Organic Chemistry
Structures
1
What do I need to know?
1. Translate between molecular, structural and ball and
stick representations of simple organic molecules
2. Describe how the functional group affects the
property of an organic compound and understand
that alkanes are unreactive towards aqueous
reagents because C—C and C—H bonds are
unreactive;
3. Write balanced chemical reactions including for
burning hydrocarbons including state symbols
2
Representations of organic molecules
• There are a number of different ways to represent
organic molecules.
• Ball and stick – this is just like molymods
3
Representations of organic molecules
• Structural formula – this is where we show the
covalent bonds between atoms as a line
• Semi-structural (molecular) – this is where we
write out the formula but do not include bonds;
these are implied eg CH3CH2OH
4
• Molecular formula – this simply counts the numbers of
each sort of atom present in the molecule, but tells you
nothing about the way they are joined together.
• Eg C2H6O
• This is the least helpful type of formula as it could be one
of two (or more) different chemicals
5
Example question
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Mark scheme
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Rules of organic molecules
Generally speaking
Carbon must make four bonds
Nitrogen must make three bonds
Oxygen must make two bonds
Hydrogen must make one bond
A double bond counts as two bonds eg C=C or C=O. A
triple bond counts as three bonds.
8
AfL - Quiz
1.
2.
3.
4.
5.
Draw the structural formula for butanol
Write the molecular formula for butanol
Draw the structural formula for hexane
Write the molecular formula for hexane
Write the molecular formula for an alkane with 25
carbon atoms.
6. How many bonds does oxygen make in methanol?
7. Give an example of a use for ethanol
8. Give an example of a use for methanol
9
1.
2.
3.
4.
5.
6.
7.
Butanol
C4H10O
Hexane
C6H14
C25H52
2
Fuel/feedstock for synthesis/solvent/used in
perfume
8. Solvent, antifreeze, feedstock for adhesives and
plastics
10
Understanding reactivity
• Alkanes are unreactive towards aqueous reagents
because C-C and C-H bonds are unreactive.
• What about organic molecules that have different
bonds?
• We call families of different types of bonded
atoms FUNCTIONAL GROUPS
• An example is the –OH group or alcohol group.
11
Different functional groups
Name
Functional group
Properties
Alkane
C-H
Relatively unreactive, burns
in air due to hydrocarbon
chain
Alkene
C=C
Used as a feedstock to make
polymers
Alcohol
-OH
Good solvent, volatile, burns
in air due to hydrocarbon
chain
Carboxylic acid
-COOH
Weak acid such as vinegar
Ester
RCOOR’
Have distinctive smells such
as fruits
12
Alkanes and combustion
• Because of the hydrocarbon chain alkanes burn
readily releasing large amounts of energy.
• Alkanes are therefore used as fuels.
• When they burn completely they make carbon
dioxide and water.
eg octane (found in petrol)
C8H18 +12 ½ O2 8CO2 + 9H2O
13
Example question
14
Mark scheme
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Example question
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Mark scheme
17
Balanced chemical equations
Write the balanced chemical equation for burning
ethanol in air as a fuel and burning pentane as a fuel
(include state symbols).
18
Answers
Ethanol
2C2H5OH(l) + 6O2(g)  4CO2(g) + 6H2O(l)
Pentane
C5H12(l)+ 8O2(g)  5CO2(g) + 6H2O (l)
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Example questions
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Mark scheme
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Alcohols and the Manufacture of
Ethanol
C7.1 and C7.5
22
What do I need to know?
1. The characteristic properties of alcohols are due
to the presence of an –OH functional group
2. Know a range of methods for synthesising ethanol
and limitations of fermentation reactions
3. Be able to explain why bioethanol is important for
sustainability
23
Functional groups - reminder
• Look back at your table of functional groups.
• Write a short paragraph to explain why different
organic chemicals have different properties in
terms of functional groups.
• Use examples such as “carboxylic acids are acidic
because they have a –COOH group”.
24
Can you recognise the functional group?
• Circle which of these are alcohols?
25
Answer
• Alcohols have an –OH group
26
Properties and uses of alcohols
Properties:
• volatile liquid (evaporates quickly at room
temperature – more than water)
• colourless
• burns readily in air because of the hydrocarbon
chain
• good solvent
27
Example question
28
Mark scheme
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Uses of ethanol and methanol
Ethanol: biofuels, solvents, feedstock for synthesis
Methanol: cleaner, feedstock for synthesis
Feedstock is the name we give to an “ingredient” on
a chemical plant
30
Reactions of different functional groups
• This is illustrated very well by comparing the
reaction of sodium with ethanol, hexane and
water.
• You have seen this reaction. Fill in the following
table and compare with the mark scheme:
31
Observations with sodium
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Mark scheme
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Comparing functional groups
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Mark scheme
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How do we make ethanol?
• Fermentation is a key process for obtaining
ethanol. It is relatively cheap and requires wheat
or beet sugar.
• The process involves the anaerobic respiration of
yeast at temperatures between 20 and 40°C and
at pH 7.
36
Conditions for fermentation
• Outside an optimum temperature the yeast does not work (high
temperatures kill the yeast).
• Outside an optimum pH the yeast does not work (extremes of pH kill the
yeast).
• To make ethanol the yeast must respire anaerobically (without oxygen).
• Eventually the ethanol concentration will be too high for the
fermentation to continue. This means only a dilute solution can be made.
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Example question
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Mark scheme
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Example question
40
Mark scheme
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Example question
42
Mark scheme
43
How do we obtain a concentrated
solution?
• Ethanol has a different boiling point to water. We
can therefore separate water and ethanol using
distillation.
44
Example question
45
Mark scheme
46
Making ethanol using ethane from crude
oil
Ethane to ethene by
CRACKING
C2H6  CH2=CH2
• zeolite catalyst OR
• heat
Ethene to ethanol by
reaction with STEAM
CH2=CH2 + H2O 
CH3CH2OH
47
• phosphoric
acid catalyst
Example question
48
Mark scheme
49
Working out masses
• We can use the useful relationship
Mass1 Mass2
=
Mr1
Mr2
• Where Mr is the molecular mass
• eg Mr of ethane C2H6 is (2 X 12) + (6 x 1) = 30
50
Example question
51
Explanation
• In this question every ethene molecule that reacts makes
one molecule of ethanol.
• We need to relate the number of molecules to mass using
our equation.
Mass1 Mass2
=
Mr1
Mr2
•
•
•
•
Mass 1 is mass of ethene = 1 tonne
Mr 1 is Mr of ethene = 28
Mass 2 is mass of ethanol = ?
Mr 2 is Mr of ethanol = 46
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Mark scheme
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Example question
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Mark scheme
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Other alternatives
• Ethanol has also been synthesised using
genetically modified e-coli bacteria and sugars
from seaweed.
• This process is sustainable as the seaweed and
bacteria are renewable sources
• Like yeast, bacteria can be killed by high
concentrations of alcohol and high temperatures
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Example question
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Mark scheme
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Ethanol – Key facts
• Ethanol is made on an industrial scale as a fuel, a solvent
and as a feedstock for other processes;
• There is a limit to the concentration of ethanol solution that
can be made by fermentation and there are optimum
conditions of pH and temperature.
• Ethanol solution can be concentrated by distillation to make
products such as whisky and brandy;
• Genetically modified E. coli bacteria can be used to convert
waste biomass from a range of sources into ethanol and
recall the optimum conditions for the process;
• Ethane from crude oil can be converted into ethanol
• Evaluating the sustainability of each process is important.
59
Bioethanol cycle
Plants
photosynthesise
•Remove CO2 from
atmosphere
Replanting
Fermentation
•Photosynthesis
removes CO2
•produces ethanol fuel
Burning
•Releases CO2 into
atmosphere
60
Balancing carbon cycle equations
• Glucose (a simple sugar) is created in the plant by
.
• Can you balance the following equation for
photosynthesis?
6 CO2 + 6 H2O → C6H12O6 + 6 O2
61
Balancing carbon cycle equations
During ethanol
, glucose is
decomposed into ethanol and carbon dioxide.
Can you balance this equation?
C6H12O6 → 2 CH3CH2OH+ 2 CO2
62
Balancing carbon cycle equations
During
ethanol reacts with oxygen to
produce carbon dioxide, water, and heat:
Can you balance this equation?
CH3CH2OH + 3 O2 → 2 CO2 + 3 H2O
63
Carboxylic acids
C7.1
64
What do I need to know?
1. understand that the properties of carboxylic acids are due
to the presence of the –COOH functional group;
2. recall the names and formulae of methanoic and ethanoic
acids;
3. recall that many carboxylic acids have unpleasant smells
and tastes and are responsible for the smell of sweaty
socks and the taste of rancid butter;
4. understand that carboxylic acids show the characteristic
reactions of acids with metals, alkalis and carbonates;
5. recall that vinegar is a dilute solution of ethanoic acid.
65
Can you recognise the functional group?
• Circle which of these is a carboxylic acid?
66
Answer
• This is a carboxylic acid
67
Methanoic and Ethanoic
Methanoic acid
Ethanoic acid (VINEGAR)
68
Acids in nature
Many acids are part of life itself, they are known as CARBOXYLIC acids
Organic or CARBOXYLIC
acids are part of life itself
and can be found in many
animals and plants.
69
Reactions of carboxylic acids
Reaction of carboxylic acids
1) Acid + metal  salt + hydrogen
Ethanoic acid + magnesium  magnesium ethanoate + hydrogen
2) Acid + metal oxide  salt + water
Ethanoic acid + copper oxide  copper ethanoate + water
3) Acid + metal carbonate  salt + water + carbon
dioxide
Ethanoic acid + sodium carbonate  sodium ethanoate + water + carbon dioxide
70
Example Question
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Mark scheme
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Example question
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Mark scheme
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Example question
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Mark scheme
76
Esters, Fats and Oils
C7.1
77
What do I need to know?
1. Recall the method for producing an ester using
reflux
2. Describe how fats and oils are all types of ester
and explain how margarine is made
3. Explain how bromine water can be used to test
whether a fat is saturated or unsaturated.
78
Making esters
What type of organic chemicals do you need to mix
together?
Can you name the ester made from ethanoic acid
and methanol?
79
Making esters
What type of organic chemicals do you need to mix
together?
• A carboxylic acid and an alcohol with an acid
catalyst
Can you name the ester made from ethanoic acid
and methanol?
• Methyl ethanoate
80
Esters
81
Example question
82
Mark scheme
83
Making esters
Reflux
84
Distillation
Purification
Drying
Reflux apparatus
85
How do I describe reflux for an exam?
1. Mixture heated in flask (1) …
2. with condenser above (1) …
3. so no liquid is lost by evaporation and allows
longer time for the reaction (1)
86
Distillation
87
Describing distillation
1. The mixture is heated
2. At the boiling point of the ester is becomes a
vapour
3. The vapour is condensed in the condenser
4. The liquid is collected
88
Purification
1. Collected ester is shaken in a separating funnel
with distilled water.
2. Impurities dissolve in the water
3. Impurities are tapped off
Ester
89
Drying
1. Solid drying agent is added to the product
2. This could be calcium chloride or sodium
sulphate
3. This removes water from the product
90
Example question
91
Mark scheme
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Example question
93
Mark scheme
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Example question
95
Mark scheme
96
Fats and oils
• These are a special type of ester made from
glycerol and fatty acids.
97
Fats and oils
• Removal of water in the condensation reaction
makes a fat or oil
98
Saturated or unsaturated?
• Have you heard these terms on the television?
• Vegetable oil is mostly unsaturated
• Animal fat is mostly saturated
99
Double bonds or not
• A saturated fat has no C=C
double bonds (alkene
functional groups) and is
usually a solid fat like
margarine or animal fat.
• An unsaturated fat has
C=C double bonds and is
usually an oil like
vegetable oil.
100
Example question
101
Mark scheme
102
Making margarine
• To make margarine we have to saturate vegetable
oil by bubbling hydrogen gas through the oil.
• This process is called hydrogenation
103
Is a fat or oil saturated or not?
• We can test for this by adding bromine water.
• If there are double bonds present the bromine
water changes from
to
.
104
Example question
105
Mark scheme
106
Hydrolysis
• When an ester is hydrolysed it goes back to an
acid and alcohol
• We can hydrolyse by adding acid or alkali (NaOH).
107
Example question
108
Mark scheme
109
Energy changes in chemistry
C7.2
110
Quiz
• When a chemical reaction takes place heat may be
given out or taken in.
1. Can you remember the word we use when heat
is given out?
2. Can you remember the word we use when heat
is taken in?
111
What do I need to know?
1. Recall and use the terms ENDOTHERMIC and
EXOTHERMIC
2. Describe examples of ENDOTHERMIC and
EXOTHERMIC reactions.
3. Use simple energy level diagrams to represent
ENDOTHERMIC and EXOTHERMIC reactions.
112
Change in energy
• Chemical reactants have a certain amount of
stored within them.
• When the reaction has taken place they have
either
within them
than before.
113
Definitions
(exothermic) then the
than they did before.
They have lost it to the surroundings.
(endothermic) then the
than they had before.
They have taken it from the surroundings.
114
Energy level diagrams
Which diagram do you think is
which is
?
Heat taken in
115
Heat given out
and
Energy level diagrams
Endothermic
Exothermic
Heat taken in
Energy level of products is higher
than reactants so heat taken in.
116
Heat given out
Energy level of products is lower
than reactants so heat given out.
Example question
117
Mark scheme
118
Bond enthalpies
C7.2
119
Quick quiz
1. Reactions where the products are at a lower energy than
the reactants are endothermic (TRUE/FALSE)
2. Activation energy is the amount of energy given out when
a reaction takes place (TRUE/FALSE)
3. A reaction which is exothermic transfers heat energy to
the surroundings (TRUE/FALSE)
4. How can we tell if a reaction is exothermic or
endothermic?
5. Sketch the energy profile for an endothermic reaction.
6. When methane (CH4) burns in oxygen (O2) bonds between
which atoms need to be broken?
120
Answers
1.
2.
3.
4.
5.
6.
Reactions where the products are at
a lower energy than the reactants
FALSE
are endothermic (TRUE/FALSE)
Activation energy is the amount of
energy given out when a reaction
takes place (TRUE/FALSE)
FALSE
A reaction which is exothermic
transfers heat energy to the
surroundings (TRUE/FALSE)
TRUE
How can we tell if a reaction is
exothermic or endothermic?
Sketch the energy profile for an Measure the
endothermic reaction.
When methane (CH4) burns in
oxygen (O2) bonds between which
atoms need to be broken?
121
temperature change
C—H bonds and O=O bonds
What do I need to know?
1. Recall that energy is needed to break chemical
bonds and energy is given out when chemical bonds
form
2. Identify which bonds are broken and which are
made when a chemical reaction takes place.
3. Use data on the energy needed to break covalent
bonds to estimate the overall energy change for a
reaction.
122
Activation energy revisited
• What is the activation energy of a reaction?
• The energy needed to start a reaction.
• BUT what is that energy used for and why does
the reaction need it if energy is given out overall?
• The activation energy is used to break bonds so
that the reaction can take place.
123
Burning methane
Consider the example of burning methane gas.
CH4 + 2O2  CO2 + 2H2O
This reaction is highly exothermic, it is the reaction
that gives us the Bunsen flame. However mixing air
(oxygen) with methane is not enough. I need to add
energy (a flame).
124
What happens when the reaction gets the
activation energy?
H H H H
Energy in chemicals
C
O
Bond
Breaking
H
C
H
H
H
O
O
O
Bond
Forming
O
O
O
O
O
H
O
Progress of reaction
125
C
O
H
O
H
H
Using bond enthalpies
By using the energy that it takes to break/make a
particular bond we can work out the overall
enthalpy/energy change for the reaction.
Sum (bonds broken) – Sum (bonds made) = Energy
change
126
BIN MIX
Breaking bonds is ENDOTHERMIC energy is TAKEN
IN when bonds are broken
Making bonds is EXOTHERMIC energy is GIVEN OUT
when bonds are made.
127
Bond enthalpies
Bond
Bond enthalpy (kJ)
Bond
Bond enthalpy (kJ)
C—H
435
Cl—Cl
243
C—C
348
C—Cl
346
H—H
436
H—Cl
452
H—O
463
O=O
498
C=O
804
C=C
614
128
Can you work out the energy change for
this reaction?
129
The answer is -120 kJ
130
Example question part 1
131
Question part 2
132
Question part 3
133
Mark scheme
134
Challenge question
• The true value for the energy change is often
slightly different from the value calculated using
bond enthalpies.
• Why do you think this is?
135
Example question
The calculated value is 120 kJ
136
Mark scheme
137
Definitions
Write each of these phrases in your book with a
definition in your own words:
•
•
•
•
•
Exothermic reaction
Endothermic reaction
Activation energy
Catalyst
Bond energy/enthalpy
138
How did you do?
Exothermic reaction
A reaction which gives energy out to the surroundings.
Endothermic reaction
A reaction which takes in energy from the surroundings.
Activation energy
The energy required to start a reaction by breaking bonds in the reactants
Catalyst
A substance that increases the rate of a reaction by providing an alternative
pathway with lower activation energy. It is not used up in the process of the
reaction
Bond energy/enthalpy
The energy required to break a certain type of bond. The negative value is
the energy given out when that bond is made.
139
Popular exam question
1. Explain why a reaction is either exothermic or
endothermic?
------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
140
Popular exam question
1. Explain why a reaction is either exothermic or
endothermic?
① In a chemical reactions some bonds are broken and some
bonds are made.
② Breaking bonds takes in energy.
③ Making bonds gives out energy.
④ If the energy given out making bonds is higher than the
energy needed to break them the reaction is exothermic.
⑤ If the energy needed to break bonds is higher than the
energy given out making them the reaction is endothermic.
141
Chemical Equilibria
C7.3 Reversible Reactions &
Dynamic Equilibria
142
What do I need to know?
1. State that some chemical reactions are reversible
2. Describe how reversible reactions reach a state
of equilibrium
3. Explain this using dynamic equilibrium model.
143
Reversible or not reversible
Until now, we were careful to say that most
chemical reactions were not reversible –
They could not go back to the reactants once
the products are formed.
144
Example
In the case of the vast majority of chemical
reactions this is true, the reaction of methane
and oxygen for example:
It is almost impossible to return the carbon
dioxide and water to the original methane and
oxygen.
145
Reversible
• Some chemical reactions, however, will go
backwards and forwards depending on the
conditions.
• CoCl2·6H2O(s)  CoCl2(s) + 6H2O(l)
pink
blue
146
How do we write them down?
• This is the symbol for used for reversible
reactions.
CoCl2·6H2O(s)
147
CoCl2(s) + 6H2O(l)
What is equilibrium?
• Reversible reactions reach a balance point, where
the amount of reactants and the amount of
products formed remains constant.
148
Dynamic Equilibrium.
• In
the forward and
backwards reactions continue at equal rates so
the concentrations of reactants and products do
not change.
• On a molecular scale there is
.
• On the macroscopic scale
. The system needs to be closed –
isolated from the outside world.
149
Example question
150
Mark scheme
151
Dynamic Equilibria
C7.3 Controlling equilibria
152
What do I need to know?
1. Recall that reversible reactions reach a state of
dynamic equilibrium.
2. Describe how dynamic equilibria can be affected
by adding or removing products and reactants.
3. Explain the difference between a “strong” and
“weak” acid in terms of equilibria
153
Position of the equilibrium
• Equilibrium can “lie” to the left or right.
• This is “in favour of products” or “in favour of
reactants”
• Meaning that once equilibrium has been reached
there could be more products or more reactants
in the reaction vessel.
154
Le Chatelier’s principle
• If you remove product as it is made then equilibrium
will move to the right to counteract the change
• If you add more reactant then equilibrium will move
to the right to counteract the change.
• In industry we recycle reactants back in and remove
product as it is made to push the equilibrium in
favour of more product.
155
Complete
When a system is at__________ to make more
product you can_________ product or add more
__________ for example by recycling them back in.
To return to reactants you ______ product or
remove_________.
[equilibrium, add, reactant, remove, product]
156
Strong and weak acids
A strong acid is one which is FULLY IONISED in water. It will
have a high hydrogen ion concentration
A weak acid is one which is NOT fully ionised and is in
equilibrium. It has a low hydrogen ion concentration
Caution – weak and strong are not the same as concentration.
157
158
Mark scheme
159
Example question
160
Mark scheme
161
Practicing definitions
Write each of these phrases in your book with a
definition in your own words:
•
•
•
•
•
Reversible reaction
Dynamic equilibrium
Position of equilibrium
Strong acid
Weak acid
162
How did you do?
Reversible reaction
A reaction that can proceed in the forward or reverse directions (represented by
two arrows in an equation).
Dynamic equilibrium
The point where the rate of the forward reaction = rate of the reverse reaction.
Position of equilibrium
The point where there is no further change in the concentration of either reactants
or products. The position can lie to the left (favouring reactants) or right (favouring
products).
Strong acid
An acid that is completely dissociated in water
Weak acid
An acid that is only partly dissociated in water because the reaction is in dynamic
equilibrium and favours the reactants (LHS).
163
Popular exam question
1. Ethanoic acid (CH3COOH) is a weak acid but
hydrochloric acid is a strong acid. Use ideas
about ion formation and dynamic equilibrium to
explain this difference.
---------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
164
Popular exam question
Ethanoic acid (CH3COOH) is a weak acid but hydrochloric acid is a
strong acid. Use ideas about ion formation and dynamic equilibrium to
explain this difference.
① Hydrochloric acid ionises completely
② So hydrogen ion concentration is high
③ Ethanoic acid only partly dissociates because the reaction is
reversible
④ Equilibrium is mainly to the left
⑤ So hydrogen ion concentration is low.
165
Analysis
C7.4 – Analytical Procedures
166
What do I need to know?
1. Recall the difference between qualitative and
quantitative methods of analysis.
2. Describe how analysis must be carried out on a
sample that represents the bulk of the material under
test
3. Explain why we need standard procedures for the
collection, storage and preparation of samples for
analysis
167
Qualitative vs. Quantitative
• A qualitative test is
. It can give vital
information without needing to wait too long for it.
• A
, for example
for a concentration in Moldm-3
tests include universal
indicator, silver nitrate for halide ions and bromine
water for unsaturation.
tests include titration,
chromatography and spectroscopy.
168
Which sample should I test?
• It is important that the sample
of the material under test.
• You may chose to
from a range of points to ensure that you have
.
• For example
? Are their pockets of
higher concentration/different composition?
169
Chemical industry
• Analysis of samples is crucial to the chemical industry
to ensure the
of the chemicals they are
manufacturing. Some are analysed numerous times a
day or even within an hour.
• To maintain consistency it is essential that we use
to:
o collect the sample
o store the sample
o prepare the sample for analysis
o analyse the sample.
170
Chromatography
C7.4 – paper chromatography
171
Chromatography
172
Solvents
1. The mobile phase is the solvent – the part that
moves
2. In paper chromatography it is water or ethanol
173
Paper/column
1. The stationary phase is the paper in paper
chromatography or the column in gas
chromatography.
2. In thin layer chromatography it is silica gel on a
glass plate
3. The stationary phase does not move.
174
How does the technique work?
In chromatography, substances are separated by
movement of a mobile phase through a stationary
phase.
Each component in a mixture will prefer either the
mobile phase OR the stationary phase.
The component will be in dynamic equilibrium between
the stationary phase and the mobile phase.
175
Substance A
• This is substance A
• Substance A prefers the stationary phase and doesn’t
move far up the paper/column.
• The equilibrium lies in favour of the stationary phase.
176
Substance B
• This is substance B
• Substance B prefers the mobile phase and moves a
long way up the paper/column
• The equilibrium lies in favour of the mobile phase
177
Using a reference
• In chromatography we can sometimes use a
known substance to measure other substances
against.
• This will travel a known distance compared to the
solvent and is known as a standard reference.
178
Advantages of TLC
TLC has a number of advantages over paper
chromatography.
It is a more uniform surface chromotograms are
neater and easier to interpret
Solvent can be used which is useful if a substance is
insoluble in water.
179
Past Paper Questions
180
Past paper question
181
Mark scheme
182
Describing how chromatography
works – exam definition
• stationary phase is paper and mobile phase is
solvent / mobile phase moves up through
stationary phase (1)
• for each compound there is a dynamic equilibrium
between the two phases (1)
• how far each compound moves depends on its
distribution between the two phases (1)
183
Using an Rf value
• In order to be more precise we can use
measurements on the TLC plate to compare the
distance travelled by our substance (the solute)
with the distance travelled by the solvent.
• The Rf value is constant for a particular
compound.
• The distance travelled however could be different
on different chromatograms.
• The Rf value is always less than 1.
184
Rf value
Distance travelled by spot
Rf value =
Distance travelled by solvent
185
Example question
This question relates to the chromatogram shown in the earlier question. Refer back…
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Mark scheme
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Example question
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Mark scheme
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Past paper question
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Mark scheme
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Gas-liquid chromatography
C7.4 GLC
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What do I need to know?
1. recall in outline the procedure for separating a
mixture by gas chromatography (gc);
2. understand the term retention time as applied to
gc;
3. interpret print-outs from gc analyses.
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Gas chromatography
• The mobile phase is an unreactive gas known as
the carrier gas this is usually nitrogen
• The stationary phase is held inside a long column
and is lots of pieces of inert solid coated in high bp
liquid.
• The column is coiled in an oven
• The sample to be analysed is injected into the
carrier gas stream at the start of the column.
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GC
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GC analysis
• Each component of the sample mixture has a
different affinity for the stationary phase compared
with the mobile phase
• Therefore each component travels through the
column in a different time.
• Compounds favouring the mobile phase (usually more
volatile) emerge first.
• A detector monitors the compounds coming out of
the column and a recorder plots the signal as a
chromatogram
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GLC Chromatograph
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Interpretation
• The time in the column is called the retention
time
• Retention times are characteristic so can identify
a compound
• Area under peak or relative heights can be used to
work out relative amounts of substances
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The key points – revise this!
• the mobile phase carries the sample (1)
• components are differently attracted to the
stationary and mobile phases (1)
• the components that are more strongly attracted
to the stationary phase move more slowly (1)
• the amount of each component in the stationary
phase and in the mobile phase is determined by a
dynamic equilibrium (1)
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Past paper question
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Mark scheme
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Titration
C7.4
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What do I need to know?
1. Calculate the concentration of a given volume of
solution given the mass of solvent;
2. Calculate the mass of solute in a given volume of
solution with a specified concentration;
3. Use the balanced equation and relative formulamasses to interpret the results of a titration;
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Concentration
• We can measure the concentration of solution in
grams/litre. This is the same as g/dm3
• 1dm3 = 1000cm3
• If I want to make a solution of 17 g/dm3 how much will I
dissolve in 1dm3.
• 17 g
• If I want to make a solution of 17g/dm3 but I only want to
make 100cm3 of it how much will I dissolve?
• 1.7g
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Making standard solutions
• For a solution of 17g/dm3
• First I will measure 17g of solid on an electronic
balance
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Making standard solutions
• Now I must dissolve it in a known 1dm3 of
water.
• I transfer it to a volumetric flask and fill up
with distilled water to about half the flask.
• I then swirl to dissolve
• Top up with a dropping pipette so that the
meniscus is on the line.
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How much to dissolve?
• Worked example:
• I want to make 250cm3 of a solution of 100g/dm3.
• How much solid do I transfer to my 250cm3
volumetric flask?
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How much to dissolve?
Worked example:
I want to make 250cm3 of a solution of 100g/dm3.
1. Work out the ratio of 250cm3 to 1000cm3
250/1000 = 0.25
2. I therefore need 0.25 of 100g in 250cm3 which is
0.25x100=25g
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General rule
3
3
Volume(cm )xConcentration(g / dm )
Mass(g) =
1000
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Practie - how much to dissolve?
• I want to make 250cm3 of a solution of 63.5g/dm3.
• How much solid do I transfer to my 250cm3
volumetric flask?
• 250/1000 x 63.5 = 15.9 g
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Practice - how much to dissolve?
• I want to make 100cm3 of a solution of 63.5g/dm3.
• How much solid do I transfer to my 100cm3
volumetric flask?
• 100/1000 x 63.5 = 6.35 g
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Concentration from mass and volume
We need to rearrange this:
3
3
Volume(cm )xConcentration(g / dm )
Mass(g) =
1000
To give
Mass(g)x1000
Concentration(g / dm ) =
3
Volume(cm )
3
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What is the concentration of?
1. 12g dissolved in 50cm3
2. 50g dissolved in 100cm3
3. 47g dissolved in 1000cm3
4. 200g dissolved in 250cm3
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What is the concentration of?
1. 12g dissolved in 50cm3
= 1000/50 x 12
= 240g/dm3
2. 50g dissolved in 100cm3
=1000/100 x 50
= 500g/dm3
3. 47g dissolved in 1000cm3
= 1000/1000 x 47
= 47g/dm3
4. 200g dissolved in 250cm3
1000/250 x 200
= 800g/dm3
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Solutions from stock solutions
Stock solution
• highest
concentration
• use to make
other solutions
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Extract a
portion of
stock solution
• as calculated
Dilute with
distilled water
• Making a known
volume of a lower
concentration
Making solutions from stock solutions
If I have a solution containing 63g/dm3, how do I make up
250cm3 of a solution of concentration 6.3g/dm3?
To make 1dm3 of 6.3g/dm3 I would need 100cm3
To make 250cm3 of 6.3g/dm3 I would therefore need 25cm3
and make it up to 250cm3 with distilled water
Final concentration(g/dm3 )
3
3
xSize
of
flask(cm
)=Amount
to
add(cm
)
3
Initial concentration(g/dm )
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Working out masses
• We can use the useful relationship
Mass1 Mass2
=
Mr1
Mr2
• Where Mr is the molecular mass
• eg Mr of NaOH is (23 + 16 + 1) = 40
• This can help us to calculate an unknown mass
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Titration calculations
• In a titration we have added a known amount of
one substance usually an acid (in the burette) to a
known amount of another substance usually an
alkali (in the conical flask).
• The amount added allows us to determine the
concentration of the unknown.
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Titration equipment
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Using a table
• It can be helpful to sketch a table to keep track of
information you know…
Value
Volume (cm3)
Mass (g)
Concentration (g/dm3)
Molecular weight (Mr)
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Acid
Alkali
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Mark scheme
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Uncertainty
• Uncertainty is a quantification of the doubt about
the measurement result.
• In a titration the uncertainty is the range of the
results.
• If results are reliable then it will be within 0.2cm3
• NOTE THAT THIS IS RELIABLE NOT NECESSARILY
ACCURATE
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C7.5 Green Chemistry
The Chemical Industry
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What do I need to know?
1. Recall and use the terms 'bulk' (made on a large
scale) and 'fine' (made on a small scale) in terms
of the chemical industry with examples;
2. Describe how new chemical products or
processes are the result of an extensive
programme of research and development;
3. Explain the need for strict regulations that
control chemical processes, storage and
transport.
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Bulk processes
• A bulk process manufactures large
quantities of relatively simple
chemicals often used as feedstocks
(ingredients) for other processes.
• Examples include ammonia, sulfuric
acid, sodium hydroxide and phosphoric
acid.
• 40 million tonnes of H2SO4 are made in
the US every year.
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Fine processes
• Fine processes manufacture smaller
quantities of much more complex
chemicals including pharmaceuticals,
dyes and agrochemicals.
• Examples include drugs, food
additives and fragrances
• 35 thousand tonnes of paracetamol
are made in the US every year.
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Example question
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Research and Development
• All chemicals are produced following an extensive
period of research and development.
• Chemicals made in the laboratory need to be
“scaled up” to be manufactured on the plant.
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Research in the lab
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Examples of making a process viable
• Trying to find suitable conditions – compromise
between rate and equilibrium
• Trying to find a suitable catalyst – increases rate
and cost effective as not used up in the process.
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Catalysts
• Can you give a definition of a catalyst?
• A substance which speeds up the rate of a
chemical reaction by providing an alternative
reaction pathway.
• The catalyst is not used up in the process
• Catalysts can control the substance formed eg
Ziegler Natta catalysts.
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Regulation of the chemical industry
• Governments have strict regulations to control
chemical processes
• Storage and transport of chemicals requires
licenses and strict protocol.
• Why?
• To protect people and the environment.
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Example question
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Process development
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Example question – part 1
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Example question part - 2
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Example question – part 3
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Factors affecting the sustainability of a
process
energy inputs
and outputs
type of waste
and disposal
environmental
impact
health and
safety risks
atom economy
renewable
feedstock
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Sustainability
social and
economic
benefits
Example question
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Mark scheme
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Atom economy
Mr of desired product
% atom economy =
x 100
Mr of reactants
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Atom economy calculation
For example, what is the atom economy for making
hydrogen by reacting coal with steam?
Write the balanced equation:
C(s) + 2H2O(g) → CO2(g) + 2H2(g)
Write out the Mr values underneath:
C(s) + 2H2O(g) → CO2(g) + 2H2(g)
12
2 × 18
44
2×2
Total mass of reactants 12 + 36 = 48g
Mass of desired product (H2) = 4g
% atom economy = 4⁄48 × 100 = 8.3%
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Example question
Example question – part 2
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