Module 2- The Acidic Environment

HSC Chemistry Summary

Module 2- The Acidic Environment

1

Robert Lee Chin

Hy

dro

gen

sulfid

e

hy

dro

cya

nic

acetic

Form

ic (Meth

an

oic)

ph

osp

ho

ric

carb

on

ic

nitrio

us

nitirc

sulfu

rou

s

sulfu

ric

hy

dro

iod

ic

hy

dro

bro

mic

hy

dro

chlo

ric

Hy

dro

fluoric

Acid

H2 S

HC

N

CH

3C

OO

H

HC

OO

H

H3 P

O4

H2 C

O3

HN

O2

HN

O3

H2 S

O3

H2 S

O4

HI

HB

r

HC

l

HF

S2

-

CN

-

CH

3C

OO

-

HC

OO

-_

PO

43-

CO

32-

NO

2-

NO

3-

SO

32-

SO

42-

I-

Br

-

Cl -

F-

An

ion

sulfid

e

cya

nid

e

aceta

te

Fo

rma

te

(meth

an

oa

te)

Ph

osp

ha

te

Ca

rbo

na

te

Nitrite

Nitra

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Su

lfite

Su

lfate

Iod

ide

Bro

mid

e

Ch

lorid

e

Flu

orid

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Zn

S

KC

N

Ag

(CH

3 CO

O)

Mg

(HC

OO

)2

Na

3 PO

4

Ca

CO

3

Na

NO

2

PB

(NO

3 )2

Na

2 SO

3

K2 S

O4

Ag

I

KB

r

Na

Cl

Ca

F2

Ty

pica

l salt



2

Robert Lee Chin

The Acidic Environment: 1. Indicators

Classify common substances as acidic, basic or neutral

Acids

Acids are substances capable of providing hydrogen ions (H+) for chemical reactions.

Free ions are only available in solutions where the proton is stabilised by a solvent

molecule. In an aqueous solution it exists as the hydronium ion:

OHOHH 32

Properties of acids:

-Sour taste

-Sting or burn skin

-Conduct electricity as an aqueous solution

-Turn blue litmus red

-React with active metals to produce a salt and hydrogen gas e.g.

4(aq)2(g)4(aq)2(s) NaSOHSOHNa

-React with many carbonates to produce salt, water and CO2 e.g.

Cl2NHCOOH2HClCO)(NH 42(g)(l)2(aq)3(s)24

-React with bases, neutralising to form water and a salt e.g.

3(l)2(aq)3(aq) NaNOOHNaOHHNO

Common

Substance

Name of acid Chemical

Formula

Uses

Cream of tartar Tartaric KHC4H4O6 Whipping eggwhites

Lemon juice Citric/2-hydroxylpropane-

1,2,3-tricarboxylic acid

C6H8O7 Flavour for food

Vinegar Acetic/ethanoic C2H4O2 Preservative, flavouring

food Fizzy drink Carbonic H2CO3 Fizzy taste

Aspirin Acetylsalicylic C9H8O4 Pain relief medicine Car battery acid sulphuric H2SO4 Car battery

Vitamin-C tablets ascorbic C6H8O6 Dietary supplement Yogurt lactic C3H6O3 Detergents, Biopolymer

precursor Wine, bananas Tartaric C4H6O6 Food flavouring

Bases

Bases are substances that react with acids to form salts or form the hydroxide ion

(OH-) in solution. A soluble base is called an alkali. Metal oxides act as bases when in

solution e.g.:

2OHCaOHCaO (aq)2

(l)2(s)

Properties of alkalis:

-bitter taste

-Soapy, slippery feel

-Conduct electricity as an aqueous solution



3

Robert Lee Chin

-Turn red litmus blue

Common

Substance

Name of base Chemical Formula Uses

Ammonia Ammonia NH3 Cleaning agent, insect

stings

Hand soap - - Cleaning Agent

Detergent - - Cleaning agent

Antacid Magnesium/Aluminium

hydroxide

Mg(OH)2

Al(OH)3

Relieve Indigestion

Bicarbonate of

soda

Sodium hydrogen

carbonate

NaHCO3 Used in baking

Lye water Sodium hydroxide NaOH Additive in some foods

cleaning agent

Neutral substances Neutral substances are neither acidic nor basic. Examples are pure water, pure alcohol

and sugar. The salts formed in neutralisation acids are neutral as are some oxides.

Common

Substance

Name of substance Chemical Formula Uses

Pure water dihydrogen oxide H2O Essential for life

Table salt Sodium chloride NaCl Food additive, preservative

sugar Sucrose C12H22O11 Food ingredient and

preservative

Pure alcohol Ethanol C2H5OH Cleaning agent, preservative

Perform a first-hand investigation to prepare and test a natural indicator

Experiment: Extracting and using a natural indicator

Aim: To prepare an indicator solution from red cabbage and test the resulting

indicator on a range of substances

Equipment:

2-3 large red cabbage leaves, shredded

500 mL beaker

250 mL beaker

Tripod, gauze mat and Bunsen burner

Test tubes & and test tube rack

Universal indicator (optional)

Approx, 2mL solution of each:

0.1 mol L-1

NaOH

0.1 molL-1

HCl

white vinegar

household ammonia

lemon juice

lemonade

bicarbonate of soda

washing powder

antacid tablet (grind into powder)

salt water



4

Robert Lee Chin

Method:

1/ Place shredded cabbage leaves in 500 Ml beaker and just cover with distilled

water (about 200 mL). Slowly boil the cabbage leaves until the water turns a

dark reddish-purple and the leaves lose most of their colour.

2/ Allow to cool and pour the liquid off into a clean 250Ml beaker. This is the

red cabbage indicator.

3/ Place 2 mL of NaOH and HCl into separate test tubes. Add a few drops of red

cabbage indicator until a definite colour is observed. Record the colour of the

indicator

4/ Repeat step 3 with the other substances and record results. Classify the

substances as acidic, neutral or basic.

5/ Optional: Test each of the solutions with universal indicator to check your

classification.

Results:

Substance Red cabbage

indicator colour

Acidic/basic/neutral Universal

indicator colour

NaOH(aq) yellow neutral Purple

HCl(aq) Red Acidic Red

white vinegar Pink Acidic Red

Household

ammonia

Dark green Basic Blue-green

Lemon juice Red Acidic Red

Lemonade Purple-magenta Acidic Red

Bicarbonate of

soda

Blue-green Basic Blue-green

Antacid Cloudy purple Slightly basic Lime green

Salt water Purple Neutral Dark green

As a generalisation, the red cabbage indicator turned acidic substances red and basic

substances blue. Neutral substances stayed the same colour as the red cabbage

indicator (purple)



5

Robert Lee Chin

Identify data and choose resources to gather information about the colour

changes of a range of indicators

Identify that indicators such as litmus, phenolphthalein, methyl orange and

bromothymol blue can be used to determine the acidic or basic nature of a

material over a range, and that the range is identified by change in indicator

colour

Indicator Highly

acidic

Slightly

acidic

Neutral Slightly

basic

Highly

basic

Litmus red red Reddish-

blue

blue blue

Phenolphthalein colourless colourless colourless pink red

Bromothymol

Blue

yellow yellow green blue blue

Methyl Orange red yellow yellow yellow yellow

Universal

Indicator

red Orange-

yellow

Green Blue Purple

Solve problems by applying information about the colour changes of

indicators to classify common substances as acidic, neutral or basic

Investigation: Testing the acidity of household substances

Aim: To determine the acidity/basicity of some household substances using some

indicators

Equipment:

Small test tubes and test tube rack

Beaker

Dropper bottles containing:

-Phenolphthalein

-Litmus

-Methyl orange

-Universal indicator

-Bromothymol Blue

Method:

Substances to be tested:

-distilled water

-drain cleaner

-ammonia

-vinegar

-lemonade

-baking soda

-shampoo

-conditioner

-egg white

-antacid

-lemon juice

Memory assist: Acids turn blue litmus red (Blue in acid goes Red- BAR)



6

Robert Lee Chin

1/ Each substance will be tested using 5 different indicators. Pour about 20 mL of

each of the substances into separate test tubes. For drain cleaner, dissolve

about a teaspoonful into 200 mL of distilled water, and then pour into the test

tubes.

2/ Add one drop of a different indicator to each one of the substances. Mix

thoroughly and record the observed colour. Repeat for each substance.

Results:

Substance Phenolphthalein Litmus Methyl

orange

Universal

indicator

Bromothymol

Blue

Distilled water Clear Purple Orange Yellow Blue-green Drain cleaner Pink Violet Yellow Turquoise Light blue

Ammonia Magenta Purple Orange Blue-grey Blue Vinegar Clear Pink Red Red Yellow

Lemonade Clear Red Red Red Yellow Baking Soda Magenta Blue Yellow Blue-

green

Blue

Shampoo White Pink-

purple

Yellow Pink Yellow

Conditioner White Violet Yellow Lime

green

Yellow

Egg white Magenta-pink Purple Orange Blue-

green

Blue

Antacid Pink Blue Yellow Red Blue Lemon juice Clear Red Red Red Yellow

As a generalisation (based on the results):

-Phenolphthalein reacts ONLY with basic substances, turning them pink, then red for

higher pH

-Blue Litmus turns stronger bases purple and strong acids red

-Methyl orange turns bases yellow and acids red

-Universal indicator turns bases blue and acids red

-Bromothymol remains blue in bases and yellow in acids

Identify and describe some everyday uses of indicators including the testing

of soil acidity/basicity

Water Testing

pH levels in swimming pools need to regularly tested and maintained between 7.2-

7.8. Above this will encourage growth of bacteria, mould and algae. Above 7.8 and

below 7.2 will cause irritation to skin and eyes. A pool pH kit is used to measure the

pH level. If it is too low, bicarbonate of soda is added. If too high, chlorine bleach

powder.



7

Robert Lee Chin

Fish in aquariums are sensitive to the pH. Too acidic or alkaline water will kill certain

fish.

Testing of soil pH

Many plants can only tolerate a certain pH range in the soil. For example, carnivorous

plants prefer acidic soils while beetroot thrives in slightly alkaline soil. To test the pH,

a white unreactive powder is first mixed with the soil to absorb moisture before

adding universal indicator.

Effluent Testing

pH can be used to assess the levels of certain types of industrial pollution. Indicators

are used to monitor the pH of waste water and natural waterways.

2. Acids in our Environment

Identify oxides of non-metals which act as acids and describe the conditions

under which they act as acids

We can distinguish whether an oxide is an acid or a base by observing its effects on

an indicator or seeing if it reacts with an acid or base

In general, oxides of metals act as bases; they turn litmus red. They react with water

to form an alkaline solution:

2(aq)2(s) Mg(OH)OHMgO E.g.

solution alkaline water oxide Metal

Basic oxides react with acids to form water and a salt:

2(aq)(l)2(s) Mg(Cl)OH2HClMgO E.g.

salt water acid oxide Metal

Oxides of non-metals act as acids; they turn litmus blue. They react with water to

form acids:

(aq)322(s)2 SOHOHSO E.g.

acid water oxide metal-Non

Acidic oxides react with bases to form water and a salt:

(aq)32(l)2(aq)2 SONaOH2NaOHSO E.g.

salt water acid oxide metal-Non

Basic oxides do not react with alkali solutions



8

Robert Lee Chin

Analyse the position of these non-metals in the Periodic table and outline the

relationship between position of elements in the periodic table of elements

and acidity/basicity of oxides.

In general, the oxides of elements on the LHS (metals) form basic oxides and the

oxides of elements on the RHS (non-metals) form acidic oxides. The noble gases do

not form oxides.

Define Le Chatelier’s principle

Revision of Equilibrium:

Many reactions are reversible reactions i.e. forwards and reverse reactions occur at the

same time. In an undisturbed, closed system, these reactions will eventually reach a

state of equilibrium

Features of a system at equilibrium:

1) It is a closed system- no energy or matter leaves or enters

2) Macroscopic properties (e.g. colour, temperature, state, pressure) remain

constant

3) Concentrations of products and reacts remain constant

4) Rate of → reaction = rate of ← reaction

5) Microscopic changes DO occur

6) There will ALWAYS be some product & reactant

Le Chatelier‟s principle applies to systems already in equilibrium that then undergo

some change.

H

C N O B F

Bi Po

Cl

Br

I

At

Te

Se

S P Si

As

Sb

Al

Ge

Sn

Pb

Zn

Zr

Be

= amphoteric oxides = basic oxides

= acidic oxides = neutral oxides



9

Robert Lee Chin

Equilibrium and Indicators:

Indicators can be written as HIn, where „H‟ is the hydrogen atom and „In” is the

indicator molecule. Indicator reactions are reversible reactions. The equilibrium

situation is:

HIn H+

+ In-

Colour 1 Colour 2

If an alkali is added, the forward reaction is favoured, so more product is formed and

colour 2 appears. If an acid is added, the reverse reaction is favoured, so more

reactants form and colour 1 appears.

Other reactions e.g. combustion reactions; reactions between acids and metals are not

reversible- they go to completion.

Identify factors which can affect the equilibrium in a reversible reaction

By changing concentration, pressure or temperature of reactants and products, we can

affect the equilibrium point.

Concentration:

Increasing concentration of reactants will drive the reaction forward, while increasing

the products will drive the reaction in the reverse direction. For example, in

A+B C+D, increasing reactants will drive the reaction forward, producing more

products, thus reducing the concentration of A and B and maintaining equilibrium.

Temperature:

Reaction Effect on equilibrium if temperature

increases

Exothermic: A+B C+D + heat Shifts left- favours reactants

Endothermic: A+B + heat C+D Shifts right- favours products

Pressure:

(For reactions involving gases) If pressure is increased, the equilibrium will favour

the side with the lower amount of substances because this will reduce the number of

particles per volume.

Le Chatelier’s principle states that if a system in equilibrium is

disturbed/changed, then the system adjusts itself to minimise this change

These changes are:

-concentration of products + reactants

-temperature (different effects for endo- and exothermic reactions)

-pressure & volume (only if gases are involved)



10

Robert Lee Chin

Describe the solubility of carbon dioxide in water under various conditions as

an equilibrium process and explain in terms of Le Chatelier’s principle.

Carbon dioxide comes from volcanic gases, burning of organic matter and respiration

of plants and animals. It exists in sea and other natural waters and forms 0.03-

0.04%/V of the atmosphere. The concentration of CO2 in the atmosphere will

continue to increase due to more animals, more machines & factories, fewer

rainforests and increase in temperature (this means CO2 is les soluble in water, so

more is released).

When CO2 dissolves in water, an equilibrium forms:

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

The solubility of carbon dioxide in water can be explained in terms of De Chatelier‟s

principle. Changing the concentration, pressure, temperature or adding chemicals that

react with products or reactants alters the equilibrium.

Concentration:

If the concentration of CO2 is increased, the equilibrium will shift to the right to use

up the extra carbon dioxide (and if CO2 concentration is decreased, it will shift to the

left to produce more carbon dioxide). If more H2CO3 is added, the equilibrium will

shift to the left to use up the extra carbonic acid (and if H2CO3 is removed, it will shift

to the right to make more H2CO3).

Pressure:

Increasing the pressure of the carbon dioxide means will force the equilibrium to use

up more CO2 so there are fewer particles. The equilibrium will move to the right, so

more carbonic acid will be formed and the solution will become more acidic.

Temperature:

The reaction is exothermic, so can be written as:

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

Increasing the temperature will cause the equilibrium to shift to the left to use up the

added heat. This is why a warm can of fizzy drink is less fizzy than a cold can- less

CO2 can be dissolved

Adding reactive chemicals:

As carbonic acid forms, it ionises, so equilibrium occurs:

H2CO3(aq) 2H+

(aq) + CO32-

(aq)

So the equation can be written as:

CO2(g) + H2O(l) 2H+

(aq) + CO32-

(aq)

If we add OH- ions, they will react with the H

+ ions, removing them from solution.

The equilibrium will shift to the right to make more H+ ions (more CO2 dissolves).

The concentration of CO2 in the atmosphere will continue to increase due to more

animals, more machines & factories, fewer rainforests and increase in temperature

(this means CO2 is les soluble in water, so more is released).



11

Robert Lee Chin

Identify natural and industrial sources of sulfur dioxide and oxides of

nitrogen

Sulfur Dioxide

Sulfur dioxide is a colourless, toxic, gas with a pungent odour. It irritates the eyes,

damages the respiratory tract and can cause asthma. Industrial sources account for

over 75% of all emissions, in particular, combustion of fossil fuels

Natural Sources Industrial Sources

1. Burning organic matter (bushfires)

2. Decay of organic matter

3. Volcanic and hot spring emissions

1. Combustion of fossil fuels (esp.

Power plants, vehicles)

2. Smelting of sulphide ores into metal

(Pb, Zn, Cu)

3. Manufacture of sulphuric acid, paper,

food processing, sewage treatment

4. Petroleum refineries

5. Burning garbage

Oxides of nitrogen

There are 3 oxides of interest, all of which cause damage to the respiratory system,

increasing the risk of respiratory infections and asthma.

Nitrogen dioxide, NO2 Dinitrogen

monoxide, N2O

Nitrogen

monoxide, NO

Names Nitrogen (IV) oxide -Nitrogen (I) oxide

-Nitrous oxide (aka

„laughing gas‟)

-Nitrogen (II) oxide

-Nitric oxide

Colour -Red brown -Colourless Colourless

pH -Acidic

-Poisonous

-neutral -Not acidic, but

reacts with oxygen,

forming acidic NO2

In the atmosphere, these oxides are oxidised to nitric acid, nitrates and nitrites which

settle or are washed away by rain. Strong sunlight causes oxides of nitrogen to react

with hydrocarbons, forming photochemical smog. This became a major problem

during the industrial revolution in the mid 20th

century.

Nitrogen oxide Natural sources Industrial sources

Nitrogen dioxide, NO2 -Action of sunlight on

nitrogen monoxide and

oxygen

-Combustion of fossil fuels

in vehicles and power

stations

Dinitrogen monoxide,

(nitrous oxide) N2O

-Produced by soil bacteria -Fuel for racing cars

-Sedative/analgesic

(„laughing gas‟)

Nitrogen monoxide

(nitric oxide), NO

-Produced by soil bacteria

-Lightning

-Burning organic matter

-Combustion of fossil fuels

in vehicles and power

stations



12

Robert Lee Chin

Describe, using equations, examples of chemical reactions which release

sulfur dioxide and chemical reactions which release oxides of nitrogen

Reactions releasing sulfur dioxide

Manufacture of Iron(II) sulfate is prepared from tanks of dilute sulphuric acid used to

clean iron sheets before plating/galvanising. Iron(II) sulfate heated above 300°C

decomposes into Iron(III) oxide, sulfur dioxide and sulfur trioxide:

3(g)2(g)3(s)2

C300

4(s) SOSOOFe2FeSO

Iron sulphide (Pyrite/Fool‟s Gold) is a source of sulfur dioxide when roasted in air.

The other product is Iron(III) oxide:

2(g)3(s)2

heat

2(g)2(s) 4SOO2Fe7O4FeS

Oxidation of hydrogen sulphide during the decay of organic matter produces sulfur

dioxide and water:

)l(2)g(2

oxidation

)g(2 OHSOSH

Smelting of metal ores (copper, lead, zinc) releases sulfur dioxide. E.g. smelting zinc

sulphide releases sulfur dioxide and zinc oxide:

(s)2(g)

heat

2(g)(s) ZnOSOOZnS

In the laboratory, sulfur dioxide is prepared by heating copper with sulphuric acid.

Copper sulfate and water are by-products:

(l)24(aq)2(g)

heat

4(aq)2(s) O2HCuSOSOSO2HCu

Sulfur dioxide is also produced when a sulphite e.g. sodium sulphite (Na2SO3) is

treated with dilute acid:

(aq)(l)22(g)(aq)3(s)2 2NaOHSO2HSONa

Reactions releasing nitrogen oxides

High temperatures (e.g. combustion engines, lightning), nitrogen and oxygen combine

to form nitric oxide:

(g)

heat

2(g)2(g) 2NOON

This nitric oxide can slowly react with oxygen to form nitrogen dioxide:

2(g)2(g)2(g) NO2O2NO

Industrially, nitric oxide is prepared by catalytic oxidation of ammonia, producing

water as a by-product:

)g(2)l(2

oxidation

)g(2)g(3 NO4OH6O5NH4

In the laboratory, nitric oxide is produced using copper and nitric acid, producing

water and copper(II) nitrate as by-products:

)aq(23)l(2)g(2)aq(3)s( )NO(CuOH2NO2HNO4Cu2



13

Robert Lee Chin

In the laboratory, nitrogen dioxide is prepared by heating lead(II) nitrate crystals. The

by-products are lead oxide and oxygen:

)g(2)s()g(2

heat

)s(23 OPbONO2)NO(Pb

When heated, nitrogen dioxide forms nitric oxide and oxygen:

)g(2)g(

heat

)g(2 ONO2NO2

Heating ammonium nitrate produces nitrous oxide and water:

)l(2)g(2

heat

)s(34 OH2ONNONH

Assess the evidence which indicates increases in atmospheric concentration of

oxides of sulfur and nitrogen

Ice Core Samples

Ice core samples show Dinitrogen monoxide (N2O) levels have increased by 10%.

Damage

Damage to buildings, forests and aquatic organisms provides the most obvious

evidence for increasing levels of sulfur and nitrogen oxides. Human health is affected

in the form of respiratory diseases.

Difficulties obtaining evidence

Unlike carbon dioxide, sulfur and nitrogen dioxide are highly water soluble, so the

validity of atmospheric measurements is questionable. These oxides also occur in

much smaller concentrations than carbon dioxide (≈380ppm). The instruments used to

measure these changes have only been available since the 1970‟s.

Explain the formation and effects of acid rain

Acid rain is acidic because it contains dissolved acidic oxides i.e. carbon, nitrogen and

sulfur dioxides). Acidic oxides are released by several pathways. Natural sources are

volcanoes & geysers; decaying vegetation. Industrial sources include the combustion

of fossil fuels in industry and vehicles. In the atmosphere, they dissolve to form weak

acids. These acidic particles can precipitate as rain, hail, snow of fog.

For example, nitrogen dioxide forms nitric and nitrous acid:

2(aq)3(aq)(l)2(g) HNOHNOH2O2NO



14

Robert Lee Chin

Formation of acid rain:

The effects of acid Rain

Acid rain causes defoliation, stunted growth and decreased ability of plants to

withstand frost. It can also leech into the soil and cause chemical reactions that affect

plants

Sulphuric acid ionises in water and removes plant nutrients. Sulfate ions leech out

calcium and magnesium ions, which are essential for healthy soil.

Normally insoluble compounds such as aluminium sulfate dissolve in acidic water,

releasing toxic aluminium ions into the soil.

Acid rain lowers pH in lakes and streams, killing aquatic life such as fish and

crayfish. Aluminium ions cause reproductive problems and clog fish gills.

Sulfate ions reduce visibility, especially in major cities in the US.

Acid rain corrodes metals, stone structures, and paint. It is especially harmful to

calcium carbonate in concrete, limestone and marble.

Inhalation of sulfate ions has contributed to chronic respiratory diseases (lung cancer,

bronchitis, asthma) in humans.

O2 O2 H2S SO2 SO3 H2SO4

H2SO3

Acid Rain

Volcanoes &

geysers

Factories and

Industries

Vehicles Decaying

vegetation



15

Robert Lee Chin

Calculate volumes of gases given masses of some substances in reactions, and

calculate masses of substances given gaseous volumes, in reactions involving

gases at 0°C and 100kPa or 25°C and 100kPa

Use the following relationships:

massmolar

massmoles ofnumber

memolar volu

volumemoles ofnumber

(L) volume)(molLion concentratmoles ofnumber 1

23106.022

particles ofnumber moles ofnumber

Steps:

1. Change the mass/volume of given substances to moles

2. Write a balanced equation for the reaction to find mole ratios

3. Change the moles you have calculated back into the units the question asks.

Examples:

1)

a) Phosphorus trioxide, P2O3, slowly reacts with water forming phosphorus acid,

H3PO3. Write a balanced equation for this reaction

)(33)(2)(32 462P aqls POHOHO

b) When phosphorus acid is heated, it decomposes into phosphoric acid, H3PO4

and phosphine, PH3. Write a balanced equation for this reaction.

c) 7.10 L of phosphine gas was collected at 25°C, 100kPa. Show that the mass of

phosphine gas is 9.72 g.

The number of moles of phosphine gas is: moles ...2864.079.24

10.7

The mass of phosphine gas is the number of moles times the molar mass of

phosphine:

g9.73607...3(1.008)][30.97moles ...2864.0 , which is approximately equal to

9.72 g.

d) What mass of pure, solid phosphorus trioxide was involved in this reaction?

1 mole of phosphine is formed by 4 moles of phosphoric acid. The number of moles

of phosphoric acid is: moles...13736.124.79

7.104

3(g)4(aq)3

heat

3(aq)3 PHPO3HPO4H



16

Robert Lee Chin

4 moles of phosphoric acid is produced by 2 moles of phosphorus trioxide i.e. molar

ratio of phosphorus trioxide to phosphoric acid is 1:2. The number of moles of

phosphorus trioxide is: moles0.56868...5.024.79

7.104

The mass of phosphorus trioxide is:

d.p.) (2 44.91g.44.90885..3(16)]30.97[moles0.56868...

Identify data, plan and perform a first-hand investigation to decarbonate soft

drink and gather data to measure the mass changes involved and calculate

the volume of gas released at 25°C and 100 kPa.

Investigation: Carbon dioxide in carbonated mineral water

Background: As heat is applied to a soft drink, the amount of dissolved and reacted

carbon dioxide decreases, and thus more and more escapes as a gas i.e. increased heat

causes CO2(g) solubility to decrease.

Equipment:

-300 mL bottle of mineral water with flat base

-500 mL beaker, gauze mat, tripod, Bunsen burner, digital balance

-marble (calcium carbonate) heating chips

Method:

1/ Weight the capped bottle of mineral water

2/ Pour 200mL of water in the beaker

3/ Uncap the soft drink, being careful not to spill any drink. Reserve cap.

4/ Rest the soft drink in the „water bath‟. Gently heat the water and let boil for 3-

5 mins.

5/ Take off heat and let bottle dry completely. Reweight combined soft drink and

cap.

Bunsen burner

Tripod & gauze

Beaker with water &

marble chips

Mineral water



17

Robert Lee Chin

Results:

Total mass before heating = 520 g

Total mass after heating = 510 g

Mass difference = 520-520 = 10g

Calculations

s.f.) (3 L 5.63 ...6340909.5

79.2422

5memolar volu CO moles CO Volume

moles22

5

)]16(212[

10

CO massmolar

difference massescaping CO Moles

2(g)2(g)

2

2(g)

Analyse information from secondary sources to summarise the industrial

origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for

concern about their release into the environment

The main industrial sources of sulfur dioxide and oxides of nitrogen are:

-combustion of fossil fuels in vehicles and power stations

-burning of organic matter

-smelting of metal sulphides

-production of sulfuric acid, paper, food production, car fuels

-petroleum refining

Reasons for concern When these acidic oxides are released into the atmosphere (air pollution) by

industries, they dissolve in the water to form acidic rain. The acidic particles produced

can fall as gaseous or solid precipitates i.e. rain, hail, snow, dew, fog. Areas most

affected are the USA, Canada and North-Western Europe. Australia has had less of a

problem due to small population, isolation from other counties, coastal winds and low

sulfur content in or fossil fuels.

Acid rain causes damage to plant-life, aquatic life and ecosystems, man-made

structures, causes respiratory diseases in humans and reduced visibility in major

cities.



18

Robert Lee Chin

3. Acids occur in many foods drinks, and even within our

stomachs

Define acids as proton donors and describe the ionisation of acids in water

For example, sulphuric acid ionises to give hydrogen ions and sulphate ions:

(aq)2

(aq)4(aq)2 SO42HSOH

This reaction is known as an ionisation reaction and is exothermic

The H+ ions do not exist alone. The attach themselves to water molecules to form a

hydronium ion, H3O+. So the ionisation of sulphuric acid can be written more

correctly:

(aq)2

(aq)3(l)4(aq)2 SO4O2HH2OSOH

For example, when potassium hydroxide ionises in water it forms hydroxide ions and

potassium ions:

(aq)(aq)(aq) OHNaNaOH

The hydroxide ions can accept H+ ions to form water:

(l)2(aq)(aq) OHHOH

Some acids, such as acetic acid (CH3COOH) are weak acids. Acetic acid has 4

hydrogen ions, but only one actually ionises in water (strong covalent bonds prevent

the other hydrogen atoms from ionising):

CH3COOH(aq) H+

(aq) + CH3COOH-

Identify acids including acetic acid (ethanoic), citric (2-hydroxypropane-

1,2,3-tricarboxylic), hydrochloric acid and sulphuric acid

Name Formula Sources Other Info.

Acetic (ethanoic) acid CH3COOH -Vinegar (4%

solution)

-Bacterial action

-Pungent, colourless

-Used to make acetates

Citric acid

(2-hydroxypropane-

1,2,3-tricarboxylic)

C6H8O7 -citrus fruits

-Antioxidant

additive

-produced

fermentation of

-Colourless, crystalline

solid

-Found in blood and

urine

-Added to jam

Acids are substances that release hydrogen ions (protons) when dissolved in water.

Acids can be defined as proton donors.

A base can be defined as a proton acceptor.



19

Robert Lee Chin

sugar by

Aspergillus fungus

-antioxidant (food

additive)

Hydrochloric acid HCl -Stomach acid

Uses:

-Industry

-Cleaning brickwork

Sulfuric acid H2SO4 -acid rain Uses:

-explosives

-fertilisers

-petroleum refining

Lactic acid C3H6O3 -stiff muscles

-yogurt, whey Uses:

-lacquers and inks

-cosmetics

-pharmaceuticals

Methanoic (formic)

acid

H2CO2 -ants Uses:

-Rubber processing

Ascorbic acid C6H8O8 -fruits and

vegetables

-blood

(metabolically

active compound)

Uses:

- antioxidant (food

additive)

-blood cell formation,

tissue growth and healing

Describe the use of the pH scale in comparing acids and bases

The pH scale is used to compare the concentration of hydrogen ions in acid and base

solutions. The following table relates pH to the concentration of hydrogen ions,

hydroxide ions and example of common substances for given pH.

pH [H+] [OH

-] Example

substance

0 100

= 1 10-14

1 M HCl

1 10-1

10-13

0.1 M HCl

2 10-2

10-12

Stomach acid

3 10-3

10-11

Lemon juice

4 10-4

10-10

Beer

5 10-5

10-9

Acid rain

6 10-6

10-8

Urine

7 10-7

10-7

Pure water

8 10-8

10-6

Sea water

9 10-9

10-5

Toothpaste

10 10-10

10-4

Detergent

11 10-11

10-3

Ammonia

12 10-12

10-2

Drain cleaner

13 10-13

10-1

0.1 M NaOH

14 10-14

100 1 M NaOH

For aqueous solutions the product of the concentration of hydrogen ions and

hydroxide ions is the same, regardless of whether the solution is an acid, base or

mixture.



20

Robert Lee Chin

[H+] x [OH

-] = ionisation constant, Kw = 10

-14 at 25°C

For acidic solutions, [H+] greater than 10

-7 molL

-1 and pH less than7

For basic solutions, [OH-] less than 10

-7 molL

-1 and pH greater than 7

Identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-

fold change in [H+]

Because pH is a logarithmic scale, a change in pH of 1 indicates the hydrogen ion

concentration has changed by a factor of 10.

Mathematically, the pH of a solution is given by:

litreper molesin ionshygrogen ofion concentrat theis ][H where],H[logpH 10

To find the pH using a calculator:

1/ Tap the minus key

2/ Type in the [H+] (e.g. 2.0 x 10

-5)

3/ Tap the log key and press enter

Process information from secondary sources to calculate pH of strong acids

given appropriate hydrogen ion concentrations

In strong acids, all hydrogen is assumed to ionise. The concentration of hydrogen

ions will depend on the number of hydrogen ions an acid can donate i.e. monoprotic

acids release one H+ per molecule, diprotic releases two and triprotic releases three.

Examples:

1) Calculate the pH of a sulfuric acid solution of molarity: 0.001 molL-1

)aq(2

4)aq()aq(42 SOH2SOH

From the above balanced equation, 1 mole of sulfuric acid produces 2 H+.

Therefore, [H+] = 2(0.001) = 2.0 x 10-3

mol L-1

.

s.f.) (2 70.2...69897.2]100.2log[]Hlog[pH 3

2) Calculate the pH of 0.02molL-1

acetic acid if 3% ionises in water.

)aq(3)aq()aq(3 COOCHHCOOHCH

.)f.s 2( 2.4...2218.4]10 x 6.0log[]Hlog[pH

L mol 10 x 6.0 = 0.002 x 0.03 = ][H

5-

-1-5+

3) Calculate the pH of the solutions produced by:

a) Dissolving 2 g of NaOH and making volume to 2L



21

Robert Lee Chin

.)f.s 3( 4.12...39794.12]100.4log[]Hlog[pH

100.4105.2

10

][OH

10]H[

10][OH]H[

molL105.2][OH

.completely ionise willsobase, strong a is hydroxide Sodium

OHNaNaOH

molL105.22

05.0

volume

molesmolarity

moles05.0)11623(

2

massmolar

massNaOH moles

13

13

2

14-

-

14-

14--

12-

)aq()aq()aq(

12

Solve problems and perform a first-hand investigation to use pH

meters/probes and indicators to distinguish between acidic, basic and neutral

chemicals

Investigation: Using pH meter to distinguish between acidic, basic and neutral

chemicals

Aim: To use pH meters to determine the pH of certain chemicals

Equipment: -pH meter

-buffer solution

-0.1 M of the following solutions: HCl, NaCl, NaOH

-test tubes and test tube rack

Method:

1/ Calibrate the pH meter using the buffer solution.

2/ Place 25mL of each solution into separate test tubes

3/ Record the pH of the each of the solutions by placing the tip of the probe into

the solution. Rinse tip of probe with distilled water in between substances.

Results:

Substance pH

HCl 1

NaCl 7

NaOH 13



22

Robert Lee Chin

Conclusion:

-HCl has a low pH and is therefore acidic.

-NaCl has a medium pH, and is therefore neutral

-NaOH has a high pH, and is therefore basic.

Plan and perform a first-hand investigation to measure the pH of identical

concentrations of strong and weak acids

Investigation: measuring the pH of identical concentrations of strong and weak

acids

Aim: To determine the pH of identical concentrations of weak and strong acids.

Equipment:

-pH meter

-buffer solution

-0.1 M of the following solutions: CH3COOH, HCl, H2SO4

Method:

1/ Calibrate the pH meter using buffer solution.

2/ Place 25mL of each solution into separate test tubes

3/ Record the pH of the each of the solutions by placing the tip of the probe into

the solution. Rinse tip of probe with distilled water in between substances.

Results:

Substance pH Degree of ionisation

CH3COOH 3 1%

HCl 1 100%

H2SO4 0.7 100%

Gather and process information from secondary sources to write ionic

equations to represent the ionisation of acids

For strong acids that ionise 100% e.g. HCl, H2SO4, the ionic equation can be written

with one arrow from left to right.

For example, the ionic equation for a strong monoprotic acid, hydrochloric acid:

)aq()aq()aq( CLHHCl

Ionic equation for a strong diprotic acid, sulphuric acid:

)aq(2

4)aq()aq(42 SOH2SOH

For weak acids, the equation will be written with reversible arrows to indicate that the

equilibrium point has a significant amount of both reactants and products.

For example, the ionisation of carbonic acid:

H2CO3(aq) 2H+

(aq) + CO32-

(aq)



23

Robert Lee Chin

For organic acids such as acetic acid and citric acid, the H+ from the =COOH group

ionises. For example, the ionisation of acetic acid:

CH3COOH (aq) CH3COO-(aq)+ H

+(aq)

Describe the difference between a strong and a weak acid in terms of an

equilibrium between the intact molecule and its ions

A strong acid is one where nearly 100% of the molecules ionise in an aqueous

solution. For example, hydrochloric acid:

)aq()aq(3)l(2)aq( ClOHOHHCl

A weak acid is one that does not fully ionise. For example, when hydrogen cyanide is

placed in water, less than 1% ionises and an equilibrium situation is set up

HCN(aq) + H2O(l) H3O++CN

-

Adding more water increases the degree of ionisation, but the concentration will not

increases because there is more solution.

Use available evidence to model the molecular nature of acids and simulate

the ionisation of strong and weak acids

Strong acids e.g. HCl, H2SO4 disassociate almost entirely in water to form

positive hydrogen ions (protons) and anions

Weak acids do not dissociate entirely in water. Most of the acid molecules

remain in the solution.

+

+ +

+ _

_ _

_ _

_

_

_ +

_

_

_

_ + Acid molecule

+

+

_

_

_

_ _

_ +

_

_ +

_

_ +

_

_



24

Robert Lee Chin

Describe acids and their solutions with the appropriate use of the terms

strong, weak, concentrated and dilute

A concentrated solution is one in which the total concentration of solute species is

high. A 10molL-1

solution would be called concentrated.

A dilute solution is one in which the total concentration of solute species is low.

A strong acid is one in which all the acid present in solution has ionised to form

hydrogen ions. There are few neutral acid molecules left.

A weak acid is one in which only some of the acid molecules present in the solution

have ionised to form hydrogen ions. Weak acids for equilibrium reactions with water

Compare the relative strengths of equal concentrations of citric, acetic and

hydrochloric acids and explain in terms of the degree of ionisation of their

molecules

If we compare different acids of equal concentrations, the pH will depend on the

number of H+ ions that ionise in solution. For strong acids, the acid will ionise 100%

e.g. ClHHCl (aq)

For weak acids such as acetic and citric acid, only 1% ionises

e.g. CH3COOH CH3COO- + H

+

Gather and process information from secondary sources to explain the use of

acids as food additives

Microorganisms such as clostridium botulism produce toxins in food, which can cause

severe food poisoning. Acids are used to preserve foods because many

microorganisms including yeasts, moulds and bacteria, are pH sensitive and are killed

when exposed to acidic conditions. Some acids act as antioxidants by retarding the

oxidation of certain chemicals in food e.g. vitamin C. The addition of acids extends

the shelf life of many processed food products including dairy, baked goods, cured

meats, fruits and vegetables. In some cases, acids also give a unique flavour to some

foods e.g. picked vegetables, sweet and sour sauces

Common acids used as preservatives include acetic acid in vinegar. Vinegar is often

used to pickle vegetables for canning. Other foods may be fermented to produce acids

by bacteria or fungi. For example, the fermentation of milk to yogurt converts lactose

to lactic acid.

Sulfur dioxide is the only acidic oxide is commonly used as a food preservative. It is

added to foods such as dried fruit and preserved deli meats because it maintains the

appearance of the food and helps prevent rotting. Other acids used as food

preservatives include phosphoric acid, citric acid, propanoic acid, benzoic acid,

sodium nitrate and sorbic acid.



25

Robert Lee Chin

Identify data gather and process information from secondary sources to

identify examples of naturally occurring acids and bases and their chemical

composition

Name/s Composition Acid/base pH in

natural form

Natural source/s

Hydrochloric

acid

HCl acid 1.0 Gastric juices

Acetic acid CH3COOH acid 2.0-3.0 Vinegar, fruits and

vegetables

Salicylic acid C6H4(OH)COOH acid 2.5 Plants e.g. rhubarb

Tartaric acid C4H6O6 acid 2.5 (2.5%

solution)

Wine fermentation

caffeine C8H10N4O2 base 8.0 Coffee, tea

Calcium

carbonate

CaCO3 base 9.0 Chalk, marine

shells, eggshells

Sodium

hydroxide

NaOH base 13 Burnt ashes, lye

water

Ammonia NH3(aq) base 11.5 All living

organisms

4. Definition of acids and bases

Outline the historical development of acids including those of:

-Lavoisier

-Davy

-Arrhenius

Antoine Lavoisier (1743-94) was a French chemist who demonstrated that

combustion reactions involved oxygen. Experimentation led him to believe that acids

were composed of two substances, one of them being oxygen. He believed oxygen

was present in all acids and as responsible for acidity

Humphry Davy (1778-94) was an English chemist who was famous for electrolysis

experiments. In 1810, he decomposed hydrochloric acid and found it was composed

of hydrogen and chlorine and did not contain oxygen. He observed that metals could

displace hydrogen from acids to form salts. He concluded that all acids contain

hydrogen.

In 1884, Swedish chemist Svante Arrhenius (1859-1927) proposed definitions for

acids and bases. He suggested that acids were neutral substances that produce

hydrogen ions as the only poitive ion in an aqueous solution and that bases are

substances that produce hydroxide ions as the only negative ion in an aqueous

solution. His theory was limited because it applied only to aqueous solutions, only

accounted for substances containing hydrogen or hydroxide ions and could not

explain amphoteric substances such as zinc oxide



26

Robert Lee Chin

Gather and process information from secondary sources to trace

developments in understanding and describing acid/base reactions

Scientist/s Acid

definition

Base

definition

Notes

Lavoisier n/A n/A Oxygen is present in all acids and

is responsible for the acidity

Davy n/A n/A Acids contain hydrogen. They do

not have to contain oxygen

Arrhenius Acid ionises

in water to

form protons

and anion

Base ionises in

water to form

hydroxide ion

and cation

-Applies only to aqueous solutions

-Only accounts for substances

containing hydrogen or hydroxide

ions

-Cannot explain amphoteric

substances

Brönsted-

Lowry

Proton

donators

Proton

acceptors

Acids must contain hydrogen

Each acid has a conjugate base

Outline the Brönsted-Lowry theory of acids and bases

An acid-base reaction is one in which a proton is transferred from an acid to a base.

An acid is defined as a proton acceptor while a base is a substance that accepts a

proton from an acid.

The Brönsted-Lowry theory is advantageous because it is able to explain:

- non-aqueous reactions

-why some salts can act as acids or bases

-why some substances are amphoteric.

It forms the basis for the qualitative treatment of acid-base equilibriums and pH

calculations

Describe the relationship between an acid and its conjugate base and a base

and its conjugate acid

Identify conjugate base pairs

Every acid has its conjugate base a substance with exactly one less proton. An acid

transfers a proton to its conjugate base in an acid-base reaction. Together, this acid

and base form a conjugate pair.

The products of an acid-base reaction are another acid and base, so there are always

two conjugate acid-base pairs in each reaction. For example:

2 acid 1 base 2 base 1 acid

OHCOOCHOHCOOHCH 33)l(2)aq(3



27

Robert Lee Chin

CH3COOH is an acid and CH3COO- is its conjugate base. Together they form a

conjugate pair. Similarly, H3O+ is an acid, and H2O is its conjugate base.

A strong acid has a weak conjugate base and a strong base has a weak conjugate

acid:

Strongest Acids Acid Base

Weakest Bases HCl Cl-

H2SO4 HSO4-

HNO3 NO3-

H3O+ H2O

HSO4- SO4

2-

H3PO4 H2PO4-

CH3COOH CH3COO-

H2CO3 HCO3-

H2S HS-

NH4+ NH3

H2O OH-

Weakest Acid HS- S

2- Strongest Acid

OH- O

2-

Identify amphiprotic substances and construct equations to describe their

behaviour in acidic and basic solutions

Amphoteric= a substance that can act as both an acid and a base e.g. zinc and

aluminium oxide

Amphiprotic (as defined by the Brönsted-Lowry theory) = an amphoteric substance

that can donate or accept protons i.e. it can act as a conjugate acid and a conjugate

base.

Amphiprotic substances include water (H2O), ammonia (NH3), hydrogen carbonate

ion (HCO3-) & phosphane (PH3).

Water acting as a proton donator: H2O(l) + NH3(g) NH4+(aq) + OH-(aq)

Water acting as a proton acceptor: H2O(l) + HCl(aq) H3O+

(aq) +Cl-(aq)

Identify neutralisation as a proton transfer reaction which is exothermic

In a neutralisation reaction, hydrogen ions and hydroxide ions form water.

Neutralisation reactions usually occur between a strong acid and a strong base.

H+

H+



28

Robert Lee Chin

For example, the reaction between hydrochloric acid and sodium hydroxide:

salt water acid base

NaClOHHClNaOH )aq()l(2)aq()aq(

The net ionic equation shows that it is a proton transfer reaction:

)l(2)aq()aq( OHHOH

Almost all neutralisation reactions are exothermic, releasing about 57 kJ of heat per

mole of water formed i.e. ΔH=-57kJ.

Perform a first-hand investigation and solve problems using titrations and

including the preparation of standard solutions, and use available evidence to

quantitatively and qualitatively describe the reaction between selected acids

and bases

Experiment: Preparation of a standard solution of hydrochloric acid

Aim: To prepare a primary standard solution of sodium carbonate and use it to

determine the concentration of a hydrochloric acid solution.

Equipment:

-Small beaker -Electronic scale-Burette

-Pipette (20mL) -3 x 250mL conical flasks

-250 mL volumetric flask -Burette clamp and retort stand

-Wash bottle with distilled water -Pipette filler

-Approx 0.1 molL-1

HCl (100mL) -Approx 2.0g dried Na2CO3

-Stirring rod -Suitable indicator (methyl red)

Safety: Wear safety glasses. Hydrochloric acid is corrosive, so avoid contact with

skin. If contact occurs, wash well with soap and water. Do NOT pipette by mouth: use

pipette filler. Many indicators are poisonous and should be handled with care

Method:

A Preparing the primary standard

1/ accurately weigh 2.0 g of sodium carbonate in a small beaker

2/ Add a small amount of distilled water to beaker and stir ti dissolve sodium

carbonate. Use a wash bottle with distilled water to wash out all the sodium carbonate

solution into the funnel.

3/ Rinse the beaker and stirring rod with small amounts of distilled water and transfer

the wash water into the flask

4/ Add distilled water to the volumetric flask until it is about two-thirds full. Fit the

stopper and shake to dissolve all the sodium carbonate. When all is dissolved, top up

the flask until the bottom of the meniscus is level with the mark.



29

Robert Lee Chin

B Standardising the hydrochloric acid/ solution

5/ Rinse the burette with distilled water and then with HCl, discarding the rinsings.

6/ Set up the burette with burette clamp and fill with HCl. Record the starting volume.

7/ Rinse the pipette with sodium carbonate solution and place in a clean 250 mL

beaker. This exact volume is known as the aliquot. Add 2-3 drops of indicator.

8/ Place the flask under the burette and run HCl into the flask, swirling continuously

until colour changes. This is the end point. Record end volume.

9/ Calculate the volume of HCl used and record

10/ Refill the burette and repeat steps 5-9 at least twice more until three precise results

are obtained.

C Standardising the 10% vinegar solution

11/ Measure 25mL vinegar using measuring flask

12/ Place into clean conical flask and fill to 250mL solution

13/ Rinse the burette with distilled water and then with 10% vinegar, discarding the

rinsings.

14/ Set up the burette with burette clamp and fill with 10% vinegar.

15/ Rinse the pipette with sodium carbonate solution and place in a clean 250 mL

beaker. This exact volume is known as the aliquot. Add 2-3 drops of indicator.

16/ Place the flask under the burette and run 10% vinegar into the flask, swirling

continuously until colour changes. This is the end point.

17/ Calculate the volume of 10% vinegar used and record

18/ Refill the burette and repeat steps 5-9 at least twice more until three precise results

are obtained.

Results:

Mass Na2CO3 used = 2.0g

Trial 1 Trial 2 Trial 3 Trial 4

Volume

Na2CO3 used

(L)

0.02 0.02 0.02 0.02

Initial burette

reading (L)

0.031 0.05 0.03 0.03

Final burette

reading (L)

0.007 0.03 0.01 0.011

Volume HCl

used (L)

0.024 0.02 0.02 0.019

Volume 10%

vinegar used

0.091 0.074 - -

Calculations:

Molarity of Na2CO3 solution:

1

32

32

32

molL0754.025.0

531

volume

molesCONa Molarity

L25.0mL250CONa Volume

moles53

1

)16(312)23(2

0.2

massmolar

massCONa moles



30

Robert Lee Chin

Moles Na2CO3 used in titration:

.)f.s 3( moles 0015.0...0015094.002.0...0754.0volumeionconcentratmoles

Average volume of acids used:

L0825.02

074.0091.0used COOHCH 10% volumeAverage

L02075.04

0.0190.020.020.024used HCl volumeAverage

3

acid. acetic moles .)2(0.0015.. neutralise will

carbonate sodium oles0.0015...m acid. acetic of moles 2 sneutralise carbonate sodium of mole 1

COOCHNa2OHCOCOOHCH2CONa

acid. ichydrochlor moles .)2(0.0015.. neutralise willcarbonate sodium

oles0.0015...m acid. ichydrochlor of moles 2 sneutralise carbonate sodium of mole 1

NaCl2OHCOHCl2CONa

)aq(3)l(2)g(2)aq(3)aq(32

)aq()l(2)g(2)aq()aq(32

Molarity of acid solutions:

.)f.s 3( molL 0.376 solution COOHCH 100% ofMolarity

.)f.s3( molL 004.0...003659.00825.0

.)2(0.0015..

volume

molessolution COOHCH 10% ofMolarity

.)f.s3( molL 015.0...14548761.002075.0

.)2(0.0015..

volume

molessolution HCl ofMolarity

1-

3

1

3

1

Describe the correct technique for conducting titrations and preparation of

standard solutions

A titration is a method used to experimentally determine the molarity of a solution. It

is a volumetric analysis technique. A solution of a known concentration called the

standard solution is added to a solution of unknown concentration until the

neutralisation reaction is complete.

Standard Solutions



31

Robert Lee Chin

There are two types of standard solutions: Primary standard and secondary standard.

Standard solutions are also known as titrants.

A primary standard is a solution that has been made by dissolving an accurately

measured mass of solute in a small amount of solvent and made to the required

volume in a volumetric flask.

A secondary standard is a solution whose concentration has been found by titrating

against a primary standard.

For a chemical to be suitable to prepare as a primary standard solution it must:

*Be a water soluble solid

*Be obtainable in pure form

*Have an accurately known formula

*Be stable in air

To prepare the standard solution:

1) Accurately weigh a calculated amount of solid

2) Dissolve it in water

3) Transferring ALL of the dissolved solid to a volumetric flask

4) Adding water to prepare a known volume of solution

The reaction is complete at the equivalence or end point. This is when the molar ratio

of H+ to OH ions is equal i.e. basebaseacidacid cvMolescvMoles . The solution

changes colour at the equivalence point.

Selecting the Indicator

The equivalence point is not always at pH=7. The salts formed by combining different

strong and weak acids have acidic or basic properties. Thus, an indicator must be

chosen that changes colour near the equivalence point

Colour in:

Indicator Acid Base pH change

litmus red blue 6.0-8.0

Bromothymol

blue

yellow blue 6.2-7.6

Methyl orange Red Yellow 3.1-4.4

Phenophalein clear pink 8.3-10.0

Strong acid and strong base:



32

Robert Lee Chin

Strong acid and weak base:

Weak acid and strong base:

Equivalence point

pH

14

7

Volume acid in base

Equivalence point

pH

14

7

Volume acid in base

pH

14

7

Volume acid in base

Equivalence point



33

Robert Lee Chin

Titration Equipment

The main equipment includes:

-pipette and burette to measure volume of reactants

-Flask to mix reactants

The pipette measures a fixed volume of solution to provide a fixed number of moles

of one reactant. Before using, it must be rinsed with distilled water, then with the

solution to be used. Rinsing with the solution removes any water which would alter

the volume and hence, number of moles of the solution being drawn. The solution

should be drawn so that the bottom of the meniscus is in line with the etched line. The

volume measured by the pipette is called an aliquot.

The flask should be rinsed only with distilled water. It does not matter if it is wet, as

this will not alter the number of moles of solution used (this has already been

accurately measured by the pipette).

The burette allows the exalt volume of the reactant required to reach the equivalence

point. Like the pipette, it must first be rinsed with distilled water, then with the

solution to be used. The volume delivered by the burette is called a titre.

Titration Procedure

1. Ensure all equipment is cleaned and rinsed with correct liquid

2. Add one solution to the burette

3. Use pipette to measure volume of other solution

4. Transfer this to a conical flask

5. Add a few drops of the suitable indicator

6. Perform a rough titration to find endpoint.

7. Repeat carefully until at least three readings within 0.1 mL of each other are

obtained

8. Perform calculations.

20mL pipette

Burette



34

Robert Lee Chin

Choose equipment and perform a first-hand investigation to identify the pH

of a range of salt solutions

Investigation: Determining the pH of salt solutions

Aim: To select equipment and perform an experiment to determine the pH of various

salt solutions

Equipment:

Test tubes and test tube rack

Universal indicator

Demineralised water

5mL of a 0.1 M solution of the following salts:

-Ammonium Chloride (NH4Cl)

-sodium carbonate (Na2CO3)

-sodium hydrogen sulfate (NaHSO4)

-potassium nitrate (KNO3)

-ammonium acetate (CH3COONH4)

Safety: Wear safety glasses

Method:

1/ Place 5mL of each salt solution and demineralised water into separate test tubes

2/ Use a few drops of universal indicator to determine the pH

Results:

Salt Formula Experimental pH Acid/neutral/base

Ammonium

chloride

NH4Cl 6.0 Weakly acidic

Sodium carbonate Na2CO3 11.0 Moderately acidic

Sodium Hydrogen

sulfate

NaHSO4 3.0 Strongly acidic

Potassium nitrate KNO3 7.5 neutral

Ammonium acetate NH4CH3COO 7.0 Neutral

Demineralised

water

H2O 7.0 neutral

basek weaacideak w

Cl NHClNH )aq()aq(4)aq(4

base strong neutral

CoNa2CONa )aq(2

3)aq()sq(32

acid strong neutral

HSO NaNaHSO )aq(4(aq))aq(4



35

Robert Lee Chin

Qualitatively describe the effect of buffers with reference to a specific

example in a natural system.

A buffer is a chemical that controls the pH of a solution. Buffer solutions are usually a

mixture of a weak acid and the salt of that acid or a weak base and the salt of that base

e.g. hydrogen carbonate ion (HCO3-) and carbonate ion (CO3

2-).

In a buffer, equilibrium is established between the weak acid and its conjugate base.

There are two reactions involved. For example, the buffer system involving hydrogen

carbonate ions and carbonate ions:

If an acid is added to the buffer, hydrogen ions are removed:

)aq(32)aq(3)aq( COHHCOH

If a base is added to a buffer, hydroxide ions are removed:

)aq(2

3)l(2)aq(3)aq( COOHHCOOH

In living organisms, blood is a buffered solution containing carbonic acid and sodium

bicarbonate:

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

(aq) + HCO3-(aq)

The more CO2 that dissolves, the more H+ will form. The equilibrium shifts to the left

to resist this change. If the pH is increasing, more carbonic acid will dissolve and the

equilibrium will shift to the right to minimise the change. Carbonic acid is a weak

acid, so a change in hydrogen concentration will not affect the pH much.

Analyse information from secondary sources to assess the use of

neutralisation reactions as a safety measure or to minimise damage in

accidents or chemical spills

Neutralisation reactions are commonly utilised for safety in laboratories where acids

and bases are used. When selecting the appropriate neutralisation reagent, the

following factors need to be considered:

-speed of neutralisation

-need for reagent that will not be harmful in excess

-safe to handle and store

-cost

-ability to use for both acids and bases i.e. amphiprotic

Common neutralising reagents include the hydrogen carbonate ion found in sodium

hydrogen carbonate.

When the carbonate ion is used for acid spills, it combines with a hydrogen ion,

forming water and carbon dioxide:

)g(2)l(2)aq(3)aq( COOHHCOH



36

Robert Lee Chin

When it is used for base spills, it combines with the hydroxide ion, forming water and

the carbonate ion: 2

3)l(2)aq(3)aq( COOHHCOOH

Identify a range of salts which form acidic, basic or neutral solutions and

explain their acidic, neutral or basic nature

A salt formed by a strong acid and a strong base is neutral e.g. NaCl:

Sodium chloride forms by reacting sodium hydroxide with hydrochloric acid.

When sodium chloride dissolves in water, the sodium chloride forms Na+ and Cl

-. The

water forms H+ and OH

-:

NaCl(aq) Na+

(aq) + Cl-(aq)

H2O(l) H+(aq) + OH

-(aq)

Sodium ions are attracted to hydroxide ions, forming NaOH, a strong base that

completely ionises. Chlorine ions are attracted to hydrogen ions, forming HCl, a

strong acid which ionises completely. Thus, the concentration of hydrogen ions equals

the concentration of hydroxide ions and the solution is neutral.

A salt formed by a strong acid and a weak base is acidic e.g. NH4Cl:

Ammonium chloride forms by reacting hydrochloric acid and ammonium hydroxide.

When ammonium chloride dissolves in water, it forms NH4+ and Cl

-. The water forms

H+ and OH-.

NH4Cl(aq) NH4+

(aq) + Cl

-(aq)

H2O(l) H+

(aq) + OH-(aq)

Cl- Cl

-

Cl-

Na+

H+

Na+

Na+

OH-

OH-

OH-

H+

H+



37

Robert Lee Chin

Ammonium ions are attracted to hydroxide ions, forming NH4OH, a weak base.

Chlorine ions are attracted to hydrogen ions, forming hydrochloric acid, a strong acid.

Thus, the concentration of hydrogen ions is greater than the concentration of

hydroxide ions and the solution is acidic.

A salt formed by a weak acid and a strong base is basic e.g. CH3COONa:

Sodium acetate forms by reacting acetic acid and sodium hydroxide. When sodium

acetate dissolves in water, the sodium acetate forms Na+ and CH3COO

-. The water

forms H+ and OH-.

CH3COONa (aq) Na+

(aq) + CH3COO

-(aq)

H2O(l) H+

(aq) + OH-(aq)

Sodium ions are attracted to hydroxide ions, forming NaOH, a strong base. Acetate

ions are attracted to hydrogen ions, forming acetic acid, a weak base. Thus, the

concentration of hydroxide ions is greater than the concentration of hydrogen ions and

the solution is basic.

Cl-

Cl-

Cl-

NH4+

H+

OH-

OH- H

+

H+

NH4OH

NH4OH

NH4+

Na+ OH

-

OH-

H+

OH-

Na+

CH3COOH

CH3COOH

CH3COO-

Na+

Na+

OH-



38

Robert Lee Chin

4. Esterification

Describe the differences between the alkanol and alkanoic functional groups

in carbon compounds

The alkanol functional group is the hydroxyl group, –OH. The alkanoic acid

functional group is the carboxyl group, –COOH. The carboxyl group makes alkanoic

acids polar molecules. Alkanoic acids form hydrogen bonds so are water soluble.

The general formula for alkanoic acids is RCOOH, where R represents the akyl chain

with the formula CnH2n+1. For example, pentanoic acid is C4H9COOH. Remember that

one C from the akyl group is already included on the COOH functional group. Thus,

for pentanoic acid (5 carbons), the number of carbon atoms in the R group is 4.

Structural formula for Pentanoic acid:

Identify the IUPAC nomenclature for describing the esters produced by

reactions of straight chain alkanoic acids from C1 to C8 and straight chain

primary alkanols from C1 to C8

Esters are named with the alkanol functional group first, replacing the suffix „anol‟

with “yl”.

When writing the chemical formula, the acid comes first, followed by the

–COO– functional group, then the alkanol group. The suffix „oic acid‟ is replaced by

„anoate‟

For example:

r wateester acid alkanoic alkanol

OH pentonoate butyl acid pentoic Butanol

OHHCOOCHCCOOHHCOHHC

2

SOH conc.

)l(2)l(9494

SOH conc.

94(l)94

42

42

Carbon

atoms

First

Part

Second

Part

1 Methyl Methanoate

2 Ethyl Ethanoate

3 Propyl propanoate

4 Butyl butanoate

5 Pentyl pentanoate

6 hexyl Hexanoate

7 Heptyl Heptanoate

8 octyl octanoate

C

-COOH functional

group

OH

O

CH2 CH2 CH2 CH3



39

Robert Lee Chin

Explain the differences in melting point and boiling point caused by straight

chain alkanoic acid and straight chain primary alkanol structures

Alkanes, alkenes and alkynes are non-polar and do not form hydrogen bonds. They

only have weak dispersion forces and thus low boiling points.

The high melting and boiling points of alkanols are due to the hydrogen bonding of

the O in one molecule, and the H from the -OH group in a nearby molecule. They also

have 1 centre of polarity, forming dipole bonding.

In alkanoic acids each carboxyl group is able to form two strong hydrogen bonds.

This is because they have two O groups and plenty of hydrogen groups. They have 2

centres of polarity and dipole bonding. This gives alkanoic acids an even higher

boiling point than their corresponding alkanol.

Identify esterification as the reaction between an acid and an alkanol and

describe, using equations, examples of esterification

Esters are produced in a condensation reaction between an alkanol and an alkanoic

acid called esterification. This is a reversible reaction that forms an equilibrium

situation.

O

H

H

O

Akyl chain

Akyl chain Hydrogen

bond

Alkanes,

alkenes &

alkynes

Alkanols

Boiling

point

Carbon atoms per molecule

Alkanoic

acids

Hydrogen

bonds

Akyl chain

OH O

C

O

C Akyl chain

HO



40

Robert Lee Chin

A molecule of water is condensed out during the reaction. Use of tracers indicates that

the OH comes from the alkanol and the O from the acid. For example:

water ethanoate methyl methanol acid icethano

Describe the purpose of using acid in esterification for catalysts

Esterification is a slow process that does not reach completion at room temperature

because it forms an equilibrium situation. Concentrated sulphuric acid is used as a

catalyst. Also, the acid is hydroscopic meaning it absorbs water, shifting the

equilibrium to the right and producing more ester.

Explain the need for refluxing during esterification

The reflux system consists of a reflux condenser fixed onto a reaction flask. The

reaction flask is heated to speed up the reaction. The reflux condenser prevents the

loss of volatile reactants (i.e. alcohol) or products during heating. It is open at the top

to avoid the dangerous build-up of pressure.

Outline some examples of the occurrence, production and use of esters

Esters occur widely in living things, esp. fruits and flowers. They are responsible for

many aromas and flavours in foods. The aroma and flavour of fruits and flowers is

from a complex mixture of esters and other compounds but esters are responsible for

the main aroma.

Fats and oils are the long chain “fatty acid” esters of the triple-alcohol molecule,

glycerol. These form the long-term energy storage sites in plants and animals.

Esters are produced via the reflux of an alcohol, alkanoic acid and a catalyst on an

industrial scale.

Name Structure Use

Ethyl ethanoate CH3COOC2H5 Nail polish remover

Ethyl butanoate CH3COOC4H9 Pineapple

Pentyl ethanoate C4H9COOC2H5 Banana

Octyl ethanoate C7H15COOC2H5 Orange

H2O

Conc. H2SO4

→ Heat

HO CH2

O

C OH CH3 + +

O

C CH3 O CH3



41

Robert Lee Chin

Uses of esters include artificial flavours for drinks and processed foods, industrial

solvent in the plastic industry and in cosmetics such as shampoos and lipstick.

Identify data, plan, select equipment and perform a first-hand investigation

to prepare an ester using reflux.

Experiment: Preparation of an ester

Aim: To prepare an ester using reflux

Equipment:

The following alkanols:

-methanol

-Ethanol

-Concentrated sulfuric acid

-1.0 molL-1

Na2CO3 solution

-retort stand and clamps

-conical flask

-Funnel

-separating funnel

-boiling chips (not marble)

-condenser with rubber tubing

-Bunsen burner

-clay triangle

-tripod

Method:

1/ Add a few boiling chips to the funnel. Place 8mL of one alkanol, 24mL of one

of the alkanoic acids and 1Ml concentrated sulfuric acid into a flask using

funnel.

The following alkanoic acids:

-butanoic acid

-Glacial acetic acid

-Salicylic acid

Safety:

Wear safety glasses at all times. Sulfuric acid is corrosive. Clean up spills

immediately and wash affected area with large quantities of water. Organic chemicals

are flammable. Do not allow liquids or vapours to come into contact with sparks or

flames and avoid inhaling vapours.



42

Robert Lee Chin

2/ Set up equipment as shown below:

3/ Connect tubing to tap and condenser and turn on water so a uniform flow is

achieved.

4/ Heat the mixtures over a steady Bunsen burner for 30 minutes (do not let

mixture boil too vigoursly) and allow to cool for 5 minutes. Turn off water.

5/ Carefully remove the flask and pour contents into a separating funnel

containing 15mL water. Stopper the funnel and shake. Allow layers to

separate, drain off and discard aqueous layer.

6/ Add 15mL sodium carbonate solution. This will neutralise the acid and

prevent the reaction from going backwards. Shake and drain the lower layer.

The ester should be in the separating funnel.

7/ Carefully smell the ester and describe the smell.

Results:

Alkanol Alkanoic acid Ester Aroma

ethanol Acetic acid Ethyl acetate Nail polish

remover; acetone

ethanol Butanoic acid Ethyl butanoate Banana, fruity

methanol Salicyclic acid Methyl

salicycoanate

Oil of

wintergreen

Retort stand

and clamps

Bunsen, tripod clay

& clay triangle

Flask with reaction

mixture and boiling

chips

water out

Condenser

water in

Documents

Module 2- The Acidic Environment