- A is the pre-exponential factor
Taking logs of both sides: ln k = + ln A (also given in data sheet)
This is in y = mx + c form, so ln k can be plotted against . The resulting graph will have a gradient of and y-intercept of ln A
Remember that the calculated activation enthalpy will be in J mol-1 and temperature will be in kelvins, so remember to convert if required

e Rate-determining steps:

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In a multi-step reaction, the rate-determining step is the step with the lowest rate. This step determines the order of the reaction
If a reactant isn't in the rate equation, it can't be in the rate-determining step
Catalysts can be in rate-determining steps (they can appear in rate equations sometimes)
If given the equation of the rate-determining step, the order of a reactant can be predicted from the number of times it appears
However, this method doesn't always work - for example with the reaction 2A + B → C it would be assumed that the reaction is second order with respect to A, but it could also be first-order if only one A is present in the rate determining step and the second A is involved in another, much quicker, step

f Factors affecting the equilibrium constant:

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Increasing the pressure of a dynamic equilibrium will move the position of the equilibrium towards the side with the fewest moles of gas. It will not affect
Changing the concentration of a reactant or product in an equilibrium will move the position, but also not affect the value of
Adding a catalyst will not affect the position of equilibrium or value of , but they do reduce the time taken for the equilibrium to form
The three values above don't affect because the proportions of reactants and products change in a way to keep it the same
However, changing the temperature of an equilibrium will affect both the position of equilibrium and value of

h Equilibrium calculations:

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The fraction used to find can be rearranged to find unknowns

Finding an equilibrium constant experimentally:

If one of the substances is coloured, colorimetry can be used, with the calibration curve being used to calculate the concentration of the coloured substance
If one of the sides of the reaction contains an acid or alkali, a pH probe can be used. If the equilibrium took a long time to form, use a titration
But do not use a titration for equilibria with a high speed of formation because adding an acid/alkali would significantly change the position of equilibrium, giving a different, inaccurate value for concentration

j Nitrogen chemistry:

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Nitrogen atoms bond diatomically to form N₂ molecules with a triple bond, therefore giving it a very high bond enthalpy
Ammonia has the structure NH₃ with a lone pair
An ammonium ion is NH₄⁺

Nitrogen oxides:

NO is nitrogen monoxide / nitrogen(II) oxide, a colourless gas
N₂O is dinitrogen monoxide / nitrogen(I) oxide, a colourless gas with a sweet smell (laughing gas)
NO₂ is nitrogen dioxide / nitrogen(IV) oxide, a brown gas with sharp odour

Testing for ammonium compounds:

Add NaOH and heat. Ammonium ions will react with the OH^- ions, releasing ammonia gas (which will turn damp red litmus paper blue)

Testing for nitrate(V) ions:

Nitrate(V) ions have the formula NO₃^-
To test for them, heat the solution with NaOH and add a piece of Devarda's alloy (mainly aluminium and copper)
The aluminium reduces the nitrate(V) ions, creating ammonia gas, water and Al(OH)₄^-

The nitrogen cycle:

N₂ and H₂ react to form ammonia
(ammonia) + H⁺ creates ammonium ions
(ammonium ions) + O₂ form nitrate(III) ions (NO₂^-)
(nitrate(III) ions) + H₂O creates nitrate(V) ions (NO₃^-)
(nitrate(V) ions) + H⁺ + e^- forms nitrogen gas again
Oxides are formed with N₂ and a multiple of O₂