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9.3 - The Acidic Environment

Section #2

The Acidic Atmosphere

While we usually think of the air around us as neutral, theatmosphere naturally contains acidic oxides of carbon, nitrogen

and sulfur. The concentrations of these acidic oxides have

been increasing since the Industrial Revolution

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Oxides are chemical compounds formed when an element reacts with oxygen.

For example, when a metal such as magnesium is burnt in oxygen, themagnesium burns with a brilliant white flame to generate a white powder(magnesium oxide).

2Mg(s) + O2(g) g 2MgO(s)

When magnesium oxide is added to water containing universal indicator theindicator changes from green to blue-violet.

This shows that magnesium oxide is a basic oxide.

This reaction occurs because of the presence of water. The magnesium oxidedissolves in the water to generate magnesium hydroxide. The hydroxide ionsthat are released cause the colour change of the indicator.

2MgO(s) + H2O(l) g 2Mg(OH)2(aq)

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Remember

The general equation for Neutralisation is

Acid + Base g Salt + Water

The salt that is formed is an ionic compound that is named after their parent

acid... with the metal ion of the base is named first, followed by the name of theanion of the parent acid.

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If we take an amount of an oxide, dissolve them in water, then the resultingsolution may display acidic or basic properties.

The general trend is

1. Metal Oxides usually form basic solutions.remember that basic aqueous solutions are called Alkalis

2. Non-Metal Oxides usually form acidic solutions.

3. Several elements that are located on the boundary between the metalsand the non-metals, form solutions that are both acid and basic in nature,and these are called Amphoteric oxides.

Oxides - Acidic or Basic?

Amphoteric = able to react chemically as either an acid or a base

4. The Inert Gases do not form oxides!

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In summary

Also there are some other trends to understand

1. The strength of the basicity of metallic oxides increases as you go down agroup, hence BaO is a stronger basic oxide than MgO.

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2. The strength of the acidity of non-metallic oxides decreases as you go down

a group (examine the group 5 N, P, As, Sb, Bi note the change fromacidic to basic oxides).

3. Generally, the higher the oxidation state of the metal or non-metal, themore amphoteric or acidic the oxide Cl2O7 is more acidic than Cl2O (infact the effect of differing oxidation states on the acidity of oxides makes it

difficult to discuss trends in acidity and basicity in the oxides of the elementsof the periodic table).

Our atmosphere is a sink for many pollutants. Some of these pollutants areoxides.

Volcanoes and geysers release many different gases into the atmosphere,including sulfur dioxide.

Oxides & Pollution

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Carbon monoxide, carbon dioxide and sulfur dioxide are also formed during thecombustion of organic matter during bushfires.

Bacterial decomposition can also release sulfur dioxide into the environment.

Lightning storms produce toxic nitrogen oxides as nitrogen and oxygen

molecules react together.

The activities of human beings result in the release of oxides into theatmosphere.

Industries generate a variety of oxides such as carbon monoxide (CO), carbon

dioxide (CO2), sulfur dioxide(SO2) and the oxides of nitrogen (NOX).

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The natural cellular respiratory processes of animals and plants produce carbondioxide, which is ultimately released into the atmosphere.

Green plants utilise the carbon dioxide for photosynthesis.

Large amounts of carbon monoxide and carbon dioxide are also producedduring bushfires, since the process of incomplete combustion produces carbonmonoxide, whereas carbon dioxide is the product of complete combustion.

The combustion of fossil fuels such as petrol, kerosene and diesel oil releaseslarge amounts of the carbon oxides into the atmosphere.

Coal-fired power plants also release vast amounts of carbon oxides into theatmosphere. The increasing levels of carbon dioxide in the atmosphere areoften linked to global warming.

Carbon monoxide does not build up over time in the atmosphere as it is rapidly

removed by the action of soil organisms or by oxidation to carbon dioxide.

i. Carbon Oxides

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The oxides of sulfur are generally irritating, poisonous gases.

These gases particularly affect people who suffer from respiratory problemssuch as asthma and emphysema.

About 50% of the sulfur dioxide that enters our atmosphere is derived from theoxidation of hydrogen sulfide (H2S) produced from bacterial decomposition.

Many oxides of sulfur are also released into the atmosphere during volcaniceruptions.

Sulfur dioxide is produced by the combustion of fuels such as coal and dieseloil. In Australia considerable quantities of coal are burnt to generate most ofour electricity (~75%).

Fossil fuels usually contain small quantities of sulfur minerals (e.g. FeS2)although the bituminous coal from eastern NSW is lower in sulfide minerals than

coal from other locations.

ii. Sulfur Oxides

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These sulfide minerals in coal are oxidised during combustion, and sulfur

dioxide is released.

Metal smelters that convert sulfide minerals into metals are also a major source

of sulfur dioxide pollution.

4FeS2(s) + 11O2(g) g 2Fe2O3(s) + 8SO2(g)

Example

The smelting of chalcopyrite (CuFeS2) during the production of copper results inthe release of sulfur dioxide. Towns near these smelters can suffer from sulfur

dioxide pollution.2CuFeS2(s) + 5O2(g) + 2SiO2(s) g 2Cu(l) + 4SO2(g) + 2FeSiO3(l)

In regions such as the lower Hunter, about 95% of SO2 emissions areassociated with power and smelting industries.

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Sulfur trioxide is produced mainly by oxidation of sulfur dioxide in the

atmosphere. Oxygen and ozone are the oxidants involved in this oxidationprocess.

2SO2(g) + O2(g) g 2SO3(g)

SO2(g) + O3(g) g SO3(g) + O2(g)

Nitric oxide (NO) and nitrogen dioxide (NO2) are the most common atmosphericoxides of nitrogen found in urban air.

In Sydney, about 86% of the total emissions of NOX comes from the engines ofmotor vehicles and other transport.

The nitric oxide (a colourless gas) is formed when nitrogen and oxygen react athigh temperatures. Nitric oxide is a neutral oxide.

iii. Nitrogen Oxides

N2(g)

+ O2(g)

g 2NO(g)

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NOX is also derived from other sources, including industries, electrical power

production and oil refining.

NO2 is of concern as it causes damage to the respiratory system of humans aswell as irritating the eyes.

Young children and older people are particularly susceptible.

Un-flued gas heaters and kerosene heaters also contribute to NO2 pollution ofthe air inside houses.


NO2 is a brown gas and is produced by oxidation of NO. The rate of NO2

formation depends on the concentration of NO in the air. Up to 10% of the totalNOX in air is NO2. Nitrogen dioxide is an acidic oxide.

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

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During the industrial revolution of the 19th century, large amounts of coal wereburnt to provide power for factories and their machines.

Vast quantities of carbon dioxide and sulfur dioxide poured into the air. Iron

smelters generated large volumes of sulfur dioxide as they produced thegrowing quantities of steel required for industry.

The atmosphere of large industrialized cities in Europe and the USA becamehighly polluted with acidic gases.

The increased use of motor vehicles in the 20th century (particularly after 1945)increased oil consumption.

Emissions of sulfur dioxide doubled in the 25-year period following World WarII.

Monitoring Atmospheric Pollution

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Adding to this pollution burden on the atmosphere was the increasing

production of nitrogen oxides in internal combustion engines.

Following numerous deaths (about 4000) in London in 1952 due to heavy acidicsmogs, pollution controls began to be introduced to clean up the air of theselarge cities.

High levels of photochemical smog (produced by the action of sunlight on aircontaining moisture, ozone, hydrocarbons and NOX) in cities such as LosAngeles and Tokyo in the 1960s accelerated the push for emission controls onmotor vehicles.

In the 1970s the development of more sensitive gas analysis technologiesallowed chemists to monitor the global increase in sulfur dioxide emissions dueto the expansion of industries in Asia.

In recent years the air quality has improved in most westernized countries. InEurope, the sulfur dioxide emissions dropped by about 45% in the 1990s.

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In the same period, nitrogen dioxide emissions dropped by about 20%.

However, due to increasing population and usage of motor vehicles, the levelsof pollutants have stabilised rather than continuing to decrease.

The EnvironmentalProtection Agency of

NSW (EPA) monitorsthe levels of pollutantgases in theatmosphere in manyregions across NSW.

The NSW Departmentof Environment andConservation preparesquarterly reports onair quality.

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This data confirms that the air quality in Sydney with respect to SO2 and NO2 ishigh when compared with NEPC standards.

The data shows no discernible trends overall across the 5-year period.

The air in Sydney compares favorably with that of much larger cities overseas.

The large industrialised cities in China, the USA and Europe have much higherlevels of atmospheric NO2 and SO2than Sydneys.

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this can be achieved if people use appliances that are more energy-efficient; asindividuals we can help by reducing our own personal use of electricity


In order to reduce emissions of sulfur dioxide and nitrogen oxides there are anumber of options available...

Reducing Atmospheric Pollution

i. Reduce the amount of coal being burnt

Power companies could switch to burning low-sulfur coal rather than high-sulfurCoal. They can also switch to natural gas, which produces very little sulfurdioxide on combustion

ii. Use Coal with lower sulfur content

The development of other forms of power production such as hydroelectricityand solar power, this reducing our need to burn fossil fuels.

iii. Reduce the reliance of Fossil Fuels

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Techniques such as collecting the sulfur dioxide produced by smelting metalsulfides and the using it to make sulfuric acid (a very important industrialchemical) could be employed.

iv. Recycling the gases

The acidic gases are passed through a slurry of a base such as calcium oxide;would cause the sulfur dioxide to react with the calcium oxide to form solidcalcium sulfite. Hence the gas is removed from the smoke that is released intothe atmosphere.

v. Flue Gas Scrubbing

Most modern cars are equipped with catalytic converters, which convert thenitrogen oxides back to nitrogen gas.

vi. Catalytic Converters

2NO(g) + 2CO(g) g N2(g) + 2CO2(g)

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When all gases are removed from pure water, its pH is 7.0 (neutral) at 25C.

However, natural water contains dissolved gases including carbon dioxide, whichmakes the water weakly acidic (pH 6.06.5) due to the presence of carbonicacid.

Acid Rain

CO2(g) + H2O(l) g H2CO3(aq)

When the atmosphere is polluted with acidic oxides such as sulfur dioxide andnitrogen dioxide, rainwater can become quite acidic (pH 4.05.0) due to thehigh solubility of these gases in water.

Rain that has such a low pH is calledAcid Rain.

Acid rain with a pH of 3.6 has been recorded in many severely pollutedindustrialised areas in Europe and the USA.

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SO3(g) + H2O(l) g H2SO4(aq)


When the oxides of sulfur and nitrogen dioxide dissolve in water they producesolutions of various acids.

Sulfur dioxide forms weak sulfurous acid (H2SO3) whereas sulfur trioxideproduces strong sulfuric acid (H2SO4).

SO2(g) + H2O(l) g H2SO3(aq)

2H2SO3(aq) + O2(g) g 2H2SO4(aq)

Sulfurous acid can be catalytically oxidised to produce sulfuric acid.

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Nitrogen dioxide produces weak nitrous acid (HNO2) and strong nitric acid(HNO3) when it dissolves in water.

2HNO2(aq) + O2(g) g 2HNO3(aq)

In the presence of water and oxygen the nitrous acid is catalytically oxidised tonitric acid.

2NO2(g) + H2O(l) g HNO2(aq) + HNO3(aq)

When you visit the ancient cities of Europe you may observe that many marblestatues and building facades are eroded.

When the calcium carbonate of the marble (& limestone) is attacked by thesulfuric acid in acid rain, the surface of the marble is converted into insoluble

calcium sulfate.

Effects of Acid Rain

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CaCO3(s) + H2SO4(aq) g CaSO4(s) + H2O(l) + CO2(g)

The wet calcium sulfate crystallises to form a porous and crumbly mineral calledgypsum (calcium sulfate dihydrate).

Over several centuries the soot from coal burning, as well as various dusts,have collected in the pores of the gypsum and turned the marble black.

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Acidic rain can also attack metallic structures composed of iron and steel. The

iron is oxidised by the hydrogen ions in the acid and becomes chemicallyweathered.

Fe(s) + H2SO4(aq) g FeSO4(aq) + H2(g)

Acid rain has had a devastating effect on many northern hemisphere forests,especially the famous Black Forest in Germany that has been significantlydamaged by acid rain.

Acid rain also affects the soil, and the acidified soils inhibit the growth of plantseedlings.

Basic minerals in the soil (such as dolomite and limestone) are attacked anddissolved by acidic water.

Many types of sandstone have grains that are cemented together with calcite(calcium carbonate), and the acid rain will dissolve this cement and so causesignificant chemical weathering and erosion.

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Equilibrium

The term Equilibrium is used in many science disciplines.

For example, in physics if a hot coin is dropped into a cup of cold water, thewater/coin system will reach thermal equilibrium when both water and coin

are at the same temperature.

Although the configuration (or large-scale properties) of a system in equilibriumdo not change over time, the small-scale, microscopic configuration of thesystem is not static & unchanging.

This is especially important in the study of chemistry since not all reactionsproceed to completion.

In many reactions, the final reaction mixture consists of both products andreactants

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Consider the following, chemical reaction

A + B g C + D

Molecules A & B react to form molecules C & D. the reactant molecules aretotally converted in to the product molecules, the reaction has gone tocompletion (combustion of ethanol).

However for many some reactions, once the molecules of C & D are formed,they can react together to produce the original molecules of A & B

C + D g A + B

The overall reaction is often written as

A + B C + D

Forward Reaction

Reverse Reaction

The system is inChemical Equilibrium.

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Microscopically the configuration of the equilibrium system is not static, becauseforward and reverse reactions are always occurring.

However since the microscopic reactions are in balance the large-scaleproperties of the system appear constant (no colour change etc).


Example

Consider the following, equilibrium system

Initially there would only be N2O4 molecules in the container.

N2O4(g) 2NO2(g)

Oxygen

Nitrogen Time

Concentratio

nN2O4

NO2

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Time

Con

centration

N2O4

NO2


Example continued

Over time, the N2O4 reacts to make NO2

Some of theNO2 molecules

will recombineto create N2O4and thesystem settlesinto anequilibrium.

Time

Concentratio

n

N2O4

NO2

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The Equilibrium Constant

Chemists use the Equilibrium Constant (Kc) as a mathematical tool to helpthem study systems in equilibrium.

Consider the following general equilibrium system

The equilibrium constant for this system is given by the formula

The equilibrium constant allows us to predict the position of the equilibrium.

aA + bB gC + dD

[C]g

[D]d

[A]a [B]bKc =

products

reactants

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very greater than one then virtually no reactants remain or the reaction goesto completion.


If the equilibrium constant (Kc) is

very less than one then virtually no products are formed or no reactionoccurs at all.

less than one then the [reactants] are favoured reverse reaction ispredominant little products are formed.

equal to one then the [reactants] & [products] are about the same neither the reactants or products are favoured.

greater than one then the [products] are favoured forward reaction ispredominant little reactants remain.

aA + bB gC + dDReactants Products

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Example

The concentrations of each of the substances in the above reaction wereanalysed

The equilibrium constant for this reaction is given by

[HI]2

[H2

] [I2

]Kc =

H2(g) + I2(g) 2HI(g)

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Example continued

H2(g) + I2(g) 2HI(g)

For experiment I:

[0.156]2

[0.0222] [0.0222]Kc =

Kc = 49.4

Repeating this calculation for the other experiments we get Exp II = 49.8, Exp

III = 49.4, Exp IV = 49.5 average for all four experiments = 49.5.Note the variancebetween values isdue to experimentalerror!!

Hence in this equilibrium system the products arefavored and there will be small amounts of thereactants remaining.

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Le Chatelier's Principle

Since no chemical equilibrium system can remain completely undisturbed by theoutside environment for a long period of time, humans began to investigate theeffect of altering systems that were at chemical equilibrium.

Le Chateliers principlegoverns disturbances to equilibrium systems

If a chemical system at equilibrium is disturbed, the equilibriummoves in the direction that tends to reduce the disturbance.

This principle allows us to predict the effect of any change we wish to make toequilibrium systems.

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There are several external factors that can potentially effect a system inequilibrium.

The most common being


Factors that Effect Equilibrium

1. TheAddition / Removal of a reactant or product.

2. Changing theVolume of the system.

3. Changing the Temperature of the system.

4. The addition of a Catalyst.

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The equilibrium will shift to consume the added reactant/product, orto replace the removed reactant/product.

1. Addition / Removal of a Reactant / Product

Example

If N2 is added

forward reaction favoured g more NH3 will be produced

N2 must be consumed

If NH3 is removed

forward reaction favoured g more NH3 will be produced

NH3 must be replaced

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

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Effect of change in concentration

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Changing the volume of a reaction system will only effect Gaseous reactants orproducts.

Remember from the preliminary course, the volume and pressure of a gas are

inversely proportional1

PV a

h Volume i Pressure i Volume h Pressure


2. Changing the Volume of the System

Also remember from the preliminary course, 1 mole of any gas has the same

volume if they are at the same temperature & pressure.

0C & 1atm = 22.71 Lmol-125C & 1atm = 24.79 Lmol-1

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So the effect of changing the volume (or pressure) of an equilibrium systemthat contains a gaseous material will depend upon the number of mole ofgas reactants or products.

The equilibrium will shift to ensure that the total number of mole ofgases will be advantageous in the altered system.


Example

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

4 mol gas reactants 2 mol gas products

IfiVolume hPressure forward reaction favoured g more NH3 will be produced

#mol of gas will want to decrease

IfhVolume iPressure reverse reaction favoured g more H2 & N2 will be produced

#mol of gas will want to increase

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Effect of change in pressure

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3H2(g) + N2(g) g 2NH3(g) DH = -46.19 kJmol-1

Remember from the preliminary course, exothermic reactions give off heatand have a negative DH value

If heat is given off then you can imagine heat as one of the products

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

2NH3(g) g 3H2(g) + N2(g) DH = +46.19 kJmol-1

Likewise, endothermic reactions absorb heat and have a positive DH value

If heat is absorbed then you can imagine heat as one of the reactants

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


2. Changing the Temperature of the System

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So the effect of changing the temperature of an equilibrium system will dependupon exothermic/endothermic nature of the forward reaction.

Decreasing the temperature of an equilibrium favour the exothermicreaction direction.


Example

If temperature is increased reverse reaction is enhanced g NH3 will be consumed

endothermic reaction is favoured

3H2(g) + N2(g) 2NH3(g) DH = -46.19 kJmol-1

If temperature is decreased

forward reaction is enhanced g NH3 will be produced

exothermic reaction is favoured

Negative DHForward Reaction is Exothermic

+ HEAT

added heat removed heat

removed heat added heat

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Effect of changing temperature

For the N2O4/NO2 system, the forward reaction isendothermic.

N2O4(g) + 57 kJ 2NO2(g)

The forward reaction therefore absorbs heat and thereverse reaction releases heat.

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Catalysts speed up reactions without being altered in the reaction.

The addition of a catalyst will speed up the forward and reversereaction rates, but will have no effect on the reactants or products ofthe reactions.

So the composition of the equilibrium mixture is unchanged, the catalyst simplyhelps the reaction reach equilibrium faster.


4. The Addition of a Catalyst

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Carbon Cycle

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Aqueous CO2 Equilibrium

NB: solubility of gases increase as temperature decreases

Carbon Dioxide Water Equilibrium

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Carbon Dioxide Water EquilibriumThe reaction between carbon dioxide and water

CO2(g) + H2O(l) H2CO3(aq)is of great practical significance.

The removal of carbon dioxide from our bodies

The transport of carbon dioxide in photosynthesisThe removal of carbon dioxide from the atmosphereThe preparation of aerated drinks.

The reaction does not go to completion, rather it is anequilibrium reaction.

CO2(g) + H2O(l) H2CO3(aq)

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Carbon dioxide is an acidic oxide, that dissolves in water


Aqueous CO2 Equilibrium

CO2(g) g CO2(aq)

to produce a solution of carbonic acid (H2CO3)

and the carbonic acid is in equilibrium with hydrogen ions and hydrogen

carbonate ions


H2CO2(aq) g H+

(aq) + HCO3-(aq)

This dissolution of CO2 in water, overall is exothermic.

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Hence soft drink tastes sour because of their acidity (pH ~ 4).


Carbonated soft drinks are manufactured by dissolving carbon dioxide in water

under pressure (~400500kPa), hence the water in the soft drink issupersaturated in carbon dioxide.

If the bottle is sealed, high pressure in the space under the cap causes...

CO2(g) g CO2(aq)


H2CO2(aq) g H+(aq) + HCO3-(aq)

carbon dioxide is dissolved

favored

carbonic acid created

favored

H+ is created increasing theacidity of the drink

favored

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When the cap of a soft-drink bottle is unscrewed, there is a rapid effervescence

observed, and overtime the soft drink eventually goes flat and tastes less sourdue to the reversal of these three equilibria.

The equilibrium of a soft drink can also be altered by changing its pH, andheating/cooling the drink.

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Gas volumes involved in chemical reactions can be calculated using very simpletechniques.


Calculating Gas Volumes

Avogadros

Number

Avogadros

Number

Gas Constant

Gas Constant

Molar Mass

Molar Mass

Volume (L)

Volume (L)

only if an aqueous solutiononly if in the gas state

# mol

mass

(g)

gas volume

(L)

# species

(atoms, molecules, ions)

concentration

(molL-1)

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Hence the volume of gases can be directly compared to the #mol of gas in a

sample, using the following constants

1 mol of any gas at 100kPa and

at 0C (273.15 K) g 22.71 L

at 25C (298.15 K) g 24.79 L

Example

If 4.5g of Sodium Carbonate is reacted with excess Hydrochloric Acid, calculatethe volume of CO2 gas that is produced (assume conditions to be 25C & 100kPa).

#mol Na2CO3 = 4.5 106 = 0.042mol

2HCl(aq) + Na2CO3(s) g 2NaCl(aq) + CO2(g) + H2O(l)

Na2CO3:CO2 g 1:1

#mol CO2 = 0.042mol

volume CO2 = 0.042 24.79

Documents

9.3 - Section 2