Analía Bellizzi – Chemistry Classes

Ronald Reagan Senior High School

## KINETICS AS

#### A review of rates from IGCSE

Rate of reactions:

• Chemical reactions occur at a definite rate (speed) determined by the reaction conditions.
• Some reactions occur very fast. Explosions are a good example of very high rates.
• Some other reactions occur at a very slow rate. Rusting normally takes a long time.

The rate of reaction is directly affected by the amount of collisions in between the reactants, so studying how different conditions affect the collisions, we can infer how the rate can be changed.

Rates can be changed by:

1. Increasing temperature:
• Increasing the temperature increases the rate of collisions between reactants, since the particles have more kinetic energy, speeding up the reaction.
2. Increasing concentration of reactants
• Increasing the concentration of the reactants increase the possibility of collisions between reactants since they are closer, so, if there is more concentration, the rate also increases.
3. Increasing surface area
• If we try to burn a log or a bunch wood splints, we know that it is easier, the smaller the particles. Increasing the surface area (or in other words, the surface of contact between the reactants increase) and the rate of the reaction increase.of wooks.

To follow the rate of a reaction, one must either measure the decrease in concentration of a reactant or the increase in concentration of a product with time.

Some techniques for doing this are:

1. Measure the volumes of gases evolved (gas syringe).
2. Volumetric analysis – samples are removed at regular intervals, the reaction stopped by cooling, and mixture analysed by titration.
3. Measuring changes in pressure (for gas reactions)
4. Colorimetry may be used if one of the constituents is coloured. The colorimeter follows the change in intensity of colour.
5. A conductivity meter may be used if there is a change in conductivity during the reaction i.e. if the number of ions present is changing. A pH meter is a special type of meter which will follow changes in H+.

#### Typical results

Consider the reaction:

#### Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

This may be followed by measuring the volume of H2 collected in the graduated cylinder at intervals of time: The rate of reaction at a particular time is given by the gradient of the tangent.

At time t: rate = change of volume / time or a / b

Rates are usually expressed by the calculus notation.

The units are cm-3 s-1 or mol-1 s-1

The rate of a reaction is fastest at t = 0 and decreases steadily as the reactants are used up. Hence, the gradient of the graph decreases with time.

When examining the effect of changing conditions (e.g. concentration) on the rate of reaction, we usually take the initial gradient, at t = 0 as a measure of rate. We can then compare the initial rates of reaction under the different conditions and determine any effect.

When working in mines where there is a lot of powder (coal for example), or flower mills, gases mix with oxygen and they can explode since the particles size are very small, so the rate of reaction increase very much.

Following the progress of a reaction:

• We can follow the change in mass by putting some chips in HCl solution and record how the mass changes with the time. (We need a digital balance). Why is a cotton wool used to cover the opening of the flask?  Observing the graph we can infer that the mass is lost faster at the beginning of the reaction. As the reactants are used up, the reaction slows down and we have no mass loss after a while. Why the mass is not going down to zero?

• We can follow the amount of gas delivered in a reaction by placing HCl + Mg, to measure the amount of H2 delivered. In this case, we have two curves, one that corresponds to the cold acid, (blue) and the other curve that corresponds to the hot acid, (red)  • We can follow the amount of solute produced by precipitation in the reaction of sodium thiosulfate with HCl by measuring the time it takes to cover the mark underneath the flask.  Since the speed of the reaction changes, so does the slope of the curve , at the beginning we have a fast reaction but as the reactants are used up, the rate decreases and so the slope of the curve. In the case of a PRODUCTION of gas, we will have a curve like the one below. If the measurement is for example the mass DECREASE, like the example of the marble chips with hydrochloric acid, we will have the curve inverted, since the mass amount decrease with the progress of the reaction. Change in Temperature:

As the temperature increases, the particles move faster and the amount of collisions also increase, which increase the rate of reaction.

So:

• The reactant particles move quicker
• They have more energy
• The particles collide more often, and more of the collisions are successful
• The rate of reaction increases

In the example below you can see that the increase in the temperature, increases the amount of gas delivered for a given reaction.

Compared to a reaction at a low temperature, the graph line for the same reaction but at a higher temperature:

• Has a steeper slope at the beginning
• Becomes horizontal sooner, showing that the reaction time is less Change in Concentration

Increasing the concentration also increases the rate of reaction since there are more chances of collision.

The rate of a chemical reaction can be changed by altering the concentration of a reactant in solution, or the pressure of a gaseous reactant.

If the concentration or pressure is increased:

• The particles in the reactant are closer.
• There is a bigger chance of collisions.
• The rate of reaction increases When we compare concentrations we normally talk about liquids. When increasing the pressure, we only refer to gases, since the pressure change does not affect the collisions when the reactans are in liquid state. Anyway, increasing the concentration of a liquid or the pressure in a gas reaction, we will notice the same consequences:

• The graph has a steeper slope at the start
• Becomes horizontal sooner, showing that the reaction finishes faster

Photochemical reactions:

All reactions require a minimum amount of energy to start. Many times we achieve this amount of energy through heat,  we call Photochemical reactions those that light provides that “activation energy”

Photosynthesis is one of those examples. If sunlight is not present, photosynthesis does not take place.

Is the reaction that take place in every leave of a plant where CO2 is taken from the air and with water it is combined to form Glucose and water in presence of Sunlight and chlorophyll.

•
•

6CO2 + 6 H2  Sunlight   > C6H12O6 + 6 O2

Another reaction or set of reactions that take place in light and not in darkness are the ones involving free radicals. Free radicals are groups of atoms which have an extra UNPAIRED electron. This electron is very reactive and can produce chain reactions as the ones that produced the ozone depletion in the ’90s.

Silver salts decompose in light to produce a dark precipitate of silver metal. the reactions taking place are decomposition reactions where the silver halides decompose into Silver metal and the Halogen. Since there is a change in the oxidation state of the silver (and the halogen), this is considered a redox reaction

2AgBr ==> 2 Ag + Br2

2AgCl ==> 2 Ag + Cl2

Oxidation numbers of silver in both reactions:            Ag +1      ==>      Ag

The silver was reduced from ox #:                                   +1         to          0

NOW, AS – STUFF ONLY

Consider the reaction:

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

This may be followed by measuring the volume of H2 collected in the graduated cylinder at intervals of time: The rate of reaction at a particular time is given by the gradient of the tangent. At time t

 Rate = change of volume / time

in the case of the graph the gradient is determined by

 Rate = a/b

The units are cm3s-1 or mols-1
The rate of a reaction is fastest at t = 0 and decreases steadily as the reactants are used up.
On the graph, you will note that the final volume does not depend on the rate of the reaction.

The collision theory applied to rates

Collision theory qualitatively explains how chemical reactions occur and why reaction rates differ for different reactions.
In the picture below, increasing the concentration, the amount of collisions increase so the rate of the reaction increases as well Key points of the theory

• Molecules must collide before they can react.
• Not all collisions are successful.
• For a collision to be successful (also called effective collisions), particles need to have
• Enough kinetic energy to break bonds: As the temperature increases, molecules move faster and collide more vigorously. Most reactions involving neutral molecules cannot take place at all until they have acquired the activation energy
• Proper orientation:  Although the kinetic energy of the molecules may be the correct one, if the molecules bounce with certain orientations, the reaction may not occur. If we analyze two chemicals, A and B, which will react, we can represent the reaction as follows: All reactions have an transition state that occurs just at the moment of impact before breaking the bonds and forming new bonds. It has a very high energy, very unstable, and it is called the “ACTIVATED COMPLEX Activation Energy:

It’s the minimum energy required to break the existing bonds so the reaction can begin to form new bonds resulting in the rearrangement of atoms. In an ENERGY PROFILE the activation energy is always measured from the reactants to the highest point in the forward reaction and from the products towards the reactant in the backward reaction. The Maxwell-Boltzmann Distribution Graph

In any system, the particles present will have a very wide range of energies. The TEMPERATURE is a measurement of the KINETIC ENERGY of the molecules. The different energies of the particles in a sample can be shown the Maxwell-Boltzmann Distribution graph which is a plot of the number of particles having each particular energy.

This graph shows us that a few particles have very low energy (on the left of the curve), most of the particles have a moderate energy (yellow peak) and only some of them have enough energy to react or more (red zone) The area below the curve shows us the total amount of particles present. Changing the temperature, will change the kinetic energy of the particles in the reaction, as shown in the graph below. BE CAREFUL….. If you observe the energy for the same particles at different temperatures, you will observe that the energy of the particles increases from left to right.

This means:

• At low temperatures (BLUE CURVE), more particles will have low energy, while in the same sample only a few will reach the activation energy (minimum to react)
• At high temperatures, (RED CURVE) less particles will have low energy and many of them will have enough energy to react.

That’s why the area under the curve is much bigger under the red curve and so, many more particles will collide with the energy necessary to react. this is a very simple explanation on why we heat up something so the reaction occurs easier.

CATALYSTS The green line in the graph above represents the minimum energy that the particles need to react. This is called ACTIVATION ENERGY.

There are certain substances that provide an alternate route with lower activation energy. These substances are called CATALYSTS.

Be careful, catalysts DO NOT LOWER the activation energy. They simply form an intermediate that needs a lower activation energy.

Looking at the energy profiles, we will see something like this: If we graph again the Maxwell Boltzman plot, we will have a lower value of activation energy for the reaction when we use a catalyst. this is shown in the graph with the purple line. It means that the particles do not need to raise the energy to the original value. They will react with a lower value of kinetic energy (Temperature). HOW CATALYSTS WORK

Heterogeneous Catalysts:     In car exhaust, incomplete combustion produces carbon monoxide, CO and organic hydrocarbons. NO is also produced from the reaction of nitrogen and oxygen gas at high temperatures. The catalyst promotes the oxidation of CO and hydrocarbons (see Equations 1 and 2, below), and the reduction of nitrogen oxides (Equation 3, below).

(1)                       2CO + O2 → CO

(2)        CxHy + O2 → CO2 + H2

(3)                         2NO → N2 + O2

Each of the reactions in Equations 1 – 3 is an oxidation-reduction reaction. In Equation 1, the carbon is being oxidized and the molecular oxygen is being reduced. (Determine the oxidation numbers of the products and the reactants to convince yourself of this.) Which species are being oxidized and which are being reduced in Equations 2 and 3?

CLOCK REACTION INTERACTIVE

http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/stoichiometry/iodine_clock.html

EXTRA NOTES

2.2 notes

EXTRA EXERCISES

2.2 exercise 1

2.2 Assessed homework (Includes 2.3 Equilibrium)

8.1 Rate of reaction

1.  explain and use the term rate of reaction, frequency of collisions, effective collisions and non-effective collisions
2. explain qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure changes on the rate of a reaction
3. use experimental data to calculate the rate of a reaction

8.2 Effect of temperature on reaction rates and the concept of activation energy

1. define activation energy, EA, as the minimum energy required for a collision to be effective
2. sketch and use the Boltzmann distribution to explain the significance of activation energy
3. explain qualitatively, in terms both of the Boltzmann distribution and of frequency of effective collisions, the
effect of temperature change on the rate of a reaction

8.3 Homogeneous and heterogeneous catalysts

1. explain and use the terms catalyst and catalysis

(a) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy
(b) explain this catalytic effect in terms of the Boltzmann distribution
(c) construct and interpret a reaction pathway diagram, for a reaction in the presence and absence of an effective catalyst