Factors Affecting the Rate of Reaction

Factors Affecting the Rate of Reaction

       1.) Temperature

·         Reaction rate increase with an increase with an increase in temperature this happens because the kinetic energies of the reactants increase with a rise in temperature.

·         The faster the reactant particles move, the greater the chance that they will collide and react with each other.

·         As general rule, a 10°C rise in temperature increases the reaction twice or thrice.

Example: We refrigerate perishable food such as milk. The bacterial reactions that lead to the spoiling of milk proceed much more rapidly at room temperature that they do at the lower temperature of a refrigerator.

       2.) Concentration of Reactants

·         An increase in Concentration means increase in the molecules or particles per unit volume and thus a decrease in space between the reacting particles. With less distance to travel inside the vessel, the more frequent collision, the faster the reaction.

  

Example: Steel wool burns with difficulty in air, which contains 20% O2, but burst into a brilliant white flame in pure oxygen.

                  In 1967, Ground Test of Apollo 1, the spacecraft exploded because there was a leak in its fuel container. Its fuel (concentrated O₂) is very combustible.

 

The rate of reaction is proportional to the concentration of the reactants raised to various powers. this is known as the Rate Law.

Rate Law - express the relationship of the rate of reaction to the rate constant and the concentration of the reactants raised to some powers.

 

 

aA + bB --> cC + dD                          where small letters - coefficients

                                                                          BIG letters - substants

 

             rate law = k [A] ^m [B] ⁿ

                          ^{where m and n are integers or fractional numbers that must be determined experimentally. In general, m and n are not equal to the stoichiometric coefficients a and b. When we know the values of m, n and k, we can use the equation above to calculate  the rate of reaction given the concentration of A and B.}

Exponents m and n specify the relationships between the concentration of reactants A and B and the reaction rate. Added together, they give us the over all reaction order, defined as the sum of the powers to which all reactant concentrations appearing in the rate law are raised. overall reaction order is (m + n).

If A is in mth order and B is in nth order, then the overall reaction order is (m + n).

*the higher the overall reaction order, the faster the reaction.

*the lower the overall reaction order, the slower the reaction.

Rate Data for Reaction Between F_{2 and} ClO₂
 

[F2](M)	[ClO2](M)	Initial Rate (M/s)
1.)              0.10	0.010	1.2 x 10-3
2.)             0.10	0.040	4.8 x 10-3
3.)             0.20	0.010	2.4 x 10-3

Computing for the rate law:

   rate law = k [F₂]^m [ClO₂]ⁿ

First, solve for the raised various powers using the formula:

m =         log rate ratio                 note: to get the experimental values of

           log concentration ratio                 exponents m and n, look at the

                                                                      concentration of the reactants

                                                                      Use the rates and concentrations of the two

                                                                       experiments where one of the concentration

                                                                       of one of the reactants is changed while the

                                                                       other is kept constant or unchanged.

 

 

Using experiments 1 and 3 for the exponent of F_2:

 

 

 m =         log [(2.4 x 10^-3)/(1.2 x 10^-3)]          

                                      _{log [(0.20)/(0.10)]}

 

_{m = 1}

 

_{Using experiments 1 and 2 for the exponent of} ClO_2:

_{n =}         log [(4.8 x 10^-3)/(1.2 x 10^-3)]          

                             _{log [(0.040)/(0.010)]}

 

_{n = 1}

 

_{therefore, the rate law can be express as:}

_{rate law =} k [F₂]¹ [ClO₂]¹    

                          or 

rate law = k [F₂] [ClO₂]  

 

Solving for k(rate constant):

manipulating the equation rate law = k [A] ^m [B] ⁿ

 

^{k =    rate law}  

        [A]^m [B]ⁿ

 

then,

k =     rate law   

       [F₂] [ClO₂] 

 

using experiment 2, we can solve for k,

 

k =      (4.8 x 10^-3)      

           (0.10)(0.040)

 

k = 1.2

 

       3.) Presence of Catalyst

·        Catalyst – a substance which controls the rate of reaction without itself undergoing a permanent chemical change. Catalysts are of two types.

·         Positive Catalyst - increases the rate of reaction by lowering the energy of activation. Thus in the presence of a positive catalyst, the greater fraction of the total molecule will posses lower energy of activation and collided successfully in a short period of time, there by increasing the rate of reaction. A positive catalyst functions by providing an alternate path to the reaction or by the formation of a transition (intermediate) compound having low energy of activation. The activation energy of this path is lower. As a result rate of reaction is increased.

·         Negative Catalysts or Inhibitors –  retards the rate of reaction. Negative catalyst do not lower the energy of activation rather they are combined with reactant molecule thus decreasing the number of colliding reactant molecules. This decreases the effective collisions, hence rate of reaction. A negative catalyst does not lower the energy of activation rather it combines with the reactant molecules, thus decreasing the number of colliding reactant molecules. This decreases the effective collisions, hence rate of reaction.

Example: The physiology of most living species depends crucially on enzymes, protein molecules that acts as catalysts, which increase the rate of selected biochemical reactions.

  4.) Nature of Reactants

·         Ionic Compound – reacts easily

·         Covalent Compound – reacts slowly

·         Polar-Polar – mixes or dissolves easily

·         Nonpolar-Nonpolar – mixes or dissolves easily

·         Polar-Nonpolar – mixes or dissolves slowly

·         Homogenous reactants reacts faster than heterogeneous reactants

Example: Reactions occurring between  oppositely charged ions are generally rapid at room temperature. In reactions involving covalent substances where there are many covalent bonds to be broken, the reaction usually takes place slowly at room temperature. the H-H and O-O bonds in the H2 and O2 molecules must be broken and then two bonds between H and O must be formed to produce the H-O-H molecule.

        5.) Surface Area

·         The greater the surface area, the faster the rate of reaction.

·         The smaller the size of particles, the larger the surface area exposed.

Example: With the same volume, between the small ice cubes and a block of ice, the ice cubes will melt faster because it has greater surface area exposed.