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keep it simple science®

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 or 10 you would havestudied some basic Chemistry, and it probablyinvolved studying acids and bases andindicators... so these are familiar terms, even ifyou’ve forgotten the details.

This topic begins by reminding you of thissimple way to classify all chemical substances:

Acid,Base,

or Neutraland how

“Indicators” canbe used to identify

them.

Then, we look at the Chemistry of Oxide Compounds,

and link that to Acidity.

You will study the sources and problems thatcan result from

Acids in theEnvironmentand, revise & practice

MoleCalculations

The details of Acid-Base Reactions will bestudied

...and you will learn the real meaning of the pHscale, used for measuring acid and basestrength.

+ +

proton transfer

ACID BASE

pH = -llog[H++]

Along the way, you willlearn some important

techniques in Chemical Analysis

and a process you might neverhave heard of before,

but it is common in nature andin the Chemical Industry:

Esterification




2

Strong v. WeakAcids

Self-ionisation0f Water

pH Scale&

pH Calculations

Dynamic Equilibrium&

Le Chatelier’s Principle

Alkanols&

Alkanoic Acids

EstersStructure & Naming

Occurrence &

Uses of Esters

EsterificationReaction

Acid-BaseBehaviour of

Water.AMPHIPROTIC

Early Ideasabout

Acids &Bases

Neutralisation&

Titration

Acidic Oxides in the

Environment

Acidic & BasicSalts.

Buffers

Acids as ProtonDonors

Acids & Alkalis.Acidic & Basic

Oxides

THE ACIDIC

ENVIRONMENT

Acids, Bases&

Indicators Acidsin our

Environment

Acids &the pH Scale

Acid-BaseTheory

&Titration

Esters

CONCEPT DIAGRAM (“Mind Map”) OF TOPICSome students find that memorising the OUTLINE of a topic helps them learn and remember the concepts and important facts. As you proceed through the topic,

come back to this page regularly to see how each bit fits the whole. At the end of the notes you will find a blank version of this “Mind Map” to practise on.

Bronsted-LowryTheory

Acid, Base or Neutral?The chemical definitions of acid and base willcome later. For now, you are reminded of whatyou may have learnt in earlier Science classes.

Acids and bases are chemical opposites; if youadd one to the other they “destroy” (neutralise)each other, and the end product is neutral.Acids and bases are the opposite ends of achemical property called “acidity”, which ismeasured by a numerical scale called “pH”.

The word “acid” comes from the Latin for“sour”, and refers to the fact that natural, acidicchemicals (e.g. vinegar) are sour-tasting.

IndicatorsIndicators are chemicals which change colouraccording to the acidity of the solution they arein.

The original indicators were natural extractsfrom plants or other living things. Some, suchas litmus, are still in use today, as well as newer,synthetically made chemicals.

Modern Laboratory IndicatorsThe syllabus requires that you are familiar withthe common laboratory indicators listed below.

You will have done experimental work, adding 2drops of indicator to test tubes of acid, base andpure water (which is neutral) and recorded thecolours produced.

Your results should have agreed with thefollowing:

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 ofacidic or basic substances. However, its colourchange is rather indistinct, and can occur overquite a range of pH values... it is not a “sharp”change.

In contrast, Phenolphthalein cannot tell you thedifference between a glass of water and sulfuricacid. However, the colour change is verydistinctive, and occurs suddenly at a veryspecific pH value... it is very sharp. This is notmuch use in general identification ofsubstances, but in certain methods of chemicalanalysis it is very important.

Everyday Uses of IndicatorsSoil Testing. Some plants grow best in acidicsoils; others need slightly alkaline (basic)conditions. Farmers and keen gardeners usesimple test kits containing an indicatorandcolour chart, to test the soil. They can thenadjust the soil pH to get the best results.

Water Testing. Swimming pools need regulartesting for acidity to better maintain their waterquality and hygiene. Aquariums must bemaintained at very specific pH levels for thehealth of their inhabitants.

Effluent Testing. Acidity is a useful way toassess the levels of certain types of pollutionfrom industries. Industry technicians andGovernment authorities use indicators tomonitor the pH of waste water and naturalwaterways.

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1. ACIDS, BASES & INDICATORSkeep it simple science

®



ACID NEUTRAL BASE

1 3 5 7 9 11 13pH

Practical Work: A Natural IndicatorYou may have done practical work to prepareand test a natural indicator.

A good exampleis the common

garden plantHydrangea.

If you collect aflower head andput it through a

blender with alittle water and

ethanol, thefiltered liquid

extract willwork as a

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 usingvarious liquid and paper indicators to classify arange of household substances and foodstuffsas either acid, base or neutral.

By using appropriate indicators (such as“Universal”) you may have even been able todifferentiate between substances that are“mildly acidic” and “strongly acidic”.

Some Typical ResultsHousehold Substances Found to be...

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

MilkLiquid detergent

Acids and bases are chemicala)................................ If you add one to the other,they b)............................... each other.

Acidity is measured by the c).............. scale. Onthis scale, a neutral substance has a value ofd)....... Values above this indicate e)......................substances, while values below indicatef)........................ substances.

Indicators are chemicals which g)............................................. according to the h).........................of the solution they are in. The originalindicators where extracts from i)..........................

Worksheet 1 Acids, Bases & Indicators Student Name..........................................Fill in the blank spaces The colour changes for the following indicators

need to be learnt.Indicator Acid Neutral Base

Litmus j)............. purple k)..............

Phenolphthalein l)........... ................ ................

Methyl orange m)............. ................ .............

Bromothymol n).............. green/blue o).............blueIn everyday situations, indicators are used forpurposes such as p).......................................... forfarming and gardening, q).....................................for pools and aquariums, and for monitoringr)............................... from industries.

Worksheet 2 Practice Questions & Test Questions section 1Student Name..........................................1.

Each solution listed below has been tested withone or more indicators, and the colour is given.For each, state if it is 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.

2.A household substance most likely to be verybasic is

A. vinegar.B. soap.C. sugar.D. milk.

3. (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 simpleindicator from a named natural substance.

Laboratory Acids and “Alkalis”You are reminded that the common laboratoryacids 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 solublehydroxide compounds which are sometimescalled “alkalis”.

Name Formula Solution ofSodium hydroxide NaOH Na+

(aq)+ OH-(aq)

Potassium hydroxide KOH K+(aq)+ OH-

(aq)

The Acid-Alkali Reaction:Neutralisation

You should be familiar with this reaction fromearlier Science 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 thepotassium and sulfate ions are spectators. Youcan leave them out to form the net ionicequation:

2H+(aq)+ 2OH-

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

H+(aq)+ OH-

(aq) H2O(l)

All the familiar laboratory acids are solutionscontaining hydrogen ions (H+). The laboratory“alkali” bases are solutions containinghydroxide ions (OH- ). These will always react toform water, so acid and base have neutralisedeach other.

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

Acid-Base Properties of the Oxide Compounds

The common laboratory acids and alkali-basesare not the whole story. Most of the oxidecompounds of the elements show some acid-base behaviour too, and the periodic tablereveals another pattern.

Basic Oxides of the MetalsMost of the oxides of the metallic elements areconsidered basic because they can neutraliseacids, forming water and a “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 neutralising it, to 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)

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2. ACIDS IN OUR ENVIRONMENTkeep it simple science

®



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

OXIDE

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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

WORKSHEET at end of sectionWORKSHEET at end of section

6

Chemical EquilibriumThe concept of “Dynamic Equilibrium” wasintroduced in a previous Preliminary topic. Theexample used then was the equilibrium betweena solid ionic lattice and dissolved ions in asaturated solution.

At equilibrium, it seems (macroscopically) thatnothing is happening. However, at the atomiclevel, solid salt is dissolving into the solution,but at the same time, some ions areprecipitation out of solution to form a solidlattice These 2 processes are occurring at thesame rate. THEY ARE IN EQUILIBRIUM.

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

This is a dynamic equilibrium situation. Evenwhen the process seems finished, there areactually 2 reactions (one forwards , onebackwards ) occurring at the same rate, sothat nothing appears to be happening.

Shifting an EquilibriumA chemical equilibrium is “dynamic”, meaningthat things are moving back and forth. This alsomeans that it is possible to “upset” anequilibrium and cause it to shift to a newbalance between reactants and products.

Equilibrium shifting to the left results in lesscarbonic acid and the formation of more CO2gas. The shift in the equilibrium was caused bythe pressure change which occurred when thedrink was opened.

Factors That Can Cause Equilibrium Shift

PressureIf a reaction involves a gas, any change in thepressure of that gas will shift the equilibrium.With soft drinks, increasing the pressure of CO2drives the equilibrium to the right; decreasingCO2 pressure, shifts it to the left.

TemperatureIf you kept the gas pressure constant, but raisedthe temperature, you would find this equilibriumwould shift left. (i.e. less carbonic acid)Lowering the temperature would shift it to theright (more carbonic acid).

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

ConcentrationThe equilibrium could also be shifted by alteringthe concentration of (say) the carbonic acid. Ifyou added extra H2CO3(aq) somehow, theequilibrium would shift left. If you reduced theconcentration, (e.g. by adding an alkali whichwould react and destroy it) the equilibriumwould shift towards the right.




e.g. 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 andmeasure/calculate themass and volume of CO2gas released.

If you weigh the soft drinkand container accurately,then release the pressureso that it goes flat (thismay take a day or more)you will have measured asmall loss of mass, due toCO2 gas escaping.

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 (assumingstandard 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 openand pour a fizzy drink it forms

bubbles in the glass.

The sealed drink containsdissolved CO2, but as soon as

it is opened, the gas beginscoming out of solution,

because the equilibrium hasshifted to the left.

Eventually a new equilibriumis established with a lot lessgas dissolved... we say it has

“gone flat”.





7

Le Chatelier’s PrincipleIn 1885, the French Chemist Henri Le Chatelierdiscovered the underlying pattern in theseequilibrium shifts. The equilibrium always shiftsin the direction which counteracts the changethat upset it in the first place.

Temperature Effect on EquilibriumThe dissolving of CO2 to form carbonic acid isan exothermic reaction, so heat energy may beconsidered as one of the products:

So, if the temperature is raised, the equilibriumshifts left

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

If the temperature is lowered, the equilibriumshifts right in an attempt to release heat energyand raise the temperature again.

Effect of ConcentrationImagine a sealed container with an equilibriummixture within.

If you injected extra product so its concentrationincreased, the equilibrium shifts left, attemptingto reduce the concentration of product

If you somehow removed product so itsconcentration was reduced, the equilibriumwould shift right in an attempt to make more.

In every case, the equilibrium shift triesto counteract the change

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

the system will adjust itself in the directionwhich counteracts 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))

H22CO33((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 necessarydiversion... now back to acids.

There are 3 main acidic oxides that are ofconcern in the environment:

All 3 are gases produced by both naturalprocesses and by human activities. All 3 reactwith water in the environment to 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 Dioxide is a part of the great“Carbon-Oxygen” cycle in nature.

CO2 is also released into the air by natural bushfires, and by volcanic eruptions, but generallythere’s a balance.

For millions of years, huge quantities of carbonhave been “locked-away” in fossil materialssuch as coal and petroleum.

For the last 100 years or so, human activity hasbeen releasing this “fossil carbon” as CO2 byburning the fossil fuels. CO2 levels have risen30% or more, upsetting the balance, world-wide.

This is believed to be causing environmentalproblems

“Global warming”due to the “Greenhouse Effect”.

CO22

O22

All living things.Respiration

Plants.Photosynthesis

Sulfur Dioxideis released into the airfrom volcaniceruptions and hot-springs, but the naturallevels are extremelylow.

Human activities canpollute the air of aregion by releasing SO2 from:

Smelting of Sulfide OresThe ores of some metals (esp. lead, zinc,copper) contain sulfide compounds. To extractthe metal, the ore is roasted with air:

lead(II) sulfide + oxygen lead + sulfurdioxide

PbS + O2 Pb + SO2

Burning of Fossil FuelsSome fuels, especially coal, contain smallamounts of sulfur-containing compounds.When the coal is burnt, the sulfur burns too:

sulfur + oxygen sulfur dioxideS + O2 SO2

The main environmental concern with SO2release is “Acid Rain”

GGeeyysseerr eerruuppttiinngg,,YYeelllloowwssttoonnee NNPP

Nitrogen Dioxide occurs naturally inextremely small quantities due to the reactionbetween 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 referredto as the “NOx” gases.

High temperaturecombustions(especially coal-burning powerstations and vehicleengines) producelarge amounts ofNOx gases.

NOx pollution cancause toxic “smog”around large cities.

Sources of Acidic Oxides in the Environment

9

Environmental ImpactsEach of these acidic oxides can cause adifferent major environmental problem:

The main problem with SO2 pollution is itsacidity.

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 SO2pollution is serious, the acidity of rainfall stingsthe eyes, corrodes metals, erodes stonebuildings and monuments, and can haveserious environmental impacts:

Lakes and wetlands can become acidic enoughto kill plants and animals and disrupt the food-chains and the normal ecological balance.

Forests can be killed by acidity of the rain, andthe leaching of the soils by acids.

In the 1960’s (when Acid Rain became known)many forests were seriously damaged, includingthe famous “Black Forest” of Germany. InCanada, 15,000 lakes were known to be “dead” inan ecological sense.

Since then, emissions of SO2 have been limitedand damage reduced, but the threat of Acid Rainis still a serious one, especially in rapidlyindustrialising countries such as China.

Evidence for Acidic Oxide Pollution

We know that CO2 levels in the atmosphere haveincreased because measurements have beencollected over many years. The measurementsof SO2 and NOx gases have not been collectedfor as long, and these gases are rapidly“washed” out of the air by rain, so the evidencefor their presence is not so certain.

What we can be certain about are the localisedeffects of SO2 pollution in places where it is, orwas, prevalent. Some examples from the 1960’swere mentioned before. Closer to home and tothe present, is the evidence of devastationaround Queenstown, Tasmania.The environmental damage around Queenstownwas caused by the release of SO2 from theMt.Lyell copper mine and smelter. Althoughpollution ceased over 20 years ago, theenvironment 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“Global Warming”

Acidity is not the main concernwith the NOx gases either. In the

next topic you will study theChemistry of “smog” and ozone

pollution. (But NO2 doescontribute 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 smeltinga copper ore containing mainly copper(I) sulfide.After concentrating the ore by froth flotation, thematerial was smelted by roasting in a furnacewith a blast of air to provide oxygen:

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

How much SO2(g) is produced from 1 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!!)


Acids and bases react with each other to forma)................... and a “b)............................” andc)................................... each other.

The oxides of most metals act asd)........................., in that they will react with anacid to form e).................. and a .......................

Oxides of non-metals mostly act asf).......................... because they will eitherdissolve in/react with water to form ang)..................... and/or they react with bases andh).......................................... them.

Chemical i).................................... occurs when areaction, and its opposite are occurring at thesame j)......................., so that the concentrationsof 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 positionaccording to n)......................................................Principle. This states that when an equilibrium isdisturbed, the system will shift in the directionwhich o)..................................................................................................................

10



Worksheet 3 Acids in the Environment Fill in the blank spaces Student Name..........................................

There are 3 acidic oxides which are of concernenvironmentally:

Carbon dioxide is a natural part of thep).....................- ............................ cycle in nature.Human activities have increased the CO2 levelsmainly from q)........................... .............................The major problem arising is “r).........................warming” due to the “s)........................................”

Sulfur dioxide occurs naturally int)............................ eruptions and u).......................Human activities which release SO2 includeburning of v)................................ and thew).............................. of some metal ores. Themain environmental problem is “x).......................................” which can cause serious ecologicaldamage to lakes and forests.

Nitrogen dioxide and nitric oxide (collectivelyknown as y)................ gases) are producednaturally in very small amounts byz)............................. Human activities whichproduce them are high temperaturecombustions in aa)............................... and................................... Environmentally, NOxgases are the main cause of toxic“ab)..................” in large cities, but also cancontribute to Acid ac).......................

Worksheet 4 Practice Questions Acid-Base Reactions Student Name..........................................

1. Acid-Alkali Reactionsa) Name the salt formed in a reaction between:i) hydrochloric acid & calcium hydroxide

ii) sulfuric acid & magnesium hydroxide

iii) nitric acid & barium hydroxide

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

ii) sulfuric acid and sodium hydroxide

iii) nitric acid and magnesium hydroxide

2. Reactions of Basic OxidesWrite a balanced equation for the reaction of:

a) sulfuric acid & iron(II) oxide

b) hydrochloric acid & magnesium oxide

c) nitric acid & copper(II) oxide

3. Reactions of Acidic Oxidesa) carbon dioxide reacts with calcium hydroxideto form water and calcium carbonate. (This isthe “limewater” reaction) Write a balancedequation for the reaction.

b) P2O5 is an acidic oxide. It reacts with water toform phosphoric acid, H3PO4. Write thebalanced equation.

c) Sulfur trioxide reacts with water to form astrong acid. Write a balanced equation, andname the acid.


11



Worksheet 5 Practice Questions Le Chatelier’s Principle Student Name..........................................1. NO2(g) reacts with itself as follows:

The reaction to the right is exothermic, so heatcan be considered as a “product”.

Note that as the reaction proceeds to the right, 2moles of gas form 1 mole of gas, so in a fixedvolume container the pressure would drop asthe reaction proceeds to the right.

Imagine a sealed container in which a mixture ofthese gases has reached equilibrium.

State which way the equilibrium would shift ifeach of the following disturbances were made tothe mixture. Explain each answer.

i) Increase in temperature.

ii) Compress the mixture, therby increasing pressure.

iii) Injecting extra N2O4, without changingpressure.

iv) Decrease the temperature.

v) Spray in a little water. (NO2 dissolves, N2O4 does not.)

vi) Decreasing the total gas pressure.

2NO2(g) N2O4(g)

2NO2(g) N2O4(g) + heat

2. If hydrogen iodide (covalent molecular) isdissolved in water, some of the moleculesionise, and an equilibrium is reached:

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, andreduces their concentration.

iii) Dissolving extra HI in the solution.

iv) It is found that raising the temperature of anequilibrium mixture has the effect of increasingthe concentration of ions. Deduce whether thereaction as written is exo- or endothermic.

3. Ammonia is manufactured from its elementsby the reaction

i) To maximize the yield of ammonia, thereaction is carried out under very high pressure.Explain how this helps.

ii) The temperature of the reaction is kept fairlyhigh to speed up the rate of the reaction. Whateffect does higher temperature have on theequilibrium?

iii) During the reaction, ammonia is constantlyremoved from the reaction vessel, and morereactant gases constantly pumped in. Explainthe effect this has.

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

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


iii) 100L of Ne at 0oC & 100kPa?

iv) 25.0mL of O2 at 0oC & 100kPa?

b) What is the volume (at 25oC & 100kPa) of:i) 100g of CO2 ?

ii) 100g of He ?

iii) 1.50g of N2 ?

iv) 1.00kg of Ar ?

3. Problems Involving Reactionsa) Carbon dioxide gas reacts with aqueouscalcium hydroxide (limewater) to form water andinsoluble calcium carbonate.

i) Write a balanced equation for the reaction.

ii) What mass of calcium carbonate would beformed by the reaction, if 1.00L of CO2 isproduced, measured at 25oC & 100kPa?

iii) What volume of CO2 (at 0oC & 100kPa) mustbe absorbed by limewater in order to precipitate1.75g of CaCO3 ?

b) In the smelting of zinc, the crushed,concentrated ore is zinc sulfide. This is roastedin a blast of air, forming zinc oxide and sulfurdioxide gas.

i) Write a balanced equation for the reaction.

ii) Calculate the volume of pure oxygen gas (at25oC & 100kPa) required for complete reactionwith 1.00 Tonne of zinc sulfide.

iii) Assuming the air used is 21% oxygen, whatvolume of air must be supplied to the smelter foreach tonne of ore?

iv) From your answer to (ii), use Gay-Lussac’sLaw to find the volume of SO2 (at 25oC &100kPa) released from one tonne of ZnS.

12



Volume of Gases under Different Conditions

You know that the volume of a gas changes atdifferent temperatures and pressures, so themolar volume changes as well.

At 25oC & 100kPa, 1 mole of any gas occupies 24.79 Litres

What about at other temperatures and pressures?

Another set of standard conditions commonlyused in Chemistry is 0oC and 100kPa.

Under these conditions, 1 mole of any gas occupies 22.71 Litres

Notes:1. You do NOT need to remember these values.They are always given in the Data Sheet intests/exams.

2. There is a formula for calculating the volume atany temperature and pressure, but its use is notrequired in this course.

Worksheet 6 Tutorial & Practice Questions Molar Gas Volumes Student Name..........................................

1. Molar Gas Volumes

a) What is the volume of:i) 2.59 mol of O2 at 25oC & 100kPa?

ii) 0.0453 mol of H2 at 0oC & 100kPa?

iii) 120 mol of CO2 at 0oC & 100kPa?

iv) 4.67 x 10-2 mol of N2 at 25oC & 100kPa?

b) How many moles of gas is:i) 12.4L of He at 25oC & 100kPa?

ii) 250mL (=0.250L) of O2 at 0oC & 100kPa?

iii) 10,000L of N2 at 25oC & 100kPa?

iv) 1.00mL of Ar at 0oC & 100kPa?

2. Mass - Volume of Gases

a) What is the mass of:i) 5.00L of CO2 at 25oC & 100kPa?

ii) 5.00L of H2 at 25oC & 100kPa?

Multiple Choice1.The salt formed when nitric acid reacts withcalcium hydroxide would be:A. calcium nitrateB. nitric hydroxideC. waterD. calcium nitric

2. An oxide compound, MO2 (M is not the correctsymbol) is found 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

3.According to Le Chatelier’s Principle, achemical system in equilibrium, which is thendisturbed, will adjust itself A. 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.

4.In the reaction

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

5.A significant natural source of sulfur dioxide inthe environment is:A. lightning storms.B. volcanic eruptions.C. burning of fossil fuels.D. forest fires.

6.The major environmental impact of NOx gasesis:A. Global Warming.B. Acid Rain.C. Ozone depletion.D. Smog.

Longer Response QuestionsMark values shown are suggestions only, and are togive you an idea of how detailed an answer isappropriate. Answer on reverse if insufficient space.

7. (7 marks)a) Write a balanced, symbol equation for the reactionbetween hydrochloric acid and magnesium oxide.

b) Explain how this equation supports theclassification of magnesium oxide as a “basic oxide”.

c) Write a balanced, symbol equation which showsthat carbon dioxide may be considered as an “acidicoxide”.

8. (8 marks)a) Explain why a sealed bottle of fizzy lemonadeshows no signs of bubbles within the liquid, yetbubbles form immediately the lid is removed.Include a relevant chemical equation in your answer.

b) The reaction of carbon dioxide with water isexothermic. Given that information, predict the effectof higher temperature on the rate of bubble formationin a just-opened bottle of soft drink. Explain yourprediction, naming the scientific principle involved.

9. (8 marks)a) Name a significant natural source of sulfur dioxidegas.

b) Use a balanced, symbol equation to explain howthe smelting of some metal ores can produce sulfurdioxide.

c) Name a serious environmental effect that can becaused by sulfur dioxide, and describe a possibleenvironmental impact it causes.

d) Calculate the volume of sulfur dioxide gas(measured at 25oC & 100kPa) that can be formedfrom 1.00 tonne of sulfur.

13




Worksheet 7 Test Questions section 2 Student Name.....................................

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

Acids as Proton DonorsSo far, you have learnt a number of things aboutacids, but not the most basic thing... thechemical definition of just what an acid is!

What defines all acids is their ability to donate ahydrogen ion (H+) to another species. Since ahydrogen ion is really just a “naked” protonfrom the nucleus of the hydrogen atom, theformal definition of an acid is:

Acids in Aqeous SolutionWhen acids dissolve in water, we often justimagine an ionisation 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”) youlearnt about how ions are “hydrated” by polarwater molecules when in solution. The dissolvingof HCl in water could be represented as:

Although there may be many water moleculesclinging to a H+ ion, for simplicity we imaginethere is just one, and it forms a special ion,called the “hydronium ion”, H3O+.

HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

HCl is an acid because it has “donated” aproton to a water molecule. The hydronium ionin the solution is an acid because it can, in turn,donate a proton to other species.

Formation of Hydronium IonsMore examples:

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

(aq) + NO3-(aq)

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

(aq) + SO42-

(aq)

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

(aq) + PO43-

(aq)

The “Organic” (Carbon-Based) Acids

Many naturally-occurring, biological moleculesare acids too. These are carbon-compounds,made by living things and most contain aspecial chemical group you should becomefamiliar with: the “-COOH” group.

Perhaps the most important is ethanoic acid,CH3COOH:

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

CH3COOH(aq)+ H2O(l) H3O+(aq)+ CH3COO-

(aq)

ethanoic + water hydronium + ethanoateacid ion ion

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

14

3. ACIDS & THE pH SCALEkeep it simple science

®



An acid is a chemical species which DONATES PROTONS

δδ−−δδ++

Cl- ++

δδ++

δδ++δδ−−δδ−−

δδ−−

δδ−−

δδ++

δδ++

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)

CC

OO-HH

OOThis group is

usuallyabbreviatedas “-COOH”

and is alwaysacidic.

This bond is polar, and ionises in water

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

WORKSHEET at end of section

15

Citric Acid is familiar as the acid in citrusfruits... oranges & lemons.

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

Molecular formula is C6H8O7, but the structuralformula is more meaningful:

Remember thatonly the -COOH

groups are acidic, and

you will see that citric acid is

triprotic

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

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

For example, HCl is a STRONG acid; whenadded to water the reaction...

HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

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

In contrast, ethanoic acid is a WEAK acid; whenadded to water, the reaction...

CH3COOH(aq)+ H2O(l) H3O+

(aq)+ CH3COO-(aq)

...reaches an equilibrium with only about 1% ofthe molecules ionised. (i.e. the equilibrium lieswell to the left, favouring the reactantmolecules.) If the concentration of the solutionwas 1 molL-1, then the hydronium ionconcentration would be only about 0.01 molL-1.

It is quite possible to have a “concentratedsolution of a weak acid” or a “dilute solution ofa strong acid”... just be careful with theseprecise meanings!

The pH Scalehas already been introduced, and you should befamiliar with 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 Ionisation in Solution

WEAK Acid = Partial Ionisation 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”.

2. “log10” means the “logarithm to base 10”. Thisis a mathematical function, best understood byexample: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 todo log 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.


16

More about pHSince the numbers on the pH scale are powersof 10, it follows that if an acid solution has a pHone unit lower than another, it is actually 10times more acidic. Two units on the pH scalerepresents 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 2units 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 of exactly the same concentration.




Measuring pHYou may have carried out some simple

laboratoryexperiments, usinga pH meter (or data-logger probe) tomeasure the pH ofvarious solutions.

You may haverecorded the pHvalue, then usedyour knowledge ofthe pH scale todecide if eachsolution 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, theionisation 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 youwill have 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)+ SO4

2-(aq)

Since H2SO4 is a STRONG ACID,the ionisation is 100%. It isDIPROTIC, so each mole of H2SO4produces 2 moles of H3O+.

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 youwill have found that a 1 molL-1

H2SO4 solution has a pH = -0.30.(below zero)

1.00 molL-1 Ethanoic Acid

When CH3COOH dissolves in water:

CH3COOH(aq)+H2O(l) H3O+

(aq)+ CH3COO-(aq)

Since this is a WEAK ACID, theionisation is incomplete.

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

and the pH can be expected to bewell above the values found for HClor H2SO4.

Sure enough, by experiment youwill have found that a 1 molL-1

CH3COOH solution has a pH ≅≅ 2.




17

The Self-Ionisation of WaterWater itself can be considered as an extremelyweak acid, since in pure water a few watermolecules ionise and donate a proton toanother molecule.

In pure water, at 25oC, it turns out that theconcentration of 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

δδ+δδ-δδ+δδ-

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 tomato sauce to ice-cream topping to driedfruits or tinned sausages, you may find “foodacid (260)” (or something similar).

“Food acid (260)” is in fact ethanoic acid, themain chemical in vinegar, and a commoningredient in many recipes.

The reasons that acids are added to many foodsare:

• Preservatives. Adding acids to a food lowers thepH and makes it more difficult for some bacteriaor fungi to grow within the food. This is the mainreason why some foods, such as tomato sauce,do not go bad even when not refrigerated. Acids most commonly used are SO2 & ethanoic.

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

• Nutrition. Some foods have vitamins added toincrease their nutritional value. Most commonlyused is ascorbic acid (vitamin C).

Some Naturally Occurring Acids & Bases

AcidsEver been bitten by a bull-

ant or stung by a bee?

You were injected with“formic acid” (methanoic

acid), HCOOH.

Ethanoic acid is the main chemical in vinegar(literally “wine-sour”) known and used forthousands of years.SUGAR (grapes)

fermentationETHANOL (wine)

oxidationETHANOIC ACID (vinegar)

BasesLime is a major ingredient of cement, and hasbeen known and used for thousands of years.Chemically it is calcium oxide, CaO.

It is prepared by roasting limestone (or evenoyster shells) which results in thedecomposition

CaCO3 CO2 + CaO

“Lye” is the name given to the alkali liquidobtained by soaking wood ashes in water. Thebasic ingredient is a mixture of hydroxidecompounds, mainly potasium and sodiumhydroxides.

Lye has been used for centuries in soap-making.

Acids can be defined as species whicha)...................... ........................... In water solutionan acid molecule donates a proton (which is ab)........................ ion) to a water molecule,forming a c).................................... ion. Acidswhich donate one proton are calledd)..................................; those that donate 2protons are e)............................., and those whichdonate 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 weakrefers to an acid which j).........................................................................

Acidity is measured by the pH scale. The pHvalue is calculated from the equation:pH = k)...............................

2. Calculating pH from [H3O+]Each of the following values is the hydronium ionconcentration of a solution, in molL-1. For each,calculate the pH, and state if the solution is acidic orbasic.

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 into accountthat some acids may be diprotic or triprotic. (assumetotal ionisation)

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

18




Worksheet 8 Acids & pH Fill in the blank spaces Student Name..........................................

pH values are l)............................. of 10, so each1 unit of pH actually means a change in acidityof m)............... times.

Acids are commonly added to foods, in order to:

• help n)............................... the food by making itmore difficult for o)................................. and............................... to grow in the food and causeit to spoil. The acids most commonly used thisway are p)........................ and ................................

• q).............................. the food, by giving it ther)......................... taste that all acids have. Mostcommonly 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).....................................................

Worksheet 9 Practice Questions Acid Ionisation & pH Student Name..........................................1. Write a balanced equation for the reaction of eachacid with water. (Assume complete ionisation in eachcase, even though some may be weak 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)

Multiple Choice1.The diagram shows a moleculeof the weak acid, methanoicacid.

When dissolved in water, one of the bondsin the molecule may ionise.

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

2. If an undissociated molecule of an acid isrepresented by “H-A”, and the ionised acid byseparate “H+” and “A-” symbols, which diagramcould show a dilute solution of a strong acid?

A. B. C. D.

3.In a solution of pH=10, the concentration ofhydronium ions is:

A. 10 molL-1. B. 1010 molL-1.C. 10-10 molL-1. D. 1 molL-1.

4.If you had 4 solutions of different acids

Hydrochloric NitricEthanoic Sulfuric

all with exactly the same molar concentration ofacid, which two only would you expect to havethe same pH?

A. sulfuric and ethanoicB. hydrochloric and nitricC. nitric and sulfuricD. ethanoic and hydrochloric


5. (7 marks)A certain diprotic acid can be represented by theformula “H2A”. (“A” is not the correct symbol)

a) Write a balanced equation for the completeionisation of H2A when added to water.Label each species in the equation as acid orbase.

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

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

6. (4 marks)Explain 2 different reasons for adding acids toprocessed foods. For each reason given, namea (different) acid used for that purpose.

19




Worksheet 10 Test Questions section 3 Student Name..........................................

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-

FOR MAXIMUM MARKS SHOWFORMULAS & WORKING,

APPROPRIATE PRECISION & UNITS IN ALL CHEMICAL PROBLEMS

Early Ideas About Acids & BasesNaturally occurring acids, like ethanoic acid invinegar, have been known for thousands ofyears, and described by their simple,observable properties such as their sour taste.

Chemistry became a modern Science just over200 years ago, and one of the earliest scientifictheories about acids was made by

Antoine Lavoisier (French, 1780’s)Lavoisier used simple plant-extract indicators (e.g.litmus) to identify acids and bases. He found byexperiment that all the non-metal oxide compoundshe tested produced acid solutions. He concluded thatacids must contain oxygen.

Humphry Davy (English, 1820’s)Forty years later and armed with more chemicalknowledge, Davy realised that Lavoisier was wrongabout acids. While some acids do NOT containoxygen (e.g. HCl), he found that all known acidscontain hydrogen.

Davy also discovered that metal oxides are basic anddescribed the patterns in the way acids react withmetals to form hydrogen gas.

These early ideas were all empirical descriptions ofproperties; they described the properties of acidsdiscovered by experiment, but did not include ageneral theory to explain or predict chemicalbehaviour.

That had to wait for Svante Arrhenius (Swedish, 1880’s)The Arrhenius Theory of acids-base behaviouris still generally used in years 7-10 Science.

The Arrhenius Theory was very successful inaccounting for simple acid-base behaviour in watersolution. It explained the neutralisation reaction, andcould explain strong and weak acids as being due tocomplete or partial ionisation.

However, the Arrhenius Theory had somedeficiences:

• It could not account for acid-base behaviourthat was not in water solution.

• It could not explain why some ions (which doNOT contain any hydrogen, hydroxide or oxide)showed acid- base behaviour in water solution.

You will soon learn that the main reason for allthese deficiencies was that Arrhenius failed toconsider the role of the solvent itself, and inwater solutions this is critical.

The Bronsted-Lowry Theorywas developed independently by JohannesBronsted (Danish) and Thomas Lowry (English)in 1923. The basic concepts of this theory havealready been used earlier in this topic:

Central to the Bronsted-Lowry (B-L) Theory isthe concept of “conjugate” species. Forexample, when ethanoic acid dissolves in water:

CH3COOH(aq)+ H2O(l) H3O+(aq)+ CH3COO-

(aq)

ethanoic + water hydronium + ethanoateacid ion ion

Acid Base Conjugate ConjugateAcid Base

The CH3COO- ion is the conjugate base ofEthanoic acid, because if this reaction was to runin reverse, the ion could act as a base and accepta proton to form ethanoic acid again.

The water molecule is acting as a base, becauseit accepts a proton to become the H3O+ ion. TheH3O+ ion is the conjugate acid because, if theequilibrium shifts left, it can donate a proton andbecome a water molecule again.

Another Example: Dissolving of HCl in water:

In this case, the reaction is very unlikely to everrun in reverse, but Cl- is still considered theconjugate base.

As before, water has acted as a base, and itsconjugate acid is the hydronium ion.

20

4. ACID-BASE THEORY & TITRATIONkeep it simple science

®



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

Acid-Base Reactions involve transfer of Protons.

An Acid is a Proton Donor.A Base is a Proton Acceptor.

Cl-

HCl((gg))

moleculeH22O((ll))

moleculeCl-((aaqq))

ionH33O++

((aaqq))

ion++

++ ++

++

Acid Base Conjugate Conjugate Base Acid





21

Water Can be an Acid, TooIn both the previous examples, the watermolecule acted as a base by accepting a protonto form a H3O+ ion. Water can also act as an acidand donate a proton.

For example, when ammonia, NH3 dissolves:

In this reaction the water molecule acts as anacid by donating a proton. Its conjugate base isthe hydroxide ion, OH-.

This equation explains why ammonia is a base,and why a solution of ammonia (NH3(aq)) can beconsidered as a solution of ammoniumhydroxide, NH4OH(aq).

Acidic and Basic SaltsOne of the weaknesses of the Arrhenius Theorywas that it could not explain the results ofsimple experiments you may 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 protons

are called “AMPHIPROTIC”

You will meet more Amphiprotic species soon

You may have done simpleexperiments using a pH meter,or Universal Indicator, to testthe pH of solutions of variousionic “salts”, none of which

seem to be obviously acids orbases, from their formulas.

You may have tested many, butthey probably included...

sodium chloride, NaCl

sodium carbonate, Na2CO3

sodium hydrogen sulfate,NaHSO4

...all at the same concentration,in pure water.

Typical Results

Salt pH Conclusion

NaCl 7 Neutral

Na2CO3 11 Basic

NaHSO4 2 Acidic

The reasons WHY some saltsshow acid-base behaviour is

explained above right.

Note: Due to CO2 from the airdissolving in the solution, the

pH of the NaCl solution may beslightly acidic.

Acidic SaltsAs some ions dissolve in water they react asacids, forcing the water molecule to be a base. Anexample 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 sulfatethere would also be some sodium ions in thesolution. In terms of acid-base behaviour, theywould be merely spectators.

Basic Salts dissolve in water and react asbases, forcing the water molecule 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 be sodium ion “spectators” in thesolution 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)


22

NeutralisationNow that you know about B-L Theory, we can goback to look again at the simple acid-alkalineutralisation.

Examplehydrochloric + sodium water + sodium acid hydroxide chloride

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

However, now we know that what makes HCl anacid is really 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 ionicequation is...

hydronium + hydroxide waterion ion

H3O+(aq) + OH-

(aq) 2H2O(l)

... and this is the net reaction for ALL the simpleacid-alkali reactions, regardless of which acidor which alkali are used.

Heat of NeutralisationIf you carried out the reactionabove in a calorimeter, youwill quickly find that thetemperature rises... thereaction is exothermic.

If measured and calculated, itis found that the value for ΔΔHis the same regardless of which acid and alkaliis used.

H3O+(aq) + OH-

(aq) 2H2O(l) ΔΔH= -56kJmol-1

Salts from Acid-Base Neutralisations

Although we might view all neutralisations asessentially the same reaction and ignore the“spectator ions” which form the “salt”, thisprevents us noticing a useful pattern which wasmentioned on the previous page.

The nature of the “salt” formed depends onwhether the acid and base were “strong” or“weak”, as follows:

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

Although most of the simple examples ofneutraliation involve the reaction at left, andusually we study reactions taking place in water,be aware that this is not always the case. Forexample, if you add together the dry gasesHCl(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 isclearly an acid-base reaction, according to theB-L definition.




ACID + ALKALI WATER + A “SALT”

The reaction involves a proton transfer

Ener

gy C

onte

nt

ΔΔHnegative

Eaa

StrongAcid

e.g. HCl,HNO3,H2SO4

WeakAcid

e.g.CH3COOH

H2CO3

StrongBase

e.g. KOH, NaOH

Weak Base

e.g. NH3

formsNEUTRAL salts

e.g. KCl, NaNO3K2SO4

formsNEUTRAL salts

e.g.CH3COONH4,

(NH4)2CO3

forms ACIDIC salts

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

forms BASIC salts

e.g. CH3COONa,Na2CO3

Why is Ammonia (NH3) a “Weak Base”?

Simple!Because when it reacts with water, the

reaction does NOT ionise fully, butreaches an equilibrium with a significantconcentration of un-ionised molecules

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.




23

TitrationTitration has been one of the most important techniques in Chemical Analysis for over acentury. 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 reaction for which youcan determine the “end-point”. Acid-base

neutralisations are ideal because indicators (or pHprobes) can identify the end-point.

Burettemeasures the volume of“standard solution” neededto reach the end-ppoint.

Burette tap controls the addition of“standard solution” to the“unknown solution”.

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

Filter Paperunder flask allows betterseeing of indicator 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 agreementwith the others). Average titre = 33.90 mL

Titration Formula: Ca x Va = Cb x Vba b

Ca, Cb = concentrations (molL-1) of acid & base solnsVa, 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.

The purpose of titration is Chemical Analysis. Itallows the concentration of an “unknown”solution to be determined by calculation, aftermeasuring the volume of a known-concentration“standard solution” which reacts with a sample toreach the “equivalence-point” (end-point). This isthe point where the reactants have beenconsumed in exactly the molar ratio specified bythe balanced equation.

Titration Technique• You must have a “standard solution” (next page)

• The burette should be rinsed with smallquantities of the standard solution, then filled withit. The level is then adjusted to the zero mark,ensuring that there are no air bubbles in the tapand delivery spout.

• An exact, measured volume of “unknown”solution is placed in the reaction flask by pipette.(“Aliquot”)

• 1-2 drops of suitable indicator are added to theflask. (See “Choosing the Indicator”, next page)

• The reaction proceeds by careful addition, withmixing, of solution from the burette. Add drop-by-drop approaching the end-point, until theindicator just changes colour.

• Record the volume of solution delivered byburette. Repeat 3 times, and average the “titres”(= burette volumes). It is common to discard anyvalues which are not in close agreement. The rest is calculation.

24

Preparing a Standard SolutionThe key to titration is to have available asuitable “standard solution”. The technique formaking a solution to an exact knownconcentration was covered in the Preliminarytopic “Water” and is revised below.

The problem is to get a substance which is ofvery high purity and stability to use as a“Primary Standard” to make the solution from.

The common acids like H2SO4, and bases likeNaOH, cannot be obtained in the pure state dueto the way they rapidly absorb water and/or CO2from the atmosphere.

Suitable primary standard substances are:

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

In the example (previous page) the standardsolution was H2SO4. This would not have been aprimary standard solution.

It may have been “standardised” by titrationagainst a primary standard Na2CO3 solution, sothat its exact concentration is known. Then itcan be used, in turn, as the titration standardsolution.

Selecting the IndicatorChoosing the best indicator to use is not just amatter of which colour change you’d preferlooking at.

Each indicator changes colour over a range ofpH values, and the trick is to choose anindicator which will change colour as close tothe “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

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

No, its not! Remember (see p.22) that the “salts”formed by different combinations of strong orweak acids and bases may have acidic or basicproperties. This means that, at the exact endpoint, the pH might not be neutrality.




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

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 lastdrop of chemical at the end point causes a hugepH change.) See graph next page.

Strong Acid - Weak BaseIf titrating (say) H2SO4 against 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”.

Add water to flask to fill it tothe mark.

(Use a dropper to avoid over-shooting)

Insert stopper & mix well.


25

Measuring pH During TitrationIf a titration is carried out as shown below, with a pHmeter or data-logger probe attached to a computer,

the pH changes occurring in the solution can berecorded and graphed.



Magnetic Stirrer

Connection tocomputer or pH meter

pH probe electrode

Burette adds standardbase solution

Flask contains“unknown” acid

solution

pH

0

2

4

7

10

1

2

Volume of Base Added

Strong Acid &

Strong Base

TTiittrraattiioonneenndd-ppooiinntt

TToo ssttaarrtt wwiitthh,,aaddddiinngg bbaassee hhaasslliittttllee eeffffeecctt oonn ppHH

NNeeaarr tthhee eenndd-ppooiinntt,,ppHH cchhaannggeess rraappiiddllyywwiitthh eeaacchh ddrroopp ooff

bbaassee aaddddeedd,, ssoo tthheeeenndd-ppooiinntt ccaann bbee

ddeetteerrmmiinneedd pprreecciisseellyy..

pH

0

2

4

7

10

1

2


Strong Acid &

Weak Base


TThhee eenndd-ppooiinntt iissaatt ppHH<<77 bbeeccaauusseetthhee ssaalltt iiss aacciiddiicc......

mmuusstt cchhoooosseemmeetthhyyll oorraannggee,, oorr

ssiimmiillaarr..

pH

0

2

4

7

10

1

2


Weak Acid &

Strong Base


TThhee eenndd-ppooiinntt iiss aatt ppHH>>77bbeeccaauussee tthhee ssaalltt iiss

bbaassiicc......

mmuusstt cchhoooossee iinnddiiccaattoorrssuucchh aass pphheennoollpphhtthhaalleeiinn

pH

2

4

7

1

0


Weak Acid &

Weak Base

TThhee eenndd-ppooiinntt iisstthheeoorreettiiccaallllyy aatt ppHH==77bbuutt iiss hhaarrdd ttoo ffiinndd bbyy

iinnddiiccaattoorr bbeeccaauussee eeaacchhddrroopp ooff bbaassee iiss ccaauussiinngg

lleessss ppHH cchhaannggee

???

In every case, the end-point is thecentre of the vertical (or near-

vertical) part of the graph.

The graph at left shows why weak acid-weak base titrations

are best avoided.

TThhee ““vveerrttiiccaall”” ppaarrtt ooff tthhiissggrraapphh eexxtteennddss ssoo ffaarr eeaacchhssiiddee ooff nneeuuttrraall,, tthhaatt tthhee

cchhooiiccee ooff iinnddiiccaattoorr iiss rreeaallllyynnoott ccrriittiiccaall...... aallll iinnddiiccaattoorrss wwiillll

cchhaannggee ccoolloouurr oonn tthhee ssaammeeddrroopp ooff aaddddeedd bbaassee..




26

A Natural Buffer SystemThe classic example of a natural buffer system is theway our blood is maintained at a pH = 7.4, despitethe fact that 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 addsa further dimension 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 besummarised as:

This system is highly effective at maintaing thepH of our blood. Be aware that it is not the onlybuffer operating.

pprroottoonn ttrraannssffeerr

H2CO3 HCO3- CO3

2-+ H++ + H++

- H++ - H++

Using Neutralisation on Chemical Spills

Whether its a small vinegar spill at home, or asulfuric acid tanker leaking after a road accident,a knowledge of neutralisation can help damagecontrol and safety.

So, if someone gets splashed with a strongbase, neutralise it by throwing a bucket of acidover them, right?

Wrong! Unless you can guarantee to apply theexact molar quantity for titration-like precision,then using an acid on an alkali spill (or vice-versa) can do more harm than good, and wouldbe extremely dangerous as a first-aid method.

The best thing to do is use an Amphiproticsubstance such as bicarbonate ion (e.g. sodiumbicarbonate, NaHCO3).

As shown by the reactions above, it can reactand neutralise either an acid or a base, and oncethe neutralisation is achieved, it stops workingand poses no further threat in its own 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

“co

nnm

an21

”BuffersA “Buffer”, or “buffered solution”, is a solution whichcan absorb significant amounts of acid or base withminimal change in pH.

Biology students will be aware that in all living thingsit is vital that the conditions within the body/cells arekept very constant. Living things cannot functionproperly if their “internal environment” undergoeslarge changes in• temperature, • water content,• salt concentration, (and many other chemicals)AND• pH level.

The fluids inside all living things are buffered, so thatpH remains remarkably constant, despite changes inthe external environment, eating acidic food,breathing, excreting, etc, all of which could alter thebody’s pH.

How do Buffers Work?All buffers are solutions containing a weak acid andits conjugate base.

CH3COOH(aq)+ H2O(l) H3O+

(aq)+ CH3COO-(aq)

weak acid conjugate base

This equilibrium constitutes a buffer solution, solong as there are significant amounts of theethanoate ion (CH3COO-) present. (e.g. by addingCH3COONa)

If acid is added (i.e. [H3O+] increases) (pH lower)then (by Le Chatelier’s Principle) the equilibriumshifts left, absorbing H3O+ and lowering [H3O+]again.

If base is added (ie [H3O+] decreases) (pH higher)then (by Le Chatelier’s Principle) the equilibriumshifts right, making H3O+ and raising [H3O+] again.

This way, the pH remains quite constant, despiteacid or base being added.

Lavoisier concluded that acids mustcontain a).......................... Later, Davydescribed acids as all containingb)............................., and being able toreact with c).............................................

The first attempt at a complete acid-base theory was made byd)........................................ in 1885.According to this theory, all acidsproduce e)..................................... insolution. Bases were defined ascompounds containing f)........................or .......................... ions. This theory wassuccessful at explaining some acid-base behaviour, but failed to take intoaccount the important role ofg)...........................................

The h).....................-....................... (B-L)Theory is much more useful forexplaining things. It defines an acid as ai)................ ........................., and a baseas a j).......................... ............................,and all acid-base reactions involve thek)....................... of one or morel)..............................

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

Water can acts as either o)....................or ..................., depending on thechemical environment. The word for thisis “p)........................................”

Many salts dissolve in water to formacidic or basic solutions, because oftheir interaction with q).........................An example of an acidic salt isr)........................................ An example ofa basic salt is s)........................................An example of a neutral salt ist)..............................

The “rules” governing this are:• Neutral salts form from theneutralisation of u)................. acids byv)......................... bases.• Acidic salts form from theneutralisation of w)..................... acidsby x).......................... bases.• Basic salts form from theneutralisation of y)..................... acids byz).......................... bases.

All neutralisation reactions (in aqueoussolution) involve the reaction ofaa).......................... and .........................ions to form ab)............................ Thereaction is always ac)..............-thermic.

ad).................................... is a techniqueused for chemical analysis. A measuredae)...................... of “unknown” solutionis reacted with a “af).......................solution” until the “ag)...........................point” is reached. The main piece ofequipment involved is aah).............................. It is important tochoose the ai)........................ which willchange aj).................. at the pH of theend-point, depending on the nature ofthe salt formed.

If the ak)...................... is measuredduring titration and graphed, the graphsgenerally show an “S” shaped curve.The end-point is located in the middle ofthe al)......................... part of the curve.

Buffers are solutions containing anequilibrium mixture of a am).......................... and its .........................................If either acid or base is added, theequilibrium an)................. (according toao)...................................... Principle) sothat the change in ap)....... is minimized.Buffers are important in allaq).......................... things, by helping tomaintain ar)................ ...............................A specific example is as).........................,where the pH is maintained by a mixtureof at)...................... and .................... ions.

27




WHEN COMPLETED, WORKSHEETSBECOME SECTION SUMMARIES

Worksheet 11 Acid-Base Theory & Titration Fill in the blank spaces Student Name..........................................

1. 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-

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)

28




Worksheet 12 Reaction of Acids & Bases With Water Practice Problems Student Name..........................................

Worksheet 13 Amphiprotic Ions, Acidic & Basic SaltsPractice Problems Student Name..........................................

For each substance below, write an equationdescribing its monoprotic reaction with water.(In brackets is a description of how thesubstance behaves 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, HPO42-, (acid)

d) hydrogen phosphate ion, HPO42-, (base)

2. Acidic and Basic SaltsEach of the following salts dissolves in water,and reacts (as shown in brackets) to form anacidic or basic 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)

1.25.00mL of an “unknown” NaOH solution was

titrated against standardised HCl, concentration= 0.09255 molL-1. The titration was carried out 4times, with the end-point titres being 22.50mL,22.45mL, 23.10mL and 22.50mL.

Average these results appropriately, then writean equation, and calculate the concentration ofthe NaOH solution.

2.Using a 0.05025molL-1 standardised solution of

NH4OH (ammonia solution), 25.00mL samples ofan unknown H2SO4 solution were titrated. Theaverage titre was 28.32mL

i) Write a balanced equation for theneutralisation.

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.

3.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 weighed out?

iii) Summarise the main steps involved inpreparing the solution.

4.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)

5.To find the concentration of an ammoniasolution (NH4OH), a student has 2 choices ofstandardised 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.

29




Worksheet 14 TitrationPractice Problems Student Name..........................................

Multiple Choice1. The hydrogen phosphate ion, HPO4

2- is anamphiprotic species. If it were to act as a base,then its conjugate acid would be:A. H2PO4

- B. H3PO4 C. PO43- D. HPO3

3-

2. An ionic “salt” is found to be acidic in watersolution. It is likely that this salt is the product ofthe reaction between: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.

3. The salt mentioned in Q2 was formed during atitration. The most appropriate indicator for thetitration would be:A. bromothymol blue.B. methyl orange.C. phenolphthalein.D. universal indicator.

4. Which of the following graphs might be the“titration curve” for the titration described in Q2?

A. B. C. D.18.

5.An appropriate material to use in case of an acidspill is:A. an amphiprotic substance,

e.g. sodium bicarbonate.B. a strong base, like sodium hydroxide.C. a buffer solution, like methanoic/methanoate

mixture.D. a weak base, like ammonia.

6.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.


7. (7 marks)a) Give the definition of an acid according to theArrhenius Theory.b) Write an equation which Arrhenius mighthave used to explain why hydrogen chloride gasis an acid when dissolved in water.c) Give the definitions for acid and baseaccording to the Bronsted-Lowry Theory.d) Write an equation to show why hydrogenchloride gas in water is an acid according to theB-L Theory.

8. (7 marks)a) A solution containing carbonate ions (CO3

2- )is found to be quite strongly basic.

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

b) Write TWO different equations to show theamphiprotic nature of the hydrogen carbonateion, HCO3

-.

9. (8 marks)A diluted vinegar sample (CH3COOH solution)was analysed by titration with a standardisedsolution of KOH with concentration of 0.008263 molL-1.

9.00mL samples of the vinegar solution wereused. The titration was done 4 times, givingburette titres of 27.35, 26.70, 26.75 and 26.65mLof KOH.

a) What is an appropriate indicator for thistitration? Justify your choice.

b) Write a balanced symbol equation for thereaction.

c) Explain how the titration measurementsshould be used.

d) Calculate the concentration of the dilutedvinegar solution.

9. (4 marks)a) What are the characteristics of a “buffersolution”?

b) In general terms, what are the usualingredients of a buffer solution? Give a specificexample.

c) Give an example of a natural buffer system.

30





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

Alkanols and Alkanoic AcidsYou were introduced to the alkanols (alcohols)in the previous topic, and to some of thealkanoic acids earlier in this 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 knowthese 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 apolar bond, and how this causes the propertiesof the alkanols to be quite different to those ofthe corresponding alkanes.

Polarity of the AlkanolFunctional Group

Now, compare the -COOH Functional Group ofthe Alkanoic Acids:

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

In an alkanol, the dipoles result in hydrogenbonding between molecules. For example, inethanol:

Molecules attractedby their dipole chargesand form hydrogen bonds.

In an alkanoic acid, such as ethanoic acid:

Each pair of molecules has 2 sets ofdipoles.

You already know that the alkanols have muchhigher m.p’s & b.p’s than the correspondingalkanes, because of the hydrogen bondingbetween molecules.

The m.p’s & b.p’s of the alkanoic acids arehigher still, due to the presence of extra dipole-dipole attractions.

For comparison:

Ethane Ethanol Ethanoic acidBoiling Pt -89 +78 +118(oC)

31

5. ESTERIFICATIONkeep it simple science

®



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 mmoossttiimmppoorrttaanntt

mmeemmbbeerr ooff bbootthhHHoommoollooggoouuss

SSeerriieess 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

δδ++

δδ++

δδ++

δδ++

δδ++

δδ-

δδ-δδ-

δδ-

δδ-

OO-HHCH

33 CH22

δδ++

δδ-

δδ++

δδ-

δδ-

δδ-

Hydrogeen bond

DDiippoollee-DDiippoolleeaattttrraaccttiioonnss

EEtthhaannooll MMooddeell EEtthhaannooiicc AAcciiddMMooddeell

32

The reaction could be described as a“condensation” because it produces a watermolecule, but it is so widespread in nature, andso important, that it rates its own name;“Esterification”.

The reaction can be visualised 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, with hydrocarbon chains on either side.

This group of atoms is polar, but both ends ofthe molecule are non-polar. Esters generallyhave low solublity in water (some exceptions),and are volatile with a strong odour... oftensweet and fruity.

As you’ll learn, this is their “job”; to give thesmells and flavours to many foods andperfumes.

Naming EstersSince the esters are made by joining together 2other molecules, it should be no surprise thatthey have a 2-part name.

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 fromthis 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((RR1 rreepprreesseennttss tthheerreesstt ooff tthhee mmoolleeccuullee))

Alkanol((RR2 rreepprreesseennttss tthheerreesstt ooff tthhee mmoolleeccuullee))

TThheessee aattoommss ffoorrmmwwaatteerr,, aanndd tthhee 22mmoolleeccuulleess jjooiinn

O

OR1 C

Water+

Ester

H OH

OH

H

C C

H HHO H

HH

H

H

H

H

C C C C

RReemmnnaanntt ooff CCHH3CCOOOOHH RReemmnnaanntt ooff CC4HH9OOHH

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 is

largely due to natural esters

The Esters“Esters” are a group of carbon compounds formed by the reaction between

an alkanol and an alkanoic acid.


33

Practical Work: Making an Ester

You may have made an ester in thelaboratory using the process of Reflux.

A simple laboratoryreflux set-up isshown.

The refluxcondenseris open to theatmosphere at thetop.

This is vital to avoidany dangerousbuild-up of pressurewhich would occurin a sealed flask.

Volatile chemicalvapours rise, but arecondensed and dripback into the flask... “reflux”.

The reaction flask is heated to speedup the otherwise very slow reaction.

As well as an alkanoic acid and analkanol, it is usual to add a catalyst,concentrated sulfuric acid, H2SO4.

alkanoic acid + alkanol ester + water

As well as speeding the reaction up, theH2SO4 catalyst also absorbs the waterproduct. This has the effect of shiftingthe equilibrium to the right (accordingto Le Chatelier’s Principle) andincreasing the yield of the ester.

Occurrence of Esters

Esters occur widely in nature, especiallyin fruits and flowers. Esters are largelyresponsible for the smell and taste ofmany foods.

There may be a complex mixture ofesters and other compounds which givethe complete smell of (say) a ripestrawberry, but there’s always an estergiving the main smell and tastesensation.

Some examples:Strawberry ethyl butanoate, C3H7COOC2H5

Orange octyl ethanoate, CH3COOC8H17

Fats & Oils are Esters, Too

Esters made from very long chain “fattyacids”, and the triple-alcohol molecule“glycerol” are used by all living thingsas high-energy foods and energystorage chemicals. We call them fats oroils, depending on their melting point.

Production and Uses of EstersAs well as their wide occurrence inliving things, artificially-manufacturedesters are important industrialchemicals.

They are produced by exactly the sameprocess shown at the left... reflux of theappropriate alcohol, acid and catalyst...but on an industrial scale, of course.

Uses include:

Artificial flavours for drinks and variousprocessed foods.

Solvents. Ethyl ethanoate is widely used as anindustrial solvent (e.g. plastics industry) and iswell known as the solvent for nail varnish.

Ingredients in many products, includingshampoo and cosmetic products, and as“plasticizers” (softening agents) in someplastics, such as “vinyl”.




EElleeccttrriiccaall HHeeaattiinngg EElleemmeenntt

ccoonncc.. HH2SSOO4


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

The Alkanoic Acids contain the functional groupg)...................... This group contains 2 polarbonds, so 2 h)............................... bonds canform between molecules. This is whyi)................................................. than the alkanols.

Esters are formed by the reaction ofj)........................... with ...........................................The other product is k).................................Esters are named by the l)........................... first(with its ending changed to m)..........), followedby the n)...................................... name (with itsending changed to o)..........................)

c) i)ii)iii)

d) i)ii)iii)

e) i)ii)iii)

f) i)ii)iii)

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)3CH3

b) CH3CH2COO(CH2)3CH3

c) CH3(CH2)3COO(CH2)4CH3

d) C4H9COOCH3

e) C5H11COOC7H15

34



Worksheet 16 Esters Fill in the blank spaces Student Name..........................................

Esters are made using the technique ofp)............................... The reaction flask is opento the atmosphere to avoid any dangerous build-up of q)................................... Volatile chemicalsvapourise, but are r)..................................... by thes)................................................ and drip backinto the flask. t).................................... is used asa catalyst, and also improves the yield byshifting the u)............................. because itabsorbs water.

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

Artificially manufactured esters are used asy)........................ .................................. inprocessed foods, as z).......................... inindustry and as ingredients in many productssuch as aa)............................. and .........................

Worksheet 17 EstersPractice Problems Student Name..........................................

1. Names of EstersName the ester formed from

a) ethanol & propanoic acid

b) propanol & ethanoic acid

c) pentanol & methanoic acid

d) methanol and pentanoic acid

e) hexanoic acid and butanol

f) 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, itmay be more convenient to place the acid remnant

first in the structural formula. This system is used here.

Multiple Choice1.

The ester shown is:

A. butyl propanoateB. ethyl butanoateC. butyl ethanoateD. propyl butanoate

2.Of these compounds, which would you expectto have the highest m.p. & b.p.?

A. ethane B. ethene C. ethanol D. ethanoic acid


3. (6 marks)a) Draw structural formulas for

i) methanol

ii) 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 tocarry out this ester preparation.

c) Name the chemical you would add to thereaction flask as a catalyst.

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

35





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36



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

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

THE ACIDIC

ENVIRONMENT


37

Answer SectionWorksheet 1a) opposites b) neutralisec) pH d) 7e) basic f) acidicg) change colour h) pH (acidity)i) plants / lichensj) pink k) bluel) clear, clear, pink/red m) red, yellow, yellown) yellow o) bluep) soil testing q) water testingr) effluents

Worksheet 21. A = acid. B = base. C = neutral. D = base. E = acid.2. B3.a) A chemical which changes colour depending onthe pH of the solution it is in.b) Soil testing, to help gardening or agriculture.c) Put Hydrangea flowers through a blender with asmall amount of water and ethanol. Filter the mixture.The liquid will act as an acid-base indicator.

Worksheet 3a) water b) saltc) neutralise d) basese) water and a salt f) acidsg) acid solution h) neutralisei) equilibrium j) ratek) products l) dynamicm) temperature, concentration or gas pressuren) Le Chatelier’so) counteracts the disturbancep) carbon-oxygen q) burning fossil fuelsr) Global s) Greenhouse Effectt) volcanic u) hot springs/geysersv) fossil fuels w) smeltingx) Acid Rain y) NOxz) lightningaa) power stations & enginesab) smog ac) Rain

Worksheet 41.a) i) calcium chloride

ii) magnesium sulfateiii) barium nitrate

b) i) HCl + LiOH H2O + LiCl

ii) H2SO4 + 2NaOH 2H2O + Na2SO4

iii) 2HNO3 + Mg(OH)2 2H2O + Mg(NO3)2

2.a) H2SO4 + FeO H2O + FeSO4

b) 2HCl + MgO H2O + MgCl2

c) 2HNO3 + CuO 2H2O + Cu(NO3)2



3. a) CO2 + Ca(OH)2 H2O + CaCO3

b) P2O5 + 3H2O 2H3PO4

c) SO3 + H2O H2SO4 (sulfuric acid)

Worksheet 51. i) Left, to try to use heat & reduce temp. ii) Right, to reduce the pressure again.iii) Left, to decrease concentration of product.iv) Right, to release heat & increase temp.v) Left, to make more reactant & increase pressure.vi) Left, to increase pressure again.2.i) Equilib. shifts left, to decrease conc. of iodide ion.ii) Equilib. shifts right, to increase conc. of hydrogenion.iii) Equilib. shifts right, to decrease conc. of HI.iv) Endothermic (heat is a reactant), since higher tempcauses shift to right to consume heat.3.i) High pressure causes equilibrium to shift right toreduce the total moles of gas and pressure.ii) Shifts equilibrium to left.iii) Removing product & adding reactant keepsshifting equilib. to right, so yield is maximised.

Worksheet 61. Molar Gas volumesa) In each case, multiply moles by 24.79 (at 25oC &100kPa) or 22.71 (at 0oC & 100kPa)

i) 64.2L ii) 1.03L iii) 2.72x103L iv) 1.16L

b) In each case, divide volume by 24.79 or 22.71.i) 0.500 mol ii) 1.10x10-2mol iii) 403 mol iv) 4.41x10-5mol

2. Mass-Vol of Gasesa) i) n(CO2) = 5.00/24.79 = 0.2017 mol

m(CO2) = n x MM = 0.2017 x 44.01 = 8.88g.ii) n(H2) = 5.00/24.79 = 0.2017 mol

m(H2) = n x MM = 0.2017 x 2.016 = 0.407g.iii) n(Ne) = 100/22.71 = 4.403 mol

m(Ne) = n x MM = 4.403 x 20.18 = 88.9g.iv) n(O2) = 0.0250/22.71 = 0.001101 mol

m(O2) = n x MM = 0.001101 x 32 = 3.52x10-2g.b) i) n = m/MM = 100/44.01 = 2.272 mol

v = 2.272 x 24.79 = 56.3Lii) n = m/MM = 100/4.003 = 24.98 mol

v = 24.98 x 24.79 = 619Liii) n = m/MM = 1.50/28.02 = 0.05353 mol

v = 0.05353 x 24.79 = 1.33L

3. Problemsa) i) CO2(g) + Ca(OH)2 (aq) H2O(l) + CaCO3(s)

ii) n(CO2) = 1.00/24.79 = 0.04034 mol∴∴ n(CaCO3) = 0.04034 mol (ratio 1:1)∴∴ m(CaCO3) = n x MM = 0.04034x100.09 = 4.04g.

iii) n(CaCO3) = m/MM = 1.75/100.09 = 0.01748 mol∴∴ n(CO2) = 0.01748 mol (ratio 1:1)

∴∴ vol(CO2) = 0.01748 x 22.71 = 0.397 L (397mL)

38

Worksheet 6 (cont.)3.b) i) 2ZnS(s) + 3O2(g) 2ZnO(s) + 2SO2(g)

ii) n(ZnS) = m/MM = 1.00x106/97.46 = 10,260 mol.∴∴ n(O2) = 10,260 x 3/2 = 15,391 mol (ratio = 2:3)∴∴ vol(O2) = 15,391 x 24.79 = 3.82 x 105 L

iii) vol(air) = 3.82 x 105 x 100/21= 1.82 x106 L.

iv) Gay-Lussac’s Law is that gases react in simplewhole number ratios. Therefore the volume of SO2will be in the same ratio as the mole ratio of O2 : SO2which is 3:2.

∴∴ vol(SO2) = 3.82 x105 x 2/3 = 2.55 x 105 L

Worksheet 71. A 2. A 3. C 4. C 5. B 6. D

7.a) 2HCl(aq) + MgO(s) H2O(l) + MgCl2(aq)

b) It has neutralised an acid, forming water and a salt.Therefore it is a base.

c) CO2(g) + 2NaOH(aq) 2H2O(l) + Na2CO3(aq)Carbon dioxide has neutralised a base, therefore it isacidic.

8.a) Carbon dioxide reacts with water as follows:

In the sealed bottle, this system is in equilibrium.When the lid is removed, the pressure of CO2 abovethe liquid drops. The equilibrium shifts left(attempting to increase the pressure by making moregas) so bubbles form as CO2 comes out of solution.

b) If exothermic, then heat is a “product” of thereaction as written above. By Le Chatelier’s Principle,higher temperature should shift equilibrium left.Therefore, a warm bottle of drink will form bubblesfaster when opened.

9.a) volcanic eruptions, hot springs, geysers.

b) example: smelting of lead sulfide ore:PbS + O2 Pb + SO2

c) “Acid Rain” can acidify lakes and streams, killingthe living things. Forests can die-back due to soilacidity.

d) S + O2 SO2

n(S) = m/MM = 1.00x106/32.07= 31,182 mol

V(SO2) = 31,182 x 24.79 = 7.73 x 105 L(773,000 Litres!)

Worksheet 8a) donate protons b) hydrogenc) hydronium d) monoprotice) diprotic f) triproticg) CH3COOHh) 2-hydroxypropane-1,2,3-tricarboxylic acidi) ionises completely in water solutionj) only partially ionises k) -log10[H3O+]l) powers (index numbers) m) 10n) preserve o) bacteria & fungip) SO2 and ethanoic q) flavourr) sour s) ethanoic & citrict) nutritional u) ascorbicv) formic (methanoic) w) calcium oxide

Worksheet 9





1. Acid Ionisation in Watera) HCl(g) + H2O(l) H3O+

(aq) + Cl-(aq)

b) HBr(g) + H2O(l) H3O+(aq) + Br-

(aq)

c) HCOOH(l) + H2O(l) H3O+(aq) + HCOO-

(aq)

d) HCN(g) + H2O(l) H3O+(aq) + CN-

(aq)

e) H2SO4(l)+ 2H2O(l) 2H3O+(aq)+ SO4

2-(aq)

f) H2CO3(l)+ 2H2O(l) 2H3O+(aq)+ CO3

2-(aq)

g) H3PO4(l)+ 3H2O(l) 3H3O+(aq)+ PO4

3-(aq)

h) C6H8O7(s)+ 3H2O(l) 3H3O+(aq) + C6H5O7(aq)

2. pH from [H3O+]use pH = -log[H3O+] in each casea) 1.17 (acidic) b) 3.08 (acidic)c) -0.398 (acidic) d) 11.6 (basic)e) 2.46 (acidic) f) 7.82 (just barely basic)

3. pH from Acid Concentrationa) [H3O+] = 0.250, so pH = 0.602b) [H3O+] = 0.0750 x 2, so pH = 0.824c) [H3O+] = 7.50x10-4 x 2, so pH = 2.82d) [H3O+] = 4.5x10-3 x 3, so pH = 1.87e) [H3O+] = 6.00 x 2, so pH = -1.08

4. [H3O+] from pHuse inverse (or 2nd function) log (-pH) on calculatora) 5.01 x 10-6 molL-1. b) 3.16 x 10-12 molL-1.c) 2.51 molL-1. d) 3.00 x 10-9 molL-1.e) 1.00 molL-1.

Worksheet 101. C 2. D 3. C 4. B

5.a) H2A + 2H2O(l) 2H3O+

(aq) + A2-(aq)

acid base conj.acid conj.base

b) if [H2A] = 0.0250, then [H3O+]= 0.0500 moL-1

pH = -log10[H3O+]= 1.30

c) H2A must be a weak acid and has only partiallyionised. The [H3O+] is lower and pH higher thanpredicted.


39

Worksheet 131. Amphiprotic Substancesa) H2PO4

-(aq) + H3O+

(aq) H3PO4(aq) + H2O(l)

H2PO4-(aq) + OH-

(aq) HPO42-

(aq) + H2O(l)

b) HCO3-(aq) + H3O+

(aq) H2CO3(aq) + H2O(l)

HCO3-(aq) + OH-

(aq) CO32-

(aq) + H2O(l)

c) HS-(aq) + H3O+

(aq) H2S(aq) + H2O(l)

HS-(aq) + OH-

(aq) S2-(aq) + H2O(l)

2. Acidic & Basic Saltsa) CH3COO-

(aq)+ H2O(l) OH-(aq)+ CH3COOH(aq)

b) NH4+

(aq) + H2O(l) NH3(aq) + H3O+(aq)

c) NO2-(aq) + H2O(l) HNO2(aq) + OH-

(aq)

d) HC2O4-(aq) + H2O(l) C2O4

2-(aq) + H3O+

(aq)

e) CN-(aq) + H2O(l) HCN(aq) + OH-

(aq)

Worksheet 141.

23.10 omittedAverage titre = (22.50+22.45+22.50)/3 = 22.48mL

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

Ca x Va = Cb x Vba b

Cb = b x Ca x Va/(Vb x a)= 1 x 0.09255 x 22.48/25.00 x1= 0.08322

∴∴ c(NaOH) = 0.08322 molL-1.

2.i) H2SO4 + 2NH4OH 2H2O + (NH4)2SO4

ii) Ca x Va = Cb x Vba b

Ca = a x Cb x Vb/(Va x b)= 1 x 0.05025 x 28.32/(25.00 x 2)= 0.02846

∴∴ c(H2SO4) = 0.08322 molL-1.iii) L is best choice because stong acid-weak basetitration has end-point at acidic pH. L changes colourat 4.2.3.i) Can be obtained in a pure, dry state. Is stable and

does not react with gases in air or absorb moisture.ii) moles required:

n = C x V = 0.02500 x 0.5 = 0.01250 mol.mass required: m = n x MM (MM= 90.04g)

= 0.01250 x 90.04mass = 1.126g.

iii) Weigh out chemical into clean, dry beaker.Add enough pure water, and stir, to dissolve itcompletely.Transfer solution into 500mL volumetric flask.Rinse beaker with small amounts of extra water andadd washings to flask.Fill flask to mark with pure water. Use a dropper at theend to fill exactly to the mark.Insert stopper and invert repeatedly to mix solutionthoroughly.



Worksheet 10 (cont.)6.Acids may be added to preserve the food. By loweringthe pH, it becomes more difficult for microbes togrow in the food and cause it to spoil. Ethanoic acid(vinegar) is commonly used for this.

Acids are also added to flavour the food by adding thesourness that improves some flavours such as fruitdrinks or jams. Citric acid is often used this way.

Worksheet 11a) oxygen b) hydrogenc) metals d) Arrheniuse) hydrogen ions f) hydroxide or oxideg) the solvent (water) h) Bronsted-Lowryi) proton donor j) proton acceptork) transfer l) protonsm) conjugate n) conjugate acido) acid or base p) amphiproticq) water r) any ammonium salts) sodium carbonate t) sodium chloride/sulfateu) strong v) strongw) strong x) weaky) weak z) strongaa) hydronium & hydroxide ab) waterac) exo- ad) Titrationae) volume af) standardag) equivalence (end) ah) buretteai) indicator aj) colourak) pH al) near-verticalam) weak acid and its conjugate basean) shifts ao) Le Chatelier’sap) pH aq) livingar) constant chemical conditionsas) in the blood at) bicarbonate & carbonate

Worksheet 12In each case the conjugate is underlined.

a) CH3COOH(aq)+ H2O(l) H3O+

(aq)+ CH3COO-(aq)

b) NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

c) HPO42-

(aq) + H2O(l) PO43-

(aq) + H3O+(aq)

d) HPO42-

(aq) + H2O(l) H2PO4-(aq) + OH-

(aq)

e) S2-(aq) + H2O(l) HS-

(aq) + OH-(aq)

f) CN-(aq) + H2O(l) HCN(aq) + OH-

(aq)

g) H2S(aq) + H2O(l) HS-(aq) + H3O+

(aq)

h) NO2-(aq) + H2O(l) HNO2(aq) + OH-

(aq)

i) NH4+

(aq) + H2O(l) NH3(aq) + H3O+(aq)

j) HSO3-(aq) + H2O(l) SO3

2-(aq) + H3O+

(aq)


40

Worksheet 14 (cont.)4.C2H2O4 + 2KOH 2H2O + C2O4K2

(or) COOHCOOH + 2KOH 2H2O + COOKCOOK

Ca x Va = Cb x Vba b


∴∴ c(KOH) = 0.06290 molL-1.

5.i) use nitric acid because it is strong acid. Weak acid-weak base titrations have indistinct end-point... bestavoided.ii) HNO3 + NH4OH H2O + NH4NO3


∴∴ c(NH4OH) = 0.9491 molL-1.

Worksheet 151. A 2. D 3. B 4. C 5. A6.a) An acid produces hydrogen ions in solution.b) HCl(g) H+

(aq) + Cl-(aq)

c) Acids are proton donors. Bases are protonacceptors.d) HCl(g) + H2O(l) Cl-(aq) + H3O+

(aq)

The HCl molecule has transferred a proton to thewater molecule, therefore it is an acid.

7.a) CO3

2-(aq) + H2O(l) HCO3

-(aq) + OH-

(aq)Water acts as an acid and donates a proton to thecarbonate ion. This forms a hydroxide ion, whichexplains why the solution is basic.

b) 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)

8.a) A weak acid-strong base titration has an end pointabout pH = 8-10, so phenolphthalein is best.

b) CH3COOH + KOH H2O + CH3COOK

c) The first titre should be discarded because it doesnot agree closely with the others. The remaining 3should be averaged. Average = (26.70+26.75+26.65)/3 = 26.70mL

d) Ca x Va = Cb x Vba b

Ca = a x Cb x Vb/(Va x b)= 1 x 0.008263 x 26.70/(25.00 x 1)= 0.008825

∴∴ c(CH3COOH) = 0.08322 molL-1.



9.a) A buffer can maintain a constant pH despiteaddition of acid or base.b) A (roughly equal) mixture of a weak acid and itsconjugate base, such as ethanoic acid plus sodiumethanoate (which provides ethanoate ions).c) Our blood is buffered by a mixture of bicarbonateions and carbonate ions. The blood pH remains quiteconstant, despite constant changes occurring asgases dissolve, food is absorbed, etc.

Worksheet 16a) alcohols b) -OHc) CnH2n+1OH d) polare) hydrogen f) alkanesg) COOH h) hydrogeni) m.p. & b.p. is even higher thanj) alkanols with alkanoic acidsk) water l) alkanolm) -yl n) alkanoic acido) -oate p) refluxq) pressure r) condenseds) reflux condenser t) Sulfuric acidu) equilibrium v) odours and tastesw) fruits x) fats and oilsy) artificial flavourings z) solventsaa) shampoo/cosmetics/plastics

Worksheet 171. Names of Estersa) ethyl propanoate b) propyl ethanoatec) pentyl methanoate d) methyl pentanoatee) butyl hexanoate f) octyl ethanoate

2. Condensed Structural Formulasa) i) CH3CH2OH ii) CH3CH2COOH

iii) CH3CH2COOCH2CH3b) i) CH3CH2CH2OH ii) CH3COOH

iii) CH3COO(CH2)2CH3c) i) CH3(CH2)3CH2OH ii) HCOOH

iii) HCOO(CH2)4CH3d) i) CH3OH ii) CH3(CH2)3COOH

iii) CH3(CH2)3COOCH3e) i) CH3(CH2)2CH2OH ii) CH3(CH2)4COOH

iii) CH3(CH2)4COO(CH2)3CH3f)i) CH3(CH2)6CH2OH ii) CH3COOH

iii) CH3COO(CH2)7CH3

3. Names from Structuresa) butyl methanoate. butanol + methanoic acidb) butyl propanoate. butanol + propanoic acidc) pentyl penanoate. pentanol + pentanoic acidd) methyl pentanoate. methanol + pentanoic acide) heptyl hexanoate. heptanol + hexanoic acid.

FOR MAXIMUM MARKS SHOWFORMULAS & WORKING,

APPROPRIATE PRECISION & UNITS IN ALL CHEMICAL PROBLEMS


41

Worksheet 181. C 2. D

3.

b) The mixture needs heating, but a closed flaskcould explode due to pressure build up. An open flaskis needed, but volatile chemicals will then evaporateaway. The reflux system is open, but condensesvapours and returns them to the reaction flask.

c) concentrated sulfuric acid.

d) Esters generally have sweet, fruity odours.



O-HH

HH

H

C

H

H

H

C

a) i) ii)

ii) methyl propanoate

CH3CH2COOCH3

H H O-HH

OH

H

H

C C C

H H O

OH

H

H

C C C

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PO Box 2575PORT MACQUARIE NSW 2444

(02) 6583 4333 FAX (02) 6583 9467www.keepitsimplescience.com.au [email protected]

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