AS Chemistry
Edexcel
Unit 2
Application
of Core
Principles
Unit 2.1 Shapes of molecules & ions
(p144 – 153 in text book)
1. Predicting Shapes of molecules
The bonding of atoms will impact on the shape of the molecule formed. The
shape of a molecule can determine both physical and chemical properties. In
particular, enzyme activity is controlled by the shape since this decides the
active site.
2. Shapes of ionic compounds
Ionic bonds are non-directional. The ions will always be electrostatically
attracted to each due to their charge. Conversely, they will also repel. As a
result a lattice structure is formed to maximise attraction and minimise
repulsion.
3. Shapes of covalent compounds
Covalent bonds are highly directional. As a result, these compounds have a
very specific shape which governs the chemical and physical properties.
The shape of a molecule depends on:
the number of electron pairs around the central atom
electron pair repulsion
The electron pairs align themselves in such a way to minimise the amount of
repulsion. The bond angles depend on the number of electron pairs around a
central atom.
(Refer to table 2.1.1 p145)
Number of electron pairs around the central atom
2
Example BeCl2
3
4
5
6
7
BF3
CH4
PCl5
SF6
IF7
Dot &
cross
diagram
Shape
Bond
angle
Name
of
shape
4. The effect of lone pairs on shapes of molecules
Some molecules have lone pairs of electrons around the central atom as well
as the bonding electrons. These lone pairs impact on the shape since they
too become involved in electron repulsion.
However, the lone pair electrons are closer to the nucleus than the bonding
electron pairs. As a result, the repulsion between a lone pair and a bonding
pair is greater than that of just 2 bonding pairs. This impacts on bond angles.
Examples
Ammonia NH3
Water H2O
Xenon tetrafluoride XeF4
5. Multiple bonds
These are treated in the same way as single bonds. The electron pairs are
still arranged to minimise repulsion.
Examples
Carbon dioxide CO2
Sulphur dioxide SO2
6. Shapes of ions
The ions are formed when electrons are lost or gained. However, electron
pair repulsion can still be used to predict the shape of the ion by viewing the
electron pairs around the central atom.
Examples
Ammonium NH4+
Sulphate SO42-
Nitrate NO3-
7. Shapes of simple organic molecules
These molecules generally have a tetrahedral structure, but features such as
double bonds will have an impact on the shape.
(Refer to p148 – 149)
a.
Alkanes
All of these molecules show a tetrahedral shape around each carbon atom
with a bond angle of 109.5°
E.g. Ethane
b.
Alkenes
These molecules have a planar structure around the double bond
E.g. Ethene
c.
Alcohols
These molecules have a tetrahedral structure, but lone pair repulsion
decreases the angle around the oxygen atom
E.g. Methane
d.
Carboxylic acids
Methanoic acid is planar. In all other carboxylic acids the functional group is
planar, with the rest of the molecule displaying tetrahedral structure.
e.
Halogenoalkanes
These molecules are tetrahedral.
E.g. Bromomethane
f.
Aldehydes & ketones
These molecules have a planar structure around the carbonyl group, with
tetrahedral structure elsewhere in the molecule.
E.g. Propanal & propanone
8. Allotropes of carbon
(Refer to p150 – 153)
Allotropes are different forms of the same element that exist in the same
physical state. This is due to different bonding between the atoms resulting
in different shapes.
a.
Diamond
Pure diamond is made from interlocking crystals of carbon covalently bonded
in a tetrahedral structure. Each carbon atom has 4 bonds.
(Refer to fig 2.1.16 p150)
This structure is giant molecular and means that diamond is extremely hard
allowing it to be used for cutting tools.
Pure diamond is colourless. However, mineral contaminates can colour them
from opaque to black. Diamonds are chemically inert and poor conductors of
electricity.
Since 1955 diamonds have been able to be manufactured synthetically. This
is done by subjecting graphite to high temperature and pressure. These were
only good enough for industrial processes until 1970. However, the synthetic
diamonds of gem quality are so expensive to make they aren’t used for
jewellery.
b.
Graphite
This is a black/grey lustrous substance that easily crumbles due to the
tendency of the layers to cleave. Three carbon atoms are bonded together in
hexagonal layers, leaving delocalised electrons to flow between the layers.
The layers are held together by weak London forces.
(Refer to fig 2.1.17 p151)
Graphite is chemically inert, but less so than diamond. However, whilst a
poor heat conductor, it is a good conductor of electricity.
c.
Amorphous carbon
This is a deep black powder, often seen as soot. It can be formed through
heating organic substances with a lack of oxygen.
It is the most reactive form of carbon and easily burns in air. It appears to be
similar in structure to graphite. It is used for black pigment in paints and
inks, along with stabilising filler for rubber and plastics.
d.
Fullerenes
Buckminster fullerene is a spherical structure with 60 atoms bonded together.
Other fullerenes exist with different numbers of carbon atoms. They have
delocalised electrons so are able to conduct electricity.
9. Nanochemistry
(Refer to p152 & 153 especially the HSW boxes)
Nanotubes were developed in 1991. These are elongated cage-like structures
which can form highly complex shapes.
Various uses are being trialled for nanotubes and nanoparticles to include;
i.
superconductors
ii.
composite materials
iii.
nanoprobes to identify cancerous cells
iv.
cosmetic preparations
Unit 2.2 Intermediate Bonding & Bond Polarity
(p154 – p159 in text book)
1.
The position of the electrons in a covalent bond
When the 2 atoms in a covalent bond are identical to each other, e.g. in a
hydrogen molecule, the bonding electrons are shared equally between the 2
atoms.
However, when the 2 atoms in a covalent bond are different to each other,
e.g. in a hydrogen chloride molecule, the bonding electrons are no longer
shared equally between the 2 atoms due to different attractions from the
nuclei.
Electronegativity is a measure of the attraction of an atom in a molecule for a
pair of electrons in a covalent bond.
2.
The Pauling scale
The most common scale used to measure electronegativity is the one
produced by Linus Pauling. It is a relative scale, so without units, between 0
and 4. The higher the number, the more electronegative the element is.
The most electronegative element is fluorine with a value of 4, while the least
electronegative element is caesium with a value of 0.7.
(Refer to fig 2.2.4 p155)
There are several trends in electronegativity;
non-metals have higher electronegativity values than metals
electronegativity increases across a period
electronegativity decrease down a group
The noble gases do not have electronegativity values since they don’t have an
affinity for electrons.
Electron density maps can be used to view the uneven sharing of electrons in
a covalent bond.
(Refer to HSW box p155)
3.
Bond character
(Refer to table 2.2.1 p156)
Ionic and covalent bonding models represent two extremes. In fact, they sit
at each end of a continuum.
By calculating the electronegativity difference in a bond, the character of the
bond can be calculated i.e. how much covalent character it exhibits and how
much ionic character it exhibits.
4.
Polar bonds
a.
Covalent bonds
(Refer to fig 2.2.7 p157)
The uneven sharing of bonding electrons results in a distortion of the electron
cloud. As a result, small charges occur at each end of the bond. These are
designated as δ- and δ+. This produces a polar bond e.g. hydrogen chloride.
b.
Ionic bonds
(Refer to fig 2.2.8 p157)
Ionic bonds may be distorted if the cation has a high charge density since it
can polarise the electron cloud of the anion. If the distortion is great enough,
the electron clouds can overlap to resemble a covalent bond.
5.
Polar molecules
It is possible to have several polar bonds in a molecule without the molecule
itself being polar.
Polar liquids can be deflected by an electric charge. This property is
dependant on the symmetry of the molecule.
Liquids deflected
Liquids undeflected
For a molecule to be polar, it needs to be asymmetrical so that there is an
uneven distribution of charge across the molecule. The charge separation
makes it a dipole. This polarity is measured as its dipole moment – the
amount of charge separation multiplied by the distance between the two
centres of charge. The unit of the dipole moment is the debye, D.
Examples
i.
Boron trifluoride, BF3
Although each B-F bond is polar, this is a non-polar molecule since the
molecule has a net polarity of zero.
ii.
Water, H2O
The molecule is v-shaped so it isn’t symmetrical. Both sets of bonding
electrons are attracted towards the oxygen atom. As a result there is an
overall dipole moment.
6.
Predicting whether a substance is polar or not
Formula
Polar
Non-polar
Description
Example
Unit 2.3 Intermolecular Forces
(p160 – p169 in text book)
1.
Polarity & boiling temperature
(Refer to fig 2.3.1 p160)
Non-polar molecules have much lower boiling points than polar molecules.
The forces between polar molecules must therefore be much bigger since it
takes more energy to break these.
2.
Strength
Melting and
Boiling
Points
Bonding
Molecules
Reason
3.
Comparing intermolecular forces
Hydrogen Bonds
London Forces
Strongest
High
Weak
Low
Hydrogen and Oxygen,
Fluorine or Nitrogen
When hydrogen bonds
with any of the above
elements the electrons
are pulled towards the
said element because of
the large
electronegativity
difference. The lone pairs
on the oxygen, fluorine
or nitrogen are then
weakly bonded to the
hydrogen atoms.
All atoms and
molecules
Electrons move
around fast in an
atom and at any
given time they are
more likely to be at
one side of an atom.
This causes a dipole
which then induces a
dipole in
neighbouring atoms
or molecules. These
dipoles are weakly
attracted.
Trends in physical properties
Permanent
Dipole-dipole
Strong
High
Polar Molecules
Some molecules
have an uneven
share of electrons in
there bonds due to
polarisation. This
creates a dipolar
molecule. The
charged ends of
these molecules
attract each other.
a.
Trends in alkanes
(Refer to table 2.3.3 p164)
The melting and boiling points of the straight chain alkanes increase as the
chain length increases. The forces between alkanes are weak but they grow
larger as the chain length increases due to a greater surface area.
(Refer to table 2.3.4 p165)
Branching of alkane chains generally lowers the melting and boiling points.
The side chains interfere with the packing of the molecules and also reduce
the number of points where London forces can occur.
b.
Comparing alkanes & alcohols
(Refer to table 2.3.5 p165)
Alcohols have higher melting and boiling points than alkanes. This is due to
the hydrogen bonding that can take place between alcohol molecules.
c.
Trends in boiling temperature of the hydrogen halides
(Refer to fig 2.3.12 p166)
Boiling temperatures decrease from HI to HCl as you move up group 7.
However, HF has a much higher boiling point than any of the other hydrogen
halides. The electronegativity difference between hydrogen and fluorine is
greater. The result is the formation of hydrogen bonds between the
molecules so more energy is required for the state change.
4.
Solubility
Water is a very good solvent since it dissolves a wide variety of substances.
However, not all substances will dissolve in water e.g. chlorophyll.
Other solvents, often called non-aqueous solvents, will dissolve many of the
substances that water won’t.
Solubility strictly means the mass of a solute that dissolves in 100g of solvent
at a particular temperature.
a.
Patterns in solubility
There are some general rules about solubility;
highly polar solids such as ionic salts dissolve in water (a polar solvent)
but not in hexane (a non-polar solvent)
polar organic substances such as glucose dissolve in water but not in
hexane
non-polar solids such as candle wax dissolve in hexane but not in
water
non-polar liquids such as petrol and diesel mix completely. They are
miscible
polar liquids such as water and ethanol are miscible
a polar liquid and a non-polar liquid such as water and oil do not mix
together. They are immiscible and form separate layers
b.
Dissolving an ionic solid in water
(Refer to fig 2.3.14 p168)
Sodium chloride dissolves in water with no apparent change in energy. The
ionic solid has a lattice energy which needs to be applied to break down its
structure. However, the polar water molecules align themselves around the
ions through the process of hydration which provides a hydration enthalpy.
In this case the lattice energy and hydration enthalpy are balanced, hence the
apparent lack of energy change.
In sodium hydroxide the lattice energy is less than the hydration enthalpy so
it dissolves in water as an exothermic process.
Lithium fluoride, however, is insoluble in water since its lattice energy is
greater than the hydration enthalpy.
c.
The solubility of alcohols in water
(Refer to fig 2.3.15 p169)
Alcohols are able to dissolve in water since they contain the polar OH group.
However, their solubility decreases as the length of the carbon chain
increases. Carboxylic acids follow the same pattern.
d.
Why are non-polar substances insoluble in water?
Non-polar substances do not dissolve in water. The London forces between
the non-polar substances are weaker than the hydrogen bonds in water. As a
result, the non-polar substances are unable to disrupt the arrangement of the
water molecules.
(Refer to fig 2.3.16 p169)
Even halogenoalkanes are unable to dissolve in water since most of the
molecule is non-polar.
e.
Mixing two organic liquids
(Refer to fig 2.3.17 p169)
Hexane and octane are 2 organic liquids that mix together. They both have
weak London forces so when they are mixed together the London forces
extend throughout the mixture.
Unit 2.8 Organic chemistry – alcohols & halogenoalkanes
(p208 – 221 in text book)
1.
Alcohols
These are very important organic compounds in industry. They are good
solvents and make suitable raw materials.
They have the general formula CnH2n+2O (also written as CnH2n+1+OH).
They have an alkyl chain with an OH group attached and are named in the
same way as other organic molecules.
(Refer to table 2.8.1 p208)
Alcohol
Methanol
Ethanol
Propan-1-ol
Butan-2-ol
2-methylbutan-2-ol
Displayed formula
Skeletal formula
2.
Types of alcohols
Alcohols can be classified as either primary, secondary or tertiary depending
on the position of the functional OH group.
(Refer to fig 2.8.2 p209)
Primary
Secondary
Tertiary
Example
Definition
3.
Reactions of alcohols
Alcohols will undergo a number of different reactions.
a.
Combustion
Just like hydrocarbons, the products of complete combustion are carbon
dioxide and water. For incomplete combustion, carbon monoxide will be
formed.
E.g. Propan-1-ol
b.
Reaction with sodium
When sodium is added to an alcohol, it reacts with the hydrogen atom of the
OH group to form H2. The other product is a salt.
The alcohols are behaving like weak acids. The strength of the ‘acid’
decreases down the homologous series.
E.g. Propan-1-ol & sodium
c.
Test for alcohols / Making halogenoalkanes
Alcohols react with PCl5 to form a chloroalkane and hydrogen chloride gas.
E.g.
CH3CH2OH(l) + PCl5(s)  CH3CH2Cl(l) + POCl3 + HCl(g)
This particular reaction results in white fumes being released (the hydrogen
chloride) which can be used as an observation to confirm the presence of an
OH group.
Similar reactions occur with other phosphorous halides e.g.
Propan-1-ol & PBr5
Phosphorous (III) halides will also react with alcohols to produce
halogenoalkanes
E.g.
3 CH3CH2OH + PF3  3 CH3CH2F + H3PO3
4.
Oxidation of alcohols
These can be used to distinguish between primary, secondary & tertiary
alcohols since they each react differently with common oxidising agents such
as acidified potassium dichromate (K2Cr2O7) and acidified potassium
permanganate (KMnO4).
a.
Primary alcohols
These are quickly oxidised to aldehydes which can be rapidly changed into
carboxylic acids. The common product is the acid unless the aldehyde is
separated during the reaction.
E.g. propan-1-ol
b.
Secondary alcohols
These are readily oxidised into ketones, but no further oxidation occurs so
only one product is made.
E.g. Propan-2-ol
c.
Tertiary alcohols
These aren’t readily oxidised by any of the common reducing agents
d.
Colour changes
(Refer to fig 2.8.6 p211)
Oxidising agent
Start colour
End colour if
oxidation occurs
Potassium dichromate
Potassium
permanganate
Colourless
5.
Extracting the products of oxidation of alcohols
a.
Reflux
(Refer to fig 2.8.8 p211)
This method is used to prepare carboxylic acids from primary alcohols. The
reflux condenser prevents vapours from leaving the flask so that a full
reaction can take place.
b.
Distillation
(Refer to fig 2.8.7 p211)
This method is used to prepare an aldehyde from a primary alcohol. The
aldehyde can be collected before it has the chance to react further to form a
carboxylic acid.
This method is also used to prepare a ketone from a secondary alcohol.
6.
Halogenoalkanes
These organic molecules are rarely found naturally occurring on Earth. Most
are synthesised. They are important in medicine, plastic production and
agriculture. However, most of them are responsible for high levels of
environmental damage, both in the atmosphere and on the ground.
At least one of the hydrogen atoms from an alkane is replaced with a halogen
to form these molecules. Consequently there are many different isomers.
(Refer to fig 2.8.10 p212)
Primary
Secondary
Tertiary
Example
Definition
7.
Uses of halogenoalkanes
There are many uses of these chemicals.
a.
Anaesthetics
(Refer to HSW box p213)
These have been developed to aid surgery over the years since it is much
easier to operate on an unconscious patient than a conscious one.
In 1841 ethoxyethane (ether) was discovered to be an anaesthetic but it is
highly flammable. In 1844 nitrous oxide was used in dentistry. Whilst not
flammable or toxic it is only a mild anaesthetic. Towards the end of the
1840s trichloromethane (chloroform) was discovered to be a better
anaesthetic than these earlier two.
The middle of the 20th century saw the discovery of 2-bromo-2-chloro-1,1,1trifluoroethane (halothane) as an effective deep anaesthetic.
b.
Dry cleaning
The most common dry cleaning fluid is tetrachloroethene.
c.
Refrigerants
(Refer to p214 & 215)
These are liquids which circulate in a refrigerator, constantly changing phase
from liquid to gas and back again in order to transfer heat from inside the
fridge to the outside.
A good refrigerant should;
CFCs (chlorofluorocarbons) have been widely used refrigerants. There are so
many of these refrigerants that they have had a separate naming system
developed.
Rightmost digit
Tens digit
Hundreds digit
Thousands digit
Any remaining bonds not accounted for are occupied by chlorine atoms
A suffix of a lower case letter a, b or c indicates unbalanced isomers e.g.
R134a
When CFCs are released into the atmosphere they begin to react once they
reach the ozone layer.
Stage 1 – homolytic breakdown under UV light e.g.
Stage 2 – reaction of a chlorine radical
Since the chlorine atom isn’t used up in the reaction, one molecule of CFC can
destroy thousands of ozone molecules.
In 1987 the Montreal Protocol set out plans for the reduction of CFC use.
Instead, HCFCs (hydrochlorofluorocarbons) have been synthesised to replace
CFCs. Although they pose a much lower threat to the ozone layer they are
very potent greenhouse gases.
d.
Fire retardants
(Refer to HSW box p216 & 217)
These were used in fire extinguishers designed to tackle electrical fires and
were identified in green containers. However, their contribution to the ozone
layer depletion and the narcotic effect of the gases has seen them withdrawn
from use. Instead they have been replaced by the blue CO2 extinguishers.
TBBPA contains approximately 59% bromine and is used to cover circuit
boards as a fire retardant.
Synthetic fibres often have some halogen content in order to reduce the size
of the flame should they catch fire. Such a benefit means that children’s
clothing is less likely to catch fire near a naked flame.
e.
Insecticides
DDT was widely used as an insecticide but causes toxic build up higher up in
food chains so its use has now been prohibited.
8.
Reactions of the halogenoalkanes
The chemistry of these compounds depends on the halogen atom present
along with its position on the alkyl chain. Most reactions are nucleophilic
substitutions but under some conditions elimination will occur.
a.
Reaction with aqueous hydroxide or water
This is a nucleophilic substitution reaction but can also be referred to as
hydrolysis. The reaction with water is slow at room temperature but is much
quicker with an OH- ion. The ideal conditions are under reflux with heat. The
main product will be an alcohol.
E.g.
Equation for chloroethane with water
Equation for bromomethane with NaOH
The rate of reaction is inversely proportional to the electronegativity
difference of the RX bond i.e. as the electronegativity of the halogen
decreases, the faster the rate of reaction.
b.
Reaction with alcoholic potassium hydroxide
In this case elimination will take place under reflux with heat. The main
product will be an alkene.
E.g. reaction of ethanolic potassium hydroxide with 2iodobutane
c.
Reaction with alcoholic ammonia
Ammonia behaves as a nucleophile to undertake a substitution reaction to
form an amine.
E.g. reaction of 1bromopropane with ammonia
However, the resulting amine also possesses a lone pair of electrons so
further reaction can take place.
This reaction must take place in a sealed tube since the ammonia would
otherwise escape through the condenser if attempted under reflux.
9.
Identifying halogenoalkanes
When these compounds react it usually results in the release of a halide ion.
This can then be tested for using silver nitrate and ammonia.
Halide ion
Observation with
silver nitrate
Observation with the
addition of ammonia
Chloride
Bromide
Iodide
Plotting the production of these halide ions can be used to determine the rate
of reaction. The further down group 7 the faster the halogenoalkane will
react. Also, the tertiary halogenoalkane reacts faster than the secondary,
which in turn is faster than the primary i.e. more branching gives a faster
reaction.
10.
Preparing halogenoalkanes
This can be done by reacting an alcohol under reflux.
(Refer to fig 2.8.18 p221)
The flask is put into a beaker of cold water and concentrated sulphuric acid is
added slowly from a dropping funnel. The flask is cooled since the reaction is
exothermic at this stage.
Refluxing on a water bath takes place for 30 minutes.
(Refer to fig 2.8.19 p221)
The mixture is then distilled with the distillate collected in two layers – an
upper aqueous layer and a lower organic layer. The aqueous layer is
discarded while the organic layer is purified and redistilled to give a pure
product.
Unit 2.10 Mass spectra and IR absorption
(p230 – 235 in text book)
1.
The mass spectrometer
This was studied in unit 1.3. For organic chemistry it can be applied in the
use of pharmacology, space research, radioactive dating and catching drug
cheats in sport.
2.
Using a mass spectrometer to analyse organic
compounds
Identifying an organic compound from the huge numbers that exist can be a
daunting and long task. The use of mass spectrometry and infrared
spectroscopy can make it much quicker.
The organic compound can form a molecular ion when bombarded with the
high speed electrons. An m/z value produces a molecular ion peak which is
usually the peak with the highest m/z value. Other peaks can occur due to
fragmentation of the molecule.
Some peaks are more difficult to interpret than others. Some can be easily
identified as being from the loss of hydrogen, a methyl group, an OH group
etc. However, even in simple molecules, it is impossible to identify all the
peaks due to complicated rearrangements of the cations inside the machine.
(Refer to specific examples of butane, ethanol, propanal and propanone on
p230 – 231)
3.
Absorbing infrared
(Refer to fig 2.10.8 p232)
If a substance is irradiated with IR its molecules will absorb some of this
radiation. This absorption is as a result of molecules vibrating due to either a
stretch or a bend in the bonds. The amount of energy required to make them
vibrate depends on;
- the bond strength
- the bond length
- the mass of each atom involved in the bond
A range of IR frequencies are passed through the sample in an infrared
spectrometer. The spectrum identifies a series of troughs where energy is
absorbed, although these are referred to as peaks. The frequency of the IR
absorption is measured in wavenumbers (cm-1) shown on the x axis. The
percentage transmission is shown on the y axis. One beam of IR is passed
through the sample while a second is passed through a reference cell to
ensure only IR absorption in the compound is observed on the spectrum.
4.
Bending & stretching
A higher frequency of IR is required to make stronger bonds vibrate. The
bonds can bend, stretch symmetrically or stretch asymmetrically. Each type
of vibration has a corresponding peak.
(Refer to fig 2.10.9 p233)
Only molecules which change their polarity will absorb IR. Consequently,
molecules such as H2 and Cl2 will not have any peaks on an IR spectrum but
HCl would. It is polar molecules in the atmosphere such as CO2, and H2O that
can absorb IR which imparts on the greenhouse effect.
Functional groups within organic molecules can be recognised from their
characteristic absorption peaks which are recorded in reference tables.
(Refer to table 2.10.3 p233)
(Refer to specific examples on p234 – 235)
Unit 2.5 The Periodic Table – Groups 2 & 7
(Refer to p178 – 195 in text book)
1.
Group 2: the alkaline earth metals
Group 2 elements behave in a similar way to group 1 elements, with some
clear patterns as you go down the group.
a.
Ionisation energies of group 2 elements
There is a large increase for all the group 2 elements between the second and
third ionisation energies. They readily lose 2 electrons to form 2+ ions, but
the large third ionisation energy makes it impossible to form a 3+ ion. Losing
2 electrons gives them noble gas configuration.
(Refer to table 2.5.1 p179)
The ionisation energy decreases down the group. There is an increasing
nuclear charge down the group, but there is also an increase in atomic radius.
This means that the outermost electrons experience a weaker attraction to
the nuclear charge as you move down the group alongside a greater shielding
effect from inner electron shells.
b.
Flame tests
These are carried out to test for the presence of specific metal ions.
Concentrated hydrochloric acid is used to clean platinum/nichrome wire which
is then used to hold a sample in a hot Bunsen flame.
The colour of the flame can be used to identify the cation.
Cation
Flame colour
Lithium
Sodium
Potassium
Magnesium
Calcium
Strontium
Barium
The flame colour is produced when the electrons in the cation return to their
ground state after being excited. However, some cations don’t produce a
colour since there is too large a gap between energy levels for the electrons
to become excited.
A flame photometer can be used to provide a more accurate result by
measuring the wavelength of light emitted by the cation. It is often used to
find sodium and potassium levels in blood samples and wine.
c.
Reactions of the group 2 metals
There are also obvious chemical trends down group 2 as well as the physical
trends.
i.
Reaction with oxygen
Reactivity with oxygen increases down the group to form a metal oxide e.g.
ii.
Reaction with chlorine
Reactivity with chlorine also increases down the group to form solid metal
chlorides e.g.
iii.
Reaction with water
The reactivity increases down the group.
(Refer to table 2.5.4 p181)
Magnesium reacts with steam to form magnesium oxide and hydrogen.
Calcium, strontium and barium react with water to form a hydroxide and
hydrogen e.g.
Group 2 oxides and hydroxides aren’t naturally occurring compounds.
However the cations are often found in carbonates which can undergo
thermal decomposition to form the oxide, which then reacts with water to
form the hydroxide.
d.
Reactions of the group 2 oxides
i.
Reaction with water
Beryllium oxide doesn’t react with water, while magnesium oxide will only
slightly react.
However, calcium, strontium and barium oxides will react vigorously with
water. The calcium hydroxide (slaked lime) produced is used in water
treatment, neutralising acidic soils and making whitewash, mortar and plaster.
ii.
Reaction with dilute hydrochloric and nitric acids
Group 2 metal oxides will react with these acids to form a salt and water e.g.
e.
Solubility of group 2 compounds
Barium sulphate is used in barium ‘meals’ or enemas. Since it is insoluble in
water it shows up well on X-rays without poisoning patients.
All the group 2 nitrates and chlorides are soluble. Group 2 salts where the
anion has a charge of -2 (such as sulphates) are generally insoluble with the
exception of a few magnesium and calcium salts.
The solubility of group 2 compounds tends to decrease down the group as
atomic number and ionic size of the cation increase. However, the solubility
of hydroxides increases.
(Refer to table 2.5.5 p183)
Barium sulphate is used to test for sulphate ions. First dilute hydrochloric
acid is added to the sample to destroy any carbonates. Then barium chloride
or barium nitrate is added. A white precipitate of insoluble barium sulphate
confirms the presence of sulphate ions.
Group 1 metals produce hydroxides, sulphates, chlorides, nitrates and
carbonates that are soluble in water. Group 2 compounds are likely to have
larger lattice energies due to the increased ionic charge, but don’t have high
enough hydration enthalpies for solubility.
f.
Thermal stability of the salts of the s-block elements
The stability of ionic compounds increases as;
- cationic radius decreases
- the charge on the ion increases
The lattice energies of the carbonates and oxides fall as you go down both
groups 1 and 2 because the cations are increasing in size. The distances
between the ions increases so the attractions become weaker.
The lattice energies of the carbonates and oxides fall at different rates due to
the different sizes of the anions. The oxide ion is relatively small whereas the
carbonate ion is much larger.
In a Bunsen flame group 1 carbonates are thermally stable with the exception
of lithium carbonate. However, the group 2 carbonates decompose to form
stable oxides. The temperature at which this happens increases down the
group. Beryllium carbonate doesn’t exist at room temperature it is so
thermally unstable.
The s-block nitrates decompose upon heating in a Bunsen flame. Group 1
nitrates (except lithium nitrate) form the corresponding nitrite e.g.
Group 2 nitrates (and lithium nitrate) form the corresponding oxide e.g.
The hydroxides of groups 1 and 2 follow the same pattern as the carbonates
and nitrates. All group 1 metal hydroxides are stable up to quite high
temperatures. All of the group 2 metal hydroxides decompose to give the
corresponding oxide e.g.
2.
a.
Group 7: the halogens
Physical properties of the halogens
The halogens are diatomic molecules which are very reactive and are strong
oxidising agents. The most common oxidation state is -1, although others
exist (except for fluorine).
Down the group they become less reactive, darker in colour and less volatile.
Halogen
Fluorine
Chlorine
Bromine
Iodine
Colour
State at room
temperature
Their solubility in water decreases down the group. Chlorine reacts in water
to form a mixture of hydrochloric acid and chloric acid.
Bromine reacts in a similar way. However iodine is almost insoluble in water
but is soluble in potassium iodide solution due to the formation of I3- ions.
(Refer to table 2.5.9 p188 for other physical properties)
Halogens are non-polar so they are more soluble in hydrocarbon solvents
than in water.
(Refer to fig 2.5.13 p189)
b.
Oxidation reactions of the halogens
i.
with metals
When halogens react with metals they become reduced to halide ions e.g.
ii.
with non-metals
This is done through covalent bonding e.g.
iii.
with iron (II) chloride solution
The halogen oxidises the pale green iron (II) chloride solution to brown iron
(III) ions e.g.
iv.
Disproportionation
The reaction of chlorine, bromine and iodine with sodium hydroxide depends
on the temperature. In cold temperatures a mixture of halide and halate (I)
ions are formed.
The halate (I) ions may decompose to form more halide ions and halate (V)
ions. The rate of decomposition depends on the halogen used and the
temperature.
Chlorine decomposes slowly at 15oC but rapidly at 70oC. For bromine both
reactions happen rapidly at 15oC and at 0oC the decomposition is prevented.
However, for iodine, both reactions are rapid at 0oC so it is difficult to
separate the products.
These are all examples of disproportionation since the halogen is both
oxidised and reduced in the same reaction.
v.
Redox reactions with the potassium halides
Potassium halides will react with another halogen in a displacement reaction
e.g.
c.
Reactions of the hydrogen halides
i.
with ammonia
This reaction produces ammonium halides e.g.
ii.
with concentrated sulphuric acid
Halides will react to form the hydrogen halide since the sulphuric acid is a
proton donor e.g.
However, with bromide and iodide compounds the sulphuric acid acts as an
oxidising agent to produce the halogen