ICT for Innovative Science Teachersictforist.oeiizk.waw.pl/upload/Chemical Reactions Module_EN.pdf · Activity 1. Rate of reaction of magnesium with hydrochloric acid In this activity

Chemical Reactions

CC 2011 ICT for IST

This project has been funded with support from the European Commission under the Lifelong Learning

Programme. This publication reflects the views only of the author, and the Commission cannot be held

responsible for any use which may be made of the information contained therein.

ICT for Innovative Science Teachers Leonardo da Vinci programme

2009-1-PL1- LEO05- 05046

Chemical reactions are all around us:

cooking, burning, rusting, souring,

photosynthesis, respiration, starting

an engine, food rotting in the fridge

and many, many others.

Chemical reactions occur at different

rates. Iron rusting is a slow reaction

which can take many years.

The fermentation of sugar to alcohol is

quite slow but you can see the carbon

dioxide bubbles forming.

Combustion reactions e.g. when a fuel

burns in air or oxygen, is a very fast

reaction.

Explosive reactions would be described

as „very fast‟.

Chemical Equilibrium Module - 2

A. Introduction

1. Background theory

1. REACTION RATE

A chemical reaction is a process in which two or more substances (reactants)

change into one or more new substances (products).

The rate of a chemical reaction indicates how fast a reaction takes place. In

other words the reaction rate is the 'rate of formation of product' or the 'rate of

removal of reactant'.

For a one-step reaction BA reaction rate at which A transforms in B, is the

change in concentration of A with time, mathematically expressed as:

t

Arate

][ or

t

Brate

][

For a general reaction with stoichiometric coefficients (represented by lowercase

letters)

qQpPbBaA

The theme of this module is chemical reactions and a chemical equilibrium.

This topic requires integrated understanding of many areas of introductory

chemistry. Many students have considerable difficulties understanding

concepts and processes involved in this topic, especially the concept of

dynamic equilibrium. The following activities are available in this module:

1. Data Video: an activity to investigate and determine rate of reaction

by measuring a change of volume of a gas which is formed during

a reaction.

2. Data logging: Three laboratory experiments:

To investigate factors effecting rates of reactions.

To determine order of a chemical reaction.

To investigate a chemical equilibrium.

3. Modelling: activities which illustrate how to model non-reversible

and reversible chemical reactions.

Chemical Reactions Module - 3

reaction rate is:

dt

Qd

qdt

Pd

pdt

Bd

bdt

Ad

arate

][1][1][1][1

with negative sign for reactants and positive signs for products.

Reaction rate can be measured by using different experimental techniques:

by measuring a change of volume of a gas which is formed during a reaction,

by measuring a change of pressure of a gas which is formed during a reaction,

by measuring a change of colour during a reaction,

by measuring a change of conductivity during a reaction etc.

2. FACTORS AFFECTING RATE OF REACTION

Every chemical reaction proceeds at its own rate. Some reactions are naturally

fast and some are naturally slow under the same conditions. However, by

varying the conditions of the reaction, the rate of almost any reaction can be

modified. There are five factors which affect the rate of a reaction, according to

the collision theory of reacting particles.

Temperature: Usually conducting a reaction at a higher temperature delivers

more energy into the system and increases the reaction rate by causing more

collisions between particles, as explained by collision theory. The influence of

temperature is described by the Arrhenius equation.

Concentration: Reaction rate increases with concentration, as described by

the rate law and explained by collision theory.

Pressure: The rate of gaseous reactions increases with pressure, which is, in

fact, equivalent to an increase in concentration of the gas.

Surface Area: In reactions on surfaces, which take place for example during

heterogeneous catalysis, the rate of reaction increases as the surface area

increases. That is because more particles of the solid are exposed and can be

hit by reactant molecules.

Catalysts: The presence of a catalyst increases the reaction rate by providing

an alternative pathway with a lower activation energy.


3. RATE LAW

Experimentally it is found that reaction rate depends on the concentration of the

species involved in the reaction. The relation between the rate and these

concentrations can be expressed mathematically in the form of an equation

called a rate law.

For a reaction qQpPbBaA , the rate equation is of the form

ba BAkr ][][

In this equation k is the reaction rate coefficient or rate constant, although it is

not really a constant, it depends on several factors, such as temperature,

surface area, etc.

The exponents a and b are called the reaction orders and depend on the reaction

mechanism. They reaction orders must be determined experimentally.

For a zero-order reaction (red line),

the rate of reaction is constant as the

reaction progresses.

For a first-order reaction (green line),

the rate of reaction is directly

proportional to the concentration. As

the reactant is consumed during the

reaction, the concentration drops and

so does the rate of reaction.

For a second-order reaction (blue line),

the rate of reaction increases with the

square of the concentration, producing

an upward curving line in the rate-

concentration plot. For this type of

reaction, the rate of reaction

decreases rapidly (faster than linearly)

as the concentration of the reactant decreases.

4. TEMPERATURE DEPENDENCE

Rate constants, and hence reaction rates, are often found to dependent strongly

on temperature. It is therefore important to quote the temperature at which any

rate constant is determined. Most commonly the rate goes up with temperature,

but this is not always the case.

Experimentally, a very large number of rate constants, are found to vary with

temperature according the Arrhenius equation:

RT

Ea

Aek


Ea is the activation energy and R is the gas constant. The values for A and Ea are

dependent on the reaction.

The equation can be manipulated into a straight line by taking natural logarithms

of both sides:

TR

EAk a 1)ln()ln(

So a plot of ln(k) against 1/T should be a straight line, such a plot is called

Arrhenius plot. The slope of the plot is -Ea/R and the intercept with the vertical

axis, when 1/t goes to zero, is ln(A).

5. REVERSIBLE REACTIONS AND DYNAMICAL EQUILIBRIUM

In reversible reactions, the reactions occur simultaneously in both directions.

The forward and backwards reactions continue up to a moment of chemical

equilibrium in which the forward and backward rates are equal. As a result, there

is no change in the concentration of any of the species, even though the

reactions are still going on.

For a reversible reaction involving two reactants and two products

aA + bB cC + dD

at equilibrium the rate of forward reaction is equal to the rate of backward

reaction

dc

b

ba

f DCkBAk ][][][][

the equilibrium constant is b

f

ba

dc

eqk

k

BA

DCK

][][

][][.

This constant provide valuable chemical information, if Keq >1 this means that

products are favoured over reactants, and if Keq < 1 this means that reactants

are favoured over products.

2. Pre-requisite

knowledge required

Chemical calculations using

the mole

Definition of concentration

3. Science concepts

developed in the

module

Reaction rate

Rate law

Rate constant

Chemical equilibrium


4. Other useful

information

Helping you to understand Chemistry

http://www.chemguide.co.uk

It is possible to find interactive

activities on the Internet about

modelling chemical reactions. Just

a few examples:

Connecting Kinetics and Equilibrium:

Reversible First-Order Reactions in

http://www.bpreid.com/applets/re

versible.html

Chemical Reaction Simulations in

http://www.science.uwaterloo.ca/

~cchieh/cact/trios/simulation.html

#reaction

Chemical Kinetics Simulation in

http://www.chem.uci.edu/undergr

ad/applets/sim/simulation.htm

B. Didactical approach

1. Pedagogical context

The activities here concern basic concepts related to chemical reaction rate and

chemical equilibrium. Gradient of sophistication is offered here, for the lower

secondary level via a qualitative approach, and for the higher secondary level via

quantitative approach and via modelling.

2. Common student difficulties

Students have difficulties in:

Interpreting graphs with time as

the independent variable plotted

on the horizontal axis

Assuming that equilibrium means

“the end of the reaction” [1]

Associate a high k value with a

very fast reaction [1]

Understanding that in equilibrium

two reactions occur at the same

time [2]

Assuming that the concentrations

of all reactants are equal at

equilibrium [3]

Assuming that all rate laws are

expressed using stoichiometric

coefficients

REFERENCES:

1. Vanessa Kind, Beyond Appearances: Students’ misconceptions about basic

chemical ideas available via

http://www.chemsoc.org/networks/learnnet/miscon.htm


2. Van Driel, J.H., (1990), “Betrokken bij Evenwicht”, Utrecht: CD Press,

dissertation didactics of chemistry education

3. Barke, H.D., Hazari, A., Yitbarek, S., (2009), “Misconceptions in

Chemistry”, Berlin Heidelberg: Springer

4. http://www. daisley.net/hellevator/misconceptions/misconceptions.pdf

3. Evaluation of ICT

The specific qualities of ICT which benefit student learning are specified below

per a type of activity.

VIDEO MEASUREMENT

Activity 1. Rate of reaction of magnesium with hydrochloric acid

In this activity the reaction rate is measured by measuring a volume of hydrogen

gas, which is formed during a reaction of magnesium with hydrochloric acid.

When a gas is formed from a solid reacting with a solution, it is collected in a

measuring cylinder filled with water. The water level in the cylinder indicates the

volume of the collected gas. A video recording is a useful method of measuring

changes in the water level as a function of time. To better visualize the water

level, lightweight brown balls are floated on the water surface in the cylinder.

DATA LOGGING

Activity 2. Factors affecting rates of reaction

In this activity students investigate several factors that affect the rate of the

reaction of magnesium metal with hydrochloric acid.

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


In this reaction, hydrogen gas is produced as a product thus pressure during the

reaction increases if the reaction container is a closed system. Pressure sensor is

used to measure the change in pressure.

This activity consists of three parts:

I - determining how changing the temperature of the reactants influence the rate

of reaction;

II – determining how changing the particle

size of one of the reactants changes the rate

of reaction;

III – determining how changing the

concentration of one of the reactants

changes the rate of reaction.

An advantage of this experiment is that the

rate of reaction is indicated immediately as a

gradient on the graph line. Only small

quantities of chemicals are needed and

results are obtained so rapidly, the

experiment is easily repeated with different

conditions.

SAFETY NOTE: The rubber bung should not

be pushed too hard into the neck of the test

tube, but should be allowed to pop out when

the pressure exceeds a safe level. Typically

this happens at less than 1.5 atmospheres.

Activity 3. Rate and order of reaction

When sodium thiosulphate reacts with a hydrochloric acid, a yellow precipitate of

sulphur is formed:

Na2S2O3 (aq) + 2HCl(aq) 2NaCl(aq) + SO2 (aq) + H2O(l) + S(s)

sodium thiosulphate + hydrochloric acid sodium chloride + sulphur dioxide + water + sulphur

In this activity the rate of this reaction is measured by using a light gate to

detect the precipitate formation. The system consists of a light beam source and

a light sensor. The reaction vessel (e.g. a small cuvette) is placed between the

light emitter and sensor. The light reading falls as the sulphur precipitate forms

and the time taken to form a certain amount of sulphur is measured.


Left: Cuvette filled with a clear solution in front of the light sensor.

Right: A precipitate of sulphur forms when the chemicals mix.

The rate of reaction for each experiment run can be expressed as

amount of formed sulphur/time

Since the absolute mass of sulphur formed is not known, the reciprocal of the

time is taken as a measure of the relative rate of reaction.

According to the rate law the rate equation for this reaction is:

r = k[ HCL] m [ thiosulphate] n

where m is the reaction order with respect to HCL, and

n is the reaction order with respect to sodium thiosulphate.

The values of the reaction orders are determined experimentally, first for sodium

thiosulphate. The experiment, in which the reaction time is determined, is

repeated with other concentrations of sodium thiosulphate.

Assuming that in each of these experiments the concentration of acid is constant

and the concentration of thiosulphate varies, the rate equation becomes:

nate][thiosulphconstr

Taking logarithms of both sides of the rate equation gives:

ate])[thiosulphln(nconst)ln(lnr

When ln(r) is plotted against ln([thiosulphate]) the result is a straight line

whatever the order of the reaction. The gradient of this line, which is n, gives

the order of the reaction with respect to thiosulphate.

In similar way the order of reaction with respect to acid can be found.

This method of finding the reaction order is called the differential method.


Activity 4. Dissolving limestone

In water a chemical equilibrium between limestone, carbon dioxide and dissolved

calcium hydrogen carbonate becomes established:

CaCO3(s) + H2O(l) + CO2(g) Ca2+

(aq) + 2 HCO3-

This equilibrium is of major geological importance. A lot of Karst landscapes form

by dissolving limestone in carbon dioxide enriched water. Coral reefs develop by

precipitating limestone.

The solution and precipitation of limestone can be studied with a conductivity

sensor. A calcium carbonate suspension in water and a carbon dioxide solution

both have a quite low conductivity.

The conductivity increases immediately after adding calcium carbonate to

mineral water or bubbling carbon dioxide gas through a calcium carbonate

suspension. If you bubble air through the solution it removes carbon dioxide and

the equilibrium shifts to the left. The conductivity decreases immediately.

MODELLING

The Modelling activities in this module illustrate how to model chemical reactions.

There are many types of chemical reactions and these activities show only a few

examples. The examples use two software systems, Coach 6 and Simulation

Insight, which illustrate two different methods of expressing models. For both

systems, ready-made models are available and students can use them as

simulations which offer the possibility of experimenting with different parameters

for investigating different conditions.

Coach 6 offers an improved graphical modelling approach specially adapted for

chemical reactions. This approach is based on kinetic graphs offers much clearer

visual representations of chemical reactions:

A „Process element‟ , used in the context of chemical kinetics, stands for the

reaction rate of a stage in a chemical reaction, which depends on the reaction

rate coefficient, the concentrations of reactants involved in the reaction step and

their stoichiometric coefficients. The In and Out coefficients (defined in the

Process properties) represent stoichiometric coefficients of the reaction stage

and determine the exact relationship between the reactants and products of the

reaction (ingoing and outgoing flows).

Activity 5. Modelling non-reversible reactions

The activities here illustrate how to model non-reversible reactions.

A. Non-reversible reaction A to B

The model given in this activity describes a simple non-reversible chemical

reaction A B. For simplicity it is assumed that the reaction is of a first order. The rate law is then r = k*[A] where k is reaction rate coefficient.


The Process element used in the model stands for the reaction rate r.

Models of reaction A B

The resulting graphs show the change of concentration of species A and B.

In the extra assignment the model has to be extended to describe two

successive reactions A B C.

B. Non-reversible reaction A B + C

C. Non-reversible reactions A + B C

These two activities illustrate how to model more complex non-reversible

reactions.

Students analyse the model and investigate the effects of changing initial

concentrations of the A, B and C species and effects of changing the reaction

rate.

As assignment in activity B students have to modify the given model and create

the model of a crack reaction in which propane is cracked and forms ethylene

and methane: C3H8 (g) C2H4 (g) + CH4 (g).

Modifying the model means changing variable names and adjusting their initial

values. The reaction rate is then automatically adjusted to r=k*[C3H8] (first order

reaction). Adjusting of In and Out coefficients of the process element is not

needed since all stoichiometric coefficients of this reaction have value 1.


Model of reaction C3H8 (g) C2H4 (g) + CH4 (g)

As assignment in activity C students have to modify the given model and create

the model of the gas-phase oxidation of nitric oxide: 2NO + O2 2NO2.

Important step here is not only adapting the variable names but also adapting:

the reaction rate to:

r=k*[NO]2*[O2] (because this is a first order reaction with respect to O2, and

a second order reaction with respect to NO], and

In and Out coefficients to stoichiometric reaction coefficients:

Coefficient [NO]=2, Coefficient [O2]=1, Coefficient [NO2]=2.

The resulting graphs of concentrations changes during reaction of 2NO + O2 2NO2

Activity 6. Modelling reversible reactions

The activities here illustrate how to model reversible reactions and introduce the

concept of dynamical equilibrium.

A reversible chemical reaction can be considered as two simultaneous reactions,

each one with its rate law and its rate constants, k1 and k2.

After some time the 'system settles down' and the net concentrations of the

reactants and products remain constant i.e. a state of concentration balance

exists.


A. Reversible reactions A B

The model in this activity describes a simple reversible reaction A B. In order

to simplify the models the reaction first order rate laws.

As assignment in this activity students have to modify the given model and

create the model of reaction N2O4 2 NO2 in which, dinitrogen tetroxide N2O4, a

colourless gas, and nitrogen dioxide NO2 a dark brown gas, exists in equilibrium

with each other.

B. Modelling reversible reactions A B + C

C. Modelling reversible reactions A + B C

These two activities illustrate how to model more complex reversible reactions.

The examples use two products in the first example and two reactants in the

second example. In order to simplify the models, all reactions have first order

rate laws.

The student is asked to:

change the rate constants and discuss how the change affects the time taken

to reach the equilibrium;

choose different initial conditions and discuss what happens, particularly

when, in the beginning, only one of the two reactants is present;

analyse specific cases of initial conditions and values for the rate constants;

predict what will happen for different initial conditions and relative values of

rate constants.

Activity 7. Crystal violet

When a solution of crystal violet (violet-coloured) reacts with sodium hydroxide

it forms a colourless compound (CVOH). A simplified version of the equation is:

CV+ + OH- CVOH

In this activity students compare the results of the given model with the results

of measurement (the activity 7. Crystal violet reaction – measurement).

The model assumes

that the reaction rate

r is proportional to

concentration of

crystal violet.

r = k*[CV]

Students may verify

this assumption by

comparison with the

measurement results.

If crystal violet is not available, phenolphthalein may be used as an alternative.


4. Teaching approaches

The three activities presented here

offer distinctive but complementary

insights into the science involved in

this topic. The Modelling activities

offer graphical models describing

process of chemical equilibriums.

These models are proposed to be

used as simulations.

For the activities to be effective for

teaching and learning, it is helpful for

teachers to consider two types of

skills in using the software tools:

Operational skills which concern

the knowledge of the features in

the software and manipulation of

the computer software.

Procedural skills which concern

the manner in which the software

tools are employed in the lesson

context for the purpose of

achieving learning benefits. A

dominant aspect of these skills is

the development of an inquiring

approach to the analysis and

interpretation of data and to

making links with previous

knowledge.

Such skills are important for the

preparation of pupils for the activities,

and the activity sheets below each

contain indications of the skills

needed for the particular activity.

For the teacher, there are further

pedagogical skills which contribute

to the effectiveness of the activities:

1. Clarity of learning objectives for

each activity.

2. Understanding of the special value

of the ICT method and exploiting

its full potential in purposeful

ways.

3. To manage the activity in a way

which promotes „appropriate‟

rather than „indiscriminate‟ use of

ICT.

4. To integrate the learning from

each activity to develop pupils‟

understanding of the topic.

The development of the last of these

is a particular aim of the ICT for IST

Project, and the activities presented

have been specially selected to

illustrate how integration might be

achieved. Comparisons of the

observations and results of each

activity form a central role in this

integration process. For example

comparing data from the model with

experimental data as well from data-

logging experiment as from video

measurement.

In these, the graph is a key tool in

facilitating comparisons and

interpretations and skills with graphs

generally provide a common thread in

exploiting ICT for IST activities.

The management of the classroom

setting also has an important

influence on the successful

integration of activities. When access

to computer equipment is scarce it is

likely that the teacher will wish to

present the activity as a

demonstration in a didactic manner.

In this mode, the teacher can give

strong guidance to pupils‟ thinking

about the comparisons between the

activities.

Alternatively, pupils could perform

the activities in small groups of three

or four pupils, each group engaged

on a different activity.

Integration might be achieved by

each group making a presentation of


their results to the whole class. In

chairing these presentations the

teacher can prompt discussion of the

significant findings of each group.

It is worth considering that all the

activities may be used in a variety of

learning contexts.

Although the activities have been

designed to provide complementary

experiences, it is not essential to use

all of them; two or three activities

might be chosen according to how

well they suit the needs of teachers

and pupils in a particular context. In

varying conditions between schools

and within schools at different times

of the year or different stages in the

curriculum, needs and

appropriateness are likely to change;

for example, an individual pupil might

need a revision or extension activity,

an enrichment activity might be

required to occupy some spare time,

a quick activity might be needed if

time is scarce. The overlapping

features, such as graphical

presentation, between the activities

allows them to be used to a certain

extent as alternatives, but their

distinctive features also allow them to

be used as complements to each

other.

The table on the next page

summarises the distinctive potential

learning benefits of each. It is a

useful guide to the special value of

each ICT activity.

Activity Potential learning benefits, ‘ICT value’

Data logging Graph of measured by sensor quantity (light intensity,

pressure) versus time is displayed during experiment.

Changes are observable immediately.

Analysis tools facilitate further investigations.

Data video Allows analysis of the water level in a measuring cylinder.

A graph of gas volume versus time is created during the

measurement.

Modelling The models illustrate the chemical reactions and „visualize‟ the

changes in equilibrium.

Models are used as simulations.

The effect of altering a parameter such as a reaction rate or

reactants/products concentrations can be investigated.

The model data can be compared with experimental data.


5. Resources for Student Activities

USING COACH 6 SOFTWARE

Activity Software

program

Files available in Coach 6 Project

Chemical Equilibrium

1. Data Video Coach 6 1. Rate of reaction of magnesium with hydrochloric

acid.cma (activity file)

1. Rate of reaction of magnesium with hydrochloric

acid.cmr (result file)

2. Data logging Coach 6 2. Factors affecting rates of reactions.cma (activity

file)

2A. Reaction - effect of temperature.cmr (result file)

2B. Reaction – effect of particle size.cmr (result file)

2C. Reaction - effect of concentration.cmr (result file)

3. Data logging Coach 6 3. Rate and order of reaction.cma (activity file)

3. Rate and order of reaction.cmr (result file)

4. Data logging Coach 6 4. Dissolving limestone.cma (activity file)

4. Dissolving limestone.cmr (result file)

5. Modelling Coach 6 5A.Non-reversible reaction A to B.cma (activity file)

5A.Non-reversible reaction A to B.cmr (result file)

5A.Non-reversible successive reactions A to B to C.cmr

(result file)

5B.Non-reversible reaction A is B + C.cma (activity

file)

5B.Non-reversible reaction C3H8 is C2H4 + CH4.cmr

(result file)

5C.Non-reversible reaction A + B is C.cma (activity

file)

5C.Non-reversible reaction 2NO + O2 is 2NO2.cmr

(result file)

6. Modelling Coach 6 6A.Reversible reactions A to B.cma (activity file)

6A.Reversible reactions N2O4_N2O.cmr (result file)

6B.Reversible reactions A is B + C.cma (activity file)

6C.Reversible reactions A + B is C.cma (activity file)

7. Modelling Coach 6 7. Crystal violet.cma (activity file)

7. Crystal violet - measurement.cmr (result file)


USING INSIGHT SOFTWARE

Activity Software

program

Files available

1. Data Video Insight iLOG 1. Rate of reaction of magnesium with hydrochloric

acid (sample data)

2. Data logging Insight iLOG 2. Factors affecting rates of reactions set-up

2A. Reaction - effect of temperature (sample data)

2B. Reaction – effect of particle size (sample data)

2C. Reaction - effect of concentration (sample data)

3. Data logging Insight iLOG 3. Rate and order of reaction set-up

3. Rate and order of reaction (sample data)

4. Data logging Insight iLOG 4. Dissolving limestone set-up

4. Dissolving limestone (sample data)

5. Modelling Simulation

Insight

5A. Reaction - non-reversible

5A+. Reaction - non-reversible A-B-C

5B. Reaction - non-reversible 2 PRODUCTS

5B+. Reaction - non-reversible 2 PRODUCTS - C3H8

5C. Reaction - non-reversible 2 REACTANTS

5C+. Reaction - non-reversible 2 REACTANTS - 2NO

+O2


Insight

6A. Reaction – reversible

6B. Reaction - reversible 2 PRODUCTS

6C. Reaction - reversible 2 REACTANTS


Insight

7. Reaction – crystal violet

7. Crystal violet experiment (sample data)


EQUIPMENT AND MATERIALS FOR DATA-LOGGING ACTIVITIES

Computer

Software – See table above

Interface (data-logger)

Light sensor, pressure sensor, conductivity sensor

Light source

Disposable cuvettes, syringes

test tubes with stoppers (one hole), hot plate, ice

stand, magnetic stirrer, aquarium air pump, a measuring cylinder

chemicals (see specifications in activities)


C. Student Activities

ACTIVITY 1. RATE OF REACTION OF MAGNESIUM

WITH HYDROCHLORIC ACID

Learning Objectives:

1. To investigate rate of reaction by using

video measurement.

Operational Skills:

Making measurements on the Video Screen

Using the Scan, Slope and Function fit options.

Procedural Skills:

Analysing data using a graph

Reading values/slopes

Evaluating measurement quality

Materials:

Video clip showing a reaction of magnesium with hydrochloric acid.

Activity method (using Coach 6):

1. In the Data-Video Window you see a video clip showing a following

experiment. In a conical flask 0.08 g magnesium reacts with dilute

hydrochloric acid of concentration 0.13 mol/dm3. During the reaction

hydrogen gas is produced. The flask is connected to an inverted measuring

cylinder in a trough of water and the produced gas is collected in the

measuring cylinder.

The initial volume of HCl is 50 cm3 and its temperature is 25oC.

2. In this video you are going to determine the volume of produced hydrogen

gas by measuring the water level in the cylinder. To indicate the water

level lightweight brown balls are used. The video is already scaled and

prepared for measurement.

3. Measure the volume of the collected gas. Plot a graph of V(H2) against

time.

APPLIED ICT TECHNOLOGY:

DATA LOGGING

STUDENT LEVEL:

AGE 15-19

RECOMMENDED SETTINGS:

STUDENT ACTIVITY IF

ENOUGH EQUIPMENT IS

AVAILABLE, OTHERWISE

TEACHER DEMONSTRATION


Questions and Assignments:

What does the graph tell you about the progress of reaction?

Determine how the reaction rate is changing in time.

What do you think what factors and how may affect the rate of reaction?

During the reaction of Mg and H+ the number of [H+] decreases. Based on

the rate of reaction and measured volume determine the change of

concentration [H+] during the reaction.

Analysing activities:

Magnesium reacts with dilute hydrochloric acid in a conical flask which is

connected to an inverted measuring cylinder in a trough of water. Students

follow the rate of reaction by measuring the volume of the produced gas.

The slope of the graph V(H2) is steepest at the beginning, this shows that the

reaction is fastest at the start. As the magnesium is used up, the rate falls, the

slope becomes less steep and then levels out when the reaction has stopped

(when no more gas is produced).

The reaction is exothermic, but the dilute acid is in excess and the rise in

temperature is not high. There is some acceleration of the reaction rate due to

the rise in temperature. Some students might notice the flask becoming

slightly warm and they could be asked how this would affect the rate of

reaction, and how they might adapt the experiment to make it a „fair test‟.

Change of

hydrogen gas

volume during

the reaction of

magnesium with

hydrochloric acid.

Additional, based on the rate of reaction and measured volume students can

determine the change of concentration [H+] during the reaction and compare

it with a model. Model from the activity 5A.Non-reversible reaction A to B

should be first modified to describe the reaction. Then the measured data can

be imported as background graph and comparison with the model data can be

performed.


ACTIVITY 2. FACTORS AFFECTING RATES OF

REACTION


1. To investigate several factors that affect

the rate of a chemical reaction.

2. To determine the effects of concentration, surface

area and temperature on the rate of a reaction.

Operational Skills:

Connecting sensors and interfaces

Choosing logging parameters

Starting and finishing real-time logging

Using the cursor tools for obtaining measurements from the graph

Changing the designation of the graph axes

Deriving secondary data by calculation

Procedural Skills:


Analysing data using graph


Materials:


Pressure sensor with connecting tube

400 ml beaker

10 ml plastic syringe

test tube (borosilicate glass - Pyrex) and test tube rack

one-hole rubber bung to fit the test tube

boiling water from kettle (hot plate), crushed ice

cleaned magnesium metal strips (0.5 by 2 cm) and granular magnesium

(2g)

1.0 M hydrochloric acid (HCl)

Weighing balance


DATA LOGGING

STUDENT LEVEL:

AGE 15-19


STUDENT ACTIVITY IF

ENOUGH EQUIPMENT IS




Activity method:

I. Study the effect of temperature on the rate of reaction

1. Prepare a clean, dry test tube (of strong borosilicate glass) with a good

fitting rubber bung containing one hole. Insert a polythene tube into the

bung and use this to connect the pressure sensor to the test tube.

2. Connect the pressure sensor to the data logger interface which is already

connected to a computer.

3. Remove the bung, measure 5.0 ml of 1.0 M hydrochloric acid and add it

to the test tube.

4. Start the program logging, drop one strip of magnesium ribbon in the

test tube and replace the bung quickly. Hydrogen gas is evolved inside

the tube and its pressure is shown on the graph. At a certain pressure,

the bung will pop out of the tube and the pressure reduces to

atmospheric pressure.

5. Save the data for later analysis.

6. Empty the contents of the test tube into the sink, wash the tube with

water and dry the inside with an absorbent cloth.

7. Place the tube in a beaker containing crushed ice and filled three-

quarters full with cold water.

8. Repeat the experiment, adding acid and dropping a strip of magnesium

in as before. Observe the different graph and save the data.

9. Empty, wash and dry the tube again and place it in a beaker containing


recently boiled water. (Take care in handling the hot water beaker.)

10. Repeat the experiment, adding acid and dropping a strip of

magnesium in as before. Observe the different graph and save the data.

II. Study the effect of particle size of one of reactants on the rate of

reaction

1. Prepare the test tube, pressure sensor, data-logger and computer as

before.

2. Measure 5.0 ml of 1.0 M hydrochloric acid into the test tube and place the

tube in the rack.

3. Cut and weigh a cleaned magnesium strip then weigh out an equal

portion of granular magnesium.

Note: If you do not have granular magnesium, use a second, the same

weight, magnesium strip and cut it into as many small pieces as possible.

4. Start the program logging, drop the magnesium in the test tube and

replace the bung quickly.

5. Observe the new graph and save the data for later analysis.

III. Study the effect of changing the concentration of one of

reactants on the rate of reaction

1. Repeat the first experiment using only 2.5 ml of acid and adding 2.5 ml of

water to reduce the concentration.

2. Observe the new graph and save the data for analysis.



I. Study the effect of temperature on the rate of reaction

How and why would the rate of a chemical reaction change as

temperature changes?

Compare your pressure versus time graphs. At which temperature did

the reaction occur fastest? How do you know?

For each graph, measure the rate of reaction using the analysis tools in

the program. Do your results support your predictions?

II. Study the effect of particle size of one of reactants on the rate of

reaction

How and why would the rate of a chemical reaction change as the

particle size of one of the reactants changes?

Compare your pressure versus time graphs. At which particle size did the

reaction occur fastest? How do you know?



III. Study the effect of changing the concentration of one of

reactants on the rate of reaction

How and why would the rate of a chemical reaction change as the

concentration of one of reactants changes?

Examine your pressure versus time graph. At which concentration of acid

did the reaction occur fastest? How do you know?




In this activity students investigate several factors that affect the rate of a

simple chemical reaction.

The reaction of magnesium metal with hydrochloric acid is observed.

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

In this reaction, hydrogen gas is produced as a product thus pressure will

increase if the reaction container is a closed system. Pressure sensor is used

to measure the change in pressure during the reaction.

This activity consists of three parts:

I - determining how changing the temperature of the reactants influence the

rate of reaction,

II – determining how changing the particle size of one of the reactants



III – determining how changing the concentration of one of the reactants


Safety

Remind students to:

follow standard laboratory safety procedures,

be careful when handling chemicals,

wear protective glasses and aprons,

add acid to water when making dilutions,

rinse any spills with copious amounts of water (have sodium bicarbonate

available for neutralization),

not use open flames due to the flammability of hydrogen gas,

not push the rubber bung too hard into the neck of the test tube. A rise

in pressure is only needed for a few seconds to demonstrate the rate of

reactions.


ACTIVITY 3. RATE AND ORDER OF REACTION


1. To introduce concept of the reaction rate.

2. To study the reaction rate during the reaction

between sodium thiosulphate and hydrochloric acid.

3. To determine the order of a reaction.

Operational Skills:







Procedural Skills:


Analysing data using a graph


Materials:


Light source

Light sensor

Disposable cuvettes

Disposable syringes

Solutions:

10 ml of hydrochloric acid HCl (0.1 M),

10 ml of sodium thiosulphate Na2S2O3 (0.1 M).


DATA LOGGING

STUDENT LEVEL:

AGE 15-19


STUDENT ACTIVITY IF

ENOUGH EQUIPMENT IS




Activity method:

1. Use a disposable cuvette (5.0ml) to mix the solutions. Solutions will be added

together with the aid of syringes.

2. Position the cuvette as close as possible to the light sensor. It is not

necessary to use a special cuvette-holder but this would simplify the

experiment. You can also use a Colorimeter sensor.

3. Connect the light sensor to an interface.

4. Before starting the first measurement you must find a suitable distance

between the light source and the sensor. Place a cuvette filled with a clear

solution (water) in front of the light sensor and move the light source until

the signal level is about 80%. When you work with a Colorimeter sensor

choose a red light and zero its reading with a cuvette of a clear solution

(water).

5. Mix 2.5ml thiosulphate-solution and 2.5ml hydrochloric acid in the cuvette.

Start recording immediately after adding the solutions.

6. Repeat the experiment with other concentrations of sodium thiosulphate.


How does the appearance of the solution change during the reaction?

What does the graph tell you about the progress of reaction?

Determine the reaction rate in which certain amount of sulphur has been

formed, in other words a certain level of turbidity is reached.

(Use '1/time' as a measure for the reaction rate, do you know why?)

Repeat the experiment with other concentrations of sodium thiosulphate.

For each reaction determine the reaction rate.

How does the reaction rate depend on the sodium thiosulphate

concentration?

Determine the order of the reaction with respect to sodium thiosulphate.

How would you find the order of reaction with respect to hydrochloric acid?


Sodium thiosulphate and hydrochloric acid react to form a precipitate.

Na2S2O3 + 2 HCl SO2(g) + S(s) + 2NaCl + H2O

To determine the reaction rate the formation of sulphur during the reaction is

used.


A special cuvette-holder with a built-in LED light source and a place for the light sensor.

Left: Cuvette filled with a clear solution in front of the light sensor.

Right: A precipitate of sulphur forms when the chemicals mix.

The reaction time is equal to the time it takes to reach a certain level of

turbidity (a certain amount of sulphur is formed).

So therefore reaction rate = amount of sulphur formed/time

Since the amount of formed sulphur is constant for each run of the experiment

then reaction rate = constant/time which means that the reciprocal of time

(1/time) may be used as a comparative measure of the rate.

For determining the order of reaction „n‟ with respect to sodium thiosulphate,

reaction rates for different thiosulphate concentrations are determined.

The rate equation is expressed as:

r =k[HCL]m[thiosulphate]n

Assuming that concentration of HCL does not change:

nate][thiosulphconstr

Taking logarithms of both sides of the rate equation gives:

ln(r) = ln(const)+n*ln([thiosulphate])

Plotting ln(r) against ln([thiosulphate]) results in a straight line which gradient

is the order of the reaction in respect to sodium thiosulphate.

If temperature isn't a variable, it must be kept constant. The simplest solution

here is to make sure all the chemicals have been standing in the laboratory

prior to the lesson.

Safety

Be careful when handling chemicals.


ACTIVITY 4. DISSOLVING LIMESTONE


1. Making visible the establishment of chemical

equilibrium.

2. Studying the effect on equilibrium by removing

one of the compounds.

3. Studying the reaction rate of the formation and

decomposition of CaHCO3.

Operational Skills:







Procedural Skills:


Analysing data using graph


Materials:


Conductivity sensor

500ml flask

calcium carbonate (powder)

carbonated mineral water

Stand, magnetic stirrer and aquarium air pump

Activity method:

1. Connect the conductivity sensor to the input 1 of an interface.

2. Switch the conductivity sensor to 0 - 2000 μS range.


DATA LOGGING

STUDENT LEVEL:

AGE 15-19


STUDENT ACTIVITY IF

ENOUGH EQUIPMENT IS




3. Measure the conductivity of distilled water, distilled water with some

calcium carbonate powder and carbonated mineral water.

4. Pour 250ml of carbonated mineral water in a 500ml flask.

5. Place the conductivity probe in the solution.

6. Start stirring and start a measurement. After about two minutes add

some calcium carbonate powder (teaspoon full) to the mineral water.

7. To remove carbon dioxide produced during the reaction after about 10

minutes, start bubbling air through the solution.

Questions and Assignments

Explain the rise in the conductivity after adding a calcium carbonate to

mineral water.

Explain the decrease in conductivity while bubbling air through the

solution.

Perform the experiment using other sources of calcium carbonate.

The logarithm of the change in conductivity decreases almost linear.

Investigate this linearity using several sources of limestone.


The conductivity increases immediately after

adding calcium carbonate to mineral water or

bubbling carbon dioxide gas through a calcium

carbonate suspension.

If air bubbles are pumped through the solution

it removes carbon dioxide and the equilibrium

shifts to the left. The conductivity decreases

immediately.

Example data, rise in conductivity after adding

calcium carbonate powder (1.00 g CaCO3) to 200ml mineral water (Sourcy)


ACTIVITY 5. MODELLING NON-REVERSIBLE

REACTIONS


1. To understand a model of a non-reversible reaction.

2. To use the model to obtain the graphs of concentrations

changes during the reaction.

3. To investigate the effect of changing initial

concentrations of reactant and product and the effect of

changing the reaction rate constant.

Operational Skills:

Manipulating model variables

Using software controls for running simulation

Modifying given model

Procedural Skills:

Analysing data using graphs

Evaluating model quality

Using models and simulations for investigations

Activities (using Simulation Insight or Insight iLog)


1. Load file 5A, run the model and compare the changes in the

concentrations of A and B as shown by the graphs.

2. Assuming that at a higher temperature the reaction goes faster, k will be

larger. Alter the model to find out how this affects the graphs.

3. Predict what happens if at t = 0 the concentrations of A and B are equal.

Check your prediction by setting the initial values: A = B = 1.


MODELLING

STUDENT LEVEL:

AGE 17


STUDENT ACTIVITY



1. Load file 5B, run the model and compare the changes in the concentrations

of A and B as shown by the graphs.



3. Predict what happens if at t = 0 the concentrations of A and B are equal.

Check your prediction by setting the initial values: A = B = 1.

C. Non-reversible reaction A + B C

1. Load file 5C, run the model and compare the changes in the concentrations

of A, B and C as shown by the graphs.

2. Compare this model with those in activities 5A and 5B and describe the

differences

3. Investigate the effect of changing the initial concentrations of A and B.



5. Extension assignment: Modify the model to represent the gas-phase oxidation of nitric oxide: 2 NO

+ O2 --> 2 NO2 (Since TWO molecules of NO react with ONE molecule of O2, the rate law

becomes r = k * [NO]^2 * [O2]).

Questions/Assignments using Coach 6:


1. Assume that at time=0 there is only reactant A, ([A]=1.0 M and [B]=

0.0 M) and A will be completely transformed into B. Right click a

respective variable and define its initial value. Coefficient k is already

defined as k= 0.4.

2. Execute your model.

3. Describe the resulting graphs.

4. The diagram shows you the decrease of [A]. Predict the change of [B].

Right click the Concentration [A] and [B] diagram and select Sketch.

Draw your prediction of [B]. Press <Esc> to finish your prediction.

Explain your prediction.

5. Display a graph of the change of [B]:

Right-click the diagram and select the Create/Edit diagram option. Click

on C3 and choose Axis, Second vertical. Click OK.

6. Does the curve agree with your prediction?

Note for this reaction: Decrease of [A] = Increase of [B]


Explain how you can "reflect" this note in the model.

7. Assume that the reaction will be faster at a higher temperature. Which

initial value needs to be changed to simulate a reaction at a higher

temperature? Explain.

8. Change the initial value you decided that needs to be changed to

simulate a reaction at a higher temperature. How do the curves show

you that the reaction takes place faster?

9. Simulate also a reaction at a lower temperature. How do the curves show

you that the reaction takes place slower?

10. After some time [A] will be half the initial value. This time is half-life time

of [A] for this reaction. Determine in the diagram the half-life time of

[A].

11. Simulate the model using different initial values of [A]. Does the half-life

time of [A] also depends on [A]?

12. What happens if at t = 0 the concentrations of A and B are equal (e.g., A

= B = 1)? Sketch your predictions of concentrations A and B as functions

of time? Check you conclusions with the model. Extra Assignment

1. Modify your model to create a model of two successive reactions A B

C, species A reacts to give B, and B reacts to give C. Assume that the concentration of species A is decreasing at a rate r ...

... and the concentration of species B is increasing at a rate r (transforming from A) and decreasing at a rate r1 (forming C)...

... and the concentration of species C is increasing at a rate r1 (transforming from B).

The second reaction, B C, follows the rate law r1 = k1*[B]. 2. Assume an initial situation in which there is only reactant A, and A will be

completely transformed into B and B into C.

3. Display the concentration [C] on the 'Concentrations' diagram.

4. Display the rate of reaction B C on the 'Rate of reaction' diagram.

5. Predict the graphs of concentration change of [A], [B] and [C]. Explain

your reasoning.

6. Check your prediction by executing your model. Were your predictions

correct? If no, do you know why?

7. What happens if at t = 0 there is any amount of the species A and B but no

C? Discuss, explain your reasoning, predict and check.

8. What happens if at t = 0 there is any amount of the species B but no A and

C? Discuss, explain your reasoning, predict and check.

9. What happens if at t = 0 there is any amount of the species A and reaction

1 occurs at a “high” rate and reaction 2 at a “low” rate (e.g., k= 1.0 and k1

= 0.1)? Discuss, explain your reasoning, predict and check.



1. Compare this model with the model from the activity 5A. Describe the

differences.

2. The progress of the reaction in time is determined by the reaction rate.

Analyse the given model. Where in the model this rate is defined? How the

reaction rate is defined?

3. It the model is assumed that the initial concentration of the reactant A is

1 M and both products 0 M. Predict and sketch the change of concentration

of [A], [B] and [C] during the reaction.

4. Check your predictions by executing the model.

5. Investigate the effect of changing the initial concentrations of A, B and C.

Use the control sliders in the Animation window. In what way does the

change of initial concentrations influence the reaction rate?

6. Investigate the effect of changing k on the reaction and reaction rate. Use

the control spinner. What happens when you set higher values of k? And

lower values of k?

7. Modify your model to create a model of a crack reaction in which propane

is cracked and forms ethylene and methane

C3H8 (g) C2H4 (g) + CH4 (g)

Following tips will help you to modify existing and create a new model.

During the reaction propane is broken down into two products, ethylene

and methane, and its initial concentration decreases.

Assume that the variable [A] becomes [C3H8].

Assume that the variable [B] becomes [C2H4].

Assume that the variable [C] becomes [CH4].

Set the initial concentration of propane to 0.5 M/L and the initial

concentration of the products to 0 M/L.

Set the k value 3.28 1/s.

Check the reaction rate (the rate should be r = k * [C3H8]).

Execute your model. Simulate the model with different values of initial

concentrations of C3H8 and CH4.


1. Compare this model with the model from the activities 5A and 5 B.

Describe the differences.

2. The progress of the reaction in time is determined by the reaction rate.

Analyse the given model. Where in the model this rate is defined? From

which variable(s) does the reaction rate depend?

3. In the model it is assumed that the initial concentrations of both reactants

is 1 M and of the product 0 M. Predict and sketch the change of


concentration of [A], [B] and [C] during the reaction.

4. Check your predictions by executing the model.

5. Investigate the effect of changing the initial concentrations of A and B. Use

the control sliders in the Animation window. In what way does the change

of initial concentrations influence the reaction rate? And in what way the

amount of C formed during the reaction?

6. Investigate the effect of changing k on the reaction and reaction rate. Use

the control spinner. What happens when you set higher values of k? And

lower values of k?

7. Modify your model to create a model of the gas-phase oxidation of nitric

oxide: 2 NO + O2 2 NO2

From the point of view of chemical kinetics this reaction can be considered

as a termolecular reaction in which two molecules of NO and one O2 collide

and form a transient complex, which in a single step forms two molecules

of NO2. Following tips will help you to modify existing and create a new model.

Assume that the variable [A] becomes [NO].

Assume that the variable [B] becomes [O2].

Assume that the variable [C] becomes [NO2].

Assume that the initial concentrations values of the reactants are equal

to 0.001 M/L and the initial concentration of the product is 0 M/L.

Set the value of k to 400 L²/(M²s).

Double-click the Process symbol and adjust its properties:

modify, in and out, stochiometric coefficients;

modify the definition of the rate of reaction.

The reaction rate is described by: reaction rate = k*[NO]2*[O2]

(this is the first order reaction with respect to O2, and the second

order reaction with respect to NO.

Execute your model (adjust your model time!).

Simulate the model with different values of initial concentration of NO

and O2. Simulate different k-values.

Analysing activities (using Coach 6):

In these Modelling activities students:

o investigate the model and fill in initial concentrations,

o investigate the effect of changing the initial concentrations and reaction

rate coefficient,

o determine the half-life time of the reaction,

o modify model to describe a real chemical reaction.



The model used in activity A is a simple non-reversible chemical reaction

A B of the first order. The Process element used in the model stands for the

reaction rate r which is proportional to concentration [A]:

r = k*[A] where k is reaction rate coefficient.

In extra assignment students

have to modify the model and

create a model of two

successive reactions

A B C.


This model is slightly modified compare

to model 5A, there is a second reaction

product added.

As the last assignment students have

to modify the model to create a model

of the gas-phase oxidation of nitric

oxide: 2 NO + O2 2 NO2.

Modifying the model means changing

variable names and adjusting their initial values. The reaction rate is then

automatically adjusted to r=k*[C3H8] (first order reaction). Adjusting of In and

Out coefficients of the process element is not needed since all stoichiometric

coefficients of this reaction have value 1.


This model is slightly modified compare

to model 5A, there is a second reactant

added.

As the assignment students have to

modify the model to create a model of

the gas-phase oxidation of nitric oxide:

2 NO + O2 2 NO2.

The adaptation needed for the Process element:

the reaction rate r=k*[NO]2*[O2] (first order reaction with respect to O2,

and the second order reaction with respect to NO);

In and Out coefficients to:

Coefficient [NO]=2, Coefficient [O2]=1, Coefficient [NO2]=2


ACTIVITY 6. MODELLING REVERSIBLE REACTIONS


1. To understand the model of reversible reactions.

2. To use the model to obtain the graphs of concentrations

changes during the reaction.

3. To investigate the effect of changing initial

concentrations of reactant and product and the effect of

changing the reaction rate constants.

4. To understand the concept of dynamic equilibrium.

Operational Skills:

Manipulating model variables


Modifying given model

Procedural Skills:


Evaluating model quality



A. Reversible reaction A to B

1. If at time t = 0 only substance B is present, will the reaction take place?

What do you expect to happen to the concentration of the substance A?

2. What happens if at t = 0 the concentrations of A and B are equal (e.g., [A]

= [B] = 1) but the rate constants are different? Can you sketch the graphs

of concentrations A and B as functions of time? Alter and run the model to

check your prediction.

3. What happens if the rate constants are equal? Use the model to check your

prediction.


1. Load file 6B and run the model to observe how equilibrium becomes

established. Change the values of k1 and k2 to find out how they affect


MODELLING

STUDENT LEVEL:

AGE 17


STUDENT ACTIVITY


the time taken to reach the equilibrium.

2. If at time t = 0 only substance A is present, will the reaction take place?

What do you expect to happen to the concentration of substances B and C?

Adjust the initial value of the concentration of B to 0 and run the model to

check your prediction.

3. What happens if at t = 0 the concentrations of A, B and C are equal (e.g.,

A = B = C = 1) but the rate constants are different? Adjust the initial

values of the model and run it to check your prediction. And what happens

if the rate constants are equal? Check and discuss.


1. Load file 6B and run the model to observe how equilibrium becomes

established. Change the values of k1 and k2 to find out how they affect


2. If at time t = 0 only substance C is present, will the reaction take place?

What do you expect to happen to the concentration of substances A and B?

Adjust the initial values and run the model to check your prediction.


A = B = C = 1) but the rate constants are different? Adjust the initial

values of the model and run it to check your prediction. And what happens

if the rate constants are equal? Check and discuss.



1. Execute the model. After how much time the equilibrium is reached?

2. Calculate the equilibrium constant for the reaction

3. Change the values of the rate constants, kforward and kbackward, by using

control spinners in the Animation window. Check how they affect the time

taken to reach the equilibrium.

4. Change the values of concentrations [A] and [B] to check how they affect


5. If at time t = 0 only species B is present, will the reaction take place? What

do you expect to happen to the concentration of the species A?

6. What happens if at t = 0 the concentrations of A and B are equal (e.g., A =

B = 1) but the rate constants are different?

Predict and sketch the graphs of concentrations A and B as functions of

time? Check you conclusions with the model.

7. And what happens if the rate constants are equal? Check and discuss.

8. Dinitrogen tetroxide N2O4 , a colourless gas, and nitrogen dioxide NO2 a

dark brown gas, exists in equilibrium with each other:

N2O4 2 NO2.


Adapt your model to create a model for this reaction. The following tips will

help you.

Assume that the variable [A] becomes [N2O4].

Assume that the variable [B] becomes [NO2].

Adjust properties of the forward and backward reaction by adjusting the

properties of respective Process symbols:

rforward= kforward*[N2O4]modify, Coefficients of forward reaction: In = 1, Out = 2,

kbackward= kbackward*[NO2]2

Coefficients of backward reaction In = 2, Out = 1

Assume that kforward = 3.14, kbackward = 1

Execute your model.

Calculate the equilibrium constant.

Investigate the effect of changing the initial concentrations and the

reaction time constants. Use the Simulate option.



2. Calculate the equilibrium constant for the reaction.




4. Change the values of concentrations [A], [B] and [C] to check how they

affect the time taken to reach the equilibrium. Use the concentration

control sliders.


A = B = C = 1) but the rate constants are different? Sketch the graphs of

concentrations A, B and C as functions of time? Check you conclusions with

the model.

6. And what happens if the rate constants are equal? Check and discuss.

Adapt your model to create a model for this reaction.



2. Calculate the equilibrium constant for the reaction.




4. If at time t = 0 only species C is present, will the reaction take place? What

do you expect to happen to the concentration of the species A and B?

Explain your reasoning and check with the model.


5. Change the values of concentrations [A], [B] and [C] to check how they

affect the time taken to reach the equilibrium. Use the concentration

control sliders.


A = B = C = 1) but the rate constants are different? Sketch the graphs of

concentrations A, B and C as functions of time? Check you conclusions with

the model. And what happens if the rate constants are equal? Check and discuss.


The students in these modelling activities are invited to:

change the rate constants and discuss how the change affects the time

taken to reach the equilibrium;

choose different initial conditions and discuss what happens;

analyse specific cases of initial conditions and values for the rate constants.


The given model describes a reversible chemical reaction

A B.

The forward reaction, A B, has a rate law rforward = kforward *[A]

and the backward, B A, has a rate law vbackward = kbackward *[B].

The equilibrium coefficient K at equilibrium is expressed as K = kforward /kbackward=[B]/[A].

As the last assignment students have to adapt the given model to describe the

equilibrium reaction: N2O4 2 NO2.


The given model describes a reversible chemical reaction A B + C. The

forward reaction, A B + C, has a rate law rforward = kforward *[A]

and the backward reaction, B + C A, has a rate law rbackward = kbackward *[B][C].

The equilibrium coefficient K at equilibrium is expressed as

K = kforward /kbackward=[B][C]/[A].


The given model describes a reversible chemical reaction A + B C.

The forward reaction, A B C, has a rate law rforward = kforward *[A][B]... and the

backward reaction, C A + B , has a rate law rbackward = kbackward *[C].

The equilibrium coefficient K at equilibrium is expressed as

K = kforward /kbackward=[C]/[A][B].


ACTIVITY 7. CRYSTAL VIOLET


1. To use a model to obtain and analyse graphs of

the change of crystal violet concentration [CV] during

the reaction between crystal violet and sodium

hydroxide.

2. To compare the results calculated by the model with

the results measured during the experiment.

Operational Skills:


Using the cursor tools for obtaining readings from the graph

Procedural Skills:



Compare the model results with the experimental measured results


1. Analyse the model. What assumption does the model make about the rate of change of [CV]?

2. Run the model and observe the changes in the concentration of [CV] and

the rate of reaction r as shown by the graphs.

3. Use the Trial fit option (Data menu) to find out which formula describes the change of [CV]. Remember this when you compare the graph with the

experimental data.

4. Looking at the model, predict the shape of graph if you plot r against [CV].

Test your prediction by adjusting the axes to show r against [CV].

5. Load the file of experimental data and compare the results as directed.


1. Analyse the given model. Which assumption does the model contain about

the rate of change of [CV]?

2. Execute the model.

3. Import the data from the Coach Result file Crystal violet - measurement.

Use the Import Background Graph option available via the Tool menu of


MODELLING

STUDENT LEVEL:

AGE 17


STUDENT ACTIVITY


the Diagram pane.

4. Does the calculated curve of [CV] looks like the measured one?

5. Which initial value are you allowed to change in the model to enhance the

fit?

6. Use the Simulate option, available via the Tool menu of the Modelling

window, to find a value for the best fitting curve.

7. Should the whole graph agree with the measurement results? (What do

you know about the colorimeter?) Determine the moments between which

both curves should fit.

8. What is your conclusion about the relationship between [CV] and the

reaction rate for the reaction between [CV] and [OH-]?

9. What is your conclusion about the relationship between [OH-] and the

reaction rate?

10. Create and describe the curve of -log([CV]) for the model results.

11. Also make a curve of -log([CV]) for the measurement results.


In this experiment 7. Crystal violet - measurement the reaction between

crystal violet (CV) and sodium hydroxide is observed. A solution of crystal

violet (deep purple/blue) reacts with OH- and forms a colourless compound

(CVOH). A simplified version of the equation is:

CV+ + OH- CVOH

(crystal violet) (hydroxide)

Fainting of the crystal violet solution is observed with the colorimeter. The

measurement results are shown below.


The model simulates the change

of [CV] during the reaction

between crystal violet and

sodium hydroxide. In the model

it is assumed that that the

reaction rate r is proportional to

[CV]:

r = k*[CV].

Students must verify this assumption according to the given measurement

results.

Documents

ICT for Innovative Science Teachersictforist.oeiizk.waw.pl/upload/Chemical Reactions Module_EN.pdf · Activity 1. Rate of reaction of magnesium with hydrochloric acid In this activity