CHEMICAL KINETICS - RATES OF REACTION -

 

Chemical Reactions

    • Occur at different speeds or rates, give some examples

Chemical Kinetics

- the study of rates of reactions and the factors which affect the rates

- starts with empirical research of reaction times

- theoretical side is not well understood.

A. Measurement of reaction rates

    • reaction rate is obtained by measuring either

- rate at which product is formed

- rate at which reactant is consumed
  • mass
  • volume
  • colour
  • pressure
  • conductivity
  • pH
  • concentration
  • temperature

 

- the rate is expressed as :

       rate = a change in property of reactant or product
                                  time

Example :

A sample of Mg weighing 0.360 g is dropped into dilute HCI. At the end of 4.00 minutes, the unreacted Mg is removed and found to weigh 0.240 g. Calculate the average rate of reaction.

 



Rate = change in mass = m1 -m2 = 0.360 g - 0.240 g
            time           t        4.00 min

= 0.0300 g/min

 

 

Questions 1-5 page 2 of Hebden, and question 6 page 3 of Hebden.

Questions 7-9 page 5 of Hebden.

 

 

 

A Special Case : Using Concentration To Measure Reaction Rates :

Concentration is often used, but we must remember that :

    • concentration changes are more rapid near the beginning of the reaction and decrease with time elapsed.

 

Example CH4(g) + Cl2(g) ® CH3Cl(g) + HCI(g)

  • the reaction rate at any given time is the slope of the curve at that point (tangent line)

  • as the reaction proceeds, slope is less, therefore rate is slower

Do questions 18-19 page 11 of Hebden

 

 

B. Factors that can Affect Reactions Rates

  • Homogeneous Reaction - a reaction in which all of the reactants are in the same phase. All factors listed below except surface area can affect homogeneous reactions.
  • Heterogeneous Reaction - reaction in which reactants are present in different phases. All factors listed below can affect heterogeneous reaction.

Factor

Affect

Temperature

As temperature increases, rate increases *

Concentration

As concentration increases, rate increases

Pressure of Gaseous Reactants (similar to concentration)
  • As pressure increases, rate increases
  • As volume increases, rate decreases

Surface Area of a solid

As surface area increases, rate increases **

 

 

Nature of Reactants
  • Few or weak bonds react faster than those with many or strong bonds.
  • Oppositely charged ions react quickly
  • Reactions in homogeneous phases are generally faster than those in heterogeneous phases.
  • Aqueous ions > gases or liquids > solids

Catalyst

Increases rate

Inhibitor

Decreases rate

Catalyst - a chemical which can be added to a reaction in increase the reaction rate. After the reaction is complete, the same amount of catalyst will be present as was originally put into the reaction.

Inhibitor - a chemical which reduces reaction rate by combining with a catalyst or one of the reactants in such a way that it prevents the reaction from occurring.

 * a general rule of thumb regarding the temperature of a SLOW reaction : an increase of 10 degrees will often double the reaction rate.

** coal dust can explode, but a lump of coal will burn for hours

 

Questions 10 - 17 pages 7-10 of Hebden.

 

C. Situations that need Controlled Reaction Rates
    • Body chemistry - temperature and enzymes
    • Fuels - oxygen concentration
    • Industrial chemical reactions
    • Rusting Cars
    • Cooking
    • Preserving food - refrigeration and freezing

 

 

D. Kinetic Collision Theory of Reaction Rates

- we know what affects rates of reactions, now we will examine a theory that explains why these factors affect reactions.

- assume particles are small hard spheres which can bounce off of each other without losing kinetic energy

- this theory assumes that:
  • reactions are the result of collisions between reactant particles
  • not all collisions are successful (ineffective collisions - particles rebound from the collision, unchanged)
  • sufficient Kinetic Energy and favorable geometry are required
  • to increase the rate, the rate of successful collisions must be increased
  • energy changes are involved in reactions as bonds are broken and formed.

- the average KE of particles is the temperature of the sample.

 

Do questions 20-22 page 12 of Hebden

 

 

E. ENTHALPY CHANGES IN CHEMICAL REACTIONS

Potential Energy - is a result of an objects position in space, and the sum of all the attractive and repulsive forces in particles that make up the object. The Ep of a chemical system is directly related to the energy of the electrons in the chemical bonds, and the number and type of atoms in the molecules.

Kinetic Energy - of a system is due to movement within the system. It may be a result of the entire system moving, or the movement of individual molecules or atoms in the system.

Bond Energy - the amount of energy required to break a bond between atoms. It takes an amount of energy equal to bond energy to break a bond, and the same amount of energy is released when that bond is formed.

Enthalpy (H) - heat content, or the total kinetic and potential energy that exists in a system at constant pressure.

D H - (Hproducts - Hreactants) - the change in enthalpy between the products and the reactants.

Hproducts - combined enthalpies of all of the products.

Hreactants - combined enthalpies of all of the reactants

Endothermic

Exothermic

Energy released when new bonds form is LESS than energy required to break bonds

Energy released when new bonds form is GREATER than energy required to break bonds

Products have MORE energy than reactants

Products have LESS energy than reactants

Heat is absorbed from the surrounding

Heat is released to the surroundings

D H is positive

D H is negative

D H is written on the reactants side

D H is written on the products side

2N2 + O2 + 164KJ® 2N2O

H2 + Cl2 ® 2HCl + 184 KJ
Potential Energy Diagram
Potential Energy Diagram

 

 

Do questions 24 - 28 on page 16 of Hebden.

 

 

Kinetic Energy Distributions

 
  • Molecules at room temperature and pressure have about 1 x 1010 collisions/second, so the lack of reactivity is NOT due to lack of collisions.
  • Since high temp favours fast reactions, the energy a molecules has must determine if it will react.
  • BUT at a given temperature, some molecules do react and some do not react.
  • The following graph gives an explanation.

 

 

  • Some molecules have a very low kinetic energy and some have a very high kinetic energy.
  • The average kinetic energy is the temperature
  • As the temperature is increased, the average energy is increased
  • The activation energy (Ea) is the minimum energy required before a molecule can react
  • The reaction rate is slow at room temp, since only a few molecules have sufficient energy to react at a given instant.
  • As the temperature increases, more and more molecules have the energy required to react, so the rate increases.
  • As temperature increases, the molecules also undergo more collisions, but this increase is relatively small. (increase of 10 degrees only increases collisions by 1%)

THE INCREASED REACTION RATE DUE TO AN INCREASE IN TEMPERATURE IS PRIMARILY DUE TO THE INCREASED NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO REACT, AND NOT TO THE INCREASED NUMBER OF COLLSIONS.

 

Do questions 29 - 32 page 19

 

 

Activation Energy :

  • When two molecules approach each other
    • Electrons in their out shells repel each other
    • The molecules slow down
    • Their Kinetic Energy is converted in Potential Energy
  • If the molecules can gain enough Potential Energy, bonds can be broken and an ACTIVATED COMPLEX is formed.
  • After the reaction the product molecules repel each other and convert Potential Energy back into Kinetic Energy.
  • Higher Activation Energy (Ea) is due to a greater number of electrons present at the 'reaction site' and greater bond strengths.
  • The lower the Ea, the faster the rate of the reaction.

Activation Energy - Ea - the minimum potential energy required to change the reactants into the activated complex

Activated complex - is arrangement of atoms which occurs when the reactants are in the process of rearranging to form products. (an intermediate molecule.)

activated complex

    • high potential energy, unstable, short lived configuration of reactant atoms.
    • it may decompose and form new lower energy, more stable products
    • or revert back to the reactant state.
  • Look at Diagram "Kinetic Energy Distribution Curve & Activation Energy" in Hebden.
    • The total energy of a system will always be KE + PE
    • The KE diagram shows how many molecules have enough energy to reach the Ea
    • Having enough energy does not mean that a certain molecule will react, it must have the proper orientation also.





Do questions 33 - 40 page 23 -24

 

 

 

 

Proper Orientation AND Sufficient Energy is Required

For an Effective Collsion

 

Reactions are Reversible

If you look at any PE diagram, their is nothing stopping the reaction from going in the reverse direction

Ea(f) is the activation energy for the forward reaction

Ea(r) is the activation energy for the reverse reaction

 

Show overheads

Do questions 41-45 on page 25 of Hebden

 

Do Experiment 18C Measuring Reaction Rate Using Volume of Gas Produced.

 

How The Collision Theory Explains :

  1. concentration

    • an increase in concentration, increases the rate of reaction.
    • the number of collisions per unit time would increase if more particles were packed into a given volume. Since more collisions occur, it is likely that more product would form.

  2. temperature

    • increasing temperature increase the reaction rate.
    • this is primarily due to the increased number of molecules with sufficient energy to react, and not to the increased number of collisions.
    • increased temp causes the particles to travel faster and therefore to hit each other harder (more likely to be successful)

  3. surface area

    • increasing surface area increases reaction rate
    • a solid can only react at its surface, so increasing surface area increases the number of reaction sites where successful collisions can occur

  4. nature of reactants

    • molecular substances and/or complex ions are often slower than reactions of simple ones because any breaking of a covalent bond requires a significantly high activation energy.
    • both the strength of the bond(s) to be broken and the location of the bond(s) in the structure affect the likelihood that any given collision is effective.

  5. catalysts

    • catalysts increase the rate of reaction, they do this by lowering the Ea by providing a different pathway for the reaction to take.
    • result in the same products overall. Remember D H is NOT changed
    • finding catalysts is often trial and error, many enzymes work by shape and orientation.

 

 

Reaction Mechanism Theory

 

·       This theory assumes that simultaneous collisions of 3 or more particles is unlikely.

·       Scientists believe that most reactions occur as a sequence of steps that involve collisions of 2 particles.

·       A reaction mechanism shows the individual steps during the process of a reaction.

 

          ie .   4HBr(g) + O2(g) à 2H2O(g) + 2 Br2

 

step 1 :       HBr(g) + O2(g) à HOOBr(g)                      slow

step 2 :       HOOBr(g)  + HBr(g)  à  2HOBr(g)           fast

step 3 :       HOBr(g)  + HBr(g)  à H2O(g) +  Br2         fast

               HOBr(g)  + HBr(g)  à H2O(g) +  Br2

 


Total :        4HBr(g) + O2(g) à 2H2O(g) + 2 Br2

 

 

·       Rate determining step – in any reaction mechanism, the overall rate of the reaction is determined by the slowest step.  (the chemical change cannot go faster than the slowest step)

·       Reaction intermediates (intermediate products) – substances formed during the reaction, but immediately react again and are not present when the reaction is complete.

 

See overhead of multi step reaction Potential Energy Diagram

 

Do questions 46 – 52 (omit 53) on page 28 or Hebden.

 

 

Energy Diagram of a Reaction Mechanism

  • each step involves an individual activated complex and activation energy
  • a three step mechanism would give three humps on the PE diagram
  • the hump with the highest Ea is the rate determining step

Show overheads

 

Do questions 54-55 page 30 of Hebden

 

The effect of catalysts on activation energy

Catalyst - a substance that provides an overall reaction with an alternative mechanism having lower activation energy.

  • D H for the overall reaction is NOT changed
  • if Ea is lowered by a catalyst, more reactant molecules can react.
  • The reverse activation energy is also lowered
  • The catalyst is an active participant which is REGENERATED in a later step of the reaction mechanism.

Show overheads

Do questions 56-61 page 34

 

Write notes on SOME USES OF CATLYSTS on page 34 of Hebden

Do questions 62, 63 page 36.