cs_unit 18(e)

156

Topic 4 Acids and Bases

UnitUnitRate of reaction1818

18.1 Fast and slow reactions

Different chemical reactions go at different rates. A variety of rapid combustion reactions take place following the ignition of fireworks. The precipitation of silver chloride from a solution also goes very fast. Others, such as the souring of milk or the rusting of iron, are much slower.

Fig. 18.1 The burning of fireworks is a very fast reaction

Fig. 18.2 Rusting of iron is a very slow reaction

There are many reasons why chemists study reaction rates, for example:

• to improve the rate of production of a chemical;

• to help understand the processes going on in our bodies or in the environment so that they can control them;

• to gain an insight into the mechanism of a reaction, i.e. the series of steps involved in a reaction.

157

Unit 18 Rate of reaction

Table 18.1 Concentrations of reactants and products as a function of time at 55 °C

Time (s)Concentration (mol dm–3)

N2O5(g) NO2(g) O2(g)

0 0.0200 0 0

100 0.0169 0.00620 0.00155

200 0.0142 0.0116 0.00290

300 0.0120 0.0160 0.00400

400 0.0101 0.0198 0.00495

500 0.00860 0.0228 0.00570

600 0.00720 0.0256 0.00640

700 0.00610 0.0278 0.00695

18.2 The rate of a reaction

During a chemical reaction, reactants are being consumed while products are being produced. To describe the rate of a reaction quantitatively, we can measure how fast the concentration (or amount) of a reactant or a product changes per unit time.

Rate = change in concentration (or amount) of a reactant or a product

time

The rate of a reaction is usually expressed in mol dm–3 s–1 (i.e. change in molar concentration per second). Other units such as mol s–1, cm3 s–1 and g s–1 are also used.

Consider the thermal decomposition of gaseous dinitrogen pentoxide, N2O5, to give the brown gas nitrogen dioxide and colourless gas oxygen:

2N2O5(g) 4NO2(g) + O2(g) colourless brown colourless

Table 18.1 shows concentrations of reactants and products as a function of time at 55 °C. Notice that the concentration of N2O5(g) decreases while those of NO2(g) and O2(g) increase.

158


We can define the rate of a reaction either as the increase in the concentration of a product per unit time or the decrease in the concentration of a reactant per unit time.

Let us first look at the formation of oxygen. In the decomposition of dinitrogen pentoxide, the rate of formation of oxygen is given by the expression:

Rate of formation of O2(g)

= concentration of O2(g) at time t2 – concentration of O2(g) at time t1

t2 – t1

= Δ[O2(g)]

Δt

In the above expression,

• square brackets surrounding O2(g) denote its concentration in mol dm–3, and Δ[O2(g)] is the change in concentration of O2(g) during the interval from t1 to t2;

• Δt is the change in time.

We can calculate the average rate of formation of O2(g) during the time period 300 s to 400 s as follows:

Average rate of formation of O2(g)

= Δ[O2(g)]

Δt

= 0.00495 mol dm–3 – 0.00400 mol dm–3

400 s – 300 s = 9.50 x 10–6 mol dm–3 s–1

Fig. 18.3 shows three curves obtained by plotting the data of Table 18.1. Look at the triangle drawn on the curve for oxygen between the time period 300 s to 400 s. Δ[O2(g)] represents the vertical side of the triangle while Δt represents the horizontal side. The slope

of the hypotenuse of the triangle is Δ[O2(g)]

Δt, i.e. the average rate

of formation of oxygen during that time period. The steeper the hypotenuse, the higher the rate.

159


Fig. 18.3 Three curves showing the data of Table 18.1

Look at the triangle drawn on the curve for nitrogen dioxide. It is defined by Δ[NO2(g)] and Δt. We can calculate the average rate of formation of NO2(g) during the time period 300 s to 400 s as follows:

Average rate of formation of NO2(g)

= Δ[NO2(g)]

Δt

= 0.0198 mol dm–3 – 0.0160 mol dm–3

400 s – 300 s = 3.80 x 10–5 mol dm–3 s–1

Now look at the triangle drawn on the curve for dinitrogen pentoxide. It is defined by Δ[N2O5(g)] and Δt. As nitrogen dioxide and oxygen form, dinitrogen pentoxide disappears. Since the concentration

of dinitrogen pentoxide decreases with time, Δ[N2O5(g)]

Δt is a negative

quantity.

N1

N1 Refer to page T55.

160


As it is usual to work with positive reaction rates, we always introduce a minus sign when calculating the rate of disappearance of a reactant. We calculate the average rate of decomposition of dinitrogen pentoxide during the 300 s to 400 s period as follows:

Average rate of decomposition of N2O5(g)

= – Δ[N2O5(g)]

Δt

= – 0.0101 mol dm–3 – 0.0120 mol dm–3

400 s – 300 s = 1.90 x 10–5 mol dm–3 s–1

When quoting a reaction rate, it is important to specify the reactant or product on which the rate is based because rates of reactant disappearance and rates of product formation may differ, as those in the above example.

18.1 Calculating the average rate of formation of a gas in mol s–1

In a chemical reaction, 0.242 g of carbon dioxide gas is formed in the first minute. What is the average rate of formation of carbon dioxide in mol s–1 for that time interval?

(Relative atomic masses: C = 12.0, O = 16.0)

Solution

Number of moles of CO2 formed = 0.242 g

44.0 g mol–1

= 0.00550 mol

Average rate of formation of CO2 = 0.00550 mol

60 s

= 9.17 x 10–5 mol s–1

∴ the average rate of formation of carbon dioxide in the first minute is 9.17 x 10–5 mol s–1.

161


1 In a chemical reaction between magnesium and an acid, 0.0160 g of hydrogen gas is formed in the first 40 seconds. What is the average rate of formation of hydrogen in mol s–1 for that time interval?

(Relative atomic mass: H = 1.0)

2 X(g) undergoes thermal decomposition to give Y(g) and Z(g).

X(g) Y(g) + Z(g)

The graph below shows the variation of concentration of X with time.

Calculate the average rate of decomposition of X(g) over the 10 s to 50 s interval.

Refer to page T56.

AS 2005 Q8(a)(i)

162


18.3 Instantaneous rate of reaction

Often, chemists want to know the rate of a reaction at a specific time rather than the rate averaged over a time interval Δt. For example, in the decomposition of dinitrogen pentoxide, what is the rate of formation of nitrogen dioxide at time t = 350 s?

Refer to Fig. 18.4 that shows only the concentration of nitrogen dioxide plotted against time when dinitrogen pentoxide decomposes at 55 °C. If we make our measurements at shorter and shorter time intervals, the triangle defined by Δ[NO2(g)] and Δt will shrink to a point, and the slope of the hypotenuse of the triangle will approach the slope of the tangent to the curve at t = 350 s.

Fig. 18.4 The concentration of nitrogen dioxide plotted against time when dinitrogen pentoxide decomposes at 55 °C

instantaneous rate 瞬間速率

The slope of the tangent to a concentration-time curve at a time t is called the instantaneous rate at that particular time.

163


Fig. 18.5 Determining the instantaneous rate of formation of nitrogen dioxide at t = 350 s

Mathematically, we write

Instantaneous rate of formation of NO2(g) at time t

= d[NO2(g)]

dt = slope of tangent to curve at time t

Notice that ‘d’ is used rather than Δ to show that the change is taking place over a very small time interval.

Instantaneous rate of formation of NO2(g) at 350 s (Fig. 18.5)

= slope of tangent at 350 s

= 0.0238 mol dm–3 – 0.0124 mol dm–3

500 s – 200 s = 3.80 x 10–5 mol dm–3 s–1

Refer to Fig. 18.6 that shows only the concentration of dinitrogen pentoxide against time when it decomposes at 55 °C.

AS 2003 Q10(a)(ii)(1)

164


From the rates of decomposition of dinitrogen pentoxide at t = 0 s, 200 s, 400 s and 600 s, we can see that the rate of decomposition is the highest at the start of the reaction. The tangent to the curve at t = 0 s is the steepest.

Time t Instantaneous rate of decomposition of N2O5(g)

0 s 4.00 x 10–5 mol dm–3 s–1

200 s 2.27 x 10–5 mol dm–3 s–1

400 s 1.88 x 10–5 mol dm–3 s–1

600 s 1.23 x 10–5 mol dm–3 s–1

Fig. 18.6 Concentration of dinitrogen pentoxide plotted against time when it decomposes at 55 °C

Instantaneous rate of decomposition of N2O5(g) at time t

= – d[N2O5(g)]

dt = – slope of tangent to curve at time t

Let us now calculate the instantaneous rate of decomposition of dinitrogen pentoxide at t = 0 s, 200 s, 400 s and 600 s.

AS 2007 Paper 2 Q2(a)(ii)

AL 2007 Paper 2 Q3(a)(ii)

165


1 On heating, chloroethane decomposes according to the following equation:

CH3CH2Cl(g) H2C=CH2(g) + HCl(g)

A chloroethane sample with concentration of 0.100 mol dm–3 is maintained at 400 °C. The way in which the concentration of chloroethane varies with time is shown in the graph below.

a) From the graph, determine the instantaneous rate of decomposition of chloroethane at

i) t = 200 s;

ii) t = 400 s.

b) Suggest how the rate of decomposition of chloroethane varies with time by comparing the two values calculated in (a).

The rate of decomposition of dinitrogen pentoxide decreases with time. The rate at t = 200 s is less than that at t = 0 s and the tangent to the curve at t = 200 s is less steep.

Refer to page T57.

166


2 An experiment is carried out to study the conversion of 2-bromo-2-methylpropane to methylpropan-2-ol according to the following equation:

(CH3)3CBr(l) + NaOH(aq) (CH3)3COH(aq) + NaBr(aq)

Equal volumes of 0.200 mol dm–3 solutions of the two reactants are mixed together and maintained at a constant temperature of 30 °C. The concentration of methylpropan-2-ol is determined at different times. The results are shown in the table below.

a) Plot a graph of concentration of (CH3)3COH(aq) against time.

b) Calculate the average rate of formation of (CH3)3COH(aq) in the first 400 seconds.

c) What is the instantaneous rate of formation of (CH3)3COH(aq) at t = 200 s?

Time (s) Concentration of (CH3)3COH(aq) (mol dm–3)

0 0

120 0.0290

240 0.0480

360 0.0650

480 0.0760

600 0.0820

18.4 How are different expressions for reaction rates related?

We can express the rate of a chemical reaction in terms of the rate of consumption of a reactant or the rate of formation of a product.

For the reaction

A + B 2C

we can express the rate of the reaction as the rate of consumption of reactant A or B, or the rate of formation of product C.

167


As 2 moles of product C are formed when 1 mole of A reacts with 1 mole of B, the rate of formation of C is twice the rate of consumption of A or B. We can express the rate of the reaction as follows:

Rate = – d[A]

dt = –

d[B]

dt =

1

2

d[C]

dt

For the thermal decomposition of dinitrogen pentoxide:

2N2O5(g) 4NO2(g) + O2(g)

When 1 mole of oxygen is formed, 4 moles of nitrogen dioxide are formed and 2 moles of dinitrogen pentoxide are decomposed. Hence the rate of formation of oxygen is one-fourth the rate of formation of nitrogen dioxide and half of the rate of decomposition of dinitrogen pentoxide.

We can express the rate of the reaction as follows:

Rate = – 1

2

d[N2O5(g)]

dt =

1

4

d[NO2(g)]

dt =

d[O2(g)]

dt

For the general reaction,

aA + bB cC + dD

Rate = – 1

a

d[A]

dt = –

1

b

d[B]

dt =

1

c

d[C]

dt =

1

d

d[D]

dt

18.2 Relating expressions for reaction rates

Ammonia reacts with oxygen in the presence of a catalyst to give nitrogen monoxide and steam.

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)

NH3(g) and O2(g) are allowed to react in a closed vessel of constant volume.

The rate of formation of NO(g) = d[NO(g)]

dt

The rate of consumption of O2(g) = – d[O2(g)]

dt

a) Express the rate of the reaction in terms of the above two expressions.

b) Find the rate of consumption of O2(g) if the rate of formation of NO(g) is 1.20 x 10–4 mol dm–3 s–1.

AS 2006 Q10(a)(i)

168


Solution

a) Rate = – 1

5

d[O2(g)]

dt =

1

4 d[NO(g)]

dt

b) Rate of consumption of O2(g) = – d[O2(g)]

dt

= 5

4 d[NO(g)]

dt

= 5

4 (1.20 x 10–4 mol dm–3 s–1)

= 1.50 x 10–4 mol dm–3 s–1

∴ the rate of consumption of O2(g) is 1.50 x 10–4 mol dm–3 s–1.

1 Ammonia is manufactured by the reaction between nitrogen and hydrogen:

N2(g) + 3H2(g) 2NH3(g)

The rate of formation of NH3(g) = d[NH3(g)]

dt

The rate of consumption of N2(g) = – d[N2(g)]

dt

The rate of consumption of H2(g) = – d[H2(g)]

dt

Express the rate of the reaction in terms of the above three expressions.

2 Consider the reaction of acidified hydrogen peroxide with potassium iodide:

H2O2(aq) + 2I–(aq) + 2H+(aq) 2H2O(l) + I2(aq)

At a particular instant, the rate of formation of iodine was found to be 10–6 mol dm–3 s–1.

a) What is the rate of consumption of H2O2(aq) at that instant?

b) What is the rate of consumption of I–(aq) at that instant?

Refer to page T59.

169


18.5 Methods for following the progress of a reaction

To determine the rate of a reaction, we need to follow the change in concentration of a reactant or a product. There are two types of methods we can use:

• methods using a variety of physical properties of the reaction mixture, such as

– measuring the volume of a gaseous product;

– measuring the loss in mass of the reaction mixture;

– measuring the pressure of the reaction mixture;

– measuring the colour intensity of the reaction mixture;

– measuring the light transmittance of the reaction mixture.

• method based on titration (i.e. titrimetric analysis).

We will discuss each method in turn.

18.6 Following the progress of a reaction by measuring the volume of a gaseous product

When a gas is evolved during a reaction, we may follow the progress of the reaction by measuring the volume of the gaseous product.

For example, when magnesium reacts with dilute hydrochloric acid, hydrogen gas is evolved.

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

We can follow the progress of the reaction by measuring the volume of hydrogen formed at regular time intervals.

Fig. 18.7 shows the set-up of the experiment. Take the steps below to follow the progress of the reaction:

1 Drop a piece of magnesium ribbon into the dilute hydrochloric acid in the conical flask.

2 Start the stop watch at the same time.

3 Record the volume of hydrogen formed every 30 seconds until the reaction stops.

titrimetric analysis 滴定分析

170


The table below shows a set of sample results.

Fig. 18.7 Experimental set-up for collecting hydrogen gas produced in the reaction between magnesium and dilute hydrochloric acid (showing the changes in the first half minute only)

Time (min) 01

21 1

1

22 2

1

23 3

1

24 4

1

25 5

1

26 6

1

2

Volume of hydrogen produced (cm3)

0 8 14 20 25 29 33 36 38 39 40 40 40 40

Fig. 18.8 Results of the reaction between magnesium and dilute hydrochloric acid

Fig. 18.8 shows the curve of volume of hydrogen formed plotted against time. Notice the following points from the experimental results:

• The rate of the reaction is the highest at the start of the reaction and the slope of the tangent to the curve at t = 0 min is the steepest.

• After 5 minutes, no hydrogen is produced. So, its volume no longer changes. The reaction is over.

N2

N2 Accounting for the change in shape of the curve of the results of the reaction between magnesium and dilute hydrochloric acid

Examination questions may ask

students to account for the change in

shape of the curve for the results of

the reaction between magnesium and

dilute hydrochloric acid or other similar

reactions.

• The rate of the reaction is highest

at the start of the reaction as the

concentration of the acid is the

highest.

• The rate of the reaction decreases as

the reaction proceeds because the

concentration of the acid decreases

during the course of the reaction.

• The reaction stops when either the

magnesium or hydrochloric acid has

been used up.

CE 2009 Paper 1 Q10(b)(i)

AS 1999 Q11(a)

AS 2000 Q7(a)

AS 2001 Q10(a)

171


• Altogether, 40 cm3 of hydrogen are produced in the reaction.

Many other chemical reactions give gaseous products as well. For example:

• carbon dioxide is formed in the reaction between calcium carbonate and dilute hydrochloric acid:

CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)

• oxygen is formed in the decomposition of hydrogen peroxide solution:

2H2O2(aq) 2H2O(l) + O2(g)

We can also follow the progress of such reactions by measuring the volumes of their gaseous products.

18.7 Following the progress of a reaction by measuring the loss in mass of the reaction mixture

Besides measuring the volume of the gaseous product, we can also follow the progress of the reaction between calcium carbonate and dilute hydrochloric acid by measuring the loss in mass of the reaction mixture.


Fig. 18.9 shows the set-up of the experiment. As carbon dioxide can escape through the cotton wool, the reaction mixture gets lighter as the reaction proceeds.

Fig. 18.9 Experimental set-up for measuring the loss in mass of the reaction mixture in the reaction between calcium carbonate and dilute hydrochloric acid (showing the changes in the first half minute only)

110.00 108.50

AL 2006 Paper 2 Q3(a)(i)

AL 1998 Paper 2 Q3(a)(iii)

AS 1999 Q11(a)

172


Take the steps listed below to follow the progress of the reaction:

1 Put a conical flask containing dilute hydrochloric acid on an electronic balance.

2 Add some calcium carbonate of known mass to the acid. Plug the flask with cotton wool immediately to prevent any acid from splashing out.

3 Record the mass of the flask plus its contents and start the stop watch at the same time.

4 Record the mass of the flask plus its contents every 30 seconds until the reaction stops.

The table below shows a set of sample results. Fig. 18.10 shows the curve of loss in mass of the reaction mixture plotted against time. The loss in mass of the reaction mixture equals the mass of carbon dioxide formed. We may also plot the mass of the flask and the reaction mixture against time (Fig. 18.11).

Time t (min) 01

21 1

1

22 2

1

23 3

1

24

Mass of flask and reaction mixture (g)

110.00 108.50 107.50 106.95 106.60 106.41 106.33 106.30 106.30

Loss in mass of reaction mixture (g)

0 1.50 2.50 3.05 3.40 3.59 3.67 3.70 3.70

Fig. 18.10 A plot of loss in mass of the reaction mixture against time

Fig. 18.11 A plot of mass of the flask and the reaction mixture against time

N3N3 Refer to page T59.

173


18.8 Following the progress of a reaction by measuring the pressure of the reaction mixture

When magnesium reacts with dilute hydrochloric acid, hydrogen gas is formed. If the reaction vessel is a closed system, the pressure inside the vessel will increase. We can follow the progress of the reaction by measuring the pressure inside the vessel with a pressure sensor connected to a data-logger interface and a computer.

Fig. 18.12 shows the set-up of the experiment. Take the steps listed below to follow the progress of the reaction:

1 Execute the data-logging software on the computer. Open a graph display with a plot of pressure against time.

2 Tilt the bottle containing the hydrochloric acid to mix the acid with the magnesium ribbon.

3 Immediately start recording the pressure inside the suction flask for about 5 minutes.

Fig. 18.12 Experimental set-up for measuring the pressure change in the reaction between magnesium and dilute hydrochloric acid

18.9 Following the progress of a reaction by measuring the colour intensity of the reaction mixture

For a chemical reaction involving a coloured reactant or product, the colour intensity of the reaction mixture will change during the course of the reaction.

For example, in the oxidation of oxalate ions (C2O42–)* by

permanganate ions (MnO4–), the intensity of the purple colour of

permanganate ions decreases as the reaction proceeds.

pressure sensor 壓強感應器 data-logger interface 數據收集儀界面

The systematic name of oxalate ion is ethanedioate ion.

Refer to the video clip ‘Using Data Loggers

in Chemistry Practical Classes’ (10 mins.):

h t tp: / /e tv.hkedci ty.net /Home/Pages/

ResourceList.aspx?catId=12173&subId=6

&specialFirst=False# N4


CE 2009 Q10(a)(ii)

AS 2000 Q7(a)

174


2MnO4–(aq) + 5C2O4

2–(aq) + 16H+(aq) 2Mn2+(aq) + 10CO2(g) + 8H2O(l) purple colourless colourless colourless

We can follow the progress of the reaction using a colorimeter. A colorimeter is an instrument for measuring the amount of light absorbed by a sample when a beam of light passes through the sample.

A colorimeter consists of a light source with filters to select a suitable colour of light which is absorbed most by the coloured species in the sample. The light passes through the sample onto a detector whose output goes to a meter or a recording device.

colorimeter 比色計 absorbance 吸光度 calibration curve 校準曲線

Fig. 18.13 The basic components of a colorimeter

In practice, we shine the light upon the sample and record the fraction of light absorbed. This fraction is called the absorbance. The absorbance is directly proportional to the colour intensity of the sample and the concentration of the coloured species in the sample.

Before following the process of the oxidation of oxalate ions by permanganate ions using a colorimeter, we need to prepare a calibration curve.

Take the steps listed below to follow the progress of the reaction:

1 Preparing the calibration curve:

a) Prepare a set of solutions of permanganate ions of known concentrations.

b) Record the absorbance of each solution using a colorimeter.

c) Plot the absorbance against concentration of permanganate ions (Fig. 18.14).

AS 2008 Paper 1 Q7(a),(b)

AL 2008 Paper 1 Q7(a)(i),(ii)

N5N5 Examination questions may show the

basic components of a colorimeter

and ask students to give the name of

the instrument and the property of the

reaction mixture measured by it.

175


Fig. 18.14 A plot of absorbance against concentration of permanganate ions

2 Following the progress of the oxidation reaction:

a) Mix solutions of oxalate ions and permanganate ions under acidic conditions.

b) Measure the absorbance of the reaction mixture using the colorimeter at regular time intervals.

c) From the calibration curve, read off concentrations of permanganate ions at regular time intervals according to the absorbance recorded.

Fig. 18.15 shows the curve of absorbance of the reaction mixture plotted against time. When the reactants are mixed, the reaction mixture becomes lighter in colour gradually as permanganate ions are consumed. Thus, the reaction mixture absorbs less and less light and so the absorbance goes down.

Fig. 18.15 A plot of absorbance of the reaction mixture against time

AS 2008 Paper 1 Q7(c)

176


Many other chemical reactions involve coloured reactants or products as well. For example,

• in the reaction between iodine and propanone, the intensity of the brown colour of iodine decreases as the reaction proceeds;

I2(aq) + CH3COCH3(aq) CH3COCH2I(aq) + HI(aq) brown colourless colourless

• in the oxidation of methanoic acid by bromine, the intensity of the yellow-brown colour of bromine decreases as the reaction proceeds.

Br2(aq) + HCOOH(aq) 2Br–(aq) + 2H+(aq) + CO2(g) yellow- colourless brown colourless

We can also follow the progress of such reactions by measuring the absorbance of their reaction mixtures using a colorimeter.

18.10 Following the progress of a reaction by measuring the light transmittance of the reaction mixture

In some chemical reactions, the light transmittance of the reaction mixture may change as the reaction proceeds.

For example, when sodium thiosulphate solution reacts with dilute sulphuric acid, a yellow precipitate of sulphur forms. This changes the light transmittance of the reaction mixture.

Na2S2O3(aq) + H2SO4(aq) Na2SO4(aq) + SO2(g) + H2O(l) + S(s)

Take the steps listed below to follow the progress of the reaction (Fig. 18.16):

1 Mark a cross on a piece of paper.

2 Put a beaker containing some sodium thiosulphate solution on top of the paper. It should be easy to see the cross through the solution from above.

3 Add dilute sulphuric acid to the beaker.

4 Start the stop watch at the same time. The cross gets fainter as the precipitate forms.

5 Stop the stop watch when the cross can no longer be seen from above.

transmittance 透光度

AS 2007 Paper 2 Q2(a)(i)

AS 2008 Paper 1 Q7(a),(b)

N6

N6 Examination questions often ask about

the oxidation of methanoic acid by

bromine.

N7N7 Besides the method described here,

the progress of the reaction can also

be followed by using a colorimeter.

This is to follow the change in light

level passing the reaction mixture as

the turbidity increases. The time taken

to reach a certain level of turbidity can

be regarded as t.

177


Fig. 18.16 Following the progress of the reaction between sodium thiosulphate solution and dilute sulphuric acid

(a) When the reactants are being mixed, the cross can be clearly seen.

(b) The cross gets fainter as more sulphur precipitate forms.

(c) The cross can no longer be seen as the solution becomes opaque.

The cross disappears when enough sulphur forms and the solution becomes opaque.

If the reaction is fast, the time to reach such a stage will be short. If the reaction is slow, the time will be long. The length of time taken to reach such a stage is thus inversely proportional to the average rate of reaction from the start to the opaque stage.

The average rate of reaction from the start to the opaque stage

∝ 1

time to reach the opaque stage

For each of the following reactions, suggest with reason(s) one method that can be used to follow the progress of the reaction.

a) Reaction between zinc and dilute sulphuric acid

Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g)

b) Displacement reaction between magnesium and copper(II) sulphate solution

Mg(s) + CuSO4(aq) MgSO4(aq) + Cu(s)

c) Decomposition of dinitrogen pentoxide in tetrachloromethane

2N2O5(in CCl4) 4NO2(in CCl4) + O2(g)

d) Redox reaction between permanganate ions and iron(II) ions

MnO4–(aq) + 5Fe2+(aq) + 8H+(aq) Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)

opaque 不透明

Refer to page T60.


178


18.11 Following the progress of a reaction using titrimetric analysis

Consider the alkaline hydrolysis of an ester, such as ethyl ethanoate*:

CH3COOC2H5(l) + OH–(aq) CH3COO–(aq) + C2H5OH(aq)

We can follow the progress of the reaction by

• first withdrawing small samples of the reaction mixture at regular time intervals;

• then determining the concentration of sodium hydroxide remaining in each sample by titrimetric analysis.

Common quenching techniques

As titrimetric analysis takes time, we must have some means to stop any further reaction in samples once they are withdrawn from the reaction mixture. The process of stopping a reaction is called quenching.

Common quenching techniques include:

• rapid cooling of the sample in ice;

• removing the catalyst;

• removing one of the reactants by adding a chemical which uses up the reactant rapidly;

• diluting the sample with a large volume of ice-cold water to lower both the temperature and concentrations of the reactants*.

For the alkaline hydrolysis of ethyl ethanoate, we can quench the reaction by running each sample into an excess of ice-cold water.

Following the progress of the alkaline hydrolysis of ethyl ethanoate

Take the steps listed below to follow the progress of the reaction (Fig. 18.17):

1 Mix solutions of ethyl ethanoate and sodium hydroxide in a flask.

2 Withdraw a small sample from the reaction mixture every 3 minutes. The amount of sample withdrawn each time should be the same.

quench 猝滅

An alkanoic acid reacts with an alkanol to form an ester. Ethyl ethanoate is formed when we warm ethanoic acid and ethanol in the presence of concentrated sulphuric acid (as a catalyst).

We will discuss the effect of change in temperature and concentration on the rate of a reaction later in this unit.

AL 1998 Paper 2 Q3(a)(iii)

AS 2009 Paper 2 Q7(c)

AL 1999 Paper 1 Q7(a)


N8


179


(a) Withdraw a small sample from the reaction mixture every 3 minutes and transfer the sample to an excess of ice-cold water

Study the following example to see how we can calculate the concentration of hydroxide ions in the reaction mixture at a particular instant.

(b) Titrate the alkali remaining in the sample of reaction mixture against standard hydrochloric acid

18.3 Following the progress of alkaline hydrolysis of ethyl ethanoate

In a certain experiment, a student followed the progress of alkaline hydrolysis of ethyl ethanoate as described below:

Step 1 Mix 100.0 cm3 of 0.0400 mol dm–3 ethyl ethanoate and 100.0 cm3 of 0.0400 mol dm–3 sodium hydroxide solution in a flask kept in a thermostat.

Step 2 Withdraw a 10.0 cm3 sample from the reaction mixture every 3 minutes and transfer it to an excess of ice-cold water in a conical flask.

Step 3 Titrate the sodium hydroxide remaining in each sample against 0.0100 mol dm–3 hydrochloric acid.

3 Transfer each sample to an excess of ice-cold water.

4 Titrate the alkali remaining in each sample against standard hydrochloric acid* and record the volume of acid required to reach the end point.

thermostat 恒温器

We have discussed titration technique in Unit 17.

Fig. 18.17 Steps to follow the progress of the alkaline hydrolysis of ethyl ethanoate

180


Time (min) 3 6 9 12 15 18

Volume of 0.0100 mol dm–3 HCl(aq) required to reach end point (cm3)

14.0 10.9 9.10 7.75 6.70 5.95

The following results were obtained:

a) Suggest the purpose of transferring the withdrawn sample to ice-cold water in a conical flask in Step 2.

b) Calculate the concentration of hydroxide ions in the reaction mixture at the 3rd minute.

c) Calculate the concentrations of hydroxide ions in the reaction mixture at other time intervals and plot the concentration against time.

Solution

a) This stops the hydrolysis in the withdrawn sample.

b) Hydrochloric acid and sodium hydroxide solution react according to the following equation:

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Number of moles of HCl used in titration = 0.0100 mol dm–3 x 14.0

1 000 dm3

= 1.40 x 10–4 mol

According to the equation, 1 mole of HCl reacts with 1 mole of NaOH.

i.e. number of moles of NaOH remaining in sample = 1.40 x 10–4 mol

Concentration of hydroxide ions in sample

= 1.40 x 10–4 mol

( 10.0

1 000 ) dm3

= 0.0140 mol dm–3

= concentration of hydroxide ions in reaction mixture

∴ the concentration of hydroxide ions in the reaction mixture at the 3rd minute is 0.0140 mol dm–3.

181


c)

Advantages and disadvantages of using titrimetric analysis to follow the progress of a reaction

Advantages

• Titrimetric analysis can be applied to most reactions.

• Only simple laboratory apparatus is required.

Disadvantages

• Samples are withdrawn from the reaction mixture for analysis. This disturbs the reaction somehow.

• Continuous measurement of concentration change is impossible.

• Titrimetric analysis takes time and thus cannot be used for rapid reactions.

Time 0 3 6 9 12 15 18

Concentration of hydroxide ions (mol dm–3)

0.0200 0.0140 0.0109 0.00910 0.00775 0.00670 0.00595

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1 Compare the titrimetric analysis with methods using physical properties for following the progress of a reaction. Identify the major differences between them.

2 Why is it better to select a method that does not involve sample withdrawal for following the progress of a reaction?

18.12 Factors affecting the rate of a reaction

The rate of a reaction depends on many different factors, including

• the concentration of a reactant;

• the surface area of a solid reactant;

• the temperature; and

• the presence of a catalyst.

We are going to consider each of these factors in turn.

18.13 Studying the effect of change in concentration of a reactant on the rate of a reaction

In each of the following activities, you are going to investigate the effect of varying the concentration of a reactant on the rate of a reaction. You will use different methods to follow the progress of the reactions.

Investigating the effect of varying the concentration of vinegar solution on the rate of its reaction with baking soda (sodium hydrogencarbonate).

18.1

A worksheet on this activity is

available for download from Jing Kung Chemistry Website.

• Refer to the following website for a

simulation of the progress of the catalytic

decomposition of hydrogen peroxide

solution:

http://www.chem.iastate.edu/group/

Greenbowe/sections/projectfolder/

flashfiles/kinetics2/rxnRate01.html

• Refer to the following website for a

simulated experiment investigating

the effect of concentration change on

the rate of reaction between sodium

thiosulphate solution and hydrochloric

acid:

http://www.hkedcity.net/resources/ires/

ires_browse.phtml?res_cat_id=2840

An assessment form on this activity is available for download from Jing Kung Chemistry Website.

Refer to page T62.

Effect of varying the concentration on the

rate of reaction

183


Investigating the effect of varying the concentration of a reactant on the rate of the reaction between permanganate ions and oxalate ions in an acidic solution.

18.2

In Activity 18.2, you investigate the effect of varying the concentration of a reactant on the rate of the reaction between permanganate ions and oxalate ions in an acidic solution.

2MnO4–(aq) + 5C2O4

2–(aq) + 16H+(aq) 2Mn2+(aq) + 10CO2(g) + 8H2O(l) purple colourless colourless

As the reaction proceeds, the concentration of permanganate ions in the reaction mixture decreases and the colour intensity of the reaction mixture also decreases. Hence you can follow the progress of the reaction by measuring how the absorbance of the reaction mixture varies with time. You use a colorimeter connected to a data-logger system in the experiment (Fig. 18.18).

Fig. 18.18 Following the progress of a reaction using colorimeter and data-logger system

Reaction mixture

Amounts of reactants in the reaction mixture

sodium oxalate

solution (cm3)

hydrochloric acid (cm3)

distilled water (cm3)

potassium permanganate solution (cm3)

1 1.0 1.0 1.0 1.0

2 1.0 1.0 1.5 0.5

cuvette 比色杯

To investigate the effect of varying the concentration of permanganate ions on the rate of the reaction, conduct the experiment using two reaction mixtures according to the following scheme:

Using a data-logger system with a light

sensor to follow the progress of the

reaction between sodium thiosulphate

solution and dilute hydrochloric acid


184


Consider the reactions between equal masses of magnesium with excess hydrochloric acid of different concentrations: 0.50 mol dm–3, 1.0 mol dm–3, 1.5 mol dm–3 and 2.0 mol dm–3. The volume of acid used is the same in each case. For a fair comparision, we must keep all the variables, except the concentration of the acid, the same.

Fig. 18.20 shows the curves of volume of hydrogen produced plotted against time. You will notice that at the start of the reaction, the more concentrated the acid, the steeper is the tangent to the curve corresponding to the acid. Furthermore, the more concentrated the acid, the less time it takes to complete the reaction. That means the reaction with a more concentrated acid is faster than that with a less concentrated acid.

In most cases, the rate of a reaction increases when the concentration of a reactant is increased.

Fig. 18.20 Results of reactions between equal masses of magnesium and excess hydrochloric acid of different concentrations

Fig. 18.19 The variation of the absorbance of reaction mixtures 1 and 2 against time

Fig. 18.19 shows the variation of the absorbance of each reaction mixture against time. The absorbance of reaction mixture 1 falls more rapidly, i.e. the reaction is faster. The concentration of permanganate ions in reaction mixture 1 is twice of that in reaction mixture 2. Thus, we can conclude that the rate of the reaction increases when the concentration of permanganate ions is increased.

N10

N9N9 Students should know the concept

of ‘fair test’ when asked to make

comparison.


185


18.4 Studying the reaction between marble and hydrochloric acid of different concentrations

In an experiment, a certain volume of 1.00 mol dm–3 hydrochloric acid was allowed to react with 10.0 g of marble (in excess). The progress of the reaction was monitored using the experimental set-up shown below.

XXX.XX

The graph shows the variation of the mass of the conical flask and its contents against time.

a) Write a balanced equation for the reaction between marble and dilute hydrochloric acid.

b) Find, from the graph, the mass of gas evolved from the reaction between marble and dilute hydrochloric acid.

186


c) Keeping other conditions unchanged, the experiment was repeated using the same volume of 0.500 mol dm–3 hydrochloric acid.

i) Determine the mass of gas evolved in this experiment.

ii) Sketch a curve on the same graph to show the variation of the mass of the conical flask and its contents against time.

d) Suggest ONE advantage of using a data-logger in this experiment.

Solution

a) CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)

b) 1.10 g

c) i) The concentration of the hydrochloric acid is halved in the repeated experiment, but the same volume is used. Therefore the number of moles of HCl in the acid is half that in the first experiment. The amount of gas evolved is also half that in the first experiment.

∴ mass of gas evolved in the repeated experiment = 1.10

2 g

= 0.550 g

ii)

d) As the change in the mass is very small in this experiment, the use of a data-logger can give more accurate results.

18.5 Studying the reaction of magnesium with different acids of the same concentration

Two pieces of magnesium ribbon of equal mass were added to 50 cm3 of 1 mol dm–3 hydrochloric acid and 50 cm3 of 1 mol dm–3 ethanoic acid separately. The acid was in excess in each case. The experimental results are shown in the table below.

CE 2005 Paper 1 Q10(b)(iii)

187


Sample

Composition of sampleReaction time t (s)Volume of 0.1 mol dm–3

Na2S2O3 (cm3)Volume of water (cm3)

Volume of 1 mol dm–3 H2SO4 (cm3)

1 4.0 0.0 4.0 37

2 3.0 1.0 4.0 50

3 2.0 2.0 4.0 76

4 1.0 3.0 4.0 154

Reaction ReactantsTime required for reaction to

complete (s)

I magnesium + 50 cm3 of 1 mol dm–3 hydrochloric acid 200

II magnesium + 50 cm3 of 1 mol dm–3 ethanoic acid ?

a) How could you know when the reactions were complete?

b) Would you expect the time required for the completion of Reaction II to be shorter or longer than that for Reaction I? Explain your answer.

Solution

a) When all the magnesium was used up and dissolved in the acid.

b) The time required for the completion of Reaction II would be longer. During the reaction between magnesium and the acids, magnesium would react with hydrogen ions in the acids.

Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)

Hydrochloric acid is a strong acid that completely dissociates in water. On the other hand, ethanoic acid is a weak acid that only partially dissociates in water. Therefore hydrochloric acid has a higher concentration of hydrogen ions than ethanoic acid. The reaction rate between magnesium and ethanoic acid is thus lower and the reaction takes a longer time to complete.

18.6 Studying how the concentration of sodium thiosulphate solution affects the rate of its reaction with dilute sulphuric acid

A student conducted an experiment at 20 °C to study how the concentration of sodium thiosulphate solution affects the rate of the following reaction:

Na2S2O3(aq) + H2SO4(aq) Na2SO4(aq) + SO2(aq) + H2O(l) + S(s)

A colorimeter connected to a data-logger was used to follow the progress of the reaction.

The following samples were used in the experiment.

CE 2005 Paper 2 Q34

AS 2006 Q7(b)

AS 2002 Q11(a)

188


a) Explain how the reaction time could be measured using a colorimeter.

b) Describe how the rate of the reaction can be obtained from the experimental results.

c) Explain why

i) water was used in samples 2 to 4;

ii) the same volume of acid was used in each sample.

d) Plot a suitable graph to show how change in the concentration of sodium thiosulphate in the sample affects the rate of the reaction.

Solution

a) The time required for the sample to reach a certain level of absorbance could be taken as the reaction time.

b) Rate of the reaction ∝ 1

t

c) i) To keep the total volume of each sample constant. Thus, the concentration of sodium thiosulphate in the sample is directly proportional to the volume of thiosulphate solution usd.

ii) To ensure that the only variable is the change in the concentration of sodium thiosulphate in the sample.

d)

It can be concluded that the rate of the reaction increases when the concentration of sodium thiosulphate in the sample is increased.

189


1 To study the rate of reaction of zinc with excess hydrochloric acid, a student carried out the following experiments at room temperature.

The results obtained in Experiment 1 are shown in the graph below.

a) Write a chemical equation for the reaction involved.

b) Draw a labelled diagram of the experimental set-up for carrying out the experiment.

c) Copy the curve for Experiment 1 and draw the expected curve for Experiment 2 next to it.

d) In another experiment, the same mass of zinc was added to 50 cm3 of 2 mol dm–3 sulphuric acid. The time required for the reaction to complete was shorter than that for Experiment 1. Explain why.

2 Acidified propanone solution reacts with iodine according to the following equation:

CH3COCH3(aq) + I2(aq) CH3COCH2I(aq) + HI(aq)

The progress of the reaction can be followed by a colorimeter. Two samples are prepared as listed below:

Experiment Reactants mixed

1 5 g of zinc + 50 cm3 of 2 mol dm–3 HCl

2 5 g of zinc + 50 cm3 of 4 mol dm–3 HCl

Refer to page T64.

190


The progress of the reaction of Sample 1 was followed by a colorimeter and the results obtained are shown in the graph.

a) Account for the change of absorbance of the sample as shown in the graph.

b) Sketch, on the same graph, the expected results of Sample 2.

SampleVolume of 1 mol dm–3

propanone (cm3)

Volume of 1 mol dm–3

sulphuric acid (cm3)

Volume of distilled water

(cm3)

Number of drops of 0.02 mol dm–3

I2

1 0.75 0.75 1.50 30

2 1.50 0.75 0.75 30

Fig. 18.21 If we break a piece of solid into 8 smaller pieces, its surface area doubles

1 piece of solid 8 small pieces of solid

Surface area of one side of the solid = 2 x 2 cm2 = 4 cm2

Total surface area (6 sides)= 6 x 4 cm2 = 24 cm2

Surface area of 1 small cube = 6 x 1 cm2 = 6 cm2 Surface area of 8 small cube= 8 x 6 cm2 = 48 cm2

The surface area is doubled!

18.14 Studying the effect of change in surface area of a solid reactant on the rate of a reaction

If we break a large piece of solid into smaller pieces, its surface area increases. Study Fig. 18.21. When we break a piece of solid into 8 smaller pieces, its surface area doubles.

191


In the following activity, you are going to investigate the effect of varying the surface area of a solid reactant on the rate of a reaction.

Investigating the effect of varying the surface area of marble chips on the rate of their reaction with dilute hydrochloric acid.

18.3

Fig. 18.22 In the reaction with dilute hydrochloric acid, powdered calcium carbonate (left) reacts faster than calcium carbonate lumps (right) do

The rate of a reaction increases when the surface area of a solid reactant is increased.

18.7 Studying how the particle size of calcium carbonate affects the rate of its reaction with dilute hydrochloric acid

A flask containing excess 1 mol dm–3 hydrochloric acid was placed on an electronic balance. Some calcium carbonate lumps were carefully added into the flask. The loss in mass of the reaction mixture is plotted against time as shown.

For example, in the reaction with dilute hydrochloric acid, powdered calcium carbonate reacts faster than calcium carbonate lumps do (Fig. 18.22).

Investigating the effect of varying the surface area of marble chips on the rate of their reaction with dilute hydrochloric acid

AS 2003 Q10(a)(iii)(1)

AS 2006 Q7(a)


192


a) Write a chemical equation for the reaction involved.

b) Why did the mass of the reaction mixture decrease during the experiment?

c) The dotted line represents the results from another experiment run under the same conditions except using calcium carbonate of a different particle size. Would you expect the particle size to be larger or smaller? Explain your answer.

Solution

a) CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

b) Carbon dioxide gas escaped.

c) The particle size of the calcium carbonate was smaller as the reaction represented by the dotted line took less time to complete, i.e. the rate of this reaction was higher.

Investigating the effect of varying the temperature on the rate of the reaction between sodium thiosulphate solution and dilute sulphuric acid.

18.4Fig. 18.23 The low temperature in

a refrigerator can slow down the reaction that makes milk sour

Investigating the effect of varying the temperature on the rate of the reaction between ingredients of Alka Seltzer tablet.

18.6

18.15 Studying the effect of change in temperature on the rate of a reaction

In each of the following activities, you are going to investigate the effect of varying the temperature on the rate of a reaction.

Investigating the effect of varying the temperature on the rate of acid hydrolysis of ethyl ethanoate.

18.5

• Effect of varying the temperature on the

rate of reaction

• Reaction rate and effective collisions

— Effect of concentration


— Effect of surface area


— Effect of temperature

An assessment form on this activity is available for

download from Jing Kung Chemistry Website.





N11


193


18.8 Studying the relationship between temperature and reaction rate

The following information is taken from the side of a carton of milk.

How long does your milk keep?

30 °C (hot summer day) 1

2 day

25 °C (room temperature) 1 day

15 °C 2 days

5 °C (refrigerator temperature) 5 days

a) Plot a graph of the time the milk keeps against the temperature.

b) Describe the shape of the curve you obtain.

c) Use your curve to predict how long the milk will keep at 20 °C.

Solution

a)

In most cases, the rate of a reaction increases when the temperature is increased.N12 Refer to page T65. N12

194


b) The graph is a decreasing curve. As temperature increases, shelf life becomes shorter.

c) 1.4 days.

Some powdered calcium carbonate are allowed to react with excess 0.50 mol dm–3 hydrochloric acid in a conical flask. The loss in mass of the reaction mixture is recorded at regular time intervals. The results are represented by curve X.

Sketch next to curve X the curve you expect for each of the following changes made:

a) using a lump of calcium carbonate of the same mass;

b) using a lump of calcium carbonate of lower mass;

c) using the same volume of 1.0 mol dm–3 hydrochloric acid;

d) increasing the temperature.

A catalyst is a substance which alters the rate of a reaction without itself undergoing any permanent chemical changes.

18.16 Studying how the presence of a catalyst affects the rate of a reaction

catalyst 催化劑


N13

Refer to page T67.

CE 2006 Paper 2 Q44

AS 2002 Q10(b)(ii)(1)

AS 2006 Q10(a)(ii)

195


Catalyzing the decomposition of hydrogen peroxide in solution.

18.7

There are two types of catalysts: positive catalysts and negative catalysts. A positive catalyst is one that speeds up a reaction. A negative catalyst is one that slows down a reaction.

One way of making oxygen in the school laboratory is by the decomposition of hydrogen peroxide, H2O2.

2H2O2(aq) 2H2O(l) + O2(g)

Under normal conditions, this reaction is very slow. Adding manganese(IV) oxide, MnO2, to the solution speeds up the decomposition. The manganese(IV) oxide is a positive catalyst of the reaction.

On the other hand, dilute hydrochloric acid is a negative catalyst of the reaction. The decomposition of hydrogen peroxide slows down when dilute hydrochloric acid is added to the solution.

In the following activity, you are going to identify catalysts for this decomposition reaction.

Saved by a very fast chemical reaction

Airbags are a familiar addition to car safety these days. But a crash happens very fast, how does the airbag inflate in time to protect the driver from injury? One type of airbag uses a very fast gas-producing chemical reaction to blow up the bag.

How does an airbag work?

The airbag is made up of three parts. Firstly, there is the bag itself, which is made of thin, nylon fabric and folded into the steering wheel or the dashboard. The second part is the crash sensor that tells the bag to inflate. Finally, there is the inflation system (Fig. 18.25).

Fig. 18.24 Airbags are blown up by a very fast gas-producing reaction

inflate 膨脹

Two related worksheets with answers are available for download from Jing Kung Chemistry Website. Worksheet ➀ consists of

additional questions for promoting active reading. (Note: Questions in this worksheet also cover those in the textbook.)

Worksheet ➁, requiring students to indicate the relationship among various

terms by drawing, may help the less able readers improve their reading skills.

CE 2009 Paper 1 Q10(a)(i)

AS 2001 Q10(a)

196


Airbags are actually inflated by nitrogen gas produced from a rapid decomposition of sodium azide (NaN3). This gas inflates the bag which bursts out. About a second later, the bag is already deflating (it has holes in it) in order to get out of the way.

This is the equation for the decomposition of sodium azide:

2NaN3(s) 2Na(s) + 3N2(g)

The process, from the initial impact of the crash to full inflation of the airbag, takes only about 40 milliseconds.

When the car decelerates very quickly, as in a head-on crash, the crash sensor would send an electric spark to the inflator. This activates the decomposition of the sodium azide sealed inside the inflator. The nitrogen gas produced inflates the airbag, cushioning the driver and reducing the likelihood of injury.

Questions

Airbags have been in the news in recent years because of deaths resulting from airbag deployments in relatively minor low-speed crashes.

1 Discuss how airbag deployments may lead to death.

2 Suggest how the airbag design may be improved so as to minimize such accidents.

You may refer to the following website for some ideas: http://www.aa1car.com/library/airbag01.htm

Fig. 18.25 In case of a collision, the crash sensor sends an electric spark to the inflator, setting off a chemical reaction that produces nitrogen gas, which inflates the airbag

Refer to page T68.

197


197

1 Rate of a reaction = change in concentration (or amount) of a reactant or a product

time

2 The instantaneous rate at time t is equal to the slope of the tangent to a concentration-time curve at that particular time t. The steeper the slope, the higher the rate.

3 For the general reaction,

aA + bB cC + dD

the rate of reaction can be expressed as:

Rate = – 1

a

d[A]

dt = –

1

b

d[B]

dt =

1

c

d[C]

dt =

1

d

d[D]

dt

198


198

4 Methods to follow the progress of a reaction include:

a) measuring the volume of a gaseous product;

b) measuring the loss in mass of the reaction mixture;

c) measuring the pressure of the reaction mixture;

d) measuring the colour intensity of the reaction mixture;

e) measuring the light transmittance of the reaction mixture;

f) titrimetric analysis.

5 In most cases, the rate of a reaction increases when the concentration of a reactant is increased.

6 The rate of a reaction increases when the surface area of a solid reactant is increased.

7 In most cases, the rate of a reaction increases when the temperature is increased.

8 A catalyst is a substance which alters the rate of a reaction without itself undergoing any permanent chemical changes.

199


199

Note: The symbol indicates the level of difficulty of a question.

Part I Knowledge and understanding

1 Construct a concept map using the following concept words. You may also add other concept words to your concept map.

reaction rate, measuring volume of gas, measuring loss in mass, measuring pressure, measuring colour intensity, measuring light transmittance, titrimetric analysis, concentration of reactant, surface area of solid reactant, temperature, presence of catalyst.

2 Ethanol is produced by fermentation of glucose. The equation is:

C6H12O6(aq) 2C2H5OH(aq) + 2CO2(g) glucose ethanol

How is the rate of formation of ethanol related to the rate of consumption of glucose?

3 A flask containing excess 1 mol dm–3 hydrochloric acid was placed on an electronic balance. Some zinc powder was added into the flask. The mass of the flask and its contents was plotted against time as shown in the graph below.

a) Write a chemical equation for the reaction in the flask.

b) Why did the mass of the flask and its contents decrease?

c) Referring to the curve, which point (A, B or C) represents the most rapid reaction? Explain your answer.

Refer to page T69.

Rate of formation of ethanol = d[C2H5OH(aq)]

dt

Rate of consumption of glucose = – d[C6H12O6(aq)]

dt

Rate = – d[C6H12O6(aq)]

dt =

1

2

d[C2H5OH(aq)]

dt

Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

Hydrogen gas escapes.

Point A represents the most rapid reaction because the tangent to the curve at point A is the steepest.

200


200

4 Consider the decomposition of a dicarboxylic acid, CO(CH2COOH)2, in aqueous solution.

CO(CH2COOH)2(aq) CH3COCH3(aq) + 2CO2(g)

Titrimetric analysis can be used to follow the progress of the reaction. Briefly describe the procedure.

5 a) Draw a labelled diagram of a colorimeter and briefly describe how a colorimeter can be used to follow the progress of a reaction.

b) Compare with titrimetric analysis, suggest ONE advantage of using physical methods to follow the progress of a reaction.

6 Some experiments were conducted to measure the time taken for whole and crushed antacid tablet to react with excess hydrochloric acid of different concentrations. The volume of acid used was the same in each case. Conditions of the experiments are shown in the table below.

Experiment number Antacid tabletConcentration of acid

(mol dm–3)Initial temperature of

acid (°C)

1 whole 1.0 23

2 whole 1.0 30

3 whole 2.0 23

4 crushed 1.5 23

5 crushed 2.0 23

a) Name one active ingredient in the antacid tablet.

b) Write a chemical equation for the reaction between the active ingredient stated in (a) with hydrochloric acid.

c) Which pair of experiments can be used to compare the effect of each of the following factors on the rate of the reaction?

i) Surface area of antacid tablet

ii) Concentration of the acid

iii) Temperature

Part II Multiple choice questions

7 Which of the following reactions is the slowest?

A Reaction between dilute hydrochloric acid and copper(II) carbonate. B Reaction between lead(II) nitrate solution and sodium chloride solution. C Burning of fireworks. D Decay of fruit.

Withdraw small samples of the reaction mixture at regular time intervals.

Quench the reaction by running each sample into an excess of ice-cold water.

Determine the concentration of dicarboxylic acid in each sample by titration with standard sodium hydroxide solution.

Refer to page T70.

Magnesium hydroxide (Other reasonable answers are acceptable.)

Experiments 3 and 5

Experiments 4 and 5

Experiments 1 and 2

Mg(OH)2(s) + 2HCl(aq) MgCl2(aq) + 2H2O(l)

201


201

8 Excess marble chips (calcium carbonate) were added to 25 cm3 of 2 mol dm–3 hydrochloric acid. Which measurement, taken at regular time intervals and plotted against time, will give the graph shown below?

A Temperature B Volume of gas produced C pH of solution D Mass of the beaker and contents

9 The graph shows the variation of concentration of a reactant with time as a reaction proceeds.

What is the average reaction rate during the first 20 s?

A 0.0025 mol dm–3 s–1

B 0.0050 mol dm–3 s–1

C 0.0075 mol dm–3 s–1

D 0.0150 mol dm–3 s–1

Calcium carbonate reacts with dilute hydrochloric acid to give carbon dioxide

gas.

As carbon dioxide escapes, the reaction mixture gets lighter as the reaction

proceeds.

Hence the mass of the beaker and contents decreases as the reaction

proceeds.

Average reaction rate during the first 20 s

= – (0.050 – 0.200) mol dm

–3

20 s

= 0.0075 mol dm–3

s–1

202


202

10 In an experiment to study the reaction between marble chips and dilute hydrochloric acid, the following graph is obtained:

Which of the following statements best explains why the slope of the graph rises gently at first and then steeply beyond point X?

A The marble chips are covered by a layer of oxide which dissolves slowly in the acid. B The carbon dioxide initially formed dissolves in the acid. C The acid is cold at the start of the experiment. D The acid dissociates completely as the reaction proceeds.

11 A student carried out an experiment to investigate the effect of particle size on reaction rate. The reaction chosen was dilute hydrochloric acid with iron(II) sulphide, a reaction producing a poisonous soluble gas. Which of the following experimental set-ups is suitable?

A B

C D

Option A — A soluble gas should NOT be collected over water.

Options B and C — Poisonous gas should NOT be released into

the laboratory.

203


203

12 Two experiments were carried out separately to investigate the rate of reaction of powdered magnesium with excess dilute sulphuric acid. The same mass of magnesium was used in both experiments. The volume of hydrogen evolved was measured at regular time intervals.

The graph shows the results of both experiments.

Which conditions were used in Experiment 2?

Temperature Concentration of acid

A Lower than Experiment 1 same as Experiment 1 B Higher than Experiment 1 higher than Experiment 1 C Same as Experiment 1 same as Experiment 1 D Same as Experiment 1 higher than Experiment 1

13 Which of the following statements concerning 20 cm3 of 1 M CH3COOH and 10 cm3 of 1 M H2SO4 is correct?

A They have the same pH values. B They have the same electrical conductivity. C They react with magnesium at the same rate. D They require the same number of moles of sodium hydroxide for complete neutralization. (HKCEE 2005)

14 Which of the following statements concerning a catalyst are correct?

(1) A catalyst can alter the rate of reaction. (2) The mass of a catalyst remains unchanged at the end of the reaction. (3) A catalyst should be in the same physical state as the reactants.

A (1) and (2) only B (1) and (3) only C (2) and (3) only D (1), (2) and (3) (HKCEE 2006)

Option A — The rate of reaction in Experiment

2 was lower when using a lower

temperature.

Option B — The rate of reaction in Experiment

2 would be higher when a higher

temperature and an acid of higher

concentration were used.

Option D — The rate of reaction in Experiment 2

would be higher when an acid of higher

concentration was used.

204


204

Part III Structured questions

15 Ammonia reacts with oxygen in the presence of platinum to give nitrogen monoxide.

Pt 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)

NH3(g) and O2(g) are allowed to react in a vessel of constant volume. Find the rate of consumption of O2(g) if the rate of formation of NO(g) is 1.24 x 10–4 mol dm–3 s–1.

(HKASLE 2006)

16 The oxidation of iodide ion by arsenic acid, H3AsO4, is described by the following equation:

3I–(aq) + H3AsO4(aq) + 2H+(aq) I3–(aq) + H3AsO3(aq) + H2O(l)

a) How is the rate of formation of I3–(aq) related to the rate of consumption of I–(aq)?

b) The rate of consumption of I–(aq) is 4.80 x 10–4 mol dm–3 s–1 at a certain time interval.

i) What is the rate of formation of I3–(aq) at that time interval?

ii) What is the rate of consumption of H+(aq) at that time interval?

17 For each of the following reactions, suggest with reason(s) one method that can be used to follow the progress of the reaction.

a) Fermentation of glucose

C6H12O6(aq) 2C2H5OH(aq) + 2CO2(g)

b) The hydrolysis of an ester 4-nitrophenyl ethanoate in an alkaline solution

O

CH3C O

O

C –OH3C

NO2(l) + 2OH–(aq)

NO2(aq) + H2O(l)

colourless

yellow

O–(aq) +

c) The thermal decomposition of sulphur dichloride dioxide

SO2Cl2(g) SO2(g) + Cl2(g)

d) The acid catalyzed hydrolysis of methyl ethanoate

CH3COOCH3(l) + H2O(l) CH3COOH(aq) + CH3OH(aq)

Refer to page T70.

Carbon dioxide gas is given off in the reaction.

Follow the progress of the reaction by measuring the volume of carbon

dioxide formed. / by passing the carbon dioxide gas into limewater and

measuring the light transmittance of the limewater.

A yellow species is formed in the reaction.

Follow the progress of the reaction by measuring the absorbance of

the reaction mixture using a colorimeter.

1 mole of SO2Cl2 gas decomposes to give 2 moles of gas.

Follow the progress of the reaction by measuring the pressure of the reaction mixture.

Withdraw small samples of the reaction mixture at

regular time intervals.

Quench the reaction by running each sample into an excess of ice-cold water.

Determine the concentration of ethanoic acid in each sample by titration with standard sodium hydroxide solution.

When the hydrolysis is complete, carry out a final titration to find out the amount of alkali needed to neutralize the acid

catalyst present.

Answers for the HKCEE and HKALE questions are not provided.

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205

The following graph was obtained from experiments to find the effect of pH on the efficiency of the catalyst.

18 The decomposition of hydrogen peroxide solution into water and oxygen can be catalyzed by a catalyst contained in potatoes.

2H2O2(aq) 2H2O(l) + O2(g)

The progress of the reaction can be followed by recording the mass over a period of time.

XXX.XX

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206

a) Calculate the average rate of reaction over the first 20 minutes, in g min–1, for the experiment at pH 10.

b) Suggest how the rate of reaction at a particular time can be determined, for the experiment at pH 7.

c) Suggest another way of following the progress of the reaction. Draw a labelled diagram of the experimental set-up.

d) Suggest ONE advantage of using a data-logger in this experiment.

19 The effect of temperature changes on reaction rate can be studied using the reaction between ethanedioic acid solution and acidified potassium permanganate solution.

5(COOH)2(aq) + 6H+(aq) + 2MnO4–(aq) 2Mn2+(aq) + 10CO2(g) + 8H2O(l)

a) Describe how the reaction time can be measured.

b) The headings for a set of results are shown below:

Temperature of reaction (°C) Reaction time (s)

Describe how the rate of reaction can be obtained from the experimental results.

18 a) Average rate of reaction = – (159.68 – 159.99) g

20 min

= 0.0155 g min–1

b) Draw a tangent to the curve at the particular time.

Determine the slope of the tangent.

c) Measure the volume of gas evolved.

hydrogen peroxide solution

potato discs

gas syringe

0

d) As the change in the mass is very small in this experiment, the use of a data-logger can give more accurate results.

The time taken for the solution to become colourless.

Rate of reaction α 1

time

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207

20 Dinitrogen pentoxide, N2O5, decomposes when heated according to the equation:

2N2O5(g) 4NO2(g) + O2(g)

The concentration of N2O5(g) at a temperature of 52 °C was measured at different times, and the results are shown in the graph below:

a) Calculate the average rate of decomposition of N2O5(g) during the time interval 20 – 30 s.

b) From the graph, determine the instantaneous rate of decomposition of N2O5(g) at 10 s.

c) Find the instantaneous rate of formation of NO2(g) at 10 s.

d) Copy the graph and sketch the variation of concentration of O2(g) with time during the experiment.

20 a) Average rate of decomposition of N2O5

= – (1.20 x 10

–2 – 1.80 x 10

–2) mol dm

–3

(30 – 20) s

= 6.00 x 10–4

mol dm–3

s–1

b) Instantaneous rate of decomposition of N2O5(g) at 10 s

= – (1.00 x 10

–2 – 3.50 x 10

–2) mol dm

–3

(26.0 – 2.5) s

= 1.06 x 10–3

mol dm–3

s–1

c) Instantaneous rate of formation of NO2(g)

= d[NO2(g)]

dt

Rate = – 1

2

d[N2O5(g)]

dt =

1

4

d[NO2(g)]

dt

∴ instantaneous rate of formation of NO2(g) at 10 s

= 2 (1.06 x 10–3

mol dm–3

s–1

)

= 2.12 mol dm–3

s–1

(26.0 – 2.5) s

(1.00 x 10–2

– 3.50 x 10–2

) mol dm–3

d) 2 moles of N2O5 decompose to give 1 mole of O2. The

following table shows the concentrations of N2O5 and O2

as a function of time.

Time(s)

Concentration

of N2O5(g)

(mol dm–3

)

Change in concentration

of N2O5(g)

(mol dm–3

)

Concentration

of O2(g)

(mol dm–3

)

0 4.00 x 10–2

0 0

5 3.30 x 10–2 4.00 x 10

–2 – 3.30 x 10

–2

= 0.70 x 10–2 0.35 x 10

–2

10 2.70 x 10–2 4.00 x 10

–2 – 2.70 x 10

–2

= 1.30 x 10–2 0.65 x 10

–2

15 2.20 x 10–2 4.00 x 10

–2 – 2.20 x 10

–2

= 1.80 x 10–2 0.90 x 10

–2

20 1.80 x 10–2 4.00 x 10

–2 – 1.80 x 10

–2

= 2.20 x 10–2 1.10 x 10

–2

30 1.20 x 10–2 4.00 x 10

–2 – 1.20 x 10

–2

= 2.80 x 10–2 1.40 x 10

–2

40 0.80 x 10–2 4.00 x 10

–2 – 0.80 x 10

–2

= 3.20 x 10–2 1.60 x 10

–2

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21 An experiment was carried out to study the rate of the following reaction:


A sample of marble chips was allowed to react with 0.1 M hydrochloric acid, which had been saturated with carbon dioxide. The graph below shows the experimental results obtained.

a) i) Suggest how hydrochloric acid can be saturated with carbon dioxide.

ii) If the hydrochloric acid used has not been saturated with carbon dioxide, different experimental results would be obtained. Copy the graph and sketch the results that would be obtained.

b) Suggest how the rate of the reaction at a particular time can be determined from the graph. (HKASLE 2003)

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209

22 X2(g) undergoes decomposition according to the following equation:

X2(g) 2X(g)

In an experiment to study the decomposition of X2(g), 0.100 mol of X2(g) was charged into a closed container of volume 1 dm3 kept at a constant temperature. The graph below shows the variation of the concentration of X2(g) in the container with time.

a) From the graph, calculate the average rate of decomposition of X2(g) in the time interval from the start of the experiment to the 40th second.

b) Copy the graph and sketch the variation of the concentration of X(g) with time during the experiment.

(HKASLE 2005)

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210

23 Bromine and methanoic acid in aqueous solution react according to the following equation:

Br2(aq) + HCOOH(aq) 2Br–(aq) + 2H+(aq) + CO2(g)

In a certain experiment, the initial concentration of Br2(aq) was 0.010 mol dm–3 and the volume of the reaction mixture is 10.0 cm3. The progress of the reaction was followed by a colorimeter. The concentration of Br2(aq) was determined at different times. The graph shows the results obtained.

a) Explain how the progress of the reaction could be followed by a colorimeter.

b) Give a reason why measuring the volume of CO2(g) is NOT a suitable method for following the progress of the reaction.

c) How could you obtain the concentration of bromine from the colorimeter reading?

d) i) How was the rate of consumption of Br2(aq) related to the rate of formation of Br–(aq)?

ii) From the graph, determine the instantaneous rate of consumption of Br2(aq) at 150 s.

iii) Find the instantaneous rate of formation of Br–(aq) at 150 s.

Refer to page T71.

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211

24 When an acid is added to bleach, chlorine gas is produced. The diagram below shows the experimental set-up for carrying out this reaction.

The reaction was carried out using an excess of acid (10 cm3) at 25 °C. The chlorine was collected and its volume recorded every half minute. The graph below shows the results obtained.

a) The reaction had not finished after 2.5 minutes. Explain how you know this is true.

b) The reaction finished within 5 minutes. Copy and continue the graph to show this. Use the graph to estimate the total time for the reaction.

c) Keeping all other conditions unchanged, the experiment was repeated

i) using 20 cm3 instead of 10 cm3 of acid; and

ii) at 50 °C instead of 25 °C

State the change in the reaction rate, if any, in each case.

25 A student performed an experiment to investigate the rate of reaction between zinc and an acid. 6 g of zinc granules was added to a conical flask containing 100 cm3 of 2 M hydrochloric acid at 20 °C. Afterwards the experiment was repeated with the following changes. In each case, state and explain whether the expected reaction rate would increase or decrease.

a) 6 g of zinc powder was used instead of zinc granules.

The volume of chlorine was still increasing after 2.5 minutes.

The rate was the same (as the concentration of the acid remained unchanged).

The rate was higher (as the temperature was increased).

The total time for the reaction is about 4.5 minutes.

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212

b) 100 cm3 of 2 M ethanoic acid was used instead of hydrochloric acid.

c) The temperature was raised to 50 °C. (HKASLE 2006)

26 A student investigated the effect of temperature on the rate of reaction. The reaction between sodium thiosulphate solution and dilute sulphuric acid was used.

Na2S2O3(aq) + H2SO4(aq) Na2SO4(aq) + SO2(aq) + H2O(l) + S(s)

The following procedure was followed.

Step 1 Measure out the following solutions.

Step 2 Heat the beaker in a water bath to about 30 °C.

Step 3 Place the beaker on a card with an X on it. Measure the exact temperature of the sodium thiosulphate solution.

Step 4 Add the acid to the sodium thiosulphate solution.

Step 5 Measure the time for the X to be obscured by the sulphur produced.

Step 6 Repeat the experiment once at about 40 °C.

Suggest THREE ways of improving the procedure for the investigation.

27 Magnesium reacts with dilute hydrochloric acid to produce hydrogen gas. Describe how you can carry out a fair experiment to investigate how the concentration of the acid affects the rate of the reaction. State the results you expect.

(For this question, you are required to give answers in paragraph form. Use equations, diagrams and examples where appropriate.)

Refer to page T72.

Refer to page T73.

cs_unit 18(e)

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