TM THE ACIDIC ENVIRONMENT - Amazon S3s3.amazonaws.com/.../Chem6.AcidicEnvironmentStudent.pdfAcid, Base or Neutral? The chemical definitions of acid and base will come later. For now,

keep it simple science TM

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Emmaus Catholic College SL#802440

1

but first, an introduction...

HSC Chemistry Topic 2

THE ACIDIC ENVIRONMENTWhat is this topic about?To keep it as simple as possible, (K.I.S.S.) this topic involves the study of:1. ACIDS, BASES & INDICATORS

2. ACIDS IN OUR ENVIRONMENT3. ACIDS & THE pH SCALE

4. ACID-BASE THEORY & TITRATION5. ESTERIFICATION

...all in the context of Chemistry in our environment and society.

Way back in years 8, 9 and 10 you would have studied somebasic Chemistry, and it probably involved studying acidsand bases and indicators... so these are familiar terms, evenif you�’ve forgotten the details.

This topic begins by reminding you of this simple way toclassify all chemical substances:

Acid,Base,

or Neutral

and how �“Indicators�”can be used to identify them.

Then, we look at the Chemistry of Oxide Compounds,

and link that to Acidity.

You will study the sourcesand problems that can result from

Acids in theEnvironmentand, revise & practice

MoleCalculations

The details of Acid-Base Reactions will be studied

...and you will learn the real meaning of thepH scale, used for measuring acidand base strength.

+ +

proton transfer

ACID BASE

pHH = -lloog[HH++ ]]

Along the way, you will learnsome important techniques inChemical Analysis

and a process you might neverhave heard of before,

but it is common in nature andin the Chemical Industry:Esterification

Phot

o by

Ken

Kis

er

HSC Chemistry Topic 2Copyright © 2006-2007 keep it simple science


www.keepitsimplescience.com.auHSC Chemistry Topic 2Copyright © 2006-2007 keep it simple science


2

Strong v. WeakAcids

Self-iionization0f Water

pH Scale&

pH Calculations

Indicators&

Colour Changes

Acid, Base& Neutral

Everyday Uses of

Indicators

DynamicEquilibrium

Alkanols&

Alkanoic Acids

Buffers

EstersStructure & Naming

Occurrence &

Usesof Esters

Esterification Reaction&

Reflux Technique

Acid-BBaseBehaviour of

Water.AMPHIROTIC

Early Ideas aboutAcids & Bases

Neutralization&

Titration

Acidic Oxides in the

Environment

Acidic & BasicSalts

Acids as ProtonDonors

Mono-,, Di- & Triprotic acids

Acids & Alkalis.Acidic & Basic

Oxides

TTHHEE AACCIIDDIICC

EENNVVIIRROONNMMEENNTT

Acids, Bases&

Indicators

Acidsin our

Environment

Acids&

the pH Scale

Acid-BBase Theory&

Titration

Esterification

CONCEPT DIAGRAM (�“Mind Map�”) OF TOPICSome students find that memorizing the OUTLINE of a topic helps them learn and remember the concepts andimportant facts. As you proceed through the topic, come back to this page regularly to see how each bit fits thewhole. At the end of the notes you will find a blank version of this �“Mind Map�” to practise on.

Bronsted-LLowryTheory

Le Chatelier�’sPrinciple

Proton Transfer&

Conjugate Species

Standard Solutions,Choosing Indicators

TitrationGraphs

Acid, Base or Neutral?The chemical definitions of acid and base will come later.For now, you are reminded of what you may have learnt inearlier Science classes.

Acids and bases are chemical opposites; if you add one tothe other they �“destroy�” (neutralize) each other, and theend product is neutral. Acids and bases are the oppositeends of a chemical property called �“acidity�”, which ismeasured by a numerical scale called �“pH�”.

The word �“acid�” comes from the Latin for �“sour�”, andrefers to the fact that natural, acidic chemicals (e.g. vinegar)are sour-tasting.

IndicatorsIndicators are chemicals which change colour according tothe acidity of the solution they are in.

The original indicators were natural extracts from plants orother living things. Some, such as litmus, are still in usetoday, as well as newer, synthetically made chemicals.

Modern Laboratory IndicatorsThe syllabus requires that you are familiar with thecommon laboratory indicators listed below.

You will have done experimental work, adding 2 drops ofindicator to test tubes of acid, base and pure water (whichis neutral) and recorded the colours produced.

Your results should have agreed with the following:

Colour inIndicator Acid Water BaseLitmus pink purple blue

Phenolphthalein clear clear red/pink

Methyl Orange red yellow yellow

Bromothymol blue yellow green/blue blue

Choosing an IndicatorWhy are so many indicators needed?

Litmus is useful for general indentification of acidic orbasic substances. However, its colour change is ratherindistinct, and can occur over quite a range of pH values...it is not a �“sharp�” change.

In contrast, Phenolphthalein cannot tell you the differencebetween a glass of water and sulfuric acid. However, thecolour change is very distinctive, and occurs suddenly at avery specific pH value... it is very sharp. This is not muchuse in general identification of substances, but in certainmethods of chemical analysis it is very important.

So, each indicator has a purpose and preferred use.

Everyday Uses of Indicators

�• Soil Testing. Some plants grow best in acidic soils; othersneed slightly alkaline (basic) conditions. Farmers and keengardeners use simple test kits containing an indicator andcolour chart, to test the soil. They can then adjust the soilpH to get the best results.

�• Water Testing. Swimming pools need regular testing foracidity to better maintain their water quality and hygiene.Aquariums must be maintained at very specific pH levelsfor the health of their inhabitants.

�• Effluent Testing. Acidity is a useful way to assess the levels of certain types of pollution from industries.Industry technicians and Government authorities useindicators to monitor the pH of waste water and naturalwaterways.

3


1. ACIDS, BASES & INDICATORS




ACID NEUTRAL BASE

1 3 5 7 9 11 13pH

Practical Work: A Natural IndicatorYou may have done practical work to prepare and test anatural indicator.

A good example isthe commongarden plantHydrangea.

If you collect aflower head andput it through ablender with a littlewater and ethanol,the filtered liquidextract will work asa simple indicator.

In acid, the Hydrangea flower extract is a bluish colour.

In a base, it turns pink-ish.




4

Practical Work: Classifying Household SubstancesYou may have done laboratory work using various liquidand paper indicators to classify a range of householdsubstances and foodstuffs as either acid, base or neutral.

By using appropriate indicators (such as �“Universal�”) youmay have even been able to differentiate betweensubstances that are �“mildly acidic�” and �“strongly acidic�”.

A Few Typical Results

Household Substances Found to be...

ACIDIC NEUTRAL BASICVinegar Salt SoapOrange juice Sugar Floor cleanerTomato juice Shampoo Drain cleaner

MilkLiquid detergent

Worksheet 1

Part A Fill in the Blanks

Acids and bases are chemical a)................................ If youadd one to the other, they b)............................................. eachother. Acidity is measured by the c).................... scale. Onthis scale, a neutral substance has a value of d)................Values above this indicate e)........................................substances, while values below indicate f)................................substances.

Indicators are chemicals which g)..................... ........................according to the h)..................................... of the solutionthey are in. The original indicators where extracts fromi).......................................................

The colour changes for the following indicators need to belearnt.Indicator Acid Neutral BaseLitmus j)................ purple k)................Phenolphthalein l)................ ............... .................Methyl orange m)............... ................ .................Bromothymol blue n)............ green/blue o)...............

In everyday situations, indicators are used for purposessuch as p).................................................. for farming andgardening, q).......................................... for pools andaquariums, and for monitoring r).............................................from industries.

Part B Practice Problems

Each solution listed below has been tested with one ormore indicators, and the colour is given. For each, state if itis acidic, neutral or basic.

Solution A: Phenolphthalein is clear.Methyl orange is red.

Solution B: Phenolphthalein is pink.Methyl orange is yellow.

Solution C: Phenolphthalein is clear.Methyl orange is yellow.

Solution D: Bromothymol blue is blue.Methyl orange is yellow.

Solution E: Phenolphthalein is clear.Litmus is pink.

Photo by Jan Friml

Laboratory Acids and �“Alkalis�”You are reminded that the common laboratory acids are:

Name Formula Solution ofHydrochloric acid HCl H+

(aq)+ Cl-(aq)

Sulfuric acid H2SO4 2H+(aq)+ SO4

-2(aq)

Nitric acid HNO3 H+(aq)+ NO3

-(aq)

The most familiar laboratory bases are soluble hydroxidecompounds which are sometimes called �“alkalis�”.

Name Formula Solution ofSodium hydroxide NaOH Na+

(aq)+ OH-(aq)

Potassium hydroxide KOH K+(aq)+ OH-

(aq)

The Acid-Alkali Reaction: NeutralizationYou should be familiar with this reaction from earlierScience studies:

Example:

sulfuric + potassium water + potassiumacid hydroxide sulfate

H2SO4(aq) + 2KOH(aq) 2H2O(l) + K2SO4(aq)

Ionic equation:2H+

(aq)+ SO42-

(aq)+ 2K+(aq)+ 2OH-

(aq)

2H2O(l)+ 2K+(aq)+ SO4

2-(aq)

If you study this equation you will see that the potassiumand sulfate ions are spectators. You can leave them out toform the net ionic equation:

2H+(aq)+ 2OH-

(aq) 2H2O(l)or more simply:

H+(aq)+ OH-

(aq) H2O(l)

All the familiar laboratory acids are solutions containinghydrogen ions (H+). The laboratory �“alkali�” bases aresolutions containing hydroxide ions (OH- ). These willalways react to form water, so acid and base haveneutralized each other.

The other product is always a soluble, ionic compound.These are known collectively as �“salts�”.

Acid-Base Properties of the Oxide Compounds

The common laboratory acids and alkali-bases are not thewhole story. Most of the oxide compounds of theelements show some acid-base behaviour too, and theperiodic table reveals another pattern.

Basic Oxides of the MetalsMost of the oxides of the metallic elements are consideredbasic because they can neutralize acids, forming water anda �“salt�”.

Examples:hydrochloric + magnesium water + magnesium

acid oxide chloride2HCl(aq) + MgO(s) H2O(l) + MgCl2(aq)

sulfuric + copper(II) water + copper(II)acid oxide sulfateH2SO4(aq+ CuO(s) H2O(l) + CuSO4(aq)

Acidic Oxides of the Non-MetalsMany of the oxide compounds of the non-metalelements are acidic because they will:

�• react with water to form an acidand/or�• react with a base by neutralizing it, and form water and

a �“salt�”.

Examples:carbon dioxide + water carbonic acid

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

carbon + sodium water + sodiumdioxide hydroxide carbonate

CO2(g) + 2NaOH(aq) 2H2O(l) + Na2CO3(aq)

5


2. ACIDS IN OUR ENVIRONMENT




ACID + ALKALI WATER + A �“SALT�” ACID + METAL WATER + A �“SALT�”OXIDE

Try the WORKSHEET at the end of sectionTry the WORKSHEET at the end of section

OOxxiiddee ooff hhyyddrrooggeenniiss wwaatteerr;; nneeuuttrraall

IInneerr

tt GGaass

eess hh

aavvee

nnooooxx

iiddee

ccoomm

ppoouunn

ddss

OOxxiiddeess ooff tthee mmeettaalss

aacctt aass baasseess......

BBASSIICC OOXXIIDEESS

Non-mmetal oxides(incl. semi-mmetals)are usually ACIDIC



6

Chemical EquilibriumThe concept of �“Dynamic Equilibrium�” was introduced ina previous Preliminary topic. The example used then wasthe equilibrium between a solid ionic lattice and dissolvedions in a saturated solution.

At equilibrium, it seems (macroscopically) that nothing ishappening. However, at the atomic level, solid salt isdissolving into the solution;

...but at the same time, some ions are precipitation out ofsolution to form a solid lattice;

...and these 2 processes are occurring at the same rate.

Equilibrium of Carbon Dioxide and WaterWhen CO2 dissolves in water, it doesn�’t merely dissolve,but reacts to form the weak acid �“carbonic acid�”.

This is a dynamic equilibrium situation. Even when theprocess seems finished, there are actually 2 reactions (oneforwards , one backwards ) occurring at the samerate, so that nothing appears to be happening.

Shifting an EquilibriumA chemical equilibrium is �“dynamic�”, meaning that thingsare moving back and forth. This also means that it ispossible to �“upset�” an equilibrium and cause it to shift to anew balance between reactants and products.

Equilibrium shifting to the left results in less carbonic acidand the formation of more CO2 gas. The shift in theequilibrium was caused by the pressure change whichoccurred when the bottle/can was opened.

Factors That Can Cause an Equilibrium Shift

PressureIf a reaction involves a gas, (as in the example at left) anychange in the pressure of that gas will shift the equilibrium.With soft drinks, increasing the pressure of CO2 drives theequilibrium to the right; decreasing CO2 pressure, shifts itto the left.

TemperatureIf you kept the gas pressure constant, but raised thetemperature, you would find this equilibrium would shiftleft. (i.e. less carbonic acid) Lowering the temperaturewould shift it to the right (more carbonic acid).

Which way an equilibrium shifts due to temperature,depends on whether the reaction is exo- or endo-thermic.This is explained later.

ConcentrationThe equilibrium could also be shifted by altering theconcentration of (say) the carbonic acid. If you added extraH2CO3(aq) somehow, the equilibrium would shift left. Ifyou reduced the concentration, (e.g. by adding an alkaliwhich would react and destroy it) the equilibrium wouldshift towards the right.



e.g. NaCl(s) Na+(aq) + Cl-(aq)

NaCl(s) Na+(aq) + Cl-(aq)


Practical Work: Mass & Volume of CO2 in a Soft DrinkYou will have done a simple laboratory exercise to �“de-carbonate�” a fizzy soft drink and measure/calculate themass and volume of CO2 gas released.

If you weigh the soft drinkand container accurately,then release the pressure sothat it goes flat (this maytake a day or more) you willhave measured a small lossof mass, due to CO2 gasescaping.

Typical Results& AnalysisFor a 300mL bottle oflemonade:

Start mass = 536.9gFinal mass = 535.8g! mass loss = 1.1g

moles of CO2:n = m/MM

= 1.1/44.01n(CO2) = 0.025 mol.

Volume of CO2 (assuming standard conditions)V = n x VM

= 0.025 x 24.8 = 0.62 L!! 0.62 L (620mL) of CO2 gas was released

For example, when you open andpour a fizzy soft drink it formsbubbles in the glass.

The sealed drink containsdissolved CO2, but as soon as it isopened, the gas begins comingout of solution, because theequilibrium has shifted to the left.

Eventually a new equilibrium isestablished with a lot less gasdissolved... we say it has �“goneflat�”.


NaCl(s) Na+(aq) + Cl-(aq)




7

Le Chatelier�’s PrincipleIn 1885, the French Chemist Henri Le Chatelier discoveredthe underlying pattern in these equilibrium shifts. Theequilibrium always shifts in the direction which counteractsthe change that upset it in the first place.

Temperature Effect on EquilibriumThe dissolving of CO2 to form carbonic acid is anexothermic reaction, so heat energy may be considered asone of the products:

So, if the temperature is raised, the equilibrium shifts left

... because shifting left would use up heat andlower the temperature again.

If the temperature is lowered, the equilibrium shifts right

... in an attempt to release heat energy and raisethe temperature again.

Effect of ConcentrationImagine a sealed container with an equilibrium mixturewithin.

If you injected extra product so its concentration increased,the equilibrium shifts left, attempting to reduce theconcentration of product

If you somehow removed product so its concentration wasreduced, the equilibrium would shift right in an attempt tomake more.

In every case, the equilibrium shift tries tocounteract the change

Le Chatelier�’s Principle:If a system in equilibrium is disturbed,

the system will adjust itself in the direction whichcounteracts the disturbance

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

H2CO3(aq) + heat energy

CO2(g) + H2O(l)

H2CO3(aq) + heat energy

CO2(g) + H2O(l)

Explaining the Pressure Effect

H22CO33((aaqq))

CO22((gg))

Initial EquilibriumSituation

Extra CO22 pumped in.Pressure (and concentration)

of CO22 increased.

CO22 pumped out. Pressure (and concentration)

of CO22 decreased.

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

H22CO33((aaqq)) concentrationincreases as equilibriumshifts right, attemptingto use up the extra CO22

H22CO33((aaqq)) concentrationdecreases as equilibriumshifts left, attempting tomake more CO22

SSeeaalleedd

CO22((gg))

H22CO22((aaqq))

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

Extra H22CO33injected intosolution

EEqquuiilliibbrriiuummSShhiifftt

The Response of an Endothermic Reaction is exactly the opposite




8

Acidic Oxides in the EnvironmentLearning about Equilibrium was a necessary diversion...now back to acids.

There are 3 main acidic oxides that are of concern in theenvironment:

All 3 are gases produced by both natural processes and byhuman activities. All 3 react with water in the environmentto form acids:

carbon dioxide + water carbonic acidCO2(g) + H2O(l) H2CO3(aq)

sulfur dioxide + water sulfurous acidSO2(g) + H2O(l) H2SO3(aq)

nitrogen + water nitrous acid + nitric aciddioxide2NO2(g)+ H2O(l) HNO2(aq)+ HNO3(aq

carbon dioxide

CCOO2

nitrogen dioxide

NNOO2sulfur dioxide

SSOO2

Carbon Dioxideis a part of the great �“Carbon-Oxygen�” cycle in nature.

CO2 is also released into the air by natural bush fires, andby volcanic eruptions, but generally there�’s a balance.

For millions of years, huge quantities of carbon have been�“locked-away�” in fossil materials such as coal andpetroleum.

For the last 100 years or so, human activity has beenreleasing this �“fossil carbon�” as CO2 by burning the fossilfuels. CO2 levels have risen 30% or more, upsetting thebalance, world-wide.

This is believed to be causing environmental problems(�“Global warming�”)

due to the �“Greenhouse Effect�”.

CO22

O22

All living things.Respiration

Plants.Photosynthesis

Sulfur Dioxideis released into the air fromvolcanic eruptions and hot-springs, but the naturallevels are extremely low.

Human activities can pollutethe air of a region byreleasing SO2 from:

Smelting of Sulfide Ores (Revise Prel. topic �“Metals�”)The ores of some metals (esp. lead, zinc, copper) containsulfide compounds. To extract the metal, the ore isroasted with air:

lead(II) sulfide + oxygen lead + sulfur dioxidePbS + O2 Pb + SO2

Burning of Fossil FuelsSome fuels, especially coal, contain small amounts ofsulfur-containing compounds. When the coal is burnt, thesulfur burns too:

sulfur + oxygen sulfur dioxideS + O2 SO2

The main environmental concern with SO2 release is�“Acid Rain�”

GGeeyysseerr eerruuppttiinngg,,YYeelllloowwssttoonnee NNPP Nitrogen Dioxide

also occurs naturally in extremely small quantities due to thereaction between O2 and N2 when lightning provides thenecessary energy:

nitrogen + oxygen nitric oxide N2(g) + O2(g) 2 NO(g)

Nitric oxide is NOT acidic, but it reacts with O2:

nitric oxide + oxygen nitrogen dioxide2NO(g) + O2(g) 2NO2(g)

The combination of NO(g) and NO2(g) is referred to asthe �“NOx�” gases.

High temperaturecombustions (especiallycoal-burning powerstations and vehicleengines) produce largeamounts of NOx gases.

NOx pollution cancause toxic �“smog�”around large cities.

Sources of Acidic Oxides in the Environment

Photo by Diana



9

Environmental ImpactsEach of these acidic oxides can cause a different majorenvironmental problem:

The main problem with SO2 pollution is its acidity.

SO2 reacts with water in the environment:

sulfur dioxide + water sulfurous acidSO2(g) + H2O(l) H2SO3(aq)

Sulfurous acid is a strong acid and where SO2 pollution isserious, the acidity of rainfall stings the eyes, corrodesmetals, erodes stone buildings and monuments, and canhave serious environmental impacts:

�• Lakes and wetlands can become acidic enough to killplants and animals and disrupt the food-chains and thenormal ecological balance.

�• Forests can be killed by acidity of the rain, and theleaching of the soils by acids.

In the 1960�’s (when Acid Rain became known) many forestswere seriously damaged, including the famous �“Black Forest�”of Germany. In Canada, 15,000 lakes were known to be�“dead�” in an ecological sense.

Since then, emissions of SO2 have been limited and damagereduced, but the threat of Acid Rain is still a serious one,especially in rapidly industrializing countries such as China.

The Evidence for Acidic Oxide PollutionWe know that CO2 levels in the atmosphere have increasedbecause measurements have been collected over manyyears. The measurements of SO2 and NOx gases have notbeen collected for as long, and these gases are rapidly�“washed�” out of the air by rain, so the evidence for theirpresence is not so certain.

What we can be certain about are the localized effects ofSO2 pollution in places where it is, or was, prevalent. Someexamples from the 1960�’s were mentioned before. Closer tohome and to the present, is the evidence of devastationaround Queenstown, Tasmania.

The environmental damage around Queenstown wascaused by the release of SO2 from the Mt.Lyell coppermine and smelter. Although it shut down over 20 years ago,the environment has still not recovered.



carbon dioxide

CCOO2

CO2 is the weakest acid of these 3,and its acidity is not the problem.You should already be aware of the�“Greenhouse Effect�” and �“GlobalWarming�”

Acidity is not the main concern withthe NOx gases either. In the nexttopic you will study the Chemistry of�“smog�” and ozone pollution. (ButNO2 does contribute to �“Acid Rain�”)

nitrogen dioxide

NNOO2

and

�“Acid Rain�”

sulfur dioxide

SSOO2

Deforestation

GGrraavveell ppllaayyiinngg ffiieelldd((ggrraassss wwoonn�’�’tt ggrrooww))

Hills bare of vegetation

How Much Gas is Produced?The Queenstown Copper Industry was smelting a copperore containing mainly copper(I) sulfide. After concentratingthe ore by froth flotation, the material was smelted byroasting in a furnace with a blast of air to provide oxygen:

Copper(I) sulfide + oxygen Copper + Sulfur dioxideCu2S + O2 2Cu + SO2

How much SO2(g) is produced from each tonne of Cu2S?

Solution: 1 Tonne = 1,000 kg = 1.00 x 106 gram

moles of Cu2S: n = m/MM (MM = 159.2g)= 1.00x106/159.2= 6.28 x 103 mol.

mole ratio in equation is 1:1! moles of SO2 produced = 6.28 x 103 mol.

mass(SO2): m = n x MM (MM = 64.07g)= 6.28x103 x 64.07= 4.02 x 105 g (402 kg)

volume(SO2): V = n x VM(at SLC) = 6.28x103 x 24.8

= 1.56 x 105 L (156,000 litres!!)

Photo by Diana

Photo by Diana

Worksheet 2

Part A Fill in the blank spaces

Acids and bases react with each other to form a)...................and a �“b)............................�” and c).............................................each other.

The oxides of most metals act as d).............................., inthat they will react with an acid to form e)...................... anda ......................................... Oxides of non-metals mostly actas f)................................... because they will either �• dissolve in/react with water to form an g)......................and/or�• react with bases and h).......................................... them.

Chemical i)....................................... occurs when a reaction,and its opposite are occurring at the same j).......................,so that the concentrations of reactant(s) and k).....................do not change. The equilibrium is said to be l)........................

An equilibrium can be �“upset�” by a change inm)...................................., .................................................. or..................................................... When a change occurs, itwill �“shift�” to a new equilibrium position according ton)...................................................... Principle. This states thatwhen an equilibrium is disturbed, the system will shift in thedirection which o)..................................................................................................................

There are 3 acidic oxides which are of concernenvironmentally:

�• Carbon dioxide is a natural part of the p)...........................-................................... cycle in nature. Human activities haveincreased the CO2 levels mainly from q)........................................................................ The major problem arising is�“r)................................. warming�” due to the�“s)..............................................................�”

�• Sulfur dioxide occurs naturally in t)............................eruptions and u)......................................... Human activitieswhich release SO2 include burning of v).................................and the w)......................................... of some metal ores. Themain environmental problem is �“x)..................... ..................�”which can cause serious ecological damage to lakes andforests.

�• Nitrogen dioxide and nitric oxide (collectively known asy)......................... gases) are produced naturally in very smallamounts by z)............................................... Human activitieswhich produce them are high temperature combustions inaa).............................................. and ............................................Environmentally, NOx gases are the main cause of toxic�“ab)...............................�” in large cities, but also cancontribute to Acid ac).......................


1. Simple Acid-Alkali Reactionsa) Name the salt formed in a reaction between:

i) hydrochloric acid & calcium hydroxideii) sulfuric acid & magnesium hydroxide

iii) nitric acid & barium hydroxide

b) Write a balanced symbol equation for the reaction of:i) hydrochloric acid and lithium hydroxide

ii) sulfuric acid and sodium hydroxideiii) nitric acid and magnesium hydroxide

2. Reactions of Basic OxidesWrite a balanced symbol equation for the reaction of:a) sulfuric acid & iron(II) oxideb) hydrochloric acid & magnesium oxidec) nitric acid & copper(II) oxide

3. Reactions of Acidic Oxidesa) carbon dioxide reacts with calcium hydroxide to formwater and calcium carbonate. (This is the �“limewater�”reaction) Write a balanced equation for the reaction.

b) P2O5 is an acidic oxide. It reacts with water to formphosphoric acid, H3PO4. Write the balanced equation.

c) Sulfur trioxide reacts with water to form a strong acid.Write a balanced equation, and name the acid.

4. Le Chatelier�’s Principlea) NO2(g) reacts with itself:

The reaction to the right is exothermic, so heat can beconsidered as a �“product�”.

Note that as the reaction proceeds to the right, 2 moles ofgas form 1 mole of gas, so in a fixed volume container thepressure would drop as the reaction proceeds to the right.

Imagine a sealed container in which a mixture of thesegases has reached equilibrium.

State which way the equilibrium would shift if each of thefollowing disturbances were made to the mixture. Explaineach answer.

i) Increase in temperature.ii) Compress the mixture, therby increasing pressure.

iii) Injecting extra N2O4, without changing pressure.iv) Decrease the temperature.v) Spray in a little water. (NO2 dissolves, N2O4 does not.)

vi) Decreasing the total gas pressure.

continued...

10





WHEN COMPLETED, WORKSHEETSBECOME SECTION SUMMARIES

2NO2(g) N2O4(g)

2NO2(g) N2O4(g) + heat

4. Le Chatelier�’s Principle (continued)b) If hydrogen iodide (covalent molecular) is dissolved inwater, some of the molecules ionize, and an equilibrium isreached:

What is the effect on this equilibrium of: (Explain each)i) adding NaI(aq) solution, which increases theconcentration of iodide ions.ii) Adding NaOH, which reacts with H+ ions, and reducestheir concentration.iii) Dissolving extra HI in the solution.

iv) It is found that raising the temperature of anequilibrium mixture has the effect of increasing theconcentration of ions. Deduce whether the reaction aswritten is exo- or endothermic.

c) Ammonia is manufactured from its elements by thereaction

i) To maximize the yield of ammonia, the reaction iscarried out under very high pressure. Explain how thishelps.ii) The temperature of the reaction is kept fairly high tospeed up the rate of the reaction. What effect does highertemperature have on the equilibrium?iii) During the reaction, ammonia is constantly removedfrom the reaction vessel, and more reactant gasesconstantly pumped in. Explain the effect this has.

5. Molar Gas Volumesa) What is the volume of:

i) 2.59 mol of O2 at SLC?ii) 0.0453 mol of H2 at STP?

iii) 120 mol of CO2 at STP?iv) 4.67 x 10-2 mol of N2 at SLC?

b) How many moles of gas is:i) 12.4L of He at SLC?

ii) 250mL (=0.250L) of O2 at STP?iii) 10,000L of N2 at SLC?iv) 1.00mL of Ar at STP?

6. Mass - Volume of Gasesa) What is the mass of:

i) 5.00L of CO2 at SLC?ii) 5.00L of H2 at SLC?

iii) 100L of Ne at STP?iv) 25.0mL of O2 at STP?

b) What is the volume (at SLC) of:i) 100g of CO2 ?

ii) 100g of He ?iii) 1.50g of N2 ?iv) 1.00kg of Ar ?

7. Problems Involving Reactionsa) Carbon dioxide gas reacts with aqueous calciumhydroxide (limewater) to form water and insoluble calciumcarbonate.

i) Write a balanced equation for the reaction.ii) What mass of calcium carbonate would be formed by

the reaction, if 1.00L of CO2, measured at SLC?iii) What volume of CO2 (at STP) must be absorbed bylimewater in order to precipitate 1.75g of CaCO3 ?

b) In the smelting of zinc, the crushed, concentrated ore iszinc sulfide. This is roasted in a blast of air, forming zincoxide and sulfur dioxide gas.

i) Write a balanced equation for the reaction.ii) Calculate the volume of pure oxygen gas (at SLC)

required for complete reaction with 1.00 Tonne of zincsulfide.iii) Assuming the air used is 21% oxygen, what volume ofair must be supplied to the smelter for each tonne of ore?iv) From your answer to (ii), use Avogadro�’s Hypothesis tofind the volume of SO2 (at SLC) released from one tonneof ZnS.

11





HI(aq) H+(aq) + I-

(aq)

NH3(g) + heatN2(g) + 3H2(g)

Volume of Gases under Different Conditions

You already know that the conditions known as �“SLC�”(Standard Laboratory Conditions) means

�• 25oCand �• 1 atmosphere of pressure ( = 100 kPa)

At SLC, 1 mole of any gas occupies 24.8 Litres

What about at other temperatures and pressures?

Another set of standard conditions commonly used inChemistry is known as �“Standard Temperature &Pressure�” (STP):

�• 0oCand �• 1 atmosphere of pressure

At STP, 1 mole of any gas occupies 22.7 Litres

Notes:1. You do NOT need to remember these values. They arealways given in the Data Sheet in tests/exams.2. There is a formula for calculating the volume at anytemperature and pressure, but its use is not required inthis course.

FOR MAXIMUM MARKS SHOWFORMULAS & WORKING,

APPROPRIATE PRECISION & UNITS IN ALL CHEMICAL PROBLEMS

Acids as Proton DonorsSo far, you have learnt a number of things about acids, butnot the most basic thing... the chemical definition of justwhat an acid is!

What defines all acids is their ability to donate a hydrogenion (H+) to another species. Since a hydrogen ion is reallyjust a �“naked�” proton from the nucleus of the hydrogenatom, the formal definition of an acid is:

Acids in Aqeous SolutionWhen acids dissolve in water, we often just imagine anionization such as this example:

Hydrochloric acid: HCl(g) H+(aq) + Cl-(aq)

However, this is NOT the full story!

In a previous topic (Preliminary topic �“Water�”) you learntabout how ions are �“hydrated�” by polar water moleculeswhen in solution. The dissolving of HCl in water could berepresented as:

Although there may be many water molecules clinging to aH+ ion, for simplicity we imagine there is just one, and itforms a special ion, called the �“hydronium ion�”, H3O

+.

HCl(g) + H2O(l) H3O+

(aq) + Cl-(aq)

HCl is an acid because it has �“donated�” a proton to a watermolecule. The hydronium ion in the solution is an acidbecause it can, in turn, donate a proton to other species.

Formation of Hydronium Ions in WaterMore examples:

Nitric acid:HNO3(l)+ H2O(l) H3O

+(aq) + NO3

-(aq)

Sulfuric acid:H2SO4(l)+ 2H2O(l) 2H3O

+(aq) + SO4

2-(aq)

Phosphoric acid:H3PO4(l)+ 3H2O(l) 3H3O

+(aq) + PO4

3-(aq)

The �“Organic�” (Carbon-Based) AcidsMany naturally-occurring, biological molecules are acidstoo. These are carbon-compounds, made by living thingsand most contain a special chemical group you shouldbecome familiar with: the �“-COOH�” group.

Perhaps the most important is ethanoic acid, CH3COOH:

When this molecule is dissolved in water the O-H bondionizes, and donates a proton to a water molecule:

CH3COOH(aq)+ H2O(l) H3O+

(aq)+ CH3COO-(aq)

ethanoic + water hydronium + ethanoateacid ion ion

The syllabus requires you to know about ethanoic acid, and�“citric acid�” (next page)

12


3. ACIDS & THE pH SCALE




An acid is a chemical species which DONATES PROTONS

""##""++

Cl- ++

""++

""++""##""##

""##

""##

""++

""++

HCl((gg))

molecule

Separate, hydrated ionsCl-((aaqq)) and H++

((aaqq))

In water solution, all acids produce hydronium ions.

Some useful words:MONOPROTIC = donating 1 proton (e.g. HCl)DIPROTIC = donating 2 protons (e.g. H2SO4)

TRIPROTIC = donating 3 protons (e.g. H3PO4)

Try the WORKSHEET at the end of section

CC

OO-HH

OOThis group isusuallyabbreviated as�“-COOH�” andis always acidic.

This bond is polar, and ionizes in water

These C-HH bonds are non-ppolar, and never ionize.



13

Citric Acidis familiar as the acid in citrus fruits... oranges & lemons.

Its correct systematic name is2-hydroxypropane-1,2,3-tricarboxylic acid

Molecular formula is C6H8O7, but the structural formula ismore meaningful:

Remember that onlythe -COOH groups

are acidic, and you will see that

citric acid is triprotic

Strong & Weak AcidsIn everyday usage, a �“strong�” solution might mean thesame as �“concentrated�” (i.e. having a lot of solute) and�“weak�” can mean the same as �“dilute�”.

FROM HERE ON, YOU MUST NOT USE THESETERMS THAT WAY. �“Strong�” & �“weak�” have particularmeanings with regard to acids.

For example, HCl is a STRONG acid; when added to waterthe reaction...

HCl(g) + H2O(l) H3O+

(aq) + Cl-(aq)

...goes fully to completion.100% of the HCl molecules are ionized.If the concentration of HCl was (say) 1 molL-1, then theconcentration of hydronium ions will also be 1 molL-1.

In contrast, ethanoic acid is a WEAK acid; when added towater, the reaction...


(aq)+ CH3COO-(aq)

...reaches an equilibrium with only about 1% of themolecules ionized. (i.e. the equilibrium lies well to the left,favouring the reactant molecules.) If the concentration ofthe solution was 1 molL-1, then the hydronium ionconcentration would be only about 0.01 molL-1.

It is quite possible to have a �“concentrated solution of aweak acid�” or a �“dilute solution of a strong acid�”... just becareful with these precise meanings!

The pH Scalehas already been introduced, and you should be familiarwith it in a descriptive way.

Where do these numbers come from?



C

C

C

H

H

H

H

COOH

COOH

COOH

H-OO

STRONG Acid = Total Ionization in Solution

WEAK Acid = Partial Ionization in Solution

ACID NEUTRAL BASE

1 3 5 7 9 11 13pH

pH = -log10[H3O+]

pH is the negative logarithm of the molar concentration of

hydronium ionsNotes:1. You must know that [square brackets] around anychemical species means �“molar concentration of �”.

2. �“log10�” means the �“logarithm to base 10�”. This is amathematical function, best understood by example:If 100 = 102, then log10(100) = 2

1,000 = 103, then log10(1,000) = 3500 = 102.699, then log10(500) = 2.699

pH values are therefore, powers of 10.

3. Different calculators may handle log functionsdifferently. You must find out or figure out how to dolog functions on your calculator.

Example Calculations: pH and [H3O+]

1. If the concentration of hydronium ions is [H3O

+] = 0.500molL-1, what is the pH?Solution: pH = -log10[H3O

+]

= -log10[0.500]= -(-0.301)

!! pH = 0.301

2. If the concentration of hydronium ions is [H3O

+] = 0.00252molL-1, what is the pH?Solution: pH = -log10[H3O

+]= -log10[0.00252]= -(-0.2.60)

!! pH = 2.60

3. If the pH = 3.75, what is [H3O+]?

Solution: pH = -log10[H3O+]

so [H3O+] = Inverse(log(-3.75))

!! [H3O+] = 0.000178 molL-1.




14

More about pHSince the numbers on the pH scale are powers of 10, itfollows that if an acid solution has a pH one unit lowerthan another, it is actually 10 times more acidic. Two unitson the pH scale represents 100 times (102) difference in[H3O

+].

Example:Acid �“P�”: [H3O

+] = 10-5molL-1. pH = 5

Acid �“Q�”: [H3O+] = 10-3molL-1. pH = 3

Acid �“R�”: [H3O+] = 10-2molL-1. pH = 2

In the examples above, acid �“Q�” has a [H3O+] value 100

times larger than �“P�”, and its pH is 2 units lower.

Acid �“R�” has a [H3O+] value 10 times higher than �“Q�”, and

its pH is 1 unit lower.

The ph scale is said to be �“logarithmic�”, because thevalues are logarithms... powers of 10.

pH of Strong & Weak AcidsRemember the special meanings of �“strong�” and �“weak�” with regard to acids.

You may have used a pH meter to measure the pH of several different acids ofexactly the same concentration.



Measuring pHYou may have carried out some simple laboratory

experiments, using a pHmeter (or data-loggerprobe) to measure thepH of various solutions.

You may have recordedthe pH value, then usedyour knowledge of thepH scale to decide ifeach solution tested was

Acidic (pH < 7)

or

Neutral (pH = 7)

or

Basic (pH > 7)

1.00 molL-1 Hydrochloric Acid

When HCl(g) dissolves in water:

HCl(g)+ H2O(l) H3O+

(aq)+ Cl-(aq)

Since HCl is a STRONG ACID, theionization is 100%.

If [HCl] = 1.00 molL-1,then [H3O

+] = 1.00 molL-1

and pH = -log10[H3O+]

= -log10[1.00]!! pH = 0

Sure enough, by experiment you willhave found that a 1 molL-1 HClsolution has a pH = 0.

1.00 molL-1 Sulfuric Acid

When H2SO4(l) dissolves in water:

H2SO4(l)+ 2H2O(l) 2H3O+

(aq)+ SO42-

(aq)

Since H2SO4 is a STRONG ACID, theionization is 100%. It is DIPROTIC, soeach mole of H2SO4 produces 2 moles ofH3O

+.

If [H2SO4] = 1.00 molL-1,then [H3O

+] = 2.00 molL-1

and pH = -log10[H3O+]

= -log10[2.00]= -(0.301)

!! pH = -0.30

Sure enough, by experiment you willhave found that a 1 molL-1 H2SO4solution has a pH = -0.30. (below zero)

1.00 molL-1 Ethanoic Acid

When CH3COOH(g) dissolves in water:

CH3COOH(aq)+H2O(l) H3O+

(aq)+ CH3COO-(aq)

Since this is a WEAK ACID, the ionizationis incomplete.

If [CH3COOH] = 1.00 molL-1,then [H3O

+] << 1.00 molL-1

and the pH can be expected to be wellabove the values found for HCl or H2SO4.

Sure enough, by experiment you willhave found that a 1 molL-1 CH3COOHsolution has a pH $$ 2.




15

The Self-Ionization of WaterWater itself can be considered as an extremely weak acid,since in pure water a few water molecules ionize and donatea proton to another molecule.

In pure water, at 25oC, it turns out that the concentrationof hydronium ions, [H3O

+] = 1.00 x 10-7 molL-1.

So the pH of pure water:pH = -log10[H3O

+]= -log10[1.00 x 10-7]= -(-7.00)

!! pH = 7.00

�• That�’s why pH = 7 is neutral on the pH scale!

�• In an acid solution, [H3O+] > 1.00 x 10-7 molL-1,

so pH < 7

�• In a basic solution, [H3O+] < 1.00 x 10-7 molL-1,

so pH > 7

""+""-""+""-

Proton transfer

++-

2 WaterMolecules

HydroxideIon

HydroniumIon

2 H2O(l) OH-(aq) + H3O

+(aq)

Acids as Food AdditivesAcids are common food additives. If you read the�“Ingredients�” listing on many processed foods, from tomatosauce to ice-cream topping to dried fruits or tinned sausages,you may find �“food acid (260)�” (or something similar).

�“Food acid (260)�” is in fact ethanoic acid, the main chemicalin vinegar, and a common ingredient in many recipes.

The reasons that acids are added to many foods are:

�• Preservatives. Adding acids to a food lowers the pH and makes it more difficult for some bacteria or fungi to grow within the food. This is the main reason why somefoods, such as tomato sauce, do not go bad even when notrefrigerated.Acids most commonly used are SO2 & ethanoic.

�• Flavour. Many foods, such as jams, sauces and fruit-flavoured drinks will taste better if they have that �“sharp�”or sour effect that acids give.Acids commonly used are citric and ethanoic acids.

�• Nutrition. Some foods have vitamins added to increasetheir nutritional value.Most commonly used is ascorbic acid (vitamin C).

Some Naturally Occurring Acids & BasesAcids

Ever been bitten by a bull-antor stung by a bee?

You were injected with�“formic acid�” (methanoicacid), HCOOH.

Ethanoic acid is the main chemical invinegar (literally �“wine-sour�”) known andused for thousands of years.SUGAR (grapes)

fermentationETHANOL (wine)

oxidationETHANOIC ACID (vinegar)

Bases

Lime is a major ingredient of cement, and has beenknown and used for thousands of years.Chemically it is calcium oxide, CaO.

It is prepared by roasting limestone (or even oystershells) which results in the decomposition

CaCO3 CO2 + CaO

�“Lye�” is the name given to the alkali liquid obtainedby soaking wood ashes in water. The basicingredient is a mixture of hydroxide compounds,mainly potasium and sodium hydroxides.

Lye has been used for centuries in soap-making.

Photo by Norbert Machmek

Worksheet 3Part A Fill in the Blank Spaces

Acids can be defined as species whicha)............................... .......................................... Inwater solution an acid molecule donates a proton(which is a b).................................. ion) to a watermolecule, forming a c)........................................ ion.Acids which donate one proton are calledd)................................................; those that donate 2protons are e)......................................, and thosewhich donate 3 are called f).......................................

Two important �“organic�”, naturally occurringacids are�• ethanoic (formula g).....................................), and �• citric acid; systematic chemical name ish).......................... ........................................................

With acids, �“strong�” & �“weak�” do NOT refer tothe concentration. A strong acid is one whichi).............................................................................,and weak refers to an acid whichj)............................................................ .......................

Acidity is measured by the pH scale. The pHvalue is calculated from the equation pH =k)...................................... pH values arel)............................. of 10, so each 1 unit of pHactually means a change in acidity of m)...............times.

Acids are commonly added to foods, in order to:

�• help n)....................................... the food bymaking it more difficult foro).................................... and ............................... togrow in the food and cause it to spoil. The acidsmost commonly used this way arep)........................... and .............................................

�• q)...................................... the food, by giving itthe r).............................. taste that all acids have.Most commonly used are s)...................................and ................................................. acids.

�• Improve the t)......................................... value. Forexample, u)................................... acid (vitamin C)is added to some foods.

A naturally occurring acid is v)..............................acid in an ant bite. A natural base is lime, whichis chemically w).....................................................


1. Acid Ionizations with WaterWrite a balanced equation for the reaction ofeach acid with water. (Assume completeionization in each case, even though some may beweak acids)

Monoprotic acidsa) HCl b) HBr (hydrobromic)c) HCOOH (methanoic) d) HCN (hydrocyanic)

Diprotic acidse) H2SO4 f) H2CO3 (carbonic)

Triprotic acidsg) H3PO4 (phosphoric) h) C6H8O7 (citric)

2. Calculating pH from [H3O+]

Each of the following values is the hydroniumion concentration of a solution, in molL-1.For each, calculate the pH, and state if thesolution is acidic or basic.

a) 0.0675 b) 0.000840c) 2.50 d) 2.50x10-12

e) 3.5 x 10-3 f) 1.50 x 10-8

3. Calculating pH from Acid ConcentrationCalculate the pH of each solution, taking intoaccount that some acids may be diprotic ortriprotic. (assume total ioniz.)

a) 0.250 molL-1 HClb) 0.0750 molL-1 H2SO4.c) 7.50x10-4 molL-1 H2SO4.d) 4.50x10-3 molL-1 H3PO4.e) 6.00 molL-1 H2SO4.

4. Calculating [H3O+] from pH.

Find the [H3O+] in each solution.

a) pH = 5.30 b) pH = 11.5b) pH = -0.40 c) pH = 4.85d) pH = 8.50 e) pH = 0

16







Early Ideas About Acids & BasesNaturally occurring acids, like ethanoic acid in vinegar,have been known for thousands of years, and described bytheir simple, observable properties such as their sour taste.

Chemistry became a modern Science just over 200 yearsago, and one of the earliest scientific theories about acidswas made by

Antoine Lavoisier (French, 1780�’s)Lavoisier used simple plant-extract indicators (e.g. litmus)to identify acids and bases. He found by experiment that allthe non-metal oxide compounds he tested produced acidsolutions. He concluded that acids must contain oxygen.

Humphry Davy (English, 1820�’s)Forty years later and armed with more chemical knowledge,Davy realized that Lavoisier was wrong about acids. Whilesome acids do NOT contain oxygen (e.g. HCl), he foundthat all known acids contain hydrogen.

Davy also discovered that metal oxides are basic anddescribed the patterns in the way acids react with metals toform hydrogen gas.

These early ideas were all empirical descriptions ofproperties; they described the properties of acidsdiscovered by experiment, but did not include a generaltheory to explain or predict chemical behaviour.

That had to wait for Svante Arrhenius (Swedish, 1880�’s)The Arrhenius Theory of acids-base behaviour is stillgenerally used in years 7-10 Science.

The Arrhenius Theory was very successful in accounting forsimple acid-base behaviour in water solution. It explained theneutralization reaction, and could explain strong and weakacids as being due to complete or partial ionization.

However, the Arrhenius Theory had some deficiences:

�• It could not account for acid-base behaviour that was notin water solution.

�• It could not explain why many ions (which did NOT contain any hydrogen, hydroxide or oxide) showed acid or base behaviour in water solution.

You will soon learn that the main reason for all thesedeficiencies was that Arrhenius failed to consider the roleof the solvent itself, and in water solutions this is critical.

The Bronsted-Lowry Theorywas developed independently by Johannes Bronsted(Danish) and Thomas Lowry (English) in 1923. The basicconcepts of this theory have already been used earlier inthis topic:

Central to the Bronsted-Lowry (B-L) Theory is the conceptof �“conjugate�” species. For example, when Ethanoic aciddissolves in water:


(aq)+ CH3COO-(aq)

ethanoic + water hydronium + ethanoateacid ion ion

Acid Base Conjugate ConjugateAcid Base

The CH3COO- ion is the conjugate base of Ethanoic acid,because if this reaction was to run in reverse, the ion couldact as a base and accept a proton to form ethanoic acid again.

The water molecule is acting as a base, because it accepts aproton to become the H3O

+ ion. The H3O+ ion is the

conjugate acid because, if the equilibrium shifts left, it candonate a proton and become a water molecule again.

Another Example: Dissolving of HCl in water:

In this case, the reaction is very unlikely to ever run inreverse, but Cl- is still considered the conjugate base.

As before, water has acted as a base, and its conjugate acidis the hydronium ion.

17


4. ACID-BASE THEORY & TITRATION




Acids produce H+ ions in water solution.Bases produce OH- ions (or oxide ions).

All Acid-Base Reactions involve transfer of Protons

An Acid is a Proton DonorA Base is a Proton Acceptor

Cl-

HCl((gg))

moleculeH22O((ll))

moleculeCl-((aaqq))

ionH33O++

((aaqq))

ion++

++ ++

++

Acid Base Conjugate Conjugate Base Acid





18

Water Can be an Acid, TooIn both the previous examples, the water molecule acted asa base by accepting a proton to form a H3O

+ ion. Watercan also act as an acid and donate a proton.

For example, when ammonia, NH3 dissolves:

In this reaction the water molecule acts as an acid bydonating a proton. Its conjugate base is the hydroxide ion,OH-.

This equation explains why ammonia is a base, and why asolution of ammonia (NH3(aq)) can be considered as asolution of ammonium hydroxide, NH4OH(aq).

Acidic and Basic SaltsOne of the weaknesses of the Arrhenius Theory was thatit could not explain the results of simple experiments youmay have done:

Practical Work: Testing the pH of Various Salt Solutions

NH33((gg)) H22O((ll)) NH44++

((aaqq))

ionOH-

((aaqq))

ion

++

++

++

Base Acid Conjugate Conjugate Acid Base

Chemical Species (like water) which canboth donate and accept protonsare called �“AMPHIPROTIC�”

You will meet more Amphiprotic species soon

You may have done simpleexperiments using a pH meter, orUniversal Indicator, to test the pH ofsolutions of various ionic �“salts�”,none of which seem to be obviouslyacids or bases, from their formulas.

You may have tested many, but theyprobably included...

sodium chloride, NaCl

sodium carbonate, Na2CO3

sodium hydrogen sulfate, NaHSO4

...all at the same concentration, inpure water.

Typical Results

Salt pH Conclusion

NaCl 7 Neutral

Na2CO3 11 Basic

NaHSO4 2 Acidic

The reasons WHY some salts showacid-base behaviour is explainedabove right.

Note: Due to CO2 from the airdissolving in the solution, the pH ofthe NaCl solution may be slightlyacidic.

Acidic SaltsAs some ions dissolve in water they react as acids, forcingthe water molecule to be a base. An example is the�“hydrogen sulfate ion�”, HSO4

-.

hydrogen + water sulfate + hydroniumsulfate ion ion

HSO4-(aq) + H2O(l) SO4

2-(aq) + H3O

+(aq)

If the salt used was sodium hydrogen sulfate there wouldalso be some sodium ions in the solution. In terms of acid-base behaviour, they would be merely spectators.

Basic Saltsdissolve in water and react as bases, forcing the watermolecule to be the acid. An example is the carbonate ion:

carbonate + water hydrogen + hydroxideion carbonate ion ion

CO32-

(aq) + H2O(l) HCO3-(aq) + OH-

(aq)

Once again, if sodium carbonate was used, there would besodium ion �“spectators�” in the solution as well.

A few worth knowing...Acidic Salts Neutral Salts Basic SaltsNaHSO4 NaCl Na2CO3NH4Cl KNO3 CH3COONaNH4NO3 Na2SO4 (sodium ethanoate)

(There is a pattern to help you remember these... later)




19

NeutralizationNow that you know about B-L Theory, we can go back tolook again at the simple acid-alkali neutralization.

Examplehydrochloric + sodium water + sodium acid hydroxide chloride

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

However, now we know that what makes HCl an acid isreally the formation of H3O

+ ions in water, and the Cl-

ions are spectators, as are the Na+ ions from the NaOH.

Leaving out the spectators, the net ionic equation is...

hydronium + hydroxide waterion ion

H3O+

(aq) + OH-(aq) 2H2O(l)

... and this is the net reaction for ALL the simple acid-alkalireactions, regardless of which acid or which alkali are used.

Heat of NeutralizationIf you carried out the reactionabove in a calorimeter, you willquickly find that the temperaturerises... the reaction is exothermic.

If measured and calculated, it is found that the value for%H is the same regardless of which acid and alkali is used.

H3O+

(aq) + OH-(aq) 2H2O(l) %%H= -56kJmol-1

Salts from Different Acid-Base Neutralizations

Although we might view all neutralizations as essentially thesame reaction and ignore the �“spectator ions�” which formthe �“salt�”, this prevents us noticing a useful pattern whichwas mentioned on the previous page.

The nature of the �“salt�” formed depends on whether theacid and base were �“strong�” or �“weak�”, as follows:

Acid-Base Reactions without H3O+, OH- or H2O

Although most of the simple examples of neutralizationinvolve the reaction at left, and usually we study reactionstaking place in water, be aware that this is not always thecase. For example, if you add together the dry gases HCl(g)and NH3(g) they will react

hydrogen + ammonia ammoniumchloride chlorideHCl(g) + NH3(g) NH4Cl(s)

Although there is no water present, and no H3O+ or OH-

ions are formed at any time, this is clearly an acid-basereaction, according to the B-L definition.



ACID + ALKALI WATER + A �“SALT�”

The reaction involves a proton transfer

Ener

gy C

onte

nt

%%Hnegative

Eaa

StrongAcid

WeakAcid

StrongBase

WeakBase

formsNEUTRAL salts

formsNEUTRAL salts

forms ACIDIC salts

forms BASIC salts

e.g. HCl,HNO3,H2SO4

e.g. KCl, NaNO3K2SO4

e.g. NH4Cl, NH4NO3(NH4)2SO4

e.g. CH3COONa,Na2CO3

e.g. CH3COONH4,(NH4)2CO3

e.g.CH3COOHH2CO3

e.g. KOH,NaOH

e.g. NH3

Why is Ammonia (NH3) a �“Weak Base�”?

Simple!Because when it reacts with water, the reactiondoes NOT ionize fully, but reaches an equilibriumwith a significant concentration of un-ionizedmolecules present.

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

Proton transfer

The information in the Right-Hand columnis not specified by the Syllabus.

It is presented here in the interests of better understanding.




20

The purpose of titration is Chemical Analysis. It allowsthe concentration of an �“unknown�” solution to bedetermined by calculation, after measuring the volume ofa known-concentration �“standard solution�” which reactswith a sample to reach the �“equivalence-point�” (end-point). This is the point where the reactants have beenconsumed in exactly the molar ratio specified by thebalanced equation.

Titration Technique

�• you must have a �“standard solution�” (see next page)�• the burette should be rinsed with small quantities of thestandard solution, then filled with it. The level is thenadjusted to the zero mark, ensuring that there are no airbubbles (especially in the tap and delivery spout).�• an exact, measured volume of �“unknown�” solution isplaced in the reaction flask by pipette.�• 1-2 drops of suitable indicator are added to the flask.

(See �“Choosing the Indicator�”, next page)�• the reaction proceeds by careful addition, with mixing,of solution from the burette. Add drop-by-dropapproaching the end-point, until the indicator justchanges colour.�• record the volume of solution delivered by burette.�• repeat 3 times, and average the �“titres�” (= burettevolumes). It is common to discard any values which arenot in close agreement. The rest is calculation.

TitrationTitration has been one of the most important techniques in Chemical Analysis for over a century.

Its use has diminished as electronic �“probes�” have become more widespread,but it remains a �“must-know�” part of Chemistry.

Titration can be used for any reactionfor which you can determine the �“end-point�”. Acid-base neutralizations areideal because indicators (or pH probes)can identify the end-point.

Burette measures the volume of�“standard solution�” neededto reach the end-point.

Burette tapcontrols the addition of�“standard solution�” to the�“unknown solution�”

Conical flask holds a measured volume of�“unknown solution�”, plus indicator

Filter Paper under flask allowsbetter seeing ofindicator colour

Example Titration CalculationTitration of an �“unknown�” solution of KOH.�“Aliquot�” (volume of samples) = 25.00 mL (by pipette)

Standard solution; H2SO4 solution, C= 0.04252 molL-1.Indicator used: Bromothymol Blue (see next page)

�“Titres�” (volumes from burette) measured (in mL): 34.25, 33.90, 33.85, 33.95.The first titre is discarded (not in close agreement with theothers). Average titre = 33.90 mL

Titration Formula: Ca x Va = Cb x Vba b

Ca, Cb = concentrations (molL-1) of acid and base solutionsVa, Vb = volumes (mL) of each solution used.a, b = molar co-efficients (balancing numbers) from

balanced equation for reaction.

Step 1: Balanced equation (ALWAYS!!)

H2SO4 + 2KOH 2H2O + K2SO4

Step 2: Titration FormulaCa x Va = Cb x Vb

a bStep 3 : Re-arrange and substitute.We are trying to find the concentration of the �“unknown�”base, so make �“Cb�” the subject:

Cb = b x Ca x Vaa x Vb

= 2 x 0.04252 x 33.901 x 25.00

= 0.1153

!! Concentration of KOH solution = 0.1153 molL-1.



21

Preparing a Standard SolutionThe key to titration is to have available a suitable �“standardsolution�”. The technique for making a solution to an exactknown concentration was covered in the Preliminary topic�“Water�” and is revised below.

The problem is to get a substance which is of very highpurity and stability to use as a �“Primary Standard�” to makethe solution from.

The common acids like H2SO4, and bases like NaOH,cannot be obtained in the pure state due to the way theyrapidly absorb water and/or CO2 from the atmosphere.

Suitable primary standard substances are:

�• Base: anhydrous sodium carbonate, Na2CO3�• Acid: oxalic acid, COOHCOOH (diprotic)

In the example (previous page) the standard solution wasH2SO4. This would not have been a primary standardsolution.

It may have been �“standardized�” by titration against aprimary standard Na2CO3 solution, so that its exactconcentration is known. Then it can be used, in turn, as thetitration standard solution.

Selecting the IndicatorChoosing the best indicator to use is not just a matter ofwhich colour change you�’d prefer looking at.

Each indicator changes colour over a range of pH values,and the trick is to choose an indicator which will changecolour as close to the �“end-point�” as possible.

Colour in ChangesIndicator Acid Base at pH

Litmus red blue 6 - 8Bromothymol

blue yellow blue 6.2 - 7.6Methyl orange red yellow 3.1 - 4.4Phenolphthalein clear pink 8.3 - 10.0

But, isn�’t the end-point always at neutral, pH = 7 ??(I hear you ask)

No, its not! Remember (see p.19) that the �“salts�” formedby different combinations of strong or weak acids andbases may have acidic or basic properties. This means that,at the exact end point, the pH might not be neutrality.

Examples:



Dissolve Solute in a smallamount of (pure) water in a

clean beaker

Calculate the mass of pure,dry Primary Standard soluterequired for solution, and

weigh out accurately

Carefully transfer solutioninto a Volumetric Flask.Rinse beaker with smallamounts of water & add

washings to flask

Add water to flask to fill it tothe mark.

(Use a dropper to avoid over-shooting)

Insert stopper & mix well.

VVoolluummeettrriicc FFllaasskk

Strong Acid - Strong BaseIf titrating (say) H2SO4 against KOH,

you can expect the salt K2SO4 to be neutral.

Therefore, choose Bromothymol blue,which changes colour near pH = 7.

(Actually, it doesn�’t really matter, because the the lastdrop of chemical at the end point causes a huge pHchange.)See graph next page. Strong Acid - Weak Base

If titrating (say) H2SO4against NH3,

you can expect the salt (NH4)2SO4 to be acidic.

Therefore, choose Methyl orange,which changes colour around

pH = 3-4.Weak Acid - Strong BaseIf titrating (say) CH3COOH against KOH,you can expect the salt (potassiumethanoate,CH3COOK) to be basic.

Therefore, choose Phenolphthalein,which changes colour around pH = 8-10.

Weak Acid - Weak BaseTitrations with this combination are to be avoided,

because the end-point is not very �“sharp�”.



22

Measuring pH During TitrationIf a titration is carried out as shown below, with a pH meteror data-logger probe attached to a computer, the pHchanges occurring in the solution can be recorded andgraphed.



Magnetic Stirrer

Connection tocomputer or pH meter

pH probe electrode

Burette adds standardbase solution

Flask contains�“unknown�” acidsolution

pH

0

2

4

7

10

1

2

Volume of Base Added

Strong Acid &

Strong Base

TTiittrraattiioonneenndd-ppooiinntt

To start with,adding base haslittle effect on pH

Near the end-point,pH changes rapidlywith each drop ofbase added, so theend-point can bedetermined precisely.

pH

0

2

4

7

10

1

2


Strong Acid &

Weak Base


The end-point is atpH<7 because thesalt is acidic...mmuusstt cchhoooossee mmeetthhyylloorraannggee,, oorr ssiimmiillaarr..

pH

0

2

4

7

10

1

2


Weak Acid &

Strong Base


The end-point is at pH>7because the salt is basic...

mmuusstt cchhoooossee iinnddiiccaattoorr ssuucchhaass pphheennoollpphhtthhaalleeiinn

pH

2

4

7

1

0


Weak Acid &

Weak BaseThe end-point istheoretically at pH=7but is hard to find byindicator because eachdrop of base is causingless pH change

???

In every case, the end-point is the centreof the vertical (or near-vertical) part ofthe graph.

The graph at left shows why weak acid-weak base titrations are best avoided.

TThhee �“�“vveerrttiiccaall�”�” ppaarrtt ooff tthhiiss ggrraapphheexxtteennddss ssoo ffaarr eeaacchh ssiiddee ooffnneeuuttrraall,, tthhaatt tthhee cchhooiiccee ooffiinnddiiccaattoorr iiss rreeaallllyy nnoott ccrriittiiccaall......aallll iinnddiiccaattoorrss wwiillll cchhaannggee ccoolloouurroonn tthhee ssaammee ddrroopp ooff aaddddeeddbbaassee..




23

BuffersA �“Buffer�”, or �“buffered solution�”, is a solution which canabsorb significant amounts of acid or base with minimalchange in pH.

Biology students will be aware that in all living things it isvital that the conditions within the body/cells are kept veryconstant. Living things cannot function properly if their�“internal environment�” undergoes large changes in�• temperature,�• water content,�• salt concentration, (and many other chemicals) AND�• pH level.

The fluids inside all living things are buffered, so that pHremains remarkably constant, despite changes in theexternal environment, eating acidic food, breathing,excreting, etc, all of which could alter the body�’s pH.

How do Buffers Work?All buffers are solutions containing a weak acid and itsconjugate base.


(aq)+ CH3COO-(aq)

weak acid conjugate base

This equilibrium constitutes a buffer solution, so long asthere are significant amounts of the ethanoate ion(CH3COO-) present. (e.g. by adding CH3COONa)

If acid is added (i.e. concentration [H3O+] increases)

(pH lower) then (by Le Chatelier�’s Principle) the equilibrium shifts left, absorbing H3O

+

and lowering [H3O+] again.

If base is added (i.e. concentration [H3O+] decreases)

(pH higher) then (by Le Chatelier�’s Principle) the equilibrium shifts right, making H3O

+

and raising [H3O+] again.

This way, the pH remains quite constant, despite acid orbase being added.

A Natural Buffer SystemThe classic example of a natural buffer system is the wayour blood is maintained at a pH = 7.4, despite the factthat we keep exchanging CO2 (acidic), excreting wastes,absorbing foods, etc.

The main chemical buffer in our blood is a solutioncontaining both the bicarbonate ion (HCO3

- ) and itsconjugate base, the carbonate ion ( CO3

2- ).

HCO3- is Amphiprotic

The bicarbonate ion is amphiprotic, which adds a furtherdimension to its buffering ability.

If the environment is acidic, HCO3- acts as a base:

H3O+

(aq) + HCO3-(aq) H2CO3(aq)+ H2O(aq)

If the environment is basic, HCO3- acts as an acid:

OH-(aq) + HCO3

-(aq) CO3

2-(aq)+ H2O(aq)

So, the total buffering system can be summarized as:

This system is highly effective at maintaing the pH of ourblood. Be aware that it is not the only buffer operating.

proton transfer

proton transfer

H2CO3 HCO3- CO3

2-

+ H++ + H++

- H++ - H++

Using Neutralization on Chemical SpillsWhether its a small vinegar spill at home, or a sulfuric acidtanker leaking after a road accident, a knowledge ofneutralization can help damage control and safety.

So, if someone gets splashed with a strong base, neutralizeit by throwing a bucket of acid over them, right?

Wrong! Unless you can guarantee to apply the exact molarquantity for titration-like precision, then using an acid on analkali spill (or vice-versa) can do more harm than good, andwould be extremely dangerous as a first-aid method.

The best thing to do is use an Amphiprotic substancesuch as bicarbonate ion (e.g. sodium bicarbonate,NaHCO3).

As shown by the reactions above, it can react and neutralizeeither an acid or a base, and once the neutralization isachieved, it stops working and poses no further threat in itsown right.

It does its job, and then stops automatically... perfect!

As you�’ll learn below,Amphiprotic substances

are useful in emergencies, too.

Phot

o by

�“con

nman

21�”

Worksheet 4Part A Fill in the blank spaces

Lavoisier concluded that acids must containa).......................... Later, Davy described acids asall containing b)............................., and being ableto react with c).............................................

The first attempt at a complete acid-base theorywas made by d)........................................ in 1885.According to this theory, all acids producee)................................................ in solution. Baseswere defined as compounds containingf)................................. or ................................ ions.This theory was successful at explaining someacid-base behaviour, but failed to take intoaccount the important role ofg)..................................................

The h)...........................-............................. (B-L)Theory is much more useful for explaining things.It defines an acid as ai)............................................................, and a base asa j)........................................................, and all acid-base reactions involve the k)....................................of one or more l)......................................................

When an acid loses a proton it forms the�“m).................................... base�”, and when a basegains a proton it forms then)...................................................

Water can acts as either o)....................... or..........................., depending on the chemicalenvironment. The word for this is�“p)....................................................�”

Many salts dissolve in water to form acidic orbasic solutions, because of their interaction withq)......................... An example of an acidic salt isr)............................................. An example of abasic salt is s)............................................. Anexample of a neutral salt ist).............................................The �“rules�” governing this are:�• Neutral salts form from the neutralization ofu)....................... acids by v)............................. bases.�• Acidic salts form from the neutralization ofw)......................... acids by x).......................... bases.�• Basic salts form from the neutralization ofy).......................... acids by z).......................... bases.

All neutralization reactions (in aqueous solution)involve the reaction of aa).............................. and......................... ions to formab).................................. The reaction is alwaysac)..............-thermic.

ad).................................... is a technique used forchemical analysis. A measured ae)........................of �“unknown�” solution is reacted with a�“af).................................... solution�” until the�“ag)......................................... point�” is reached.The main piece of equipment involved is aah).............................. It is important to choose theai)................................... which will changeaj)............................ at the pH of the end-point,depending on the nature of the salt formed.

If the ak).............. is measured during titration andgraphed, the graphs generally show an �“S�”shaped curve. The end-point is located in themiddle of the al)................................ part of thecurve.

Buffers are solutions containing an equilibriummixture of a am)................................... and its............................................. If either acid or base isadded, the equilibrium an)................. (according toao)...................................... Principle) so that thechange in ap).................. is minimized. Buffers areimportant in all aq)............................. things, byhelping to maintain ar)...............................................A specific example is as)..........................................,where the pH is maintained by a mixture ofat)........................... and .................................... ions.


1. Reactions of Acids & Bases with WaterFor each substance below, write an equationdescribing its monoprotic reaction with water. (Inbrackets is a description of how the substancebehaves in each case)

For each equation, identify the conjugateacid/base of the named species.

a) methanoic acid, HCOOH (acid)b) ammonia, NH3 (base)c) hydrogen phosphate ion, HPO4

2-, (acid)d) hydrogen phosphate ion, HPO4

2-, (base)e) sulfide ion, S2- (base)f) cyanide ion, CN- (base)g) hydrogen sulfide, H2S (acid)h) nitrite ion, NO2

- (base)i) ammonium ion, NH4

+ (acid)j) hydrogen sulfite ion, HSO3

- (acid)(continued...)

24






Amphiprotic

Worksheet 4 Part B (continued)

2. Amphiprotic SubstancesEach of the substances below is amphiprotic.For each, write TWO equations to show itsreaction

i) with H3O+

ii) with OH-

a) dihydrogen phosphate ion, H2PO4-

b) hydrogen carbonate ion, HCO3-

c) hydrogen sulfide ion HS-

3. Acidic and Basic SaltsEach of the following salts dissolves in water, andreacts (as shown in brackets) to form an acidic orbasic solution.Write a net ionic equation (leave out spectators)to show the reaction with water.

a) potassium ethanoate, CH3COOK (basic)

b) ammonium nitrate, NH4NO3 (acidic)

c) sodium nitrite, NaNO2 (basic)

d) potassium hydrogen oxalate, KHC2O4 (acidic)

e) lithium cyanide, LiCN (basic)

4. Titrationsa) 25.00mL of an �“unknown�” NaOH solutionwas titrated against standardized HCl,concentration = 0.09255 molL-1. The titrationwas carried out 4 times, with the end-point titresbeing 22.50mL, 22.45mL, 23.10mL and 22.50mL.

Average these results appropriately, then write anequation, and calculate the concentration of theNaOH solution.

b) Using a 0.05025molL-1 standardized solutionof NH4OH (ammonia solution), 25.00mLsamples of an unknown H2SO4 solution weretitrated. The average titre was 28.32mL

i) Write a balanced equation for theneutralization.

ii) Find the concentration of the acid.

iii) For this titration there were 3 indicatorsavailable:

Indicator pH of Colour ChangeJ $ 8.7K $ 6.8L $ 4.2

Which indicator is most appropriate for thistitration? Explain your answer.

c) A student wishes to prepare 500mL of a0.02500molL-1 solution of oxalic acid(COOHCOOH), from the solid chemical.

i) What are the characteristics that qualify thischemical as a �“primary standard�”?

ii) What mass needs to be accurately weighedout?

iii) Summarize the main steps involved inpreparing the solution.

iv) The solution, when prepared, was used todetermine the concentration of a KOH solution.25.00mL aliquots of KOH required an averagetitre of 31.45mL. Find the concentration of theKOH. (care: oxalic acid is diprotic)

d) To find the concentration of an ammoniasolution (NH4OH), a student has 2 choices ofstandardized solutions she could use:

�• 0.7438 molL-1 HNO3 solutionor �• 0.8863 molL-1 CH3COOH solution

i) Which solution should she use in the titration?Explain.

ii) Using 10.00mL samples of the ammoniaunknown, and titrating with the appropriate acid,the average titre was found to be 12.76mL.Calculate the concentration of the ammoniasolution.

25







Alkanols and Alkanoic AcidsYou were introduced to the alkanols (alcohols) in theprevious topic, and to some of the alkanoic acids earlier inthis topic.

Alkanols Alkanoic Acids

Methanol Methanoic acidCH3OH HCOOH

Ethanol Ethanoic acidCH3CH2OH CH3COOH

Propanol Propanoic acidCH3CH2CH2OH CH3CH2COOH

Butanol Butanoic acidCH3(CH2)2CH2OH CH3(CH2)2COOH

Pentanol Pentanoic acidCH3(CH2)3CH2OH CH3(CH2)3COOH

...and so on.

The syllabus requires you to know these as far as 8-carbon molecules; octanol and octanoic acid.

Differences Between these Functional Groups

You learnt in the previous topic how the alkanol�“functional group�” (the -OH group) contains a polar bond,and how this causes the properties of the alkanols to bequite different to those of the corresponding alkanes.

Polarity of the AlkanolFunctional Group

Now, compare the -COOH Functional Group of theAlkanoic Acids:

Because of the presence ofanother electronegativeoxygen atom, there are two sets of dipoles on the molecule.

In an alkanol, the dipoles result in hydrogen bondingbetween molecules. For example, in ethanol:

Each pair ofmolecules can form one hydrogen bond

In an alkanoic acid, such as ethanoic acid:

Each pair ofmolecules canform twohydrogen bonds.

You already know that the alkanols have much higherm.p�’s & b.p�’s than the corresponding alkanes, because ofthe hydrogen bonding between molecules.

The m.p�’s & b.p�’s of the alkanoic acids are higher still,due to the presence of twice as many hydrogen bonds.

For comparison:Ethane Ethanol Ethanoic acid

Boiling Pt (oC) -89 +78 +118

26


5. ESTERIFICATION




General FormulaCnH2n+1OH(starting at n=1)

General FormulaCnH2n+1COOH(starting at n=0)

O-HH

OH C

O-HH

HH

H

C

H O-HH

H

H

H

H

C C

O-HH

CH3 CH2

H O-HH

OH

H

C C

O-HH

OCH3 C

O-HH

O

CH3CH H O-HH

HHH

H

H

C C CH H O-HH

OH

H

H

C C C

TThhee mmoossttiimmppoorrttaannttmmeemmbbeerr ooff bbootthhHHoommoollooggoouussSSeerriieess iiss tthhee 22-ccaarrbboonn �“�“EEtthh-�”�”ccoommppoouunndd

CCoommmmoonn nnaammee�“�“FFoorrmmiicc aacciidd�”�”

CCoommmmoonn nnaammee�“�“AAcceettiicc aacciidd�”�”

CH2

O HThis bond is polar

CO

O

HBoth thesebonds arepolar

""++

""++

""++

""++

""++

""-

""-""-

""-

""-

OO-HHCH

33 CH22

""++

""-

""++

""-

""-

""-

Hydrogeen bond

HHyyddrrooggeenn bboonnddss



27

The Esters�“Esters�” are a group of carbon compounds formed by thereaction between an alkanol and an alkanoic acid.

The reaction could be described as a �“condensation�”because it produces a water molecule, but it is sowidespread in nature, and so important, that it rates its ownname; �“Esterification�”.

The reaction can be vizualized as follows:

Example: If ethanoic acid reacted with butanol...

...the ester formed would look like this:

All esters contain this chemical �“functional group�”inside the molecule, withhydrocarbon chains oneither side.

This group of atoms is polar,but both ends of the molecule are non-polar. Estersgenerally have low solublity in water (some exceptions), andare volatile with a strong odour... often sweet and fruity.

As you�’ll learn, this is their �“job�”; to give the smells andflavours to many foods and perfumes.

Naming EstersSince the esters are made by joining together 2 othermolecules, it should be no surprise that they have a 2-partname.

�• Alkanol name first.Drop-off �“-ANOL�”, and add �“-YL�”.

In the example at lower left, �“butanol�” becomes �“butyl�”.

�• Alkanoic acid name second.Drop-off �“-IC ACID�”, and add �“-ATE�”

In the example lower left, �“ethanoic acid�” becomes�“ethanoate�” (the same as the ion from this acid).

The name of the ester in the photo is�“butyl ethanoate�”

ethanoic + butanol butyl ethanoate + wateracid

CH3COOH + C4H9OH CH3COOC4H9 + H2O



H-OO R2

H

HC

R2

H

HC

O-HH

O

OH

H

R1 C

Alkanoic acid(R1 represents therest of the molecule)

Alkanol(R2 represents therest of the molecule)

These atoms formwater, and the 2molecules join

O

OR1 C

Water+

Ester

H OH

OH

H

C CH H

HO HHH

H

H

H

H

C C C C

Remnant of CH3COOH Remnant of C4H9OH

C CO

O

You need to be able to name any esterwith up to 8 carbons on either end

(nothing bigger than �“octyl octanoate�”)


When the label says �“Flavour�”,it means �“Esters�”

The delicious smell andtaste of ripe fruit islargely due to naturalesters



28

Practical Work: Making an EsterYou may have made an ester in the laboratory using theprocess of Reflux.

A simple laboratory reflux set-up is shown.

Reflux condenseris open to the atmosphere atthe top.

This is vital to avoid anydangerous build-up ofpressure which would occurin a sealed flask.

Volatile chemical vapours rise,but are condensed and dripback into the flask... �“reflux�”.

Reaction Flaskis heated to speed up theotherwise very slow reaction.

As well as an alkanoic acid and an alkanol, it is usual toadd a catalyst, concentrated sulfuric acid, H2SO4.

alkanoic acid + alkanol ester + water

As well as speeding the reaction up, the H2SO4 catalystalso absorbs the water product. This has the effect ofshifting the equilibrium to the right (Le Chatelier�’sPrinciple) and increasing the yield of the ester.

Occurrence of EstersEsters occur widely in nature, especially in fruits andflowers. Esters are largely responsible for the smell andtaste of many foods.

There may be a complex mixture of esters and othercompounds which give the complete smell of (say) a ripestrawberry, but there�’s always an ester giving the main smelland taste sensation.

Some examples:Strawberry ethyl butanoate, C3H7COOC2H5

Orange octyl ethanoate, CH3COOC8H17

Fats & Oils are Esters, TooEsters made from very long chain �“fatty acids�”, and thetriple-alcohol molecule �“glycerol�” are used by all livingthings as high-energy foods and energy storage chemicals.We call them fats or oils, depending on their melting point.

Production and Uses of EstersAs well as their wide occurrence in living things, artificially-manufactured esters are important industrial chemicals.

They are produced by exactly the same process shown atthe left... reflux of the appropriate alcohol, acid andcatalyst... but on an industrial scale, of course.

Uses include:

�• artificial flavours for drinks and various processed foods.

�• solvents. Ethyl ethanoate is widely used as an industrial solvent (e.g. plastics industry) and is well known as thesolvent for nail varnish.

�• ingredients in many products, including shampoo and cosmetic products, and as �“plasticizers�” (softening agents)in some plastics, such as �“vinyl�”.



HHeeaattiinngg EElleemmeenntt

ccoonncc.. HH2SSOO4

Worksheet 5Part A Fill in the Blanks

The Alkanols, also called a)..............................., allcontain the functional group b).................. andhave the general formula c)......................................The -OH group contains a chemical bond whichis d)..............................., and allowse).......................................... bonding betweenmolecules. This is why the alkanols have m.p�’s &b.p�’s much higher than the correspondingf).........................................

The Alkanoic Acids contain the functional groupg)...................... This group contains 2 polarbonds, so 2 h)............................... bonds can formbetween molecules. This is whyi)................................................................ than thealkanols.

Esters are formed by the reaction ofj)............................... with ...........................................The other product is k)................................. Estersare named by the l)........................... first (with itsending changed to m)..........), followed by then)...................................... name (with its endingchanged to o)..........................)

Esters are made using the technique ofp)............................... The reaction flask is open tothe atmosphere to avoid any dangerous build-upof q)...................................... Volatile chemicalsvapourize, but are r)..................................... by thes).......................... .............................................. anddrip back into the flask. t).................................... isused as a catalyst, and also improves the yield byshifting the u)......................... because it absorbswater.

Esters occur widely in nature, being responsiblefor many of the v)......................... and.................................. of foods, especiallyw)................................ Long-chain esters ofglycerol are the x)............................... and............................

Artificially manufactured esters are used asy)........................ .................................. in processedfoods, as z).......................... in industry and asingredients in many products such asaa)............................. and ..........................................

Part B Practice problems

1. Names of EstersName the ester formed from

a) ethanol & propanoic acidb) propanol & ethanoic acidc) pentanol & methanoic acidd) methanol and pentanoic acide) hexanoic acid and butanolf) ethanoic acid and octanol

2. Condensed Structural FormulasFor each of the compounds above, give thecondensed structural formula for the

i) alkanolii) alkanoic acid

and iii) ester

The first has been done for you as an example.Answera) i) ethanol = CH3CH2OH

ii) propanoic acid = CH3CH2COOHiii) ester = CH3CH2COOCH2CH3

Note: although the alkanol comes first in naming,it may be more convenient to place the acidremnant first in the structural formula. Thissystem is used here throughout.

3. Names from StructuresFor each of the following esters:

i) give the name of the esterii) name the alkanol and acid use to make it

a) HCOO(CH2)3CH3b) CH3CH2COO(CH2)3CH3c) CH3(CH2)3COO(CH2)4CH3d) C4H9COOCH3e) C5H11COOC7H15

29








30 www.keepitsimplescience.com.au


TTHHEE AACCIIDDIICC

EENNVVIIRROONNMMEENNTT

CONCEPT DIAGRAM (�“Mind Map�”) OF TOPICSome students find that memorizing the OUTLINE of a topic

helps them learn and remember the concepts and important facts.Practise on this blank version.

Practice QuestionsThese are not intended to be "HSC style" questions, but tochallenge your basic knowledge and understanding of thetopic, and remind you of what you NEED to know at theK.I.S.S. Principle level.

When you have confidently mastered this level, it is stronglyrecommended you work on questions from past exampapers.

Part A Multiple Choice

1.The colour changes for some common indicators areshown:

Colour in ChangesIndicator Acid Base at pHBromothymol

blue yellow blue 6.2 - 7.6Methyl orange red yellow 3.1 - 4.4Phenolphthalein clear pink 8.3 - 10.0

In a solution with pH= 8.0 the colours of phenolphthalein,bromothymol blue and methyl orange (in order) would be:A. pink, blue, yellowB. clear, blue, redC. clear, yellow, yellowD. clear, blue, yellow

2.A household substance most likely to be very basic isA. vinegar.B. soap.C. sugar.D. milk.

3.The salt formed when nitric acid reacts with calciumhydroxide would be:A. calcium nitrateB. nitric hydroxideC. waterD. calcium nitric

4.An oxide compound, MO2 (M is not the correct symbol) isfound to react as follows:MO2 + 2NaOH(aq) 2H2O(l) + Na2MO3(aq)

It would be true to say that MO2 is A. an acidic oxideB. a conjugate baseC. amphiproticD. a basic oxide

5.According to Le Chatelier�’s Principle, a chemical system inequilibrium, which is then disturbed, will adjust itselfA. so that the disturbance is amplified.B. in the direction that releases energy.C. so that the disturbance is counteracted.D. in the direction that releases the pressure.

6.In the reaction

the change that would shift the equilibrium to theright would be to:A. increase the pressure.B. increase the concentration of hydrogen.C. increase the temperature.D. remove NH3 from the equilibrium mixture.

7.A significant natural source of sulfur dioxide in theenvironment is:A. lightning storms.B. volcanic eruptions.C. burning of fossil fuels.D. forest fires.

8.The major environmental impact of NOx gases is:A. Global Warming.B. Acid Rain.C. Ozone depletion.D. Smog.

9.The diagram shows a molecule ofthe weak acid, methanoic acid.

When dissolved in water, one of thebonds in the molecule may ionize.

The bond most likely to ionize is theA. H-C bond B. C=O bondC. O-H bond D. C-O bond

10.In the reaction

HBr + NH3 NH4+ + Br-

it is true to say that:A. NH3 is the base, because it has accepted a proton.B. HBr is the acid, because it has accepted a proton.C. NH4

+ is the conjugate acid of HBr.D. Br- is the conjugate acid of HBr.

11.If an undissociated molecule of an acid is represented by�“H-A�”, and the ionized acid by separate �“H+�” and �“A-�”symbols, which diagram could show a dilute solution of astrong acid?

A. B. C. D.

31





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

CC

OO-HH

OO

HH

H-AA

H-AA

H-AA

H-AA

H-AA

H-AA

H-AA

H++H++

H++

H++

H++H++

H++

H++

H++

A-

A-

A-

A-

A-

A-

A-A-

12.In a solution of pH= 10, the concentration of hydroniumions is:A. 10 molL-1. B. 1010 molL-1.C. 10-10 molL-1. D. 1 molL-1.

13.If you had 4 solutions of different acids

Hydrochloric NitricEthanoic Sulfuric

all with exactly the same molar concentration of acid,which two only would you expect to have the same pH?A. sulfuric and ethanoicB. hydrochloric and nitricC. nitric and sulfuricD. ethanoic and hydrochloric

14.The hydrogen phosphate ion, HPO4

2- is an amphiproticspecies. If it were to act as a base, then its conjugate acidwould be:A. H2PO4

- B. H3PO4 C. PO43- D. HPO3

3-

15.An ionic �“salt�” is found to be acidic in water solution.It is likely that this salt is the product of the reactionbetween:A. a strong acid and a strong base.B. a weak acid and a weak base.C. a weak acid and a strong base.D. a strong acid and a weak base.

16.The salt mentioned in Q15 was formed during a titration.The most appropriate indicator for the titration would be:

(hint: refer to the list given for Q1)A. bromothymol blue.B. methyl orange.C. phenolphthalein.D. universal indicator.

17.Which of the following graphs might be the �“titrationcurve�” for the titration described in Q15-16?

A. B. C. D.

18.An appropriate material to use in case of an acid spill is:A. an amphiprotic substance, like sodium bicarbonate.B. a strong base, like sodium hydroxide.C. a buffer solution, like methanoic/methanoate mixture.D. a weak base, like ammonia.

19.The ester shown is:A. butyl propanoateB. ethyl butanoateC. butyl ethanoateD. propyl butanoate

20.Of these compounds, which would you expect to have thehighest m.p. & b.p.?A. ethane B. ethene C. ethanol D. ethanoic acid

Part B Longer Response QuestionsMark values shown are suggestions only, and are to give youan idea of how detailed an answer is appropriate.

21. (4 marks)a) What is meant by an �“acid-base indicator�”?b) Identify an everyday use of an indicator.c) Describe how you could prepare a crude indicator froma named natural substance.

22. (7 marks)a) Write a balanced, symbol equation for the reactionbetween hydrochloric acid and magnesium oxide.b) Explain how this equation supports the classification ofmagnesium oxide as a �“basic oxide�”.c) Write a balanced, symbol equation which shows thatcarbon dioxide may be considered as an �“acidic oxide�”.

23. (8 marks)a) Explain why a sealed bottle of fizzy lemonade shows nosigns of bubbles within the liquid, yet bubbles formimmediately the lid is removed.Include a relevant chemical equation in your answer.

b) The reaction of carbon dioxide with water is exothermic.Given that information, predict the effect of highertemperature on the rate of bubble formation in a just-opened bottle of soft drink. Explain your prediction,naming the scientific principle involved.

24. (8 marks)a) Name a significant natural source of sulfur dioxide gas.b) Use a balanced, symbol equation to explain how thesmelting of some metal ores can produce sulfur dioxide.c) Name a serious environmental effect that can be causedby sulfur dioxide, and describe a possible environmentalimpact it causes.d) Calculate the volume of sulfur dioxide gas (measured atSLC) that can be formed from 1.00 tonne of sulfur.

32





Vol Base

22

pp

HH

77

1122

Vol Base

22

pp

HH

77

1122

Vol Base

22

pp

HH

77

1122

Vol Base

22

pp

HH

77

1122

25. (7 marks)A certain diprotic acid can be represented by the formula�“H2A�”. (�“A�” is not the correct symbol)

a) Write a balanced equation for the complete ionization ofH2A when added to water.Label each species in the equation as acid or base, includingconjugate species.

b) If H2A is a strong acid, calculate the pH of a solutionwith concentration [H2A] = 0.0250 molL-1.

c) In fact, when this exact concentration solution wastested, it was found to have a pH = 2.50.What do you conclude from this?

26. (4 marks)Explain 2 different reasons for adding acids to processedfoods. For each reason given, name a (different) acid usedfor that purpose.

27. (7 marks)a) Give the definition of an acid according to theArrhenius Theory.b) Write an equation which Arrhenius might have used toexplain why hydrogen chloride gas is an acid whendissolved in water.c) Give the definitions for acid and base according to theBronsted-Lowry Theory.d) Write an equation to show why hydrogen chloride gas inwater is an acid according to the B-L Theory.

28. (7 marks)a) A solution containing carbonate ions (CO3

2- ) is found tobe quite strongly basic.

Write an equation to explain why, and state the role of thewater molecule in this reaction.

b) Write TWO different equations to show the amphiproticnature of the hydrogen carbonate ion, HCO3

-.

29. (8 marks)A diluted vinegar sample (CH3COOH solution) wasanalysed by titration with a standardized solution of KOHwith concentration of 0.008263 molL-1.

25.00mL samples of the vinegar solution were used. Thetitration was done 4 times, giving burette titres of 27.35,26.70, 26.75 and 26.65mL of KOH.

a) Choose the most appropriate indicator for this titrationfrom the list shown in Q1, and justify your choice.

b) Write a balanced symbol equation for the reaction.

c) Explain how the titration measurements should be used.

d) Calculate the concentration of the diluted vinegarsolution.

30. (4 marks)a) What are the characteristics of a �“buffer solution�”?

b) In general terms, what are the usual ingredients of abuffer solution? Give a specific example.

c) Give an example of a natural buffer system.

31. (6 marks)a) Draw structural formulas for

i) methanolii) propanoic acid

iii) the ester formed by reaction of the above,and give the name of this ester.

b) Explain the need for using a reflux system to carry outthis ester preparation.

c) Name the chemical you would add to the reaction flaskas a catalyst.

d) Predict, in general terms, the odour of the ester.

33







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TM THE ACIDIC ENVIRONMENT - Amazon S3s3.amazonaws.com/.../Chem6.AcidicEnvironmentStudent.pdfAcid, Base or Neutral? The chemical definitions of acid and base will come later. For now,