6. chemical industry

rate of reaction

change in amount of reactants or products per unit time

continuous monitoring

a way of measuring rate of reaction over the complete course of the reaction

methods of measuring rate of reaction

chosen based on what changes in the reaction:

• pH measurement

• gas volume

• loss of mass

• colour change

• titration

pH measurement

used if a product as an acid or base, can be measured with pH meter or a pH probe with a data logger, can be converted to concentration by [H+] = 10^-pH, gives mol dm-3 time -1

gas volume

if produced it can be collected in a gas syringe and volume recorded at regular intervals, gives volume time-1

loss of mass

measure mass remaining of a system if a gas is given off, measured with a balance, gives mass time-1

colour change

if a reactant or product is coloured, the colour change of the reaction can be tracked with a colorimeter measuring absorbance (more concentrated means more absorbance)calculate concentration with calibration curve

titration

small sample are taken at intervals and titrated, the samples are slowed down by diluting with deionised water, cooling or stoped with a chemical

concentration time graph

a graph that can be used to work out rate of reaction by determining gradient at any point with a tangent

initial rate

the rate of reaction at the start

= amount of reactant used or product formed ÷ time

initial rate method

the time for a set amount of product to form (eg gas) is recorded and used to calculate initial rate,

repeated serval times and the initial concentration of a reactant is changed

assumptions of initial rate method

• concentration of other reactants isn’t changing much- typically done by having them in excess

• temperature stays constant

• reactant hasn’t proceeded to far when measurement is taken

clock reaction

a initial rates reaction where you measure time taken for a set amount of product to form and concentration of a reactant is changed, but has an early observable end point like a colour change, quicker colour change means faster rate of reaction

iodine clock reaction

• small amount of sodium thiosulfate solution and starch are added to an excess of hydrogen peroxide and iodine ions in acid solution

• when the sodium thiosulfate is added it reacts instantaneously with any iodine that forms

• initially all iodine made is used up immediately, but once sodium thiosulfate is used any more iodine stays in solution storing starch blue black (starch turns blue back in presence of iodine so acts as an indicator), this is the end point

• changing the concentration of iodine ions or hydrogen peroxide will give different times for the colour change

rate equation

tells you how rate is affected by concentration of reactants

rate = k[A]m[B]n

rate constant

k

bigger it is the water the reaction

remains the same for a reaction at a particulate temperature

unit vary

= rate ÷ [A]m[B]n

orders

tell you how a reactants concentration affects the rate

determined only by reactions

zero order

order if concentration of the reactant changes and rate does not

first order

order of concentration of the reactant changes and the rate changes proportionally

eg if conc doubles then rate doubles

second order

oder if concentration of reactant changes and rate changes proportionally to [A]^2

eg if conc doubles then rate will be 2^2 so 4 times faster

overall order

sum of individual orders of reactants, gives order of reaction

determine order

by:

• constructing a rate concertation graph

• comparing initial rate for different concentration

• constructing a conception time graph and comparing half lives

rate concentration graph

made by constructing a concentration time graph and calculating rate at various concentrations and plotting it on a new graph

units of rate constant

calculated by substituting units for rate and concentration in to rate equation

done by cancelling out units

eg units of k = mol dm-3 s-1 ÷ (mol dm-3)2 (mol dm-3) = mol-2 dm6 s-1

half life

the time taken for a reactant to halve in quantity (half to be used up)

easily calculated on concentration time graph

half life increases

half life of a zero order reaction where the rate doesn’t change

half life is the same

half life of a first order reaction where rate is proportional

half life increases

half life of a second order reaction where rate is proportion to x2

word out rate constant from half life

using equation k = ln2 ÷ half life time

units for each are: no units ÷ s = s-1

works for first order reaction with equal half life

temperature

effects rate of reaction as:

• increases collisions

• increases reactants with activation energy

• more reactants have the right orientation

• more kinetic energy

higher temperature

higher rate constant

higher rate of reaction

Arrhenius equation

links rate constant and activation energy and temperature

k = Ae^-Ea/RT given on data sheet

Ea = activation energy J mol-1

T = temperature K

R = gas constant 8.31 J K-1 mol-1

A = pre-exponetial factor (a constant)

e = exponential relationship (button on calculator)

activation energy

as it gets bigger, k gets smaller

high activation energy

means there will be a slow rate and k will be smaller

high temperature

as it increases, k rises as more reactants have the reaction energy

logarithmic form

allows Arrhenius equation to calculate activation energy

ln k = -(Ea/RT) = ln A

ln = logarithmic button

Arrhenius plot

made by plotting ln k again 1/T

produces a graph with a gradient of -Ea / R and a y-intercept of ln A

used to find activation energy and pre-exponetial factor (A)

work out gradient then times it by R (8.31 J K-1 mol-1) to give activation energy

find y intercept and use e, this works out A, or substitute in values to equation

rate determining step

slowest step in a multistep reaction

at least one reactant must appear in rate equation

used to predict rate determining step

affect the rate

reactants that appear on rate equation

order of a reactant

tells you how many molecules of that reactant are involved in rate determining step

mechanism

can be predicted with the rate equation, based on what appears in rate determining step

decomposition

can mean that there can be one reactant with 2 moles in rate determining step but its not second order as it breaks down by itself instead of reacting together

costs involved in producing a chemical

raw materials

fuel/energy

overheads/fixed costs

disposal costs

raw materials

chemicals needed for the reaction

cheap, widely available ones are best

fuel/energy

reactions needing Hugh temperature or pressure use more energy

transporting chemicals uses energy

overheads/fixes costs

costs that need to be met regularly, including staff wages, rent, taxes, insurance, bills, etc

disposal costs

any unwanted by-products need to be disposed of safely

temperature

high means faster rate of reaction by more expensive

pressure

higher makes gaseous reaction faster but uses lots of energy so expensive, also very dangerous

catalyst

used to speed up reactions allowing for lower temperatures, but can be expensive and may have to be separated form products if in same state

good investment as don’t get used up

equilibrium constant

ratio of products and reactants at dynamic equilibrium , represented by Kc

Kc = [D]d[E]e / [A]a[B]b

can also be used to determine concentration of things given Kc and some concentrations

dynamic equilibrium

when the forwards and backwards reaction cancel each other out in a reversible reaction

alter Kc

temperature changes do but pressure changes do not

effect of temperature on Kc

if increased reaction shifts endothermic direction

if decreased reaction shifts to exothermic direction

if more product formed then Kc rises and the opposite the other way round

effect of pressure on Kc

as the reaction always reverses the change there is not effect on Kc

effect of catalysts on Kc

the have no effect as they do not increase yield they only mean equilibrium is approached faster

how economical a reaction is

determined by position of equilibrium and the yield of the reaction

generally a compromise of conditions

harder process

carried out at 400˚c and 200 atmospheres of pressure

higher pressure makes it faster but very expensive

high temperature increases rate but decreases yield as forward reaction is exothermic so equilibrium gets shifted to left, so compromise between speed and yield

risks of production

some chemicals are high flammable and explosive eg pressured hydrogen

some chemicals are toxic/harmful to health eg chlorine

some chemicals damage the environment eg acid rain

Kc units

put the units into Kc equation and cancel them out

(mol dm-3) (mol dm-3) / (mol dm-3) = mol dm-3

find concentrations at equilibrium

by using a technique that won’t disturb position of the equilibrium eg colorimetry or pH

colorimetry

can be used to determine concentrations of a reaction if it is coloured, find absorbance and use a calibration curve to determine concentration

pH

if one side of the reversible reaction contains an acid or base at equilibrium the a pH probe can be used and concentration calculated

never use titration

as it would cause any reactants/products to to react and from a salt and water decreasing the concentration therefore shifting the equilibrium, this means it cannot be used to find equilibrium concentrations

ice tables

method to work out initial concentrations and equilibrium concentrations

initial eg 1

change eg -0.5

equilibrium eg 0.5

diatomic molecule

nitrogen, top of group 5, electronic configuration of 1s2 2s2 2p3 so there’s 5 electrons inter shell they pair up to share 3 of them making a triple bond, difficult to break so there unreactive

ammonia NH3

formed form 1 N and 3H, forms 3 covalent bonds and has one lone pair

can form hydrogen bonds so is very soluble in water

acts as a by the lone pair forming a dative covalent bond with complex ions with transition metals

lone pair also causes It to act as a base as it can from date covalent bond with H+

nitrogen oxides

NO, N2O and NO2

NO

nitrogen monoxide, nitric oxide, or nitrogen (ii) oxide, colourless gas

N2O

dinitrogen monoxide, nitrous oxide or nitrogen (i) oxide (laughing gas) sweet smell and colourless

NO2

nitrogen dioxide or notion (IV) oxide, down gas, sharp order and is toxic

sodium hydroxide

added to a mixture to test for NH4+ (ammonium ions), gently heated, causes ammonia gas to be produced which is alkaline so turns damp red litmus paper blue

aluminium

used to test for NO3- (nitrate (V) ions) by warming solution with sodium hydroxide solution (devardas alloy also used, contains aluminium), the nitrate ions are reduced in the presence of the alkali, ammonia gas is produced, turns damp red litmus paper blue

nitrogen cycle

process where nitrogen goes to ammonia, ammonium ions, nitrate (III) ions, nitrate (V) ions and back to nitrogen to to nitrogen oxides

reduce half equations

involved in nitrogen cycle