1.1 rate of reaction

Chapter 1

Rate of Reaction

Chapter 1: Rate of ReactionChapter 1: Rate of Reaction1.1 Rate of Reaction

Which reaction is faster?

The meaning of rate of reactionThe meaning of rate of reaction

1 A chemical reaction takes place when reactants are mixed together under suitable conditions.

2 The speed of a chemical reaction

is called the rate of reaction.

3. The rate of reaction is inversely proportional to the time taken for the reaction to be completed.

Rate of reaction

The reaction is fast if it takes a short time to complete. Conversely, the reaction is

slow if it takes a long time for the reaction to complete.

takenTime

4 During a chemical reaction, the reactants are used up as the products are formed. Thus, the amounts of reactants decrease (Figure 1.1(a)) while the amounts of products increase as the reaction proceeds (Figure 1.1(b)).

5 Therefore, the rate of reaction can be determined in one of the following ways:

(a) the rate of disappearance of a reactant, or (b) the rate of formation of a product

6 (a) The rate of reaction can be defined as the amount of a reactant used up per unit time.

Rate of reaction = TakenTime

upusedtreacofAmount tan

6 (b) The rate of reaction can also be defined as the amount of a product obtained per unit time.

Rate of reaction =TakenTime

obtainedproductofAmount

7 Methods of measuring reaction rates (a) The amount of a reactant used up or a product

obtained can be measured in terms of the mass of the substance or the concentration of the substance.

7 Methods of measuring reaction rates (b) For chemical reactions that produce gases, the

rate of reactions can be measured as the volume of gas produced per unit time.

Rate of reaction =TakenTime

producedofVolume

8 Reaction between calcium carbonate and dilute hydrochloric acid.

(a) The reaction between calcium carbonate (marble chips) and dilute hydrochloric acid can be represented by the equation:

(b) During the reaction, the following observable changes take place.

(i) The mass of calcium carbonate (the reactant) decreases.

(ii) The concentration of hydrochloric acid (the reactant) decreases.

(iii) The volume of carbon dioxide (the product) produced increases.

(c) Thus, the rate of reaction between calcium carbonate and hydrochloric acid can be determined by measuring

(i) the decrease in mass of calcium carbonate per unit time, or

(ii) the increase in volume of carbon dioxide per

unit time.

9 (a) Besides changes in the mass of the reactants or the changes in the volume of the gaseous products, other changes that can be observed are as follows:

(i) changes in colour (ii) formation of precipitates

(b) Therefore, the time taken for the colour of a reactant to change or a certain amount of precipitate to form can be used to measure the rate of reaction.

10 The units used for the rate of reaction will depend on the changes measured.

For example (a) cm3 per unit time (second or minute) for a

gas evolved

For example (b) g per unit time or mol per unit time for a

solid reactant

For example (c) mol dm-3 per unit time for a reactant in

aqueous solution

11 Different chemical reactions take place at different rates. Some reactions occur very rapidly and some very slowly. Table 1.1 shows some examples of very fast reactions and very slow reactions.

Fast SlowCombustion RustingPrecipitation of silver chloride PhotosynthesisReaction between reactive metal and water

Fermentation

Measuring reaction ratesMeasuring reaction rates

1. The rate of reaction can be express in two ways:

(a) the average rate of reaction over a period of time, or

(b) the rate of reaction at any given time.

2 The average rate of reaction is the average of the reaction rates over a given period of time. We can measure the average rate of reaction by measuring the change in amount (or

concentration) of a reactant or a product over a period of time.

Example 1 A piece of magnesium ribbon weighing 0.1 g is

added to dilute hydrochloric acid. After 5 seconds, all the magnesium had dissolved. What is the average rate of reaction?

average rate of reaction=10.1

Example 2 Calcium carbonate reacts with dilute hydrochloric

acid according to the equation: CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) +

CO2(g) After 1.2 minutes, the volume of gas produced is

100 cm3. Calculate the average rate of reaction in the units of

(a) cm3 min-1, (b) cm3 s-1.

Example 2Solution:

min83.3cm (min)2.1

)100(cm

takentime

produced CO of volume ratereaction average

Example 2Solution:

takentime

producedCOofvolumerateAverage 2

39.172

100 scms

3 The rate of reaction at any given time is the actual rate of reaction at a given time. The reaction rate at any given time is also known as the instantaneous rate of reaction.

4. The rate of reaction at a given time can

determined by measuring the gradient of the

graph mass of reactant against time (Figure1.2)

Determining the gradient of the tangent time t: Rate of reaction at time t (gs-1) = Gradient of the graph

Solving numerical problems Solving numerical problems involving rate of reactioninvolving rate of reaction

Example 3: Hydrogen peroxide decomposes according to the

equation: 2H 2O2 (aq) 2H2O(l) + O2(g) The results of an experiment on the decomposition

of hydrogen peroxide are given below.

Calculate the rate of reaction at 40 seconds in the units of (a) cm3 s-1, (b) cm3 min -1.

Time(s) 0 15 30 45 60 90Volume of O2(cm3)

0 16 30 40 48 56

Example 3: Solution (a) The rate of reaction at 40 s = Gradient at 40 s

1370.01858

Example 4: Solution (b) time =

Rate in cm3 min-1

Rate of reaction in cm3 min-1

= min3

228 3cm

= 42cm3 min-1

(58 18) 40 2min min min

60 60 3

5 Analysing a reaction rate curve: (a) The steeper the gradient, the faster the

rate of reaction.

Steeper

5 Analysing a reaction rate curve:

(b) Figure 1.5 shows that the rate of reaction at t2,

is slower than the rate of reaction at t1. The

gradient at t2 is less steep than the gradient at t1. As the reaction proceeds, the curve becomes less steep because the rate of reaction decreases.

6 Comparing reaction rates at a give time

Activity: To find the reaction rates at (a) 90 s, (b)180 s a (c) the average rate of the reaction between zinc and dilute sulphuric

Apparatus: Conical flask, measuring cylinder, delivery

tube, burette, basin, retort stand, retort clamp and stopwatch.

Materials: Granulated zinc and 0.3 mol dm-3 sulphuric acid.

Procedure: 1 The burette is filled with water and inverted over a

basin of water.

Procedure: 2 Using a measuring cylinder, 20.0 cm3 of 0.3 mol

dm-3 sulphuric acid is measured out and poured into a conical flask.

Procedure: 3 5.0 g of granulated zinc are then added to the

sulphuric acid in the conical flask.

Procedure: 4 The conical flask is then closed and the hydrogen

gas produced is collected in the burette by the displacement of water as shown in Figure 1.5.

Procedure: 5 The stopwatch is started immediately. 6 The volume of hydrogen gas collected in the burette is

recorded at 30-second intervals.

Results:

Calculation: (a) The rate of reaction at 90 s

= slope at 90 s

)30180(

)2052( 3

13213.0 scm

Calculation: (b) The rate of reaction at 180 s = slope at 180 s

)18240(

)3048( 3

13081.0 scm

Calculation: (c) The average rate of reaction

157.0300

taken timetotal

produced H of volumetotal

Conclusion:

The rate of reaction decreases as the reaction proceeds.

Activity: To measure the rate of reaction between calcium carbonate (CaCO3) and excess hydrochloric acid

Apparatus: Conical flask, electronic balance, measuring

cylinder and stopwatch.

Material: Calcium carbonate (CaCO3) pieces, 2.0 mol dm-3

hydrochloric acid and cotton wool.

Procedure 1 Using a measuring cylinder, 50 cm3 of 2 mol dm-3

hydrochloric acid are measured out and poured in a dry conical flask. The mouth of the conical flask covered with some of cotton wool. The cotton wool is inserted into the mouth of the conical flak is prevent liquid from splashing out during the reaction.

Procedure 2 The conical flask is placed on the electronic balance

as shown in Figure 1.6.

Procedure 3 The mass of conical flask and its contents (calcium

carbonate, hydrochloric acid and cotton wool) is recorded.

Procedure 4 The calcium carbonate is then transferred to the

hydrochloric acid in the conical flask and the stopwatch is started immediately.

Procedure 5 The mass of the conical flask (and its contents) is

recorded at one-minute intervals.

Results:

Time(s) 0 1 2 3 4 5 6 7 8Mass of conical flask + content(g)

60.0 59.1 58.3 57.9 57.4 57.0 56.8 56.5 56.3

Based on the experimental results, a graph of the mass of conical flask and its contents against time is plotted (Figure 1.7)

Calculation The average rate of reaction for the first minute. Decrease in mass = mass of carbon dioxide produced = (60.0-59.1) g = 0.9g Average rate of reaction for the first minute

12 min9.0min0.1

takentime

produced CO of mass gg

Time(s) 0 1 2 3 4 5 6 7 8Mass of conical flask + content(g)

60.0 59.1 58.3 57.9 57.4 57.0 56.8 56.5 56.3

Calculation the average rate of reaction between 1.4 minutes and 2.2

minutes Rate of decrease in mass

1min625.04.12.2

3.588.58

Calculation (c) The reaction rate at the 5th minute =Gradient of the graph at 5.0 minutes =

1min306.03.6

1.1Gradient

minutes6.34.30.71.14.565.57

Conclusion: The rate of reaction decreases as the reaction

proceeds. Finally, the reaction will stop when all the calcium carbonate added have reacted.

Example 4: 3.0 g of excess marble (CaCO3) are added to 100 cm3 of

dilute hydrochloric acid. Figure 1.8 shows the graph of volume of carbon dioxide produced against time.

Example 4: Calculate (a) the average rate of reaction (b) the concentration of hydrochloric acid in mol dm-3. (1 mole of any gas occupies 24 dm3 at room conditions).

Example 4: Solution (a) Total volume of carbon dioxide evolved = 360 cm 3

Time taken = 8.0 minutes Average rate of reaction

= 13 min458

360 cm

Example 4: Solution (b) Number of moles of CO2 evolved

CaCO3 + 2HCl CaCl2 + H2O + CO2

Mole ratio of HCl : CO2 =2: 1

mol015.0100024

Example 4: Solution (b) According to the equation, number of moles of

hydrochloric acid used

CaCO3 + 2HCl CaCl2 + H2O + CO2

2 moles HCl produces 1 mol CO2

? moles HCl produces 0.015 mol CO2

= 2 x 0.015 = 0.03 mole

Example 4: Solution (b) According to the equation, number of moles of

hydrochloric acid used = 2 x 0.015 = 0.03 mole

Concentration of hydrochloric acid, M

3.01.0

)dmin volume(

moles ofnumber

V = 100 cm3 =0.1 dm3

1.1 rate of reaction

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