Chemistry Helpers
Kinetics and Equilibrium


Kinetics and Equilibrium
I:   Definitions
1)    Activation Energy: the minimum amount of energy needed to produce an activated complex
2)    Heat of Reaction: the amount of heat released or absorbed in a reaction        
3)    Exothermic Reaction:      a reaction that releases heat energy      
4)    Endothermic Reaction:  a reaction that absorbs heat energy         
5)    Activated Complex:         the temporary, unstable, intermediate union of reactants
6)    Equilibrium:  a dynamic chemical condition in which opposing reactions are proceeding at equal rates, producing an apparent constant condition.              
II: Kinetics
A.   Chemical Kinetics deals with:
1)          The rates of chemical reactions                                          
2)          The pathway by which the reaction occurs                         
B.   A Reaction Rate depends on several factors: in order for particles to react, they must collide with each other. Each collision must have enough energy and hit in the right direction-they must be effective.
1)          The nature of the reactants: solutions of ionic solids tend to react more quickly than non-ionic solids because ions are created. Polar reacts more quickly than non-polar.               
2)          The concentration of the reactants: an increase in the concentration will increase the rate of the reaction. For gases, an increase in pressure brings the particles closer together and acts like an increase in concentration.                         
3)          Surface Area: an increase in surface area will increase the rate of the reaction. A powdered substance will always react faster than a solid of the same amount.                                      
4)    The temperature of the system: an increase in temperature will increase the rate of a reaction. Why? Because more energy speeds up particles and increases the number of collisions. The more collisions, the more effective collisions, the faster the reaction.      
5)    Presence of a Catalyst: a substance that increases the rate of a reaction without chemical alteration to itself. **Will not begin the reaction, it just lowers the required activation energy by providing a new pathway for reacting particles.     
C.   Roles of Energy in Reactions
1)    Activation Energy: the minimum energy needed to cause a reaction to begin         
2)    Heat of Reaction: the difference between the potential energy of products and of the reactants        Formula: ∆H = H products – H reactants
a) Exothermic Reaction: reactions that release energy. The potential energy of the products is lower than the potential energy of the reactants. The sign for the Heat of Reaction (∆H) is negative.
b) Endothermic Reaction: reactions that  absorb energy. The potential energy of the products is greater than the potential energy of the reactants. The sign for the Heat of Reaction (∆H) is positive.
III: Equilibrium
A.   Dynamic Equilibrium: When the forward and the reverse rates of a reaction are equal.
B.   Phase Equilibrium: In a closed container, when the rate of melting is equal to the rate of freezing; or when the rate of evaporation is equal to the rate of condensation.         H2O (s) ↔ H2O (l)    or    H2O (l) ↔ H2O (g) 
C.   Solution Equilibrium: In a saturated solution (in a closed container), when all that can be dissolved in liquid has been reached, the rate of dissolving and the rate of recrystallizing are equal;  C12H22O11 (s) ↔ C12H22O11 (aq)   or      CO2 (g) ↔ CO2 (aq)        
1)    Solubility: The maximum amount of a substance that can dissolve in a certain amount of a solvent at a particular temperature.
a)    Solute: the dissolved portion of a solution;  the one present in smaller amount
b)     Solvent: the part of a solution in which the solute is dissolved;  the one present in greater amount (usually water) 
**We will use Table G to determine solubility**
D.   Chemical Equilibrium:      when the forward and reverse reactions occur at equal rates.  ↔ The products can become the reactants again. **The concentration of the reactants and the products are kept constant but do not have to be equal!!!    
E.    Le Chatelier’s Principle: changes to concentration, pressure or temperature are known as  applied stresses. This principle describes what happens to a system under such stresses.
1) Concentration Changes: If there is an increase in concentration, the system will want to use up what has been increased, therefore the reaction will shift away from the increase. If there is a decrease in concentration the system will want to create more of what was removed, therefore the reaction will shift towards the decrease. At equilibrium, if you add more reactants, equilibrium is upset and must return to the value so equilibrium will shift to  the product side of the reaction
*put in more = shift to the other side     *take out = remain on the same side
2) Pressure Changes: **effects GASES** 
a) Increase the pressure, the system will move to the side with  the fewest number of molecules (count the coefficients).
b) Decrease pressure, the system will move to the side with     the most number of molecules (count the coefficients).
3) Temperature Changes: 
a)     If the system is exothermic:
                   -an increase in temperature will shift to  reactant side
                   -a decrease in temperature will shift to    product side
b)   If the system is endothermic:
                   -an increase in temperature will shift to  product side
                   -a decrease in temperature will shift to    reactant side
4) Effect of a Catalyst: since catalysts will increase both the forward and the revere reaction, there is  no change in equilibrium
F.     Spontaneous Reactions: one that occurs in nature under a given set of conditions.
Gibbs Free Energy Change: ∆G tendency of a reaction to proceed to a minimum energy and a maximum entropy.   Whether or not reactions proceed seems to depend on the balance of two basic principles:
1)    The drive toward greater stability (reduced potential energy)      -∆H
2)    The drive toward less organization (increased entropy)  +∆S     
a)    Entropy: the amount of randomness or disorder in a system; the symbol for the change in entropy is ∆S
∆S solid > ∆S liquid > ∆S gas
          Gibbs Free Energy Change Reaction: ∆G = ∆H - T∆S
-∆G is Spontaneous
For a reaction to occur spontaneously – more likely exothermic reactions with an increase in entropy.
          -∆G = a spontaneous reaction
+∆G = a non spontaneous reaction
∆G is 0 = a reaction at equilibrium.

Write Your Questions Below

while a sample of liquid is freezing, its average kinetic energy is doing what?
November 6, 2014

 During a phase change (melting/freezing, vaporizing/condensing) the kinetic energy remains the same and the potential energy changes. Remember that the definition of temperature is the average kinetic energy of a system so if the temperature is changing so is the kinetic energy. Only one energy can change at a time.

November 10, 2014 -  Replied By Expert

. The reaction: 2 NO (g) + Cl2 (g) ⇔ 2 NOCl (g) is allowed to reach equilibrium. Analysis of the equilibrium mixture shows [NO] = 0. 574, [Cl2] = 0.226, and [NOCl] = 0.148. Calculate Kc for this reaction at this temperature.
December 5, 2011

Whenever you are calculating Kc for an equation you must divide the concentration of the products(raised to any coefficients present) by the concentration of the reactants (raised to any coefficients present).  For example:


2H2 + O2 → 2H2O      Kc = [H2O]2 / [H2]2 x [O2]

I am sorry that we are unable to answer exact homework questions, we are only able to guide you but if you input your data into the formula noted above you can calculate the Kc for this reaction.

Kc = [product]coefficient x [product]coefficient / [reactant]coefficient x   [reactant]coefficient



December 5, 2011 -  Replied By Expert

Copyright © 2011, all rights reserved.