module 2- the acidic environment
TRANSCRIPT
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
te
Su
lfite
Su
lfate
Iod
ide
Bro
mid
e
Ch
lorid
e
Flu
orid
e
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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)
HSC Chemistry Summary
Module 2- The Acidic Environment
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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)
HSC Chemistry Summary
Module 2- The Acidic Environment
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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.
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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)
HSC Chemistry Summary
Module 2- The Acidic Environment
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).
HSC Chemistry Summary
Module 2- The Acidic Environment
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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.
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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.
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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.
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[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
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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
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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)
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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
+
+
_
_
_
_ _
_ +
_
_ +
_
_ +
_
_
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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.
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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
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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
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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+
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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.
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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
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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
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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:
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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
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
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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+
HSC Chemistry Summary
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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-
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
Module 2- The Acidic Environment
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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
HSC Chemistry Summary
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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.
HSC Chemistry Summary
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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