Chemistry – The Acidic Environment The Acidic

Transcription

Chemistry – The Acidic Environment
The Acidic Environment
1. Indicators were identified with the observation that the colour of some
flowers depends on soil composition
Classify common substances as acidic, basic or neutral
Acids: are substances capable of providing H+ ions
H++H2O
H3O+͢
Bases: are substances that give rise to hydroxide ions OH Substances
Acid
Vinigar
Acetic
Lemon Juice
Citric
Asprin
Acetyl salicyle
Car batteries
Ascorbic
Rust converters
Sulfuric
Cloud ammonia
Phosphoric
Whashing soda
Base
Neutral
-
Antacid tablets
Sodium carbonate
Cleaners
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Lime
Sodium hydroxide
Pure water
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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
©
(2012) All Rights Reserved
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Indicators: change colour depending on how acidic or basic the solution is
Universal indicator is a mixture of litmus, phenolphthalein, methyl orange and
bromothymol blue.
Each of these can also be used independently as an indicator
Indicators were first identified by the observations that the colour of some
flowers depends on the soil composition
Indicator
Colours
Highly acidic Slightly
acidic
Neutral
Slightly
alkaline
Highly
alkaline
Methyl orange
Red
Yellow
Yellow
Yellow
Yellow
Bromothymol
blue
Yellow
Yellow
Green
Blue
Blue
Litmus
Red
Red
Purple
Blue
Blue
Colourless
Colourle
ss
Colourle
ss
Purple/p
ink
Phenolphthale Colourless
in
Identify and describe some everyday uses of indicators including the testing of
soil acidity/basicity
In chemistry:
Used regally in laboratories to:
• Moniter how acidic or alkaline solutions from industries are before they
are pumped into a river or sea- this ensures that we can monitor the
affect they have on the environment and make sure we minimise this
affect
Not in chemistry:
• Testing the water in aquariums- some aquatic organisms can only
survive withing a very narrow pH, it is therefore imperative that we
measure this pH to ensure their survival
• Testing the water in swimming pools
• Testing soil samples- if the pH of soil is too great or small this may
result in the death of some plants
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Perform a first-hand investigation to prepare and test 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 Red cabbage leaves
• 2 beakers
• Distilled water
• Measuring cylinder
• Tripod, gauze mat and
Bunsen burner
• Test tubes
• Test tube rack
• Dropper
- Range of household
substances (Eg: ammonia,
lemon juice, bicarb soda,
washing powder, white vinegar
- Procedure:
1. Put cabbage leave sin
beaker and slightly cover
with water
2. Slowly boil until water turns
dark and leaves loose their
colour
3. Allow to cool and decant the
liquid (this is the plant
indicator)
4. Spread plant indicator
between test tubes (Having
ONE control)
5. Add a household substance
to each test tube and record
the colour change
6. Classify substances as
acidic, basic or neutrel
Identify data and choose resources to gather information about
the colour changes of a range of indicators
Solve problems by applying information about the colour
changes of indicators to classify some household substances as
acidic, neutral or basic
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- Jif
- Phenolphth
alein
- Magenta
- Universal Indicator
- Purple
- Methyl Red
- Bromothyothl
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©
- Yellow
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- Facial wash
- Shampoo
- No change
- No change
- Orange
- Red
- No change
- Red
- Yellow
- Yellow
Discussion and conclusionThe results are were relatively accurate, although some of the
substance mixed together making the results inaccurate.
Through following the method it was discovered that Jif, the
washing powder, is alkali, the facial wash is slightly acidic and
the shampoo is acidic. While using the Phenolphthalein and
Methyl red there was no change when it came into contact with
two substances
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•
ConclusionWe have achieved discovering the pH of household substances
using different indicators.
2.
While we usually think of the air around us as neutral, the atmosphere
naturally contains acidic oxides of carbon, nitrogen and sulfur. The
concentrations of these acidic oxides have been increasing since the
Industrial Revolution
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Identify oxides of non-metals which act as acids and describe the
conditions under which they act as acids
Analyse the position of these non-metals in the Periodic Table and
outline the relationship between position of elements in the Periodic
Table and acidity/basicity of oxides
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Acidic Oxides:
•
React with water
to form an acid
And/or
React with bases
to form salts
Are generally
oxides of nonmetals (covalent
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oxides)
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Basic Oxides:
React with acids to form salts
Does not react with basic (alkali solutions
Are generally oxides of metals (ionic oxides)
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Amphoteric oxides:
React with acids to form salts
React with alkalis eg, ZnO ,PbO, Al2O3
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Neutral oxides:
Do not react with acids
Do not react with bases
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Define Le Chatelier’s principle
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Le Chatelier’s principle: if a system at equilibrium is disturbed, then the
system adjusts itself so as to minimise the disturbance
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Identify factors which can affect the equilibrium in a reversible reaction
Temperature change
Pressure change
Concentration change
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If one of these changes occur the system will adjust as to minimise this
change, by reacting in one specific way.
The effect of changing temperature and pressure can be regarded as
moving the ‘position of equilibrium’
‘Position of equilibrium’ or ‘equilibrium position’ means the extent to
which the reaction has gone forward or reversed in direction
If the equilibrium position ‘lies to the left’, then only some of the
reactants have reacted to form the products
If the equilibrium position ‘lies to the right’, then most of the reactants
have reacted to form the products
Describe the solubility of carbon dioxide in water under various
conditions as an equilibrium process and explain in terms of Le
Chatelier’s principle
CO2(g)+H2O(l)
Endothermic
H2CO3(aq)
Exothermic
This equilibrium is disturbed if:
• The concentration or pressure changes of one species is changed
• The total pressure has changed
• The temperature is changed
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e.g. what would raising the temperature of the system do?
(constant pressure is assumed)
It would move to the left
Identify natural and industrial sources of sulfur dioxide and oxides of
nitrogen
describe, using equations, examples of chemical reactions which
release sulfur dioxide and chemical reactions which release oxides of
nitrogen
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
Sulfur Dioxide:
Industrial sourcesThe burning of fossil fuels, this is because most coal contains sulfur
and it is not practical to extract the sulfur prior to burning it. The sulfur
is therefore converted to sulfur dioxide when it is burnt with the coal, as
is shown in the following reaction:
S + O2(g)
SO2
Crude oil also contains sulfur compounds. These can be extracted and
sold to sulfuric acid manufacturers, although the extraction process
results in the release of some sulfur into the atmosphere. Also, some
sulfur can be left in the gas or crude oil and is later released into the
atmosphere.
Also the extraction of metals from sulfide ores. Many metals, such as
Cu, Zn, Ag and Ni, are found in sulfide ores. The first step in the
extraction process involves roasting the sulfide in air, resulting in the
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release of sulfur dioxide. For example in the extraction of Cu from it’s
ore: 2ZnS(s) + 3O2(g)
2ZnO(s) + 2SO2(g)
The manufacturing of sulphuric acid
Refining of petroleum
Natural Sources- About two thirds of the sulfur dioxide released world
wide originates from:
Geothermal hot springs
Volcanoes
Decay of organic material
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Oxides of Nitrogen:
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Nitric Oxide (NO) and Nitrogen dioxide (NO2)
Natural Sources• NO is formed as a result of the high and concentrate temperatures
formed as a result of lightning. As is demonstrated by the following
reaction:
O2(g)+N2(g)
2NO(g)
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The NO then gradually reacts with the oxygen found in the
atmosphere to
form NO2: 2NO(g)+O2(g)
2NO2(g)
Industrial sources• NO and NO2 are both formed as a result of combustion, from both cars
and power stations. Due to the high temperature the oxygen and
nitrogen found in air to combine to form NO, just as they do due to
lightning in the environment. Also, just as stated above, the NO is
gradually changed to NO2.
Nitrous oxide
Natural sources• Bacteria acts upon nitrogenous materials in the soil in a way that
results in the formation of nitrous oxide
• Industrial sources• The use of nitrogenous fertilizer adds to the amount of nitrous oxide
produced by bacteria, therefore increasing the amount of nitrous oxide
released into the atmosphere
•
SO2 and NO2 can have a negative impact on barley, impacting upon
its growth patterns. This could have a negative impact on food
production and damage the economy. NO2 can cause irritation to the
lungs as well as lowering the body’s resistance to respiratory infections
such as influenza. Regular exposure in children may result in acute
respiratory illness. SO2 also has negative health affects, including
respiratory illness. Asthmatics and those with chronic lung disease or
cardiovascular disease, as well as children and elderly, are particularly
receptive to SO2.
•
Assess the evidence which indicates increases in atmospheric
concentration of oxides of sulfur and nitrogen
Evidence of the change in atmospheric sulfur oxides and nitrogen
oxides is not easy to find, due to the following:
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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
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n= L/MolL-1
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Mole: is the quantity of a substance that contains 6.022 x 1023
particles
Molar Mass: the formula mass of a substance in grams
Molar Volume: volume in litres, occupied by one mole of any gas at a
particular temperature and pressure.
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Formulas
n=m
M
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n = Volume (L)
22.71 or 24.79 (Molar volume)
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n= concentration (mol/L) x volume (L)
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n= N
NA
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©
There is only a very small amount of these gasses, 0.001 ppm in
populated areas. The instruments required to measure these very low
concentrations of the gasses have only been available since the 1970s
When SO2 dissolves in water it forms sulfate ions and NO2 forms
nitrate ions. Sulfate and nitrate ions are usually soluble in water, they
circulate the biosphere and hydrosphere being chemically changed in
the process. This means that SO2 and NO2 are very hard to measure
from water samples.
Despite this there is some evidence to suggest the atmospheric
concentration of oxides of sulfur and nitrogen has increased. This
evidence comes from analysis of air bubbles that have been trapped in
Antarctic ice; measurement of carbon isotopes in old trees and grass
seeds found in museums, as well as calcium carbonate found in coral.
The amount of acid oxides released into the atmosphere has increased
dramatically since the industrial revolution. Large amounts of SO2
and NO2 dissolve in rainwater and are as a result washed out of the
atmosphere with rain. As a result there is not a significant build up of
these gasses over time, compared with CO2 and Nitrous oxide which
have built up to a huge extent. Overall, the evidence is inaccurate due
to the difficulties involved in collecting data. Due to technological
advances, the evidence collected in recent times is considerably more
reliable, however, there is a lacking of reliable past data to compare it
to in order to see changes over time.
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Gas pressure is due to constant bombardment of the walls of the
container by moving molecules. A rise in temperature increases the
kinetic energy of the molecules, which results in more collisions
therefore a rise in gas pressure.
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Pressure equals force per unit area. SI unit is Pascal (Pa)
Pressure can be used using a manometer. A manometer used only to
measure atmospheric pressure is called a barometer.
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Molar Volume of a Gas
According to Avogadro, any gas occupies the same volume at the
same temperature and pressure. The molar volume of a gas is the
volume of one mole of any gas.
At 100 kPa one mole of any gas will occupy 22.71 litres at 00C and at
250C it will occupy 24.79 litres
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• Examples:
1) Calculate the mass of 1 litre of NO gas at 00C and 100kPa.
- n=V
- n=m
22.71
M
- =1
= 0.044 x (14.01 +16)
- 22.71
- m = 1.32 g
- n = 0.044 moles
2) Calculate the volume occupied by 2.0g of CO2 gas at 250C and 100 kPa.
- n=m
- n=V
=2
24.79
- (12.01 + 32)
- V = 0.045 x 24.79
- n = 0.045 moles
- V = 1.127 litres
3) S (s) + O2 (g)  SO2 (g)
• What volume of sulfur dioxide forms when 50g of sulfur is burnt at
250C and 100kPa?
- n=m
- n=V
M
24.79
- = 50
- V = 1.559 x 24.79
32.07
= 38.65 litres
- =1.559 moles
Therefore 1.559 moles of SO2 are formed.
4) a) How many litres of CO2 gas form at room temperature and pressure in
the complete combustion of 0.5 moles ethane?
• C2H6 + 3.5O2  2CO2 + 3H2O
• Therefore 1 mole of ethane gives 2 moles of CO2.
• n=V
•
24.79
• V = 1 x 24.79
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V = 24.79 L
b) What mass of water vapour would be produced?
• n=m
•
M
• m = 1.5 x (1.008 x 2 +16)
• m = 27.024 g
• 5) Sulfur dioxide is produced by the smelting of copper sulfide ores.
• a) Write a symbolic equation to show the smelting of copper sulfide.
• CuS (s) + O2 (g)  SO2 (g) + Cu (s)
b) Calculate the number of moles present in 980kg of copper sulfide.
• n=m
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M
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= 980 000
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(63.55 + 32.07)
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= 10 248.9 moles
c) Calculate the number of moles produced by the smelting of 980kg of
copper sulfide.
• 10 248.9 moles of CuS (s)
• Therefore SO2 = 10 248.9 moles
d) What volume would this mass sulfur dioxide occupy at 250C?
• n=V
•
24.79
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V= 10 248.9 x 24.79
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= 254 070 L
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Explain the formation and effects of acid rain
Acid rain is rain that has a higher hydrogen ion concentration than
normal, unpolluted rain. Unpolluted rainwater therefore has a pH of
about 5.6, due to atmospheric CO2 dissolving in the rain water to form
carbonic acid. Acid rain is formed as a result of the exposure of
rainwater to, CO2, NO and SO2. The acid that is generally present in
acid rain is sulfuric acid and nitric acid.
NO slowly reacts in air to form NO2. NO2 then reacts with water to
form nitric acid, HNO3. As is demonstrated through the following
reaction:
3NO2(g) + H2O(l)
2HNO3(aq) + NO(g)
The nitric acid then disassociates in water, forming nitrate and
hydrogen ions.
SO2 reacts with water to produce sulphurise acid (it is an acid-oxide)
and can be oxidised to form sulphur trioxide:
SO2(g) + H2O(l)
H2SO3(aq)
2SO2(g) + O2(g )
2SO3(g)
SO3(g) + H2O(l)
H2SO4(aq)
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Sulfuric acid, being a strong acid, disassociates in water to give of
hydrogen ions:
H2SO4(aq)
HSO4- + H+
HSO4SO42- + H+
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There are many negative affects of acid rain both on society and the
environmentSociety:
Acid rain causes the erosion of marble and limestone, as both marble
and limestone contain CaCO3. The following reaction shows how
CaCO3 is eroded by sulfuric acid:
CaCO3(s) + H2SO4(aq)
Ca2+(aq) + SO4-2(aq) + H2O(l) +
CO2(g)
This results in the damage of building surfaces and decorations.
Environment:
Acid rain has lead to an increase in the acidity of many lakes, having
negative affect on organisms that inhabit these areas. It may also
release toxic ions such as Al3+ into the lakes or rivers.
Many pine forests in North America and parts of Europe have been
damaged as a result of acid rain. Acid rain has caused damage to
vegetation in areas around mines and smelter sites as they release
CO, NO and NO2.
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Aim: To decarbonate soft drink, measure the mass changes involved
and calculate the volume of gas released at 250C at 100kPa.
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Method:
1. Weigh the soda water bottle before opening as well as the beaker.
2. Open the soda water bottle and pour it into the beaker.
3. Weigh the empty bottle.
4. Place approximately the same amount of water into another beaker
and weigh (the water in the beaker is to be used as a control so that
the weight loss due to the evaporation so that the water can be
estimated)
5. Leave the two beakers out in the lab covered lightly with a piece of
towel so there is no chance of inaccuracy (eg: Flies falling in)
6. Weigh the beakers at regular intervals over several days.
3.
Acids occur in many foods, drinks and even within our stomachs
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©
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 100kPa
Decarbonation is the removal of dissolved CO2 from solution (soft
drink).
Define acids as proton donors and describe the ionisation of acids in
water
Acids are proton donors because they give of protons. In spolution they
give of H+, which contain one proton, hens they are proton donors. The
H+ ion then combines with the H2O in water to form hydronium- H3O+.
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When an acid dissolves in water it ionises, i.e. it breaks up into ions.
For example:
HCl(g) + H2O(l)
H3O+(aq) + Cl-(aq)
It is the presence of the hydronium ion which makes the solution acidic
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Water is a very weak electrolyte because it has slight ionisation.
H2O
H3O+ + OHThis is called the self ionisation of water.
It is an equilibrium reaction that lies well to the left.
This self ionisation also occurs in aqueous solutions such as an acid,
base or salt solution.
At any given temperature the product of the hydrogen and hydroxide
ions is constant. At 250C it is always 1 x 10-14 (mol/L)2 OR 10-14 and
the symbol used is KW. This is called the ionic product or ionisation
constant of water.
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Identify acids including:
Acetic (ethanoic)
CH3CHOO
This is a weak acid
Citric (2-hydroxypropane-1,2,3-tricarboxylic)
This is a weak acid
Hydrochloric
HCl
This is a strong acid
Sulfuric acid
H2SO4
This is a strong acid
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Acid
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Formula
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Found In
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Used For
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Hydrochlor
ic Acid
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HCl
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Stomach
(used to
digest
protein),
also man
made
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Making
dyes
Cleaning
bricks
To digest
protein
Food
Metallurg
y and oil
industrie
s
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Sulfuric
Acid
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H2SO4
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Man
made but
found
naturally
in acid
rain
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Car
batteries
Fertiliser
s
Petroch
emicals
Metal
treatmen
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Acetic Acid
Citric Acid
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CH3COOH
C6H8O7
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Vinegar
Citrus fruit
(oranges
and
lemons)
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Phosphoric
Acid
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H3PO4
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Nitric Acid
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HNO3
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t
Cooking
Food
Preserva
tion
Plastics
Solvents
Eating
-Food
flavourin
g
Cleanin
g
products
Fertiliser
s
Metal
treatmen
t
Cola
drinks
Fertilise
rs
Explosiv
es
Dyes
Describe the use of the pH scale in comparing acids and bases
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pH = power
hydrogen
• <7= acid
>7= base
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• It is a numerical scale
which allows comparison between acids and bases to be made
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Describe acids and their solutions with the appropriate use of the terms
strong, weak, concentrated and dilute
WeakDo not disassociate in water to from hydrogen ions
StrongDisassociate in water to from hydrogen ions
Monoprotic- releases one proton (H+) e.g. HCl
Diprotic- releases two protons (H+) e.g. H2SO4
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Triproticreleases
three protons
(H+) e.g.
H3PO4
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Concentrate
d-
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There is a large amount of the acid
Dilute- there is a small amount of the acid
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Identify pH as -log10 [H+] and explain that a change in pH of 1 means
a ten-fold change in [H+]
Process information from secondary sources to calculate pH of strong
acids given appropriate hydrogen ion concentrations
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pH = -log10(H3O+)
(H3O+) = 10-pH
At 25o C: (H3O+2) (OH-) = Kw = 1.00 X10-14
Kw = the ionic product of water
Using this equation, if the (H+) is known, then this can be used to
calculate the (OH-)
A change of 1 in pH means a ten-fold change in [H+]
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H
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[H+]
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[OH-]
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[H+] x [OH-]
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0
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100 = 1
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10-14
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10-14
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1
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10-1
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10-13
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10-14
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2
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10-2
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10-12
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10-14
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3
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10-3
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10-11
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10-14
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4
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10-4
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10-10
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10-14
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5
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10-5
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10-9
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10-14
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6
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10-6
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10-8
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10-14
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7
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10-7
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10-7
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10-14
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8
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10-8
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10-6
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10-14
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9
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10-9
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10-5
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10-14
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10
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10-10
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10-4
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10-14
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11
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10-11
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10-3
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10-14
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12
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10-12
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10-2
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10-14
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13
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10-13
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10-1
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10-14
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14
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10-14
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100 =1
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10-14
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
HCl is a strong acid- close to 100% is ionised in solution:
HCl(aq)
H3O+(aq) + Cl(aq)
Citric acid and acetic acid are weak acids, they are only 1% ionised in a
solution:
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Describe the difference between a strong and a weak acid in terms of
an equilibrium between the intact molecule and its ions
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Strong acid- all molocules present ionise when placed in water
Weak acid- only a small amount of the molocules ionise when placed in
water
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If you prepare 0.1 M solutions of each you should find the hydrochloric
acid has a pH of 1 corresponding to a [H+] = 10-1 = 0.1 M.
That is, practically every HCl molecule has ionised, producing a H+.
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By contrast a 0.1 M solution of acetic acid will have a pH close to 3
indicating a [H+] close to 10-3 = 0.001 M. Only about 0.001 / 0.1 = 1%
of the acetic acid molecules have ionised producing a H+.
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Solve problems and perform a first-hand investigation to use pH
meters/probes and indicators to distinguish between acidic, basic and
neutral chemicals
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Aim- to determine the acidity or otherwise of an acid, base and a
neutral solution using a probe/meter
Results•
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HCl
NaOH
Distilled
H2O
Unknown
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1.35
12.61
6.80
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12.58
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pH
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DiscussionOur results were relatively accurate because a data logger is
moderately accurate, also we rinsed the data logger in distilled water
between each measurement
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ConclusionHCl is acidic, NaOH is alkali, distilled water is slightly acidic, the
unknown substance is alkali.
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Plan and perform a first-hand investigation to measure the pH of
identical concentrations of strong and weak acids
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Aim: To compare the pH of different acids of equal and varied
concentrations.
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Equipment:
- 25mL pipette and filler
- 6 x 250mL volumetric flask
- I mol/L HCl 25mL
- I mol/L CH3COOH 25mL
- pH meter
- Distilled water
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Method:
1) Fill the pipette with 25mL of HCl.
2) Pour into volumetric flask 1and make up to 250mL. Measure pH
(0.01 mol/L)
3) Fill pipette with 25mL of solution from volumetric flask 1.
4) Pour into volumetric flask 2 and use distilled water to make up to
250mL. Measure the pH (0.001 mol/L)
5) Fill pipette with 25mL of solution from volumetric flask 2.
6) Pour into volumetric flask 3 and use distilled water to make up to
250mL. Measure pH (0.0001 mol/L)
7) Repeat steps 1-6 with acetic acid
8) Tabulate and compare results
Note: Don’t forget to rinse the pipette and pH meter after every use
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•
•
•
©
•
•
Safety:
HCl is corrosive, thus avoid contact with skin thus wear gloves,
goggles, and lab coats.
•
•
•
•
•
•
•
Variables:
Independent: concentration of each acid
Dependent: pH
Controlled:
- Molarity of original acids
- pH meter used
- Volume of flask used
15 of 36
Safety-
•
©
•
Identify
•
Describe
•
Explain
•
Take care when
diluting an acid.
•
When diluting an
acid, always add
acid to water,
NEVER add
water to acid.
•
Dilution of acids
is an exothermic
reaction. A great
deal of heat can
be released. If
water is being
added to acid,
this heat may
make th acid boil
and spit.
Spattering the
person with
concentrated
acid which is
corrosive and
thus can burn
skin and
benches. If acid
is being added to
water, the water
will boil and spit.
•
Cover up when
using acids,
especially
concentrated
acids
•
Acids are
corrosive and
may get on skin
or in eyes.
•
Acids may
splash into eyes
or be spilled on
skin or clothing
and because it’s
corrosive, it may
burn. Always
wear protective
clothing,
including safety
glasses when
using acids in
order to prevent
acids coming
into contact with
skin or eyes.
16 of 36
•
Gather and process information from secondary sources to write ionic
equations to represent the ionisation of acids
•
A strong acid ionises completely, donating protons freely.
Eg: HCl(g) + H2O(l)
H3O+(aq) + Cl-(aq)
All of the hydrochloric acid is present as ions; there are no neutral acid
molecules present.
•
•
•
A weak acid does not ionise completely; it is not a good proton donor.
Eg: CH3COOH(aq) + H2O(l)
H3O+(aq) + CH3COO-(aq)
This reaction is an equilibrium reaction. It does not go to completion.
Only a small percentage of the molecules ionise; the rest remain as
molecules of acetic acid.
•
Use available evidence to model the molecular nature of acids and
simulate the ionisation of strong and weak acids
•
Here you can build models of acid molecules and remove the
appropriate H to simulate ionisation. Note that in organic acids, the H
that ionises comes from a –COOH group. The two electronegative
oxygen atoms attract the pair of electrons in the covalent bond
between the O and the H. This weakens the bond so that sometimes
the H can break away as a H+ ion, leaving an electron from the H on
the –COO– , thus providing the charge of –1.
•
Gather and process information from secondary sources to explain the
use of acids as food additives
•
Acids such as vinegar have been used as food preservatives for
thousands of years, acid oxides such as SO2 are also widely used.
They prevent the growth of microbes, as they can not grow in acidic
environments.
(Any product with food codes 220 to 225 or 228 has one of these acidic
substances added: SO2 and H2SO3 or a sulfite, bisulfite or
metabisulfite salt)
•
•
•
•
•
•
•
The concentration of the additive decreases with time and this limits
the shelf life of the food.
Examples of food preservatives:
Sulfur dioxide SO2
• Sorbic acid
CH3CH=CHCH=CHCOOH
Lactic acid CH3CHOHCOOH
• Acetic acid CH3COOH
Nitrous acid HNO2
• Propionic (propanoic) acid
Formic acid HCOOH
CH3CH2COOH
Benzoic acid C6H5COOH
Everyday examples:
Dried apricots 220
Soy sauce (contains sulfites)
Apricot flavoured bars 220
•
Antimicrobials
•
•
•
•
•
•
•
©
17 of 36
©
•
Preservatives that prevent spoilage of food from some microorganisms:
bacteria, moulds, fungi and yeast. They can extend shelf life of a food
product and can protect the natural colour and flavour of food.
•
Identify data, gather and process information from secondary sources
to identify examples of naturally occurring acids and bases and their
chemical composition
•
Name
•
Stucture
•
Acid/Base
•
•
Sulfuric
Acid
•
H2SO4
•
Acid
•
•
Acetic
Acid
•
CH3COOH
•
Acid
•
•
•
Ammonia
•
NH3
•
Base
•
18 of 36
pH in
natural
occurring
form
•
Sodium
carbonate
Na2CO3
•
•
Base
•
•
©
•
Acid
•
Formula
•
Anion
•
Anion
name
•
Hydrofluoric
•
HF
•
F-
•
Fluoride
•
Hydrochloric
•
HCl
•
Cl-
•
Chloride
•
Hydrobromi
c
•
HBr
•
Br-
•
Bromide
•
Hydroiodic
•
HI
•
I-
•
Iodide
19 of 36
©
•
Sulfuric
•
H2SO4
•
SO42-
•
Sulfate
•
Nitric
•
HNO3
•
NO3-
•
Nitrate
•
Sulphurous
•
H2SO3
•
SO32-
•
Sulfite
•
Nitrous
•
HNO2
•
NO2-
•
Nitride
•
Carbonic
•
H2CO3
•
CO32-
•
Carbonat
e
•
Phosphoric
•
H3PO4
•
PO43-
•
Phosphat
e
20 of 36
©
•
Formic
(Methanoic)
•
HCOO
H
•
HC00-
•
Formate
(Methano
ate)
•
Acetic
(Ethanoic)
•
CH3CO
OH
•
CH3COO-
•
•
Acetate
(Ethanoat
e)
•
Hydrocyanic
•
HCN
•
CN-
•
Cynide
•
Hydrogen
Sulphide
•
H2S
•
S-2
•
Sulfide
21 of 36
•
Acid
•
Formula
•
Natural Form
•
Formic
Acid
•
HCOOH
•
Ant, bee stings. It is
used in textile
treatments and
tanning hides.
•
Ascorbic
Acid
•
C6H8O6
•
Vitamin C. It is used
as antioxidant in food
preservation.
•
Base
•
Formula
•
Natural Form
•
Calcium
Carbonate
•
CaCO3
•
Limestone
•
Nicotine
•
C8H14N2
•
Liquid found in
tobacco leaves.
•
©
22 of 36
•
4. Because of the prevalence and importance of acids, they have been
used and studied for hundreds of years. Over time, the definitions of
acid and base have been refined
•
Putline the historical development of ideas about acids including those
of: Lavoisier Davy and Arrhenius
Outline the Brönsted-Lowry theory of acids and bases
•
•
©
Chemist
•
Time
•
Definition
of an acid
•
Problems
•
Lavoisier
•
1780
•
Acids are
substance
s that are
formed
from nonmetals and
contain
oxygen
•
Some acids,
like HCl, do
not contain
oxygen
•
Davy
•
1815
•
Acids are
substance
s that
contain
replaceabl
e
hydrogen
(hydrogen
that could
be
partly/totall
y replaced
by a metal)
Bases
were
substance
s that
reacted
with acids
to form
salt and
water.
•
This worked
well for most
of that century
23 of 36
•
Arrhenius
•
1884
•
•
•
BronstedLowry
•
1923
•
•
•
©
24 of 36
Acids are
substance
s that
provide
hydrogen
ions as the
only
positive
ion in an
aqueous
solution
Basis are
substance
s that
provide
hydroxide
ions as the
only
negative
ion in a
solution
•
This restricted
reactions to
those in an
aqueous
solution
•
This failed to
take into
account
insoluble
hydrogen
compounds
•
This did not
include
substances
like
carbonates
which react to
form
hydroxide ions
in water
An acid is
a
substance
that, in
solution,
tends to
give up
protons
(hydrogen
ions), and
a base is a
substance
that tends
to accept
protons
A
substance
s acidic
properties
are
relative to
those of
the solvent
•
If the
substance
HA, has a
greater
tendency
to give up
protons
than the
solvent,
then that
substance
in that
solvent will
be an acid
•
•
The Brönsted-Lowry Theory of Acids and Bases:
Acid: a substance that donates a proton (H+).
•
•
Conjugate Acid-Base Pairs
The term conjugate acid base pair connects two species that differ by
one proton, H+.
In the Brönsted-Lowry theory every acid has a conjugate base (a
species that has one proton less than the acid.)
Example:
NH4+(aq) + HCOO-(aq) NH3(g) + HCOOH(aq)
Acid
Base x
Conjugate Conjugate
Base
Acid
 A strong acid has a weak conjugate base
 A strong base has a weak conjugate acid
 A moderately weak acid has a weak conjugate base.
•
•
•
•
•
•
•
•
•
•
•
•
Describe the relationship between an acid and its conjugate base and
a base and its conjugate acid
•
•
•
An acid gives up a proton form a conjugate base:
HCl + H2O
Cl- + H3O+
Acid
base
Every acid has a conjugate base.
•
When a base accepts a proton, it forms a conjugate acid:
•
Together, the acid and base are called a conjugate pair
•
Identify a range of salts which form acidic, basic or neutral solutions
and explain their acidic, neutral or basic nature
•
Salt ions that are formed from weak acids or bases can react with
water to reform the acid or base. In this hydrolysis reaction they
release OH-or H+ ions, which can produce basic or acidic salt solutions.
•
©
Also as part of the Bronson-lowry theory:
Monoprotic acids (HCl) donate one proton per molecule
Diprotic acids (H2SO4) donate two protons per molecule
Triprotic acids (H3PO4) donate three protons per molocule
25 of 36
•
•
This then result in the reaction: H3O+ + OH- 2H2O
This causes a neutral solution
•
Ammonium salt solutions are acidic. This is because:
NH4+ + H2O
NH3 + H3O+
Sodium Chloride solution is neutral, because Na + and Cl- do not under
go hydrolysis
Sodium carbonate solution is basic, because the carbonate ion from
the weak acid carbonic acid can hydrolyse: CO32- + H2O HCO3- + OH-
•
•
•
•
Potassium acetate solution is basic: CH3COO- + H2O
OH-
•
Identify conjugate acid/base pairs
•
Some examples are:
•
©
Acid
•
Base
•
HCN
•
CN-
•
HCl
•
Cl-
•
HSO4-
•
SO4-2
CH3COOH +
26 of 36
NH4+
•
NH3
•
HClO4
•
ClO4
•
Identify amphiprotic substances and construct equations to describe
their behaviour in acidic and basic solutions
•
•
•
An amphiprotic molecule or ion can donate or accept a proton.
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
©
•
Water is an amphiprotic molecule:
Water as an acid: H2O
H+ + OH- or more fully H2O +H2O H3O+
+ OHWater as a base: H+ + H2O
H3O+ or more fully H3O++H2O H2O
+ H3O+
Bicarbonate ions are amphiprotic ion:
As an acid HCO3- H+ + CO32- or more fully HCO3-+H2O
H3O+
+ CO32As a base H+ + HCO3H2CO3 or more fully H3O++ HCO3H2O +H2CO3
Identify neutralisation as a proton transfer reaction which is exothermic
Neutralisationacid + base
salt + water
e.g. HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
In this example the acid HCl transfers a proton to the base NaOH,
specifically the OH-part of it
H3O+ + OH-(aq)
2H2O
This is an exothermic reaction
Equilibrium in acid base reactions
A strong acid or base is completely ionised and is therefore not an
equilibrium reaction.
e.g. HCl + H2O
Cl- + H3O+
27 of 36
•
•
•
Describe the correct technique for conducting titrations and preparation
of standard solutions
•
A volumetric analysis is a form of chemical analysis in which the
concentration or amount of substance A is determined by measuring
the volume of solution of a known concentration of another substance
B which is just sufficient enough to react with all of the sample A.
Titration is the process the ‘just sufficient’ volume
•
•
•
The equivalence point: the point of a chemical reaction at which the
amount of two reactants are just sufficient to cause complete
consumption of two reactants.
This point is often determined by using an indicator. The indicator must
be chosen so that the pH at the end point (when the indicator changes
colour) is as close as possible to the equivalence point.
•
•
•
•
•
•
•
•
•
•
For a strong acid-strong base titration:
Litmus (6 to 8)
Bromothyle blue (6.2 to 7.6)
For a strong acid- weak base titration:
Methyl orange (3.1 to 4.4)
Bromophenol blue (3.0 to 4.6)
Bromocresol green (3.8 to 5.4)
For a weak acid-strong base titration:
Phenolphthalein (8.3 to 10.0)
Thymol blue (8.0 to 9.6)
•
•
Method in acid-base titrations:
Wash the burette with water and then rinse with a little of the solution to
be placed in it
Fill the burette with a solution of known concentration (solution A) and
adjust the level of solution to the zero mark
Wash the pipette with water and then rinse with a little of the solution
(solution B) to be placed in it
Wash the conical flask with water
Use the pipette to transfer an exact volume, and hens an exact number
of moles of solution B into the conical flask
Add one or two drops of indicator to the conical flask
Slowly run solution A from the burette into the conical flask with
continuous swirling and washing of the inside of the conical flask until
the indicator just changes colour
Accurately record the volume delivered from the burette into the flask
Calculate the concentration of solution B
•
•
•
•
•
•
•
•
•
©
In contrast, a weak acid is only partly ionised and is there fore at
equelibrium.
e.g. CH3COOH + H2O
CH3COO- + H3O+
A primary standard used in volumetric analysis is a substance of
exactly known concentration, it must:
28 of 36
•
•
•
•
•
•
•
Qualitatively describe the effect of buffers with reference to a specific
example in a natural system
•
A buffer solution is usually a mixture between a weak acid and it’s
cunjugate base. For example: HCO-3 and CO3-2
•
A buffer controls the pH of a solution. If an acid or base is added to a
buffer solution there is hardly any change in pH.
•
•
•
•
•
©
Be obtainable in pure form
Cannot absorb water or other substances from the atmosphere during
weighing or when in use
It must be a crystalline solid and hence able to accurately be weighed
It must have a relatively high molar mass and so the experiment is very
accurate and the % error is very small
Sodium carbonate and oxalic acid make suitable primary standards for
titrations.
Sodium hydroxide and hydrochloric acid are not suitable primary
standards. NaOH absorbs water and carbon dioxide from the air. HCl,
obtained commercially from muriatic acid, is approximately 10 molar.
If an acid is added to a buffer the hydrogen ions are removed by:
2H(aq)+ + CO3(aq)-2
H2CO3(aq)
If a base is added to a buffer the hydroxide ions are removed by:
OH- + HCO3H2O +CO3-2
•
Gather and process information from secondary sources to trace
developments in understanding and describing acid/base reactions
•
Choose equipment and perform a first-hand investigation to identify the
pH of a range of salt solutions
•
•
Salts: ionic compounds.
Salt solutions can be acidic, basic or neutral depending on the
tendencies of the ions in the salt to undergo hydrolysis.
•
•
•
•
Aim:
To identify the pH of salt solutions.
Risk Assessment:
Salt solutions may be acidic as well as corrosive to skin, thus wear
safety goggles and gloves. They are also very dangerous to marine
organisms as they are acidic, thus don’t tip chemicals down the sink.
Salts are poisonous so don’t swallow them.
29 of 36
•
Perform a first-hand investigation to determine the concentration of a
domestic acidic substance using computer-based technologies
•
A data logger stores data input electronically (data can be displayed on
a graphics screen or transferred to a computer.) Probes are available
for attachment to many data loggers to measure pH and electrical
conductivity.
These can be used to measure the pH changes during a titration and
as they are far more reliable than a indicator
•
©
•
Indicator
•
Colour on
Acidic
Side
•
Range of
Colour
Change
•
Colour on
Basic Side
•
Methyl
violet
•
Yellow
•
0 - 1.6
•
Violet
•
Bromoph
enol blue
•
Yellow
•
3 - 4.6
•
Blue
•
Methyl
orange
•
Red
•
3.1 – 4.4
•
Yellow
•
Methyl
red
•
Red
•
4.4 –
6.26
•
Yellow
30 of 36
Litmus
•
Red
•
5–8
•
Blue
•
Bromothy
mol blue
•
Yellow
•
6 – 7.6
•
Blue
•
Phenolpht
halein
•
Colourless
•
8.3 – 10
•
Pink
•
Alizarin
yellow
•
Yellow
•
10.1 - 12
•
Red
•
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 widely used for safety in laboratories and
factories where acids or bases are used. Because many acids and
alkalis are very corrosive, it is important to neutralise any spills of these
substances quickly.
•
•
•
•
Sodium carbonate is widely used to neutralise acidic spills because:
It is a stable solid, which is easily and safely handled and stored
It is the cheapest alkali available
If too much of it is used there is less danger than from excess sodium
hydroxide or lime (calcium hydroxide)
As it is amphiprotic it can be used for acid and base spills
•
©
•
31 of 36
•
•
•
•
•
•
•
Factors that need to be considered in choosing a substance to
neutralise acid or alkali spills in factories or laboratories, need to
consider:
The speed of the reaction for neutralising the spill material
The need for a regent that will not have any harmful effect if any excess
of it is used (since it is hard to determine exact quantities, for
neutralising spills)
The safety in handling and storing the reagent
The cost of the reagent
The possibility of one reagent being able to neutralise both acid and
alkali spills
Therefore, neutralisation is valuable but only when you have the
neutraliser as you’re creating an environmentally friendly substance
(salt and water) which is easy to get rid of compared to the acid base.
5. Esterification is a naturally occurring process which can be performed in
the laboratory
•
Describe the differences between the alkanol and alkanoic acid
functional groups in carbon compounds
Identify the IUPAC nomenclature for describing the esters produced by
reactions of straight-chained alkanoic acids from C1 to C8 and straightchained primary alkanols from C1 to C8
•
Functional group- gives them there distinctive proporties
•
•
•
•
•
•
Alkanol: Contains a OH functional group
Naming- according to the chain length
PropertiesThe polar –OH group makes them polar
Hydrogen bonding can occur between the different molocules
Soluble in water, but this solubility decreases with the increasing
carbon chain length
Higher meling and boiling points than there corresponding alkanes and
alkenes
They react with alkanoic acids to make esters
•
•
•
•
•
•
•
•
•
©
Alkanoic acid- contains a -COOH functional group
Naming- they are named accoding to the number of carbon atoms.
Formic acid is the first. When they are names the ‘e’ of the
corresponding alkane or alkene is crossed of and ‘oic’ is added
PropertiesThey are polar
Strong hydrogen bonds between near by alkanoic molecules
They are soluble in water, but this decreases with the carbon chain
length
32 of 36
•
•
•
•
•
•
•
Esters- contain the ______ functional group
Naming- the first part of the name comes from the alkanol, the second
comes from the acid.
GIVE EXAMPLES
•
Explain the difference in melting point and boiling point caused by
straight-chained alkanoic acid and straight-chained primary alkanol
structures
•
The first few members of the alkane or alkene homologous series are
gases at room temperature. Intermolecular bonds do not form to the
same extent as they do for other substances of similar masses that are
liquid or solid at room temperature.
Alkanols contain C-O and O-H bonds with are polar. Also alkanols can
form hydrogen bonds with other alkanols. So there are strong
intermolecular forces in alkanols. As a result there are no alkanols that
are gas at room temperature
Alkanoic acid contains polar C=O bonds. The intermolecular forces are
even greater than that of the corresponding alkanol, therefore, they
have higher boiling points.
Alkanoic acids contain polar group, C=O. They have higher boiling
points than the corresponding alkanols
•
•
•
•
Identify esterification as the reaction between an acid and an alkanol
and describe, using equations, examples of esterification
•
An ester is formed when an alkanol and an alkanoic acid join:
Alkanol + Alkanoic Acid Ester + water
•
•
•
•
•
Example:
Ethanol + Ethanoic acid/formic acid Ethyl ethanoate/ Ethyl formate +
water
HCOOH + CH3CH2OH HCOOCH2CH3 + H2O
The underlined atoms ate the ones that in equation join together to
form water
•
Describe the purpose of using acid in esterification for catalysis
•
Esterification is an equilibrium reaction, which occurs relatively slowly
at room temperature.
A few drops of a concentrated mineral acid, such as H2SO4 can be
added as a catalyst.
It speads up the reaction by:
•
•
©
They have a higher melting and boiling point than there corresponding
alkanes, alkenes and alkanols
They are weak acids
They react with alkanols to form esters
Salts of long chain alkanoic acids are soaps
33 of 36
•
•
•
•
•
•
Absorbing the water formed, as a result forcing the equilibrium to the
right.
A strong acid may donate a proton to the unshared oxygen electron
pairs of either the acid or the alkanol. This makes the alkanol and the
alkanoic acid more reactive. The proton is eventually returned to the
solution. This reaction can be written asAlkanoic acid + alkanol (H+, above the arrow) ester + water
Acid-catalysed esterification is reversible and generally at equilibrium
there are appreciable quantities of both the ester and alkanol present
Explain the need for refluxing during esterification
•
As an alternative for the use of a catalyst during esterification the
reaction mixture can be heated above 100o C, this removes the water
as it is forms, forcing the equilibrium to the right.
This heating is nessesary to cause the reaction to proceed to any
appreciable extent. The heating gives the reacting partivles a greater
kinetic energy and therefore more successful collisions, but in the
process, volatile substances may be lost.
•
Outline some examples of the occurrence, production and uses of
esters
•
•
•
•
•
•
Esters occur naturally as:
Flavouring agents and scents in plants and fruits
Animal fats and plant oils.
The production of esters involves either:
Extraction from natural sources.
Manufacture in an industrial environment, which is generally far less
expensive (esterification)
Uses of esters include:
Flavouring agents and scents in foods.
Solvents and thinners.
Medications.
Plasticisers.
•
•
•
•
•
•
•
•
•
•
ExampleOctyl ethanoate:
It is associated with orange flavour, as it is the main ester present in
oranges.
Production- reacting 1-octanol with ethanoic acid while concentrated sulfuric
acid acts as a catalyst and increases the yield
• Uses- as a flavouring agent in foods, such as orange-flavoured
confectionary.
•
©
Note- see http://www.ausetute.com.au/esters.html for production
process
34 of 36
•
Process information from secondary sources to identify and describe
the uses of esters as flavours and perfumes in processed foods and
cosmetics
•
Esters occur naturally in living things as animal fats and oils and give
perfume and taste to flowers and fruits.
Esters are used in processed foods and in cosmetics (eg: ethyl acetate
in nail polish remover) as they have strong flavours and odours. They
are also useful as solvents.
Examples of common natural flavours:
Apple: methyl butanoate
Orange: octyl thanoate
Pear: pentyl ethanoate
•
•
•
•
•
©
35 of 36
©
•
Product
•
Ester
•
Use of Ester
•
Nail polish
remover
•
Ethyl acetate
•
‘pear drop’
smell, solvent
•
Lipstick
•
2-propyl myristate
•
Reduce
stickiness
•
Perfume
•
Benzyl acetate
•
Jasmine scent
•
Honey flavoured
sweets
•
Methyl
pheylacetate
•
Honey flavour
36 of 36

Chemistry – The Acidic Environment The Acidic

Transcription

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CHEM-1105 THERMOCHEMISTRY 1.

EnJuvenate Pina Colada