Rates of Reaction

Define a catalyst

Increases rate of reaction by lowering activation energy (min E required to start reaction) that is an alternative reaction pathway. Catalysts are not part of the chemical reaction and thus not altered by the end of it.

Distinguish between homoegenous and heterogenous catalysts

Homoegenous: Catalysts is in same phase as reactants (eg. both liquid)
Heterogenous: Catalysts is in different phase as reactants (eg. reactant gas but catalyst liquid)

List examples of catalysts and their uses

• Biological catalysts: controls biochemical reactions within cells. Important industry to allow industrial reactions to happen at lower temperatures and pressures than usually needed (saves time + $)

• Transition metals: Able to form more than one stable oxidation state

Detail the reaction rate graph and what each features tells about the rate of reaction

Slope: how fast the rate is. Steeper slope = quicker reaction rate. Always a positive value.

Y-intercept: inital concentration of reactants/products at 0 sec

List some ways to measure the rate of a reaction

• Mass loss

• Gas production

• Colorimetery

What is colorimetery and how does it work?

• Colorimetry measures the light intestity of light passing through a sample

• If solution changes colour = changes rate.

• Intensity of light reaches detector that measures every few seconds plots data on graph to show how the concentration of the reactants and products changes with time.

• Light intensity is related to concentration.

What is measuring rate of reaction using changes in mass and how does it work?

• When gas is produced in reaction = escapes from setting/reaction vessel = mass decreases = measured

• Ex: Calcium carbonate and hydrochloric gas produces CO₂

• Measuring changes in mass using a balance and cotton wool blocks the neck of flask so only gas escapes.

• Limitation is that mass of gas is too small to measure on 2/3 decimal place balance

What is measuring rate of reaction using changes in volume of gas and how does it work?

• Method (flask): Gas syringe traps gas by

• Ex: magnesium with hydrochloric acid produces hydrogen gas

• Method (inverted measuring cylinder):

What is measuring rate of reaction by titration with concentration changes and how does it work?

What is measuring rate of reaction by using conductivity how does it work?

What is a clock reaction and how does it work?

List the factors of collision theory

• Collision frequency= increase = more chances of preferred orientation of substrate to catalyst

  • More particles = more frequent collisions/collision per unit time

• Collision energy= Sufficient energy = successful collision (combined energy of colliding particles = collision energy)

• Activation energy

  • Min energy required for colliding particles need to to react.

  • Collision E < Activation E = unsucessful collision

  • Collision E ≥ Activation E = successful collision

  • Activation E can be changed by addition of catalyst

• Collision geometry

  • Right orientation = successful

Define collision theory

Explain why chemical reactions occur and at what rate. Reactions must happen when reactant molecules collide with each other. Collisions must have suffieint energy (activation energy), propert orientation for reactants to rearrange and form new bonds.

Distinguish between successful and unsuccessful collision

Successful: Particles colide in correct orientation AND has sufficient energy for chemical reaction to occur

Unsuccessful: Particles collide in wrong orientation OR doesn’t have enough energy and bounces off eachother without causing chemical reaction

Outline how pressure, concentration, volume, temperature, surface area, and use of catalysts affect collision rates.

Pressure (gas only) : Increase pressure = less volume but same number of particles = less space in which they can move = succesful collisions because increased collision frequency

Concentration(no solid): More concentrated solution = greater number of particles in given V of solvent = increased collision frequency/collision per unit of time = more frequency of successful collision

Volume(gas only): Lower

Temperature: High temperatures = increase of kinetic energy = more collision = higher frequency proportion

Surface area (heterogenous reactions (reactants are diff state) only). Increased SA=

Catalysts:

Define a Maxwell-Boltzmann distribution curve

• Graph that shows distribution of energies at certain temperature

• Small proportion of particles in sample have enough energy for an effective successful collision and for chemical reaction to take place

• Most probable energy of particle = highest point of curve (E_MP)

• Area under curve = 100% of particles of solution (doesn’t change even if temperature changes)

List the effects of temperature on Maxwell-Boltzmann distribution curve

• Increased T = more kinetic kinetic energy = more frequent collisions = higher proportion of particles posses min activation energy compared to decreased T

• Wider graph. peak shifts to right

• Activation energy = same

List the effects of catalyst on Maxwell-Boltzmann distribution curve

• Activation energy