Rate of reaction Flashcards
Collision theory states that for a reaction to occur reactants must :
- COLLIDE
- In the correct ORIENTATION
- With energy GREATER than the ACTIVATION energy , which is the minimum energy molecules must have to react.
Rate of reaction
- How FAST a reaction is taking place
Indicates how fast either:
- the concentration of the reactants decrease
- the products increase
( at a certain point during the reaction)
FASTEST RATE ~ BEGINNING :
the concentration of reactants are at their highest values and it slows down as the reactants are used up.
- The reaction STPS when ONE of the reactants has fully reacted
Effect of TEMPERATURE
An INCREASE in temperature results in an INCREASE in the rate of reaction because:
- KINETIC ENERGY of the reactants increase
- COLLIDE more frequently
- GREATER PROPORTION of the reactants have ENERGY GREATER than the ACTIVATION ENERGY
- MORE SUCCESSFUL COLLISIONS
Effect of CONCENTRATION
By INCREASING the concentration of the reactants:
- MORE reactants PER UNIT VOLUME
- MORE COLLISIONS per unit time because the reactants are CLOSER TOGETHER
- A greater PROPORTION of collisions will have energy GREATER than the ACTIVATION ENERGY
- Rate of reaction INCREASES
Effect of PRESSURE
An INCREASE in pressure results in :
- More GASEOUS reactants PER UNIT VOLUME
OR - The volume of space occupied by the reaction mixture DECREASED.
- More COLLISIONS per unit time because the reactants are CLOSER TOGETHER
- A greater PROPORTION of collisions will have energy GREATER than the ACTIVATION energy
- Rate of reaction INCREASES
Effect of SURFACE AREA
For the SAME MASS, a REDUCTION in particle size leads to an INCREASE in surface area:
SA : powder > lumps/ribbons
An INCREASE in surface area:
- A GREATER number of COLLISIONS to take place with GASEOUS or LIQUID reactions per unit time
- A greater PROPORTION of collisions will have energy GREATER than the ACTIVATION energy
- Rate of reaction INCREASES
Graph showing REACTANT concentration plotted against TIME
BEGINNING:
- The rate at t=0 is called the INITIAL RATE
- The concentration of the reactants are at their HIGHEST at t=0
- The gradient of the curve and rate are at their highest
MIDDLE:
- The concentration of the reactants has FALLEN
- FEWER successful collisions taking place
- The gradient of the curve and rate DECREASES as the reaction PROCEEDS
END:
- The concentration of the reactants is
ZERO
- The gradient of the curve and rate both become ZERO
Graph showing PRODUCT concentration plotted against TIME
BEGINNING:
- The concentration of the product is at its lowest
- The reactants have just started to react
MIDDLE:
- The concentration of the product continues to INCREASE
- The reactants are being CONVERTED into products
- The rate is DECREASING as reactants are being used up
END:
- ALL of the reactants have been converted to products
- The rate LEVELS OFF and become ZERO
Calculating GRADIENT ( rate of reaction )
ROR= change in concentration/ time taken for change
(This formula will just give the correct value or magnitude of the rate)
CATALYSTS & their PROPERTIES
CATALYSTS ~ increase the rate of reaction by providing an alternative route with a lower activation energy.
- NOT used up in a chemical reaction
Either:
- React with a reactant to form an
INTERMEDIATE
- Provide a SURFACE on which the reaction
can take place
- They are REGENERATED by the end of the reaction without being permanently changed
Phase
- If a BOUNDARY can be seen between two different components in a mixture
- They are in DIFFERENT PHASES
- Like physical states, the types of phases are S , L , G & AQ.
Homogenous Catalysis
- The catalyst is in the SAME phase or state as the reactants
- Usually in the aq , l or g phase
- The catalyst reacts with the reactant to form an INTERMEDIATE which breaks down to give the products and is REGENERATED
EXAMPLE:
- H2SO4 when making esters
- ethanol and ethanoic acid are in the same phase as the H2SO4(aq)
Heterogeneous Catalysis
- The catalyst is in DIFFERENT phase to the reactants
- Usually catalyst is a SOLID and reactants are either GASES or LIQUIDS
- The reactants are ADSORBED ( weakly bonded ) onto the surface where the reaction takes place
- After the reaction , the products leave the surface of the catalyst via DESORPTION
INDUSTRIAL PROCESSES using Heterogeneous catalysts
MAKING NH3 ( HABER PROCESS):
Catalyst ~ Fe(s)
Equation N2(g) + 3H2(g) —– 2NH3(g)
PRODUCING SO3 FOR MAKING H2SO4 ( CONTACT PROCESS):
Catalyst ~ V2O5(s)
Equation ~ 2SO2(g) + O2(g) —- 2SO3(g)
HYDROGENATION OF ALKENES :
Catalyst ~ Ni(s)
Equation ~ H2C=CH2(g) + H2(g) ——CH3CH3(g)
SUSTAINABLE & ECONOMIC IMPORTANCE of catalysts
- Increase the rate of reaction by providing an alternative route with a lower activation energy
- The TEMPERATURE and ENERGY requirements are consequently reduced
- LESS electricity and fossil fuel are used
- The RUNNING COSTS are reduced
- The products are made faster and PROFITS increase
- Benefits OUTWEIGH cost of developing catalytic process
- Industries aim to use processes with HIGH ATOM ECONOMIES and minimise the products POLLUTANTS
- The emission of CO2 which is linked to GLOBAL WARMING, is REDUCED
The BLOTZMANN DISTRIBUTION of molecular energies in a gas, liquid or solution at a constant temperature
KEY FEATURES
W :
- The PEAK of the curve represents the AVERAGE ENERGY of the sample
- The DISTRIBUTION is NOT SYMMETRICAL as most particles have INTERMEDIATE energies
X - Curve starts at ORIGIN as NO particles have ZERO energy
Y - The total AREA under the curve is equal to the total NUMBER OF PARTICLES
Z - NO INTERCEPT of the energy axis by the curve at high energy
Boltzmann Distribution ~ TEMPERATURE INCREASE
Average energy:
- INCREASES
- Each particle will gain a certain amount of energy
- The peak moves to the RIGHT
Distribution:
- BROADEN
- Each particle will gain a certain amount of energy
- The peak of the curve LOWERS
Area under the curve:
- STAYS THE SAME
- The total number of particles does not change in the sample when the temperature increases
Boltzmann Distribution : TEMPERATURE DECREASE
- The peak of the curve shifts to a lower energy (LEFT)
- The curve becomes NARROWER and more POINTED due to the smaller spread of energies
- FEWER particles that are able to react at a lower temperature
- Area under the curve remains the SAME as number of particles remains unchanged.
Boltzmann Distribution : ACTIVATION ENERGY at two different temperatures
ACTIVATIN ENERGY~ the minimum energy particles must have for a reaction to take place.
- The SHADED area under the curve beyond Ea indicates the number of molecules with the activation energy
- An increase in TEMPERATURE leads to a GREATER PROPORTION of particles having an energy greater than the Ea.
- This leads to more FREQUENT & SUCCESSFUL collisions
- Rate of reaction INCREASES
Boltzmann Distribution : CATALYST
CATALYST ~ provides an alternative route for the reaction with a lower activation energy.
- The activation energy will be to the LEFT of the activation energy for the uncatalyzed reaction
- The DISTRIBUTION of molecular energies remains UNCHANGED
The use of a CATALYST leads to :
- a GREATER AREA under the curve beyond the activation energy
- A GREATER PROPORTION of molecules have energy greater than the ACTIVATION ENERGY
- MORE SUCCESSFULL collisions
- Rate of reaction INCREASES
