Concept+of+the+Rate+of+Reaction

__**Chemical Kinetics**__ is the area of chemistry that deals with the concept of the rates of reactions. Another way to phrase the idea of rates of reactions is to consider the rate to be the same thing as the speed of a reaction, or how fast it happens. Reactions can differ greatly in their speeds or rates. The basis for how fast a reaction occurs is how often and with how much energy the reactant molecules collide (though they do need a certain amount of energy to react at all). As with other aspects of a chemical reaction, its rate can be affected by other factors, including:
 * The Physical State of the Reactants
 * reactants can be in either solid, liquid, or gaseous form, and the speed of their reactions depends on which form they are in
 * in the gaseous state, molecules move faster, so the reaction will occur faster because the molecules will collide more often
 * if the reactants are in different phases, the reaction is limited to the space that the two reactants share
 * The Concentration of the Reactants
 * the rate of a reaction will be faster with a larger concentration of reactants because with a larger concentration, there are more reactant molecules that are able to collide more often with each other
 * The Temperature of the Reaction
 * as with increased concentration, the rate of the reaction will increase with a higher temperature
 * if a temperature is higher, the energy of the molecules will be greater, causing them to collide quicker and therefore react at a higher speed
 * Catalysts
 * a catalyst is an agent that will speed up a chemical reaction and is never used up
 * a catalyst is able to speed up the chemical reaction by reducing the activation energy which is the minimum amount of energy needed to start a chemical reaction
 * an example is an enzyme which is a protein molecule

The unit for describing the rate of a reaction is Molarity per second (M/s). This is because we measure the speed of a reaction by determining the change of a reactant's concentration over time passed. The rate of reaction can be expressed in terms of the disappearance rate of a reactant, or the appearance rate of the product. For example, if molecule X reacts with molecule Y, than the rate of disappearance of X will equal the rate of disappearance of Y. This, however, is not true for reactions that contain reactants and products not in one-to-one ratios. That is where the stoichiometry below explains the rate in terms of different molar ratios.


 * The dots do not imply any information, they are only in place for the spacing of words and equations.

The rate of disappearance of X is equal to __ Δ X (change in X's concentration)__ ..........The negative sign is there because rates are always Δ time (change in time) ......................expressed as __positive quantities__.

The rate of appearance of Y is equal to __ Δ X (change in X's concentration)__ Δ time (change in time)

Often, the rate of a reaction will decrease over time and is shown through a grapzh such as The rate of the reaction of at a particular moment in time is known as the __instantaneous rate.__ The instantaneous rate is determined by the slope of the curve between the two periods of time. For example, if we want the rate of a reaction at 60 seconds, we draw a tangent line form 0 to 60 seconds and use the equation: __ Δ in reactant __ or __ ( [the concentration of the reactant at 60 seconds] - [the concentration of the reactant at 0 seconds])__ Δ in time ...................................................................(60 seconds-0 seconds)


 * __ Stoichiometry and Its Effects __**

When determining the rate of reactions in which the stoichiometry of the reactants and products differ, the rates fo appearances and disappearances will no longer be the same, For example: 2 HI(g) ---> H2(g) + I2(g) rate = __-1__ __Δ[HI]__ = __Δ__ __[H2]__ = __ Δ[I2] __ ..............................................2.. Δt........ Δt ........Δt
 * Here the rate of the disappearance of HI(g) will occur twice as fast as appearance of H2(g). This is due to only one mole of H2 and I2 being made per 2 moles of HI.
 * The overall idea of stiochiometry in determing the rate of reactions is expressed in the general equation of:

a A + b B ---> c C + d D rate = __-1__ __Δ[A]__ = __-1__ __Δ[B]__ = __1__ __ Δ[C] __ = __1__ __ Δ[D] __ .............................................. a ..Δt.........b ..Δt ......c ..Δt ....d ..Δt **__Rate Law__** The rate law of a reaction demonstrates how the rate of reaction is directly proportional and dependent on the concentrations of the reactants. Usually, the form of this law is:

Rate = //k//[A]m[B]n *m and n are exponents

//k// is the rate constant, which can be determined if we know the initial rate law of a reaction and its rate for a set of reactant concentrations. Just substitute data, solve for //k//, and then use //k// to solve for the rate law of any other reactant concentrations.

Rate = //k//[A]m[B]n //k// = Rate/[A]m[B]n Rate of new concentrations = //k//[C]d[E]f and then calculate

The unit of the rate constant //k// is L/ mole sec. However, //k// is also dependent on the reaction order because  there may be fewer liters per moles per second or more depending on if the reaction of an order is second, first, or zero.  The exponents in the rate laws are called reaction orders and the sum of all the exponents of a rate law is the overall reaction order. For example, if the rate law = //k//[A][B], than the overall reaction order is 2. The reaction orders show how the concentrations of the reactants affect the rate law. For example, a reactant is first order (exponent of 1) if the rate doubles every time the concentration doubles and it is second order (exponent of 2) is the rate quadruples as the concentration doubles. The reaction order may be zero order (exponent of 0) if it has no effect on the rate. The values of the exponents do not always mirror the coefficients in a chemical formula and must be determined experimentally, just like the data used in rate laws.  For more examples or explanations, visit f acstaff.bloomu.edu/eschultz/CHEM%20116%20SPRING%202007/ 116sp07%20kinetics1B.ppt