jacaraonda chapter 17

7/21/2019 jacaraonda chapter 17

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COMBUSTION

AND REACTION

KINETICS&,Introduction

Combustion is an important chemical process. The burning

of fuels such as coal, oil and natural gas provides heat to

keep us warm and to power machines. Combustion of fuels

requires oxygen. If oxygen supply is restricted, gases such

as carbon monoxide, which pollute the environment, are

released. In this chapter, we will investigate combustion

and the rate at which combustion reactions and otherreactions occur.

In this chapter

17.1 Combustion page 314

17.2 Reaction kinetics page 327

8]VeiZg

Figure 17.1

Combustion reactions require fuel, oxygen and an

ignition source.


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314 ENERGY

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Chemical changeLet us consider the indicators of chemical change and the chemical bonds

that are broken and formed as chemical reactions proceed.

Indicators of chemical change

In this course, you have investigated many different types of chemicalreaction and physical change. Physical changes do not involve theproduction of new materials. For example, when wax melts, the liquid

wax is still composed of the same molecules as the solid wax. This physicalchange is readily reversed by removing heat so that the solid wax reforms.Melting wax requires much less heat than would be required for a chemicalreaction involving the wax molecules.

Chemical changes involve the production of new materials. When waxis burnt in air (as in a burning candle), new chemical compounds, such ascarbon dioxide and water, are formed. Mixing carbon dioxide and waterdoes not cause wax and oxygen to re-form.

In general, the indicators of chemical change are:

Reactants are permanently converted into new products with a differentappearance or property (e.g. different colour).

Chemical reactions are difficult to reverse.

Large energy changes occur.

Chemical reactions involve breaking

and making bonds

Chemical changes occur because some or all of the chemical bonds in

the reactants are broken; the released atoms, molecules or ions combine

to form new products by forming new bonds. Breaking chemical bondsrequires energy, which comes from the thermal motion of the particles.

Energy is transferred from one particle to another during particle

&,#&

Remember

Before beginning this section,you should be able to: identify and use the IUPAC

nomenclature for describingstraight-chained alkanes andalkenes from C1 to C8

recall the safety issuesassociated with the storageof alkanes from C1 toC8 in view of their weakintermolecular forces(dispersion forces).

Key content

By the end of this section, youshould be able to: describe the indicators of

chemical reactions identify combustion as an

exothermic chemical reaction

outline the changes inmolecules during chemicalreactions in terms of bond-breaking and bond-making

explain that energy is requiredto break bonds and energyis released when bonds areformed

describe the energy neededto begin a chemical reactionas activation energy

describe the energy profilesfor endothermic andexothermic reactions

explain the relationshipbetween ignition temperatureand activation energy

identify the sources ofpollution that accompanythe combustion of organiccompounds, and explain howthese can be avoided

describe chemical reactionsby using full balanced chemicalequations to summariseexamples of complete andincomplete combustion

solve problems and performa first-hand investigation tomeasure the change in masswhen a mixture such as woodis burnt in an open container

identify the changes of stateinvolved in combustion of aburning candle

perform first-handinvestigations to observeand describe examples ofendothermic and exothermicchemical reactions.

Figure 17.2 Burning wood is an example of a chemical change.


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CHAPTER 17 COMBUSTION AND REACTION KINETICS 315

collisions. When new bonds are formed, energy is released to the systemand increases the kinetic energy of the particles.

The strength of a chemical bond can be measured by the bond energyrequired to break the bond. An OH covalent bond is stronger than anHH bond. The bond energy of OH is 463 kJ per mole of bonds whilethe bond energy of HH is 436 kJ/mol. Triple carboncarbon bondsare much stronger than single carboncarbon bonds. Table 17.1 lists the

bond energies of some common covalent bonds.

Table 17.1Bond energies of some covalent bonds

Bond Bond energy (kJ/mol)

CC 346

C=C 614

C C 839

CH 414

HH 436

OH 463

O=O 498

CO 358

C=O* 745

(*for C=O in carbon dioxide)

Bond energy tables can also be used to determine the amount of energyreleased when covalent bonds form. Thus, when one mole of HH bondsform, 436 kJ of energy is released.

Calculate the energy change per mole of oxygen when hydrogen gasburns in oxygen gas to form water vapour.

Step 1:Write the balanced equation for the combustion reaction.

2H2(g) + O2(g) 2H2O(g)

Step 2:Rewrite the equation to show the covalent bonds of thereactants to be broken and bonds of products to be formed.

HH + HH + O=O HOH + HOH

Step 3:Identify the bonds to be broken and those to be formed. Bonds broken: 2 HH bonds

1 O=O bond

Bonds formed: 4 OH bondsStep 4:Use table 17.1 to calculate the energy required to break the

bonds in the reactants: 2 bond energy of HH = 2 436 = 872 kJ 1 bond energy of O=O = 498 kJ

Total energy to break bonds = 872 + 498 = 1370 kJ

Step 5:Use table 17.1 to calculate the energy released when theproducts form.

4 bond energy of OH = 4 463 = 1852 kJ

SOLUTION

SAMPLE PROBLEM 17.1

bond energy:the enthalpychange associated with breakingcovalent bonds in one mole ofa gaseous substance to producegaseous fragments

bond energy:the enthalpychange associated with breakingcovalent bonds in one mole ofa gaseous substance to producegaseous fragments


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316 ENERGY

Step 6:Calculate the overall energy change.

The process of breaking the bonds of the reactants isendothermic. The process of forming bonds in the products isexothermic. By convention, the energy released in anexothermic process is assigned a negative value.

Thus, the energy change = (+1370) + (1852) = 482 kJ

Thus, 482 kJ of energy is released per mole of oxygenreacting in the combustion of hydrogen in oxygen to formwater vapour.

Figure 17.3Energy changes during bond

breaking and bond formation in the

combustion of hydrogen

4H(g) + 2O(g)

2H2(g) + O2(g)

+1370 kJ 1852 kJ

2H2O(g)

1000

0

1000

2000

H= 482 kJEnt

halpy(H)(kJ)

The processes of breaking and making bonds also occur in reactionsinvolving metals and ionic compounds. Because these materials consistof infinite crystalline lattices rather than individual molecules in thegaseous state, other ways are used to describe the energy changes involved.For example, the strength of ionic bonds in metallic oxides, sulfides andchlorides can be compared using data on the amount of energy required todissociate one mole of an ionic compound into its gaseous ions. This energyis called the lattice enthalpy. Table 17.2 compares the lattice enthalpies

of various ionic compounds. It suggests that the strength of ionic bondsvaries, and that metal oxide bonds are stronger than metal sulfide ormetal chloride bonds. We can also conclude that a large amount of energyis released when these metal and non-metal ions react in the gaseousphase to form ionic bonds in a crystal lattice.

Table 17.2Lattice enthalpies for ionic compounds (kJ/mol)

O2 S2 Cl

Na+ 2488 2199 788

K+ 2245 1986 718

Rb+ 2170 1936 693

Activation energyWhen a fuel such as hydrogen gas is mixed with oxygen gas at room tem-perature there is no observable reaction. This mixture of gases is stableat 25 C. However, if a spark or a flame is supplied, the mixture reactsexplosively. Why doesnt the reaction occur at 25 C? We can explain thisdifference in behaviour of the reacting particles in terms of energy.

The term activation energy is used to explain why many reactionsdo not proceed at low temperatures. Activation energy is the minimum

activation energy:the minimumenergy required by reactants inorder to react

activation energy:the minimumenergy required by reactants inorder to react


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amount of energy needed by the reactants to turn into products. In orderfor chemicals to react, they must collide with sufficient kinetic energy sothat bonds can be broken. At low temperatures, the reactants have lowkinetic energy. As the temperature rises, the particles move faster andfaster. Eventually, they have sufficient kinetic energy to react.

We can visualise activation energy as an energy barrier that separates thereactants and products. Sometimes we refer to this barrier as the activation

energy hill. Figure 17.4 shows an energy profile illustrating this ideafor an exothermic reaction. Reactants without sufficient energy to over-come this energy barrier will not react. A spark or flame in a hydrogenoxygen mixture would provide sufficient energy for some of the moleculesto cross the barrier and form products. When this happens, they releaseheat energy that raises the kinetic energy of other reactant molecules.Very soon, almost all the reactant molecules have sufficient energy tocross the barrier. The reaction becomes self-sustaining (a flame or sparkis no longer required) and the reaction occurs rapidly with considerableevolution of heat.

Energy profiles have several other features worth noting.

Endothermic and exothermic reactionsThe energy profiles for exothermic and endothermic reactions aredifferent. If the products have more energy than the reactants, the reactionis endothermic and the enthalpy change ($H) is positive. If the productshave less energy than the reactants, the reaction is exothermic and theenthalpy change is negative. These features are shown in figure 17.5.

Figure 17.5

Enthalpy change, activation

energy and activated complex

for (a) endothermic and

(b) exothermic reactions

(b) Exothermic reaction

Enthalpy

(H)

Reactants

Reaction coordinate

Products

Activated

complex

Products

Reactants

EA

H< 0Enthalpy

(H)

Reactants

Reaction coordinate

Products

Activated

complex

Products

Reactants

EA H > 0

(a) Endothermic reaction

Exothermic reactions are self-sustaining because the liberated heatprovides the activation energy required by other reactants in the mixture.

Endothermic reactions are not self-sustaining and energy is needed tokeep the reaction going. This may come from the kinetic energy of the

molecules in the system; this causes the reaction to slow down and the tem-perature of the system to decrease.

Activated complex

The activated complexis a transition state that exists at the top of the acti-vation energy hill. In this state, the bonds that held the reactants togetherare partially broken and new bonds holding the products together arepartially formed. This complex of atoms exists for a very short time beforeit starts to break up. At this point, there are two possible ways the reactioncould proceed:

Figure 17.4Energy profile showing the

activation energy barrier for an

exothermic reaction

Energy

Reaction coordinate

Activation energybarrier for the

forward reaction

Reactants

Products

Figure 17.4Energy profile showing the

activation energy barrier for an

exothermic reaction

Energy

Reaction coordinate

Activation energybarrier for the

forward reaction

Reactants

Products

activated complex:the (unstable)transition state formed duringa reaction (at the top of theactivation energy hill) thatbreaks down to form products

activated complex:the (unstable)transition state formed duringa reaction (at the top of theactivation energy hill) thatbreaks down to form products


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318 ENERGY

1. No reaction occurs because the activated complex breaks up and thereactants re-form.

2. Reaction occurs because the activated complex separates, formingproduct molecules.

A reaction has a greater chance of occurring when the system is suppliedwith energy in excess of the activation energy.

Activation energy for the forward and reverse reactionsNot all reactions occur in one direction only. In many reactions, theproducts recombine and the reactants re-form. To do this, the productmolecules must have sufficient energy to overcome the activation barrierfor the reverse reaction. The size of this barrier is different from the sizeof the barrier in the forward reaction, as shown in figure 17.6.

Figure 17.6

Energy profile showing the

activation energy for forward and

reverse reactions

Enthalpy

(H)

A+ B C+ D

C+ D

A+

BE

A(f)

EA(f) = Activation

energy for

forward reaction

EA(r) = Activation

energy for

reverse reaction

EA(r)

Combustion and ignition temperatureHeptane is a component of petrol. Mixtures of 17%v/v of heptanevapour in air are combustible; if there is too much heptane or too little air,the mixture will not combust. A combustible mixture of heptane vapour

and air is stable at room temperature. However, if it is heated, it willeventually reach a temperature at which combustion begins. Thisminimum temperature is referred to as the ignition temperature.

Once combustion of heptane begins, it quickly becomes self-sustainingas the heat liberated raises the temperature of the rest of the mixture abovethe ignition temperature. The energy equation for such a combustionreaction is:

C7H16(l) + 11O2(g) 7CO2(g) + 8H2O(l) $H= 4817 kJ/mol

Figure 17.7

Energy profile for the combustion

of heptane

5000

4000

3000

2000

1000

0

Enthalpy

(H)(k

J)

Reaction coordinate

Products

Reactants100

EA= +350 kJ

H= 4817 kJ

The activation energy for this reaction is 350 kJ. This means that thereactants must be heated to their ignition temperature or higher to provide

ignition temperature:theminimum temperature at whicha combustible fueloxidisermixture ignites spontaneously

ignition temperature:theminimum temperature at whicha combustible fueloxidisermixture ignites spontaneously

Exothermic

and endothermic

reactions

17.1 PRACTICALACTIVITIES

TEMPERATURE AND

REACTION RATE
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at least 350 kJ of energy. For this particular reaction, the ignition temp-erature is approximately 215 C. The enthalpy of combustion of heptane is4817 kJ/mol. Figure 17.7 shows the energy profile for the combustion ofheptane.

Ignition temperature is a measure of the activation energy of a reaction.The higher the activation energy, the higher the ignition temperature.Thus, fuels such as butane with a higher ignition temperature (405 C) than

heptane, also have a higher activation energy.Table 17.3 lists typical ignition temperatures of some fuels. Ignition tem-

perature depends on the methods used to measure it and on the composi-tion of the fuel (e.g. natural gas composition is quite variable).

Table 17.3Typical values for ignition temperatures of fuels

Fuel Ignition temperature (C)

hydrogen 585

natural gas 540560

petrol 390420

kerosene 380

diesel 300350

methane 580

butane 405

pentane 260

hexane 225

heptane 215

octane 206

Petrol has a higher ignition temperature range (390420 C) thandiesel (300350 C). The diesel vapourair mixture in a diesel engine iscompressed until its temperature reaches the ignition temperature. Com-pression heating is not used in a petrol engine, however, as the ignitiontemperature is too high. In a petrol engine, the petrol vapourair mixtureis ignited by a spark.

A mixture of hexane vapour (5 mL) and air (95 mL) is placed in a100 mL vessel and the temperature raised to 200 C. Use the followinginformation to determine whether the fuel vapouroxidiser mixturewill ignite in the absence of a flame at this temperature.

Flash point of hexane = 23 CIgnition temperature of hexane = 225 CCombustible mixture = 1.27.5% hexane in air

Step 1:Determine the % composition of the mixture.

% hexane vapour = 5/100 10 = 5%

Step 2:Determine whether this mixture is a combustible mixture.The composition lies between 1.2% and 7.5% so it is acombustible mixture.

SAMPLE PROBLEM 17.2

SOLUTION


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320 ENERGY

Step 3:Compare the ambient temperature with the ignitiontemperature and determine whether the mixture will ignite.

The ambient temperature is above flash point, the minimumtemperature at which a combustible fuel vapourair mixture willform. However, the ambient temperature is below the ignitiontemperature so the mixture will not ignite.

Incomplete combustion and pollutionThe release of waste and poisonous substances pollutes the environment,affecting all living things. The products of incomplete combustion of fuelsis an important example of pollution. Production of carbon dioxide is alsoof concern as there is considerable evidence that increasing carbon dioxidelevels in the atmosphere contributes to global warming by enhancing thegreenhouse effect.

Some pollutants released by the combustion of organic compounds(such as coal, oil and natural gas) include:

sulfur dioxide.Coal and oil often contain sulfur minerals. When these

fuels are burnt, the sulfur minerals also burn and release sulfur dioxide,a colourless, choking gas:

S(s) + O2(g) SO2(g)

Sulfur dioxide can be oxidised in the atmosphere to form sulfur trioxide.These oxides of sulfur combine with moisture in the air, forming acidrain, which can damage living things as well as the built environment.

nitrogen oxides.When fuels are burnt in air at a high temperature (e.g.in a petrol engine or coal-fired power station), the nitrogen in the airalso reacts with oxygen to form oxides of nitrogen, such as NO, NO2and N2O5:

N2(g) + O2(g) 2NO(g)

2NO(g) + O2(g) 2NO2(g)

Oxides such as NO2contribute to the formation of acid rain as well ascontributing to the formation of photochemical smog in cities.

carbon monoxide and unburnt hydrocarbons. When organic compounds,such as hexane, are burnt in a limited air supply, the combustionproducts include carbon monoxide as well as with the exhaust gases:

2C6H14(l) + 13O2(g) 12CO(g) + 14H2O(l)

Carbon monoxide is a very poisonous gas. It combines more readilywith haemoglobin in the blood than oxygen, leading to suffocation.

particulates. Coal-fired power stations and car engines emit ash andother fine particles, such as soot (carbon particles), into the air.

Burning natural gas in a Bunsen burner

You would have used a Bunsen burner many times in your chemistry lessons.The collar at the base of the burner can be rotated to allow differentamounts of air to be drawn into the gas before the mixture is ignited.When the hole in the collar is closed, a yellow, safety flame is produced.As the hole is opened gradually and more air mixes with the naturalgas, the flame changes from a luminous yellow to a mauve and, finally,

Figure 17.8

Black smoke is a sign of incomplete

combustion.

Figure 17.8

Black smoke is a sign of incomplete

combustion.


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Figure 17.9

Bunsen burner flames

Narrow mauve

outer zone

Yellow flame

Open air holeClosedair hole

Mauve

outer zone

Blue luminous

inner zone

Dark inner zone

(unburnt gas

and air)

a blue flame. The different coloured flames show whether combustion ofthe natural gas is complete or incomplete.

To achieve complete combustion, air must be drawn in through theopen hole in the collar as the gas enters the base of the burner through anarrow jet. The gas and air mixture moves through the barrel to the top ofthe burner where it is ignited. The resulting flame has a mauve outer zone(mantle), a central blue zone with a dark inner zone. The dark inner zonecontains unburnt gas and air and is quite cool. The surface of the bright

blue zone is the hottest region and is where complete combustion occurs.Complete combustion of methane produces carbon dioxide and watervapour. The mauve outer mantle is a cooler region where a variety of reac-tions occur; these reactions produce less heat than that produced by thecomplete combustion of methane in the bright blue zone. For example, inthe outer mantle, carbon monoxide formed from the incomplete combus-tion of methane reacts with oxygen to form carbon dioxide.

Complete combustion of methane:

CH4(g) + 2O2(g) CO2(g) + 2H2O(l) $cH= 892 kJ/mol

where $cHis the enthalpy of combustion.If the hole at the base of the barrel is closed, the flame becomes a

luminous yellow. The oxygen required for combustion diffuses into themethane from the surrounding air. This process of forming a combustiblemixture is not as efficient as that produced with the hole open. Theyellow flame indicates poor or incomplete combustion. It is not as hotas the blue or mauve flames because the methane burns to form carbonparticles (soot) and water vapour, as well as some carbon monoxide andcarbon dioxide. The carbon particles become incandescent and colourthe flame yellow, escaping from the combustion zone before they can allburn. The yellow flame has a narrow, blue-mauve outer edge where mostof the carbon monoxide burnt.

Incomplete combustion of methane:

2CH4(g) + 3O2(g) 2CO(g) + 4H2O(l) $cH= 609 kJ/mol

CH4(g) + O2(g) C(s) + 2H2O(l) $cH= 498 kJ/mol

2CO(g) + O2(g) 2CO2(g) $cH= 283kJ/mol

Combustion of solid and liquid fuels

Common solid fuels used in industry and power generation include coaland coke. Coke is made by heating coal in the absence of air to drive offvolatile materials. These solid fuels are not very volatile so they have to beheated to high temperatures to produce ignition. A continuous streamof air or oxygen helps to promote complete combustion and to raise theflame temperature.


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322 ENERGY

Complete combustion of solid fuel:

C(s) + O2(g) CO2(g) $cH= 394 kJ/mol

If the supply of oxygen is reduced, incomplete combustion results andthe flame temperature is much lower.

Incomplete combustion of solid fuel:

C(s) + O2(g) 2CO(g) $

cH= 111 kJ/mol

Waxy solids, such as paraffin wax, are used to make candles. Wax doesnot burn without a fibre wick to help produce a combustible mixtureof wax vapour and air. Practical activity 17.2 on page 337 deals with thechemistry of a burning candle.

Volatile, low molecular weight, liquid fuels (such as petrol) readilyvaporise to produce combustible mixtures with air. When sparked, themixture burns completely or incompletely, depending on the ratio of fuelvapour to air. Modern fuel injection systems in cars ensure the optimumfuelair ratio for maximum power and minimum pollution by carbonmonoxide emission. Catalytic converters in the exhaust system also reducecarbon monoxide emission by converting CO to CO2.

Complete combustion of octane (a petrol component):

2C8H18(l) + 25O2(g) 16CO2(g) + 18H2O(l) $cH= 5470 kJ/mol

Incomplete combustion of octane:

2C8H18(l) + 17O2(g) 16CO(g) + 18H2O(l) $cH= 3253 kJ/mol

Heavier liquid fractions, such as kerosene, do not ignite when a litmatch is applied to their surface unless a wick is present. Thus, kerosenelamps and home heaters use fibre wicks with high surface areas to producesufficient vapour to form combustible fuel vapourair mixtures.

Diesel fuel is also not as volatile as petrol. Therefore, to reduce airpollution from incomplete combustion of diesel, the fuel to air ratio

in diesel engines is kept lower than the (stoichiometric) ratio requiredby the balanced equation. The additional oxygen helps to ensure morecomplete combustion. Diesel-powered cars that emit black smoke fromtheir exhausts are poorly adjusted and produce less power. The followingequation shows a possible combustion reaction in such an engine.

Incomplete combustion of hexadecane (a component of diesel):

2C16H34(l) + 25O2(g) 16C(s) + 16CO(g) + 34H2O(l)

Exhaust gases from a poorly tuned car under different driving conditionsproduce the following relative levels of pollutants. Account for theseobservations.

Table 17.4

Pollutants

Level of pollutants when

cruising at constant speed

Level of pollutants when

stationary with engine running

nitrogen oxides high low

carbonmonoxide

low high

hydrocarbons low high

SAMPLE PROBLEM 17.3

Figure 17.10

The burning wick of a candle


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Insufficient oxygen mixes with unleaded fuel vapours when a car isidling (e.g. at traffic lights). This is less likely in a fuel injectionsystem but was quite common in an old carburettor system.Incomplete combustion results, producing less energy, and leavingunburnt fuel in the exhaust. The following equation illustrates poorcombustion of octane.

2C8H18(l) + 17O2(g) 16CO(g) + 18H2O(l)

When cruising along a highway, the ratio of air to fuel is optimal.With more available oxygen, the fuel vapours burn more completely.Complete combustion increases the temperature in the engine,promoting unwanted side reactions of air components to producenitrogen oxides. The following equations show reactions in theengine when fuel combustion is complete.

2C8H18(l) + 25O2(g) 16CO2(g) + 18H2O(l)

N2(g) + 2O2(g) 2NO2(g)

SOLUTION

Combustion

PRACTICALACTIVITIES

17.2

SYLLABUS FOCUS

19. USING INSTRUCTION TERMS

CORRECTLY

When answering questions, it is important to knowwhat the instruction terms (verbs) require you todo. Here are some examples.

Analyse

This instruction term requires you to identify

components and the relationship between them, or

to draw out and relate implications.

Example:

When a fuel burns in oxygen, the temperature of

the flame is higher if the reaction has a large heat

of combustion and the number of moles of gaseous

products that must be heated (per mole of fuel)is low. Analyse the following equations to predict

which fuel would have the higher flame temperature.

Identify the implications of this prediction.

Acetylene: C2H2(g) + 5O2(g)

2CO2(g) + H2O(g) $H= 1213 kJ/mol

Propane: C3H8(g) + 5O2(g)

3CO2(g) + 4H2O(g) $H= 2044 kJ/mol

Answer:

Each equation shows the combustion of 1 mole

of fuel. The acetylene reaction releases 1213 kJ of

energy per 3 moles of product gases (404 kJ per

mole of products). The propane reaction releases2044 kJ of energy per 7 moles of gaseous products

(292 kJ per mole of products). Thus, according to

the relationship between flame temperature, heat

of combustion and moles of products:

$H= mC$T

the acetylene flame should be hotter. In fact, an

acetylene flame is very hot and useful in high-

temperature welding and cutting of metals.

COMPLETE AND

INCOMPLETE

COMBUSTION

Evaluate

This instruction term requires you to make a

judgement based on standards or criteria, or to

determine the value of a proposal or idea.Example:

Evaluate the use of hydrogen as an alternative to

petrol for powering cars.

Answer:

Hydrogen has a much higher specific energy

(143 MJ/kg) than petrol (~47 MJ/kg). Thus, on a

weight basis, much more energy is available from

hydrogen gas. The gas needs to be compressed,

however, to provide sufficient energy density

(amount of energy stored in a given volume or

mass) to compete with petrol, which is condensed

at room temperature. Hydrogen is a clean fuel as it

burns to form non-polluting water; petrol burns to

form carbon dioxide (a greenhouse gas), carbon

monoxide (a toxic pollutant) and some carbon

(soot) that contaminates the environment. However,

hydrogen currently costs much more to manufacture

than petrol and greater safety precautions must be

taken due to the explosive nature of hydrogen

oxygen mixtures in the presence of a spark.


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324 ENERGY

Combustion

17.2 PRACTICALACTIVITIES

17.1 QUESTIONS

1. Which of the following isnotan indicator ofchemical change in the specified example?A Large amounts of heat are produced

when sodium hydroxide and

hydrochloric acid are mixed.B A colourless gas is evolved when copper

(II) carbonate is heated.C A white, waxy solid turns into a clear

liquid when gently heated.D A copper (II) sulfate solution turns a

very deep blue when ammonia solution isadded.

2. Identify the molecule that requires thegreatest input of energy to convert onemole of the gaseous compound into gaseousatoms. (Refer to table 17.1 on page 315.)

A C2H4B C2H6C C2H2D CH4

3. Identify which of the following reactions isendothermic.A Electrolysis of waterB Precipitation of lead (II) iodideC Neutralisation of zinc oxide by

hydrochloric acidD Hydration of plaster of Paris

4. The activation energy for the completecombustion of heptane is 350 kJ/mol. Theheat released when one mole of heptaneburns in excess oxygen to form carbondioxide and water is 4817 kJ/mol. Calculatethe activation energy for the reversereaction in which carbon dioxide and watercombine to form heptane and oxygen.A 350 kJB 4817 kJC 4467 kJD 5167 kJ

5. Select the correct statement about ignitiontemperatures.

A Octane has a lower ignition temperaturethan methane.

B Petrol has a lower ignition temperaturethan diesel.

C The ignition temperature of a fuel islower than its flash point.

D Compression heating is used in a petrolengine because the ignition temperatureof petrol is quite low.

6. Sulfur burns in air with a beautiful mauveflame, forming poisonous sulfur dioxide.The energy profile for the combustion ofsulfur in oxygen is shown in figure 17.11.The reaction is:

S(s) + O2(g) SO2(g)

(a) Mixing sulfur with air at roomtemperature does not lead to anyreaction. When the sulfur is heated,however, it eventually melts and starts toburn. Explain why sulfur does not burnat room temperature.

(b) The enthalpy change for thecombustion of sulfur is 300 kJ/mol.(i) Classify this reaction as

endothermic or exothermic.(ii) Use figure 17.11 to identify

which of the following algebraic

quantities equates to$H

:H

AH

C,HCHA, HAHB, HBHA, HCHB.(iii) Identify the algebraic quantity in

(ii) that equates to the activationenergy of the forward reaction.

(c) Explain why it is difficult to convertsulfur dioxide back into sulfur andoxygen.

HA

HB

HC

Enthalpy

(H)

Reaction coordinate

Products

Activated

complex

Reactants

Figure 17.11 Energy profile of the combustion of sulfur

7. A mixture of hydrogen and carbonmonoxide was prepared in the mole ratio3 : 1. (This mixture is sometimes calledtown gas.)(a) Write balanced equations for the

combustion of (i) hydrogen and(ii) carbon monoxide in oxygen.

(b) Use the following data to calculate theenergy released in the combustion of12 moles of the gaseous mixture.

Heat of combustion for H2= 286 kJ/mol

Heat of combustion for CO = 283 kJ/mol
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8. Nitrogen dioxide gas is brown but itsdimer, N2O4, is colourless. The proportionof each gas in a mixture is dependent ontemperature. Figure 17.12 shows the energyprofile for the reaction:

2NO2(g) N2O4(g)

Unstableactivated

complex

Reactants

24

79

0

Enthalpy

(H)(kJ)

Reaction coordinate

Products

Figure 17.12 Energy profile for the dimerisation of

NO2. Dimerisation is the process where a molecule (or

monomer) combines with another similar molecule.

(a) For the reaction above, determine the: (i) activation energy (EA)(ii) enthalpy change ($H).

(b) As a mixture of NO2and N2O4cools, itbecomes paler (less brown). Classify thisreaction as endothermic or exothermic.

(c) Calculate the activation energy for thedecomposition of 1 mole of dinitrogen

tetroxide.(d) Identify whether the activated complex

in the decomposition of dinitrogentetroxide has more or less energy thanthe reactant or the product.

9. When ammonia decomposes, it formsnitrogen gas and hydrogen gas.

100

200

300

400

500

0Enthalpy

(H)(kJ)

Reaction coordinate

Products

Unstable

activated

complex

Reactants

Figure 17.13 Energy profile for decomposition

of ammonia

Figure 17.13 shows the energy profile forthe decomposition of 2 moles of ammoniagas:

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

Use figure 17.13 to calculate the activationenergy for the formation of 1 mole ofammonia from its elements.

10. The volatility of a fuel is measured by itsequilibrium vapour pressure, which isusually measured at a standard temperatureof 25 C. Table 17.5 provides volatility datafor three common fuels.

Table 17.5

Fuel

Vapour pressure

(kPa) at 25 C Boiling point (C)

ethanol 10.3 78

octane 1.9 126

dodecane


14/32

326 ENERGY

11. A combustible mixture of diesel vapour andair is placed in a vessel at 250 C. Use thefollowing information to determine whetherthis fuel vapour oxidiser mixture wouldignite in the absence of a flame at thistemperature.

Flash point of diesel = 65 C

Ignition temperature of diesel = 350 C

12. Methanol has an ignition temperatureof 385 C and a flash point of 13 C.Combustible methanol air mixtures(636%v/v) can be quite dangerous as theyignite readily in the presence of an ignitionsource. The mixture burns with a pale blueflame, which is difficult to see in brightsunlight at a racetrack; racing car drivershave been burnt in such fires.

Table 17.6

Sample Volume of methanolvapour (mL) Volume of air (mL)

A 30 70

B 70 30

Samples A and B with different methanolvapour air compositions were prepared asin table 17.6.

Explain which sample would ignite whenheated to 450 C in the absence of flamesand sparks.

13. Figure 17.14 shows the apparatus used by

a student to collect and analyse some ofthe gases produced by the combustionof a candle. Use this figure to answer thefollowing questions.(a) Explain why the end of the U-tube on

the right is connected to a water pump.

(b) The blue cobalt chloride paper in theU-tube on the left turns pink. Identifythe combustion product indicated bythis test.

(c) In the U-tube on the right, thelimewater turns milky white. Identify thecombustion product indicated by thistest.

(d) Some black specks are observed on thewalls of the funnel and U-tube on theleft. Identify this black product.

(e) Classify the combustion of candle wax ascomplete or incomplete combustion.

14. The following information represents theresults of a student experiment in whichthe heat of combustion of butane gas froma portable lighter was measured usinga calorimetry experiment. Analyse thedata to determine a value for the heat of

combustion per mole of butane.Mass of water heated bythe butane flame = 50.0 g

Initial temperature of waterin glass calorimeter = 19.7 C

Final temperature of waterin glass calorimeter = 49.7 C

Mass of butane burnt = 0.25 g

Specific heat capacityof water = 4.2 J/K/g

15. LPG (liquefied petroleum gas) is largely

composed of propane. Evaluate a proposalto promote the use of LPG in suitablymodified motor vehicles as a replacementfor petrol.

&UNNEL

"LACK

SOLID

#ANDLE

)CE BLOCKS

"LUE COBALTCHLORIDE PAPER

7HITEPRECIPITATE

4

PUMP

,IMEWATER

Figure 17.14 Collecting combustion products of a candle


15/32


2%!#4)/. +).%4)#3

Particle collisions and reaction rateWhen you light a candle or burn natural gas, you are observing a chemicalreaction that occurs at a rapid rate or speed. Some reactions are so fastthat their rates are affected only by how fast we can mix the reactants.

The reaction ratesof strong acids and bases is an example of this. Otherreactions can be very slow and may take hours, days or years to produceappreciable quantities of products. The study of the rates of chemicalchange is called chemical kinetics.

The effect of temperature on reaction rate

We know from everyday experience that food cooks faster if we supply moreheat. The same principles apply to chemical reactions in the laboratory.Generally, as the temperature increases, the reaction rate increases.Example:Magnesium and hydrochloric acid

The reaction between magnesium strips and dilute hydrochloric acidproduces magnesium chloride and hydrogen gas.

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

The rate of the reaction can be monitored by measuring the volume ofhydrogen released as a function of time. The effect of increasing temperatureon the rate of hydrogen evolution can be determined experimentally.In such an experiment, the mass of magnesium (1.0 g), and the volume(100 mL) and molarity (1.0 mol/L) of the acid are held constant whilethe temperature is varied.

reaction rate:the change inconcentration of a reactant orproduct per unit time

chemical kinetics:the study ofthe rates of chemical change

reaction rate:the change inconcentration of a reactant orproduct per unit time

chemical kinetics:the study ofthe rates of chemical change

&,#'

Remember

Before beginning this section,you should be able to: explain that energy is required

to break bonds and energyis released when bonds areformed

describe the energy neededto begin a chemical reactionas activation energy

describe the energy profilesfor both endothermic andexothermic reactions.

Figure 17.15

Heating zinc metal in acid increases

the reaction rate.

Volume of gas measuredin gas syringe

Magnesium metal

Hydrogen evolved

Acid

Figure 17.16

Apparatus for measuring the volume

of gas evolved versus time


16/32

328 ENERGY

Figure 17.17 shows that the rate of the reaction (as measured by the initialslope of each line) increases as the temperature increases. Closer analysisshows that the rate approximately doubles for each 10 C rise in temperature.

The effect of concentration on reaction rate

The effect of increasing the concentration of one or more reactants onthe rate of the reaction can also be investigated experimentally. In theexample of the reaction between magnesium and hydrochloric acid,the effect of varying the concentration of hydrochloric acid could beinvestigated at constant temperature. Figure 17.18 shows the results

of three experiments in whichthe acid concentration isincreased from 1 to 3 molar. Itshows that the initial reaction rateincreases as the concentrationof acid increases.

Reactions involving gases alsoshow an increase in reaction rateas the pressure or concentration ofthe reactant gases increases. Thiscan be achieved by reducing thesize of the vessel so the reactantmolecules have less free space, orby adding more reactant gases to avessel with a fixed volume.

The effect of surface area

If 1.0 g of magnesium metal is ground into a fine powder and allowed toreact at room temperature with 100 mL of 1.0 mol/L hydrochloric acid,the rate of the reaction is very much greater than when using strips of

magnesium. The reaction is so rapid that the mixture becomes very hot.The powdered magnesium has a much greater surface area for the acid toattack, which increases the reaction rate.

Using particle models to explain reaction rates

Reactions occur because particles collide with sufficient energy to breakbonds and allow new bonds to form. Svante Arrhenius (18591927)proposed a model to explain why increasing the temperature increasesthe reaction rate. He suggested that, at room temperature, only a fewmolecules in the reaction mixture have enough kinetic energy to react,so few collisions between reactant molecules result in reaction products.Further, in dilute solutions or in low-pressure gas systems, these few

energetic molecules would rarely meet, so the reaction would be slow.Arrhenius reasoned that a reaction would occur if the sum of the kineticenergies of the reactants is greater than the minimum required energy(activation energy, EA). If the reaction mixture is heated, a greaterproportion of reactants would have kinetic energy greater than EA, somore successful collisions would occur.

Later mathematical theories of reaction rates showed that the averagekinetic energy, KE , of the colliding particles was directly proportional tothe absolute (kelvin) temperature, T, of the system.

KEsT

Key content

By the end of this section, youshould be able to:

describe combustion in terms ofslow, spontaneous and explosivereactions and explain theconditions under which theseoccur

explain the importance ofcollisions between reactingparticles as a criterion fordetermining reaction rates

explain the relationship betweentemperature and the kineticenergy of particles.

describe the role of catalystsin chemical reactions, using anamed industrial catalyst as anexample

explain the role of catalysts in

changing the activation energyand hence the rate of a chemicalreaction

solve problems, identify data,perform first-hand investigationsand gather first-hand data, whereappropriate, to observe theimpact on reaction rates of: changing temperature changing concentration size of solid particles adding catalysts

process information fromsecondary sources to investigatethe conditions under whichexplosions occur and relatethese to the importance ofcollisions between reactingparticles

analyse information and usethe available evidence to relatethe conditions under whichexplosions occur to the needfor safety in work environmentswhere fine particles mix with air

analyse information fromsecondary sources to developmodels to simulate the role ofcatalysts in changing the rate ofchemical reactions.

Figure 17.17

Effect of temperature on the rate of

the reaction between magnesium and

hydrochloric acid

Volum

eofhydrogen(mL)

40 C30 C

20 C

Initial slope = rate of reaction (R)

R30= 2 R20R40= 2 R30

VolumeofH2(mL)

Time (min)

Initial slope = rate of reaction (R) concentration

Figure 17.18

Effect of the concentration of hydrochloric

acid on the rate of its reaction with

magnesium


17/32


Figure 17.20

Effect of a catalyst on the activation energy of a reaction

Products

ReactantsReaction coordinate

Products

EA

EA

Enthalpy(H)

(EAhigher)

With catalyst

(EAlower)

ReactantsH

Thus, if the system is heated from 27 C (300 K) to 37 C (310 K), the averagekinetic energy of the reactant particles should increase (accordingto the above relationship) by about 3%. However, this small increase inkinetic energy of the particles does not explain why the reaction ratedoubles when the temperature rises by 10 C. The explanation lies in thenon-uniform spread of kinetic energies in the population of molecules.Some molecules have much more kinetic energy than others and these more

energetic molecules have energy greater than EA. Raising the temperatureby 10 C doubles the number of these more energetic molecules, so thereaction rate doubles. Figure 17.19 illustrates this concept.

Figure 17.19

Effect of temperature on the

proportion of particles with sufficient

kinetic energy to react

Proportionofparticleswith

kineti

cenergy(KE)

Kinetic energy of particles (KE)

Activation

energy

Higher proportion of

particles with KE> EA

at high temperature

Smaller proportionof particles with

KE> EAat low

temperature

Low

temperature

High

temperature

EA

In summary, the rate of the reaction depends on:

collision rate, which depends on concentration

the proportion of collisions with energy greater than EA, which depends

on temperature.

CatalysisA catalyst is a chemical substance that increases the rate of a chemicalreaction. Catalysts interact with one or more reactants or intermediatesso that chemical bonds can be broken and re-formed more easily. Animportant characteristic of catalysts is that they are re-formed afterthe reaction is complete. The reaction pathway (or mechanism of thereaction) is altered by the catalytic interaction.

Catalysed reactions have lower activation energy, EA, than similaruncatalysed reactions. Thus, at any specific temperature, a greaterproportion of reactant particles have more energy than the lower activationenergy. Figure 17.20 shows the effect of a catalyst on the energy profile ofa reaction. Note that the enthalpy change for the reaction is not affectedby the presence of a catalyst.

Experimental research is used to identify catalysts. For a substance to beclassified as a catalyst, it must:

increase the reaction rate compared with an uncatalysed control

be re-formed at some later stage in the reaction.

catalyst:a substance that altersthe rate of a reaction without achange in its own concentration

mechanism:the steps involved ina chemical reaction

catalyst:a substance that altersthe rate of a reaction without achange in its own concentration

mechanism:the steps involved ina chemical reaction


18/32

330 ENERGY

An experiment investigated the effect of copper on the rate ofthe reaction between zinc granules and dilute hydrochloric acid.Figure 17.21 shows the experimental observations and numericaldata for the reaction.

Figure 17.21 Effect of copper on the rate of the reaction between zinc

and hydrochloric acid

3LOW

EFFERVESCENCE

M, (

M, (#L

#OPPER PRESENT

-ASS OF COPPER DOES NOT CHANGE DURING REACTION

#OPPER PRESENT

:INC G

M, (

.O (EVOLVED

4UBE 4UBE 4UBE CONTROL

2APID

EFFERVESCENCE

M, (#L M, (#L

#OPPER IN

CONTACT WITH

ZINC G

#OPPER

NO ZINC

Use figure 17.21 to determine whether copper is acting as a catalystin the reaction.

Tubes 1 and 2 both contain zinc and produce hydrogen. However,tube 2 also contains copper and produces hydrogen more rapidly

than tube 1. This satisfies the criterion that a catalyst increases thereaction rate.

The mass of copper recovered at the end of the experiment is the sameas the initial mass of copper. Thus, no copper reacted permanently,satisfying the second criterion for a catalyst.

Thus, copper acted as a catalyst in this reaction.

SAMPLE PROBLEM 17.4

SOLUTION

Adding a catalyst to a reacting system cannot result in any more productthan allowed by the stoichiometry of the reaction (as determined by thebalanced chemical equation). The product accumulates, however, at a fasterrate than without a catalyst. Figure 17.22 shows that, in the reaction between

zinc and dilute hydrochloric acid, the volume of hydrogen produced is notaffected by the presence of copper. It also shows that the rate of productionof hydrogen is much faster in the presence of the copper catalyst.

Catalysts in industryLet us examine some industrial uses of catalysts.

Manufacture of nitric acid

The manufacture of nitric oxide, NO, by the Ostwald process is animportant step in the industrial manufacture of nitric acid. The first step

Figure 17.22

Volume of hydrogen evolved

on reaction between zinc and

hydrochloric acid, with and without

the copper catalyst

01

2

20

40

6080

100

120

VolumeofH2(mL)

Time (min)

Copper Samevolume ofhydrogenevolved

Initial slope = rate of reaction (R);

No copper

R2> R

1

Figure 17.22

Volume of hydrogen evolved

on reaction between zinc and

hydrochloric acid, with and without

the copper catalyst

01

2

20

40

6080

100

120

VolumeofH2(mL)

Time (min)

Copper Samevolume ofhydrogenevolved

Initial slope = rate of reaction (R);

No copper

R2> R

1


19/32


of this process used a platinumrhodium catalyst. Because the catalystand reactants are in different phases, this process is an example ofheterogeneous catalysis.The initial reaction is catalytic oxidation of ammonia:

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

Figure 17.23 shows a simplified view of the reactions that break and

re-form bonds on the surface of the metallic catalyst in this initial step.

Figure 17.23

Reactions at the surface of the

platinumrhodium catalyst

HH

H

HH

HH

HH

H HH

HH

H

HH

H

HH

H

HH

H

H H

H H

HH H H

HH

HH

(1)Ammonia moleculesmove onto thecatalyst surface.

(3)Molecules break upinto atoms on surface.

(5)New products formas atoms recombine.

(2)Oxygen moleculesadsorb onto catalystsurface.

(4)Atoms migrate onsurface.

(6)Product moleculesescape from surface ofcatalyst.

The sum of the remaining reaction steps is:

4NO(g) + 3O2(g) + 2H2O(l) 4HNO3(aq)

So, the overall reaction is:

NH3(g) + 2O2(g) HNO3(aq) + H2O(l)

Manufacture of ethyl acetate

Ethyl acetate, CH3CO2C2H5, is an important industrial solvent. It is

manufactured by the reaction of acetic acid and ethanol in the presenceof a concentrated sulfuric acid catalyst.

CH3CO2H(l) + C2H5OH(l) CH3CO2C2H5(l) + H2O(l)

Because the acid catalyst and the reactants are all in the same phase, thisreaction is an example of homogeneous catalysis. Even though the reaction isquite slow and the yield of product is less than 100%, the catalyst ensuresthat equilibrium is reached in the shortest possible time.

Rates of combustion reactionsNot all combustion reactions proceed at the same rate. Some are quite slowand others are so rapid that an explosion results. Fuels such as kerosene

undergo rapid combustion.

Spontaneous combustion

Some fuels have a high ignition temperature, which prevents spontan-eous combustion at room temperature. However, some materials with lowactivation energy and ignition temperature can combust spontaneouslyon exposure to air or oxygen. Some examples of where spontaneouscombustion can occur are:

brown coal deposits.When mined brown coal is exposed to air, it canoxidise, producing heat. The temperature of the coal deposit increases

17.3

Reaction rates

PRACTICAL

ACTIVITIES

DATAANALYSISCatalysis

17.4

CATALYSIS:

HYDROGENATION

OF ETHYLENE


20/32

332 ENERGY

and it can combust spontaneously. This may be prevented by sprayingwater onto it to cool it.

haystacks.The decomposition action of microbes (e.g. bacteria, mouldsand protozoans) on moist, mown grass can build up heat in thecentre of a stack. If the stack is quite large, the heat builds up fasterthan the stack can radiate heat. Spontaneous combustion can thenoccur. Modern farming practices have led to new methods of reducingthis danger.

oily, cotton rags.Cotton rags are used to mop up oil spills in some work-shops. However, the cotton fibres provide a large surface area for the oilto make contact with the air. If the rags become hot enough, spontaneouscombustion can result.

white phosphorus. An allotrope of phosphorus is called white phos-phorus, P4. It must be stored underwater as it will combust spontaneouslyon contact with air at quite low temperatures.

Slow combustion

Two examples of slow combustion are:

rusting.When iron is exposed to air or oxygen in a moist environment,it oxidises slowly to form a substance called rust. Other metals alsoundergo surface corrosion. These reactions release heat but the rise intemperature is quite small as the reaction occurs only on the surfacelayers.

burning of wood, coal and coke.When coal or coke or logs of wood areburnt in a fireplace, the combustion is much slower than when paperor twigs burn. These combustion reactions are slow because the fuelhas a small surface area. If the surface area of the fuel is increased (e.g.by powdering the coal), the reaction can become very fast and possiblyexplosive.

Explosive combustionFigure 17.25 compares the rates of a normal reaction and an explosivereaction as a function of temperature. In explosions, there is a rapidincrease in reaction rate with increasing temperature. This may occurif there is a rapid increase in temperature due to a chain reaction thatproduces an exponential growth in new reactants.

Examples of explosive combustion reactions include:

hydrogen and oxygen.When mixtures of hydrogen gas in air are sparked,the hydrogen burns explosively to form water. This reaction is used in acontrolled way in rocket engines; the water vapour formed in the reac-tion is expelled at high pressure and temperature to provide thrust.

hydrogen and chlorine. Chlorine, like oxygen, can act as an oxidiser.Hydrogen gas and chlorine gas mixtures are stable in the dark but,on exposure to light, chlorine molecules absorb photons and breakdown to form reactive chlorine radicals that rapidly attack hydrogenmolecules. A chain reaction begins that leads to an explosion, forminghydrogen chloride gas.

dust explosions.Fine dust particles have very high surface areas. Whenmixtures of air and particles of a combustible material (e.g. carbondust, flour dust, carbon toner dust, cellulosic dust) are sparked, theycan explode. It is important in grain silos and mines to control the dust

Figure 17.24

The rusting of iron exposed to air

is an example of a slow combustion

reaction.

Figure 17.24

The rusting of iron exposed to air

is an example of a slow combustion

reaction.

Figure 17.25

Relationship between reaction rates

and temperature for explosive and

non-explosive reactions

Reaction

rate

Temperature

Non-explosive reactions

Rate increases

with increasing

temperature.

Explosive reactions

Ignition

temperature

Explosion beginsas rate increases

very rapidly.

Reactionrate

Temperature

Figure 17.25

Relationship between reaction rates

and temperature for explosive and

non-explosive reactions

Reaction

rate

Temperature

Non-explosive reactions

Rate increases

with increasing

temperature.

Explosive reactions

Ignition

temperature

Explosion beginsas rate increases

very rapidly.

Reactionrate

Temperature


21/32


22/32

334 ENERGY

7. Figure 17.26 shows the volume of carbondioxide released when a fixed mass ofcalcium carbonate reacted with excesshydrochloric acid at constant temperature.In one experiment, calcium carbonate wassupplied as marble chips and, in the other,as a powder.

0

0

20

40

60

80

100

20 40 60 80 100 120 140

120

Volumeofcarbon

dioxide(mL)

Time (sec)

Powder

Chips

160 180

(a) Explain why the reaction with thepowdered calcium carbonate was muchfaster than with marble chips.

(b) Explain why the same volume ofcarbon dioxide was released in eachexperiment.

8. Jesse winds strips of magnesium arounddeflagrating spoons. He ignites the ribbonsin a Bunsen burner flame and then allowsthe strips to continue to burn in two

separate gas jars. Jar A contains air and jar Bcontains pure oxygen. Compare the rate ofcombustion of magnesium in the gas jars.

9. Rust is a serious problem in our builtenvironment. The rusting of iron requiresthe presence of both water and oxygen.Explain how the rate of this reaction couldbe increased.

10. Figure 17.27 shows a flask containing

chips of calcium carbonate and dilutehydrochloric acid. Explain how thisapparatus could be used to measure therate of the reaction between the acid andcalcium carbonate.

11. Explain how haystacks can suddenly burstinto flame.

12. Methaneair mixtures are explosive whenthe methane concentration is between 5 and15%v/v. An experimentalist investigates thebehaviour of burning methane. He takes atin can with a pressdown lid (e.g. Milo can).He makes a 10 mm hole in the bottom ofthe can and a 5 mm hole in the lid. The canis filled with methane gas and placed on atripod. A flame from a candle is applied tothe hole in the lid and the methane startsto burn with a blue-yellow flame, whichquickly disappears. The experimentalistwaits for a while at a safe distance behind aspecial screen. For a while, nothing seemsto happen. Then, there is an explosion thatcauses the lid to be thrown upwards.

Account for this behaviour of the burningmethane.

#ALCIUM CARBONATE

)NITIAL LEVEL

OF LIQUID IN

MANOMETER

(EIGHT DIFFER

IN LIQUID IS

MEASURED AS A

FUNCTION OF TIME

#OLOURED LIQUID

IN MANOMETER(YDROCHLORIC

ACID

#ARBON DIOXIDE

EVOLVED

H

Figure 17.27 Reaction between calcium carbonate and hydrochloric acid

Figure 17.26


23/32


SUMMARY

Hydrocarbons such as petrol, kerosene and diesel are commonfuels. How readily fuels combust depends on several factorsincluding ignition temperature, volatility and flash point.

Energy is released when hydrocarbons burn. The energy released(per mole) is called the heat of combustion or enthalpy ofcombustion. Combustion can be described as slow, spontaneousor explosive.

Chemical reactions involve bond-breaking (endothermic) andbond-forming (exothermic) processes.

Pollution can result from incomplete combustion of fuels.Carbon monoxide and soot are common pollutants produced byincomplete hydrocarbon combustion.

Chemical reactions vary in their rate. Catalysts can be used toincrease the rate of reactions. Catalysts have an important role in

industry.

Various factors, including temperature, pressure, concentration,surface area and catalysts, affect the rate of a reaction. Changesin reaction rate with temperature can be understood in terms ofchanges in the kinetic energy of the colliding molecules.

Catalysts are important in industry as they increase the rate atwhich products are produced.

The energy required to begin a chemical reaction is called theactivation energy. If molecules do not have sufficient kineticenergy to overcome the activation energy barrier, no reactionoccurs.

MODULE 4REVISION

www.jaconline.com.au/chemnsw/chemistry1
http://cdrom/module-revision/core-04.pdfhttp://www.jaconline.com.au/chemnsw/chemistry1


24/32

336 ENERGY

17.1EXOTHERMIC AND

ENDOTHERMIC REACTIONS

Aim

To investigate and classify a range of chemicalreactions as endothermic or exothermic

Safety issues Wear safety glasses throughout this experiment.

Take particular care with bases such as sodiumhydroxide that they do not come in contactwith your eyes or skin.

Avoid breathing ammonia vapours.

Identify other safety issues relevant to thisexperiment by reading the method.

Materials clean test tubes test-tube rack

Pasteur pipettes

10 mL measuring cylinder

glass stirring rod

alcohol thermometer (10 C to +100 C range)

2 mol/L HCl

2 mol/L NaOH

universal indicator solution

copper (II) carbonate solid

plaster of Paris 2 mol/L calcium chloride solution

2 mol/L sodium carbonate solution

magnesium ribbon (5cm)

hydrated barium hydroxide solid

ammonium thiocyanate solid

Method1. Table 17.7 describes seven experiments. These

should be done using small quantities of

PRACTICALACTIVITIES

PRACTICALACTIVITIES

Table 17.7

Experiment Reaction Procedure Observations Classification

1 Hydrochloricacid + sodiumhydroxide solution

Pour 2 mL of 2 mol/L NaOH into a clean test tubeand add 3 drops of universal indicator. Slowly add3050 drops of 2 mol/L HCl from a Pasteur pipette.

2 Copper (II)carbonate +hydrochloric acid

Put about one third of a scoop of copper (II)carbonate into a test tube and add 2030 drops ofhydrochloric acid using a Pasteur pipette.

3 Dissolution ofammonium nitratein water

Add half a scoop of ammonium nitrate crystals to2 mL of water in a test tube.

4 Hydration ofplaster of Paris

Place half a scoop of plaster of Paris in a dry testtube and add drops of water until no furtherchange occurs.

5 Sodium carbonatesolution + calciumchloride solution

Add drops of 2 mol/L calcium chloride solution to2 mL of 2 mol/L sodium carbonate solution in atest tube.

6 Magnesium +hydrochloric acid

Put a 5 cm coiled strip of magnesium in 2 mL of2 mol/L hydrochloric acid.

7 Hydrated bariumhydroxide +ammoniumthiocyanate

Mix half a scoop of hydrated barium hydroxidewith half a scoop of ammonium thiocyanate. Stirthe solid mixture gently with a glass rod. (Do NOTbreathe the vapours.)


25/32


chemicals in clean test tubes, with an alcoholthermometer inserted to measure changes intemperature.

2. Copy and complete table 17.7 as youconduct the experiments. Measure theinitial temperature and the highest or lowesttemperature reached during the reaction.Record these values, together with otherobservations in the Observations column. Inthe Classification column, write exothermicor endothermic to classify the observedreaction.

QuestionsAnswer the following questions in your report onthis experiment.

1. Write balanced equations for each of thereactions performed.

(Note:In experiment 4, plaster of Paris hasthe formula 2CaSO4.H2O and the gypsumthat forms has the formula CaSO4.2H2O. Inexperiment 7, the formula of ammoniumthiocyanate is NH4SCN and barium hydroxidehas 8 molecules of water in its crystal. Theproducts of the reaction are ammonia, water

and barium thiocyanate.)2. Explain why the temperature of the system

decreases during an endothermic reaction.

Conclusion

Briefly describe the outcome of your investigation.

17.2COMBUSTION

Part A: Combustion of wood

and magnesiumAimTo compare the change in weight on thecombustion of wood and magnesium

Safety issuesWear safety glasses throughout this experiment.


Materials wooden sticks (e.g. paddle-pop sticks broken

in half)

10 cm magnesium ribbon

tongs

evaporating basins

electronic balance

Bunsen burner

Method1. Determine the mass of a clean evaporating

basin. Determine the mass of a wooden stick.

2. Hold the end of the wooden stick in a pairof tongs and ignite the wood using a Bunsenflame. Allow any wood ash to fall into thebasin. Release the end of the wood into thebasin and allow the remainder of the wood tofinish burning in the basin.

3. When cool, reweigh the basin and ash.Determine the mass of the ash.

4. Weigh a new basin and a 10cm strip ofmagnesium ribbon.

5. Hold the ribbon over the basin and ignite thefree end with a Bunsen flame. Allow the ash tofall into the basin. Allow the remainder of theribbon to finish burning in the basin.

6. When cool, reweigh the basin and ash.Determine the mass of the ash.

PRACTICALACTIVITIES

PRACTICALACTIVITIES


26/32

338 ENERGY

ResultsRecord your observations in a suitable table.

QuestionsAnswer the following questions in your report onthis experiment:

1. Compare the mass of the wood with the massof the wood ash. Has there been an increaseor decrease in mass?

2. Compare the mass of the magnesium and theash. Has there been an increase or decrease inmass?

3. Explain the difference in results between the

two experiments.Part B: Combustion of a candle

AimTo observe the combustion of a candle.

BackgroundA candle is composed of a fibre wick surrounded bywax. The wax consists of long-chain hydrocarbons;it has a low volatility so combustible mixtures of itsvapour with air will form only if the wax is hot. Thewick provides a large surface area on which hot,liquid wax molecules can vaporise. The wick andany wax on its surface are ignited initially by the

flame of a burning match. The candle producesa flame around the wick due to the presence of acombustible mixture of hydrocarbon vapours andair. The flame is the reaction zone that radiatesheat outwards. Some of this radiant heat meltsthe wax and creates a pool of melted wax that canmove upwards through the pores of the wick bycapillary action to provide a continuous supply ofwax vapour for the flame. The narrow zone of waxvapour around the wick helps to prevent it burningaway too rapidly. As the candle burns down, the tipof the wick becomes exposed to the air and charsas it burns. Candle wax burns incompletely due tothe slow diffusion of air into the flame zone. Theblackening of a beaker placed above the flameshows the presence of unburnt carbon in the flame.



Materials candle (e.g. a birthday candle)

Petri dish

electronic balance

matches

Method1. Stand the candle in the Petri dish and weigh

them.

2. Light the candle and make observations as thecandle burns at the wick. Note the colours ofthe wax and the zones of the flame. Is thereany black smoke? Note how the wax melts in

some areas and re-solidifies in others.3. Continue the observations until the candle

has burnt down 12 cm. Draw a labelleddiagram of the flame, wick and top of thecandle, noting any regions of different colour.

4. Blow out the flame and reweigh the dish andcandle.

ResultsRecord your observations in a suitable way.

PRACTICALACTIVITIES

Figure 17.28


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1. Explain why the candle loses mass on burning.

2. Identify changes in state during thecombustion of a candle.

3. The innermost zone of the flame near thewick is black. In this region, there are unburntwax vapours. Predict whether this zone will becool or hot.

4. A yellow, luminous middle zone of the flameindicates a region of incomplete combustion

where the flame is moderately hot. The yellowcolour indicates the presence of glowing sootparticles as a combustion product. Explainwhy the wax vapours are only partially burnt inthis zone.

5. A blue, non-luminous flame indicates regionsof complete combustion of the hydrocarbonvapour. Explain why this blue region is at thebase and edge of the flame zone and why thispart of the flame is very hot.

6. Explain (in terms of ignition temperature andcombustible mixtures) why the combustion of

the wax vapour is self-sustaining after the litmatch is removed.

ConclusionBriefly describe the outcome of your investigation.

17.3REACTION RATES

Part A: Permanganateoxalic acidreactionAimTo investigate the effect of different variableson the rate of the reaction between potassiumpermanganate and oxalic acid

BackgroundThe reaction between pink-purple permanganateions and colourless oxalic acid in an acidicmedium is shown by the equation:

2MnO4(aq) + 5C2O4H2(aq) + 6H

+(aq)

2Mn2+(aq) + 8H2O(l) + 10CO2(g)

The average rate of the reaction can bedetermined by the time taken for the solution tochange from pink-purple to colourless.


Oxalic acid is poisonous. Do not ingest it. Potassium permanganate can stain hands and

clothing.


Materials large test tubes

burettes

conical flasks

thermometers

hotplate

stopwatches water bath

250 mL of 0.10 mol/L oxalic acid,C2O4H2.2H2O

250 mL of 4.0 mol/L sulfuric acid

250 mL of 0.005 mol/L potassiumpermanganate, KMnO4

50mL of 0.1 mol/L manganese sulfate solution

deionised water

PRACTICALACTIVITIES

PRACTICALACTIVITIES


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340 ENERGY

Method

Part 1: Effect of concentration of oxalic acid

1. All experiments should be done in large testtubes immersed in a constant-temperaturewater bath at about 30 C. Note thetemperature in your results.

2. In clean test tubes, prepare the mixtures intable 17.8.

Table 17.8

Mixture Tube 1 Tube 2

Final concentration

of oxalic acid in

reaction mixture(mol/L)

1 4.0 mLKMnO4

4.0 mLoxalic acid,

2.0 mL sulfuricacid

4.0/10.0 0.10 =0.040

2 4.0 mLKMnO4

3.0 mLoxalic acid,

1.0 mL water,2.0 mL sulfuric

acid

3 4.0 mL

KMnO4

2.0 mL

oxalic acid,2.0 mL water,

2.0 mL sulfuricacid

3. For each mixture, equilibrate the two testtubes in the water bath. Tip the contents oftube 2 into tube 1 and start timing. Agitatethe reaction mixture and record the time (inseconds) when the pink-purple colour firstdisappears.

4. Calculate the concentration of oxalic acid in

each final reaction mixture. The total volumeof each reaction mixture is 10.0 mL. The firstcalculation has been done for you.

5. Combine all class results to improve reliability.

6. Record your observations in a suitable tablewith the headings Mixture, Final oxalic acidconcentration, and Reaction time.

7. Calculate the mean reaction time anddetermine the deviation from the mean foreach mixture.

Part 2: Effect of temperature

1. Follow the same general procedure as in part1. However, this time, keep the concentrationsconstant and vary the temperature.

2. Use the mixture of reactants shown intable 17.9.

Table 17.9

Tube 1 Tube 2

4.0 mL KMnO4 4.0 mL oxalic acid2.0 mL sulfuric acid

3. Do the experiment as in part 1 at thefollowing water bath temperatures: 30 C,40 C, 50 C.

4. Record the time for decolourisation as in part1. Tabulate your results. Combine class resultsto improve reliability.

5. Calculate the mean reaction time anddeviation from the mean for eachtemperature.

Part 3: Effect of a catalyst1. Follow the same general procedure as in

part 1. However, this time, a small volume ofmanganese sulfate solution is added to testits catalytic properties. All other variables areconstant.

2. Test the reaction mixtures in table 17.10 at30 C.

3. Record the time for decolourisation as before.Tabulate your results. Combine class resultsto improve reliability. Calculate the meanreaction time and the deviation from themean for each mixture.

Table 17.10

Mixture Tube 1 Tube 2

1 4.0 mL KMnO4 3.0 mL oxalic acid2.0 mL sulfuric acid

1.0 mL water

2 4.0 mL KMnO4 3.0 mL oxalic acid2.0 mL sulfuric acid

0.5 mL water0.5 mL manganese sulfate

PRACTICALACTIVITIES


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1. Explain how the reaction time is related to thereaction rate.

2. Discuss the effect on the reaction ratedecreasing the concentration of oxalic of acid.

3. Discuss the effect on the reaction rate ofincreasing the temperature.

4. Discuss the effect on the reaction rate of thepresence of manganese sulfate. What otherinformation would be needed to prove that

manganese ions were acting catalytically?

Part B: Acid on calcium carbonate

AimTo investigate the effect of particle size on therate of the reaction between hydrochloric acidand marble.

Safety issues Wear safety glasses throughout this experiment.


Materials 2 measuring cylinders

two 50 mL beakers

stirring rods

electronic balance

stopwatch

5.0 g marble chips

5.0 g powdered calcium carbonate

1 mol/L hydrochloric acid

Method1. Weigh 5.0 g each of marble chips and

powdered calcium carbonate into separatesmall beakers.

2. Measure 50 mL of 1 mol/L hydrochloric acidinto separate measuring cylinders.

3. Add the acid to each beaker and stir themixtures. Record the time for the mixture ineach beaker to stop fizzing.

4. Combined all class results. Compare thereaction times for large and small particles.

ResultsRecord your observations in a suitable table.

QuestionsAnswer the following questions in your report onthis experiment:

1. Compare the surface area of 5 g of marblechips to 5 g of powdered calcium carbonate.

2. Explain why the reaction rate is much higherfor powdered calcium carbonate than formarble chips.

3. Determine which reactant was in excess in thisexperiment.

Conclusion

Briefly describe the outcome of yourinvestigations.

PRACTICALACTIVITIES


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342 ENERGY

17.4CATALYSIS

Model construction

1. For each step in the following reactions, use amolecular model kit to make models of eachreactant and catalyst molecule.

2. Draw a diagram of the models at each step.

Example 1: Decomposition of ozone

1. Make a model of the ozone molecule, O3. It isan angular (bent) molecule.

2. Make a model of a nitric oxide molecule, NO.3. NO reacts with O3to form NO2and O2. Model

this reaction.4. NO2then reacts with reactive oxygen atoms

(free radicals) to form nitric oxide and O2.Model this reaction.

5. Identify the catalyst, which is the atom ormolecule used up in one step and re-formed ina later step.

6. Write these reaction steps as a series of

balanced equations. Sum these reactions stepsto obtain the overall ozone decompositionreaction.

Example 2: Hydrogenation of ethylene

Hydrogenation of ethylene is catalysed by finelydivided palladium or platinum. This reactioninvolves a heterogeneous catalyst. After thegaseous molecules adsorb onto the catalystssurface, bonds break and new bonds form. Theproduct molecules then desorb from the surface.

1. Let the desktop represent the surface of thecatalyst.

2. Make a model of ethylene, C2H4, and a modelof H2gas. Place both molecules on the catalystsurface. They are now adsorbed onto thesurface.

3. Break the bond of the hydrogen moleculeto form two reactive hydrogen atoms (freeradicals). These are bound to the metal atomsbut they can move across the surface of thecatalyst towards the ethylene.

4. One reactive hydrogen atom attacks a carbonatom of the ethylene molecule, causing the

double bond to break and forming a singlecarboncarbon bond. Model this reaction.

5. The other hydrogen atom bonds to the secondcarbon atom to form ethane. Model thisreaction.

6. The ethane molecule then desorbs from thecatalysts surface. The surface is now free tocatalyse a new reaction.

7. Compare the strength of the metalhydrogenbond on the catalyst surface with thehydrogenhydrogen bond in H2.

DATAANALYSIS

DATAANALYSIS


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17.5EXPLOSIONS AND SAFETY

Part A: Hydrogen gas explosionThe reaction that occurs when a mixture ofhydrogen and oxygen is sparked is an exampleof a reaction in which the rate of formation offree radicals determines whether the mixtureexplodes or not. Changes in parameters such as

temperature, pressure and solid surfaces alter theprogress of the reaction.

The box below provides some examples ofreaction steps that occur in the combustion ofhydrogen. Steps A to G are listed in randomorder. (Note:In equations, free radicals are shownby placing a dot next to the molecular formula torepresent an unpaired electron.)

uOH + H2 H2O + Hu (A)

HO2u+ H2 H2O + uOH (B)

H2 Hu+ Hu (C)

Hu+ Hu H2 (D)

Ou+ H2 uOH + Hu (E)

Hu+ O2 uOH + Ou (F)

Hu+ O2 HO2u (G)

Answer the following questions using informationfrom the box.

Questions1. The first step in the combustion of hydrogen

is to break the HH bond to form reactivehydrogen radicals. This is achieved by

sparking so that the temperature is highenough to overcome the activation energybarrier. Use the code letter to identify theequation that represents this process.

2. Once the hydrogen free radicals form, theycan react with oxygen molecules to generateoxygen and hydroxyl free radicals. Identify theequation that represents this reaction.

3. Once oxygen free radicals form, they canattack hydrogen molecules to form more

hydrogen and hydroxyl radicals. Identify theequation that represents this reaction.

4. Above a temperature of 890 K (617 C),the reaction is always explosive because theproduction of radicals described in (2) and(3) occur so rapidly that a chain reactionresults. Below this temperature, the walls ofthe vessel can absorb and deactivate radicals.Three of the listed equations describe thepropagation of radicals in the gas phase andone describes the termination of radicals toform a stable molecule.

(a) Identify the three equations involvingpropagation and formation of newradicals.

(b) Identify the equation that shows a radicaltermination step.

5. It is known that some substances, such as KCl,on vessel walls promote the termination ofradicals by allowing HO2uand Huradicals tocombine. What molecule forms when thesetwo radicals collide?

6. Explain why hydrogenoxygen mixtures aresafe at room temperature.

7. Identify the safety precautions necessary to

store and handle hydrogen safely.

Part B: Dust explosionsRead the following information and answer thequestions that follow:

As long ago as 1785, it was recognised that certaintypes of dust can lead to explosions. Flour dustcontains starch and cellulose, which burn undercertain conditions. In 1785, flour dust in a bakersstoreroom in Italy exploded when it came incontact with an oil lamp.

Dust explosions in industry must be avoidedas they can injure or kill workers and damageequipment and buildings.

Various dust clouds in air can form flammablemixtures. If an ignition source is present,an explosion can occur in which the flamepropagates rapidly through the dust cloud. A flashfire results if the flame propagates in an openspace. In a closed space, the pressure wave leadsto an explosion. The maximum pressure in a dustexplosion is typically 5001000 kPa.

DATAANALYSIS

DATAANALYSIS


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The severity of the explosion is related to theheat released in the combustion reaction. Someexamples are shown in table 17.11.

The conditions that lead to dust particleexplosions are:

suspended dust particles less than 0.1 mm indiameter

suspended dust concentration between 50 and3000 g/m3

moisture content of dust less than 11%

oxygen concentration greater than 12%

presence of an ignition source (e.g. openflame, hot surface or electrical discharge)

presence of a combustible mixture in a closedspace.

The types of dust that form combustible mixtureswith air include:

metals (e.g. Al, Mg, Zn and Fe)

non-metals (e.g. C and S)

natural organic materials (e.g. starch, celluloseand coal)

synthetic organic materials (e.g. pigments and

plastics).Various industrial standards for dust level are setto maintain safety. The Australian government hasset the industrial hygiene range at about110 mg/m3.

Some common procedures used to prevent dustexplosions include:

using pneumatic dust collectors to keep dustlevels low at all times

venting of dust

adding inert dust to reduce the concentration

of combustible dust (e.g. rock dust in coalmines)

using unreactive gases (N2, CO

2, inert gases) in

reaction vessels to prevent oxygen reaching thedust

regular monitoring and maintaining electricalequipment to avoid sparking.

Primary dust explosions can also lead tosecondary explosions. The primary blast wave cancause dust in other areas to become suspendedand burn explosively as the primary flame frontpropagates outwards.

Go to:

www.jaconline.com.au/chemnsw/chemistry1

and click on the Dust Explosion link for thischapter to view videos of dust explosions.

Questions1. Explain why flour dust can be dangerous in

bakeries.

2. A closed storage room contains suspendeddust from sugarcane fibres. The dust particlesare 0.01 mm in size and the concentrationof dust is 1000 g/m3. Identify the additionalconditions required to cause the explosion ofthis dust suspension.

3. A dust of finely powdered aluminium iscreated in a metal grinding factory. The sparksfrom the grinding caused the dust suspensionto explode.

(a) Write a balanced equation for the reactionthat occurred.

(b) Explain how the factory managers couldprevent such explosions in the future.

4. Identify which of the following produces themost severe explosions: carbohydrate dusts,active metal dusts, non-metal dusts.

5. Explain how secondary explosions can result

from primary explosions in suspended dust.

DATAANALYSIS

Table 17.11

Dust sulfur carbon starch zinc aluminium magnesium

Heat released per mole O2(kJ) 297 394 470 700 1117 1204
http://www.jaconline.com.au/chemnsw/chemistry1

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jacaraonda chapter 17