Elements of the Sea
Ionic equations
o If 2 solutions react to form a solid, a precipitation reaction occurs.
Eg. Ag+(aq) + No3-(aq) + Na+(aq) + Br-(aq) goes to AgBr(s) +
Na+9aq0 + NO3-(aq)
Then cancel spectators
Ag+(aq) + Br-(aq) goes to AgBr(s)
o Always include state symbols
o Only split into ions if aqueous and ionic
Ionic solids
o Cations and anions are held together by strong electrostatic
attractions in a giant ionic lattice.
o Each sodium ion is surrounded by 6
chloride ions 6:6 coordination. It is a
simple cubic structure.
o Ionic solids are hard and have a high
melting point due to strong
electrostatic attractions between
ions.
o If ionic crystals contain water, the
water molecules fit into the lattice
the same way as regular ions. This is called the water of
crystallisation. These crystals are said to be hydrated.
o If these hydrated crystals are heated, the water molecules are
driven off and you are left with an anhydrous solid.
Ionic substances in solution
o As ionic substances dissolve in water, the ions become
surrounded by water molecules and they spread throughout the
solution. The hydrated ions are randomly arranged and behave
independently.
o Water surrounds the ions in this way because it is polar and bent.
The partially + H ions are attracted to the partially negative ions.
o In positive ions, the water molecules surround the ion with the
oxygen atoms pointing inwards, and with the hydrogen atoms
pointing inwards in negatively charged ions.
Rules on solubility
Soluble salts
All sodium, potassium and
ammonium salts.
Nitrates are always soluble
Most chlorides, bromides and
iodides.
Most sulphates.
Insoluble salts
All carbonates are insoluble except
for sodium, potassium and
ammonium carbonates.
Silver and lead chlorides, bromides
and iodides.
Calcium, strontium and barium
sulphate.
Charges on ions
Positive ions (cations)
Hydrogen H+
Sodium Na+
Silver Ag+
Potassium K+
Lithium Li+
Ammonium NH4+
Barium Ba2+
Calcium Ca2+
Copper(II) Cu2+
Zinc Zn2+
Magnesium Mg2+
Lead Pb2+
Iron(II) Fe2+
Iron(III) Fe3+
Aluminium Al3+
Negative ions (anions)
Chloride ClBromide BrFluoride FIodide IHydrogencarbonate HCO3Hydroxide OHNitrate NO3Oxide O2Sulphide S2Sulphate SO42Carbonate CO32-
Atoms and ions
First ionisation enthalpy
o The energy needed to remove one electron from every atom in a
mole of gaseous atoms.
o One mole of gaseous ions with one positive charge is formed.
X(g) goes to X+(g) + eo Units are kJmol-1 and FIE is always positive as energy must be put
in to remove the electron as it is attracted to the positive nucleus.
o The peaks are elements in
group 0 (noble gases). It is
difficult to remove an
electron from these atoms
because they have full
outer shells and are
unreactive.
o The troughs are elements in
group 1. These only have 1
outer shell electron so are
easy to ionise and are very reactive.
o It becomes harder to remove an electron across the period ( as
atomic number increases) due to the increasing nuclear charge.
o It becomes easier to remove an electron down the group as the
electron is further from the nucleus and more shielded by the
inner shells of electrons.
o It becomes harder to remove each successive electron because
once an electron is removed from the atom, the atom becomes
positively charged so the electrons left on (the ones you are trying
to remove) are very strongly attracted to the atom.
o There is a big jump in IE when an electron is removed from a full
outer shell eg. Aluminium has 3 electrons in its outer shell so the
4th IE is much bigger than the 3rd because an electron is being
removed from a full shell.
First IE: X(g) goes to X+(g) + eSecond IE: X+(g) goes to X2+(g) + eThird IE: X2+(g) goes to X3+(g) + eFourth IE: X3+(g) goes to X4+(g) + e-
Redox
OILRIG
Oxidation is losing (electrons) reduction is gaining (electrons).
Oxidation is an increase in oxidation states
Reduction is a decrease in oxidation states
Oxidising agents are reduced (decrease in oxidation states)
Reducing agents are oxidised (increase in oxidation states)
o ‘Oxyanions’ are negative ions that contain oxygen and another
element. Their names end in ‘ate’ eg.sulfate. They should include
an oxidation state in their name.
o Oxidation states should be included in names when the elements
in the compound can exist in more than one oxidation state. Eg.
FeO is called iron(II) oxide because O’s oxidation state is -2 so Fe’s
oxidation state must be +2.
o Half-equations for redox are as follows.
o If the element has been oxidised the electron goes on the
right
o If it’s been reduced, the electron goes on the left.
o Think reduced, right. But then switch it around!
A higher (more reactive) halogen oxidises a lower (less reactive) halide.
Eg. Cl2(g) + 2I-(aq) goes to 2Cl-(aq) + I2(aq)
This can be split into 2 half equations: Cl2(aq) +2e- goes to 2Cl-(aq)reduction
2I-(aq) goes to I2(aq) + 2e- oxidation
Disproportionation
In some redox reactions, it is the same element that is simultaneously
oxidised and reduced. This is called disproportionation. This occurs
when the products consist of 2 forms of the disproportionised element.
The Halogens
Physical properties
Flourine: Pale yellow gas, reacts with water, soluble in organic solvents.
Chlorine: Pale green gas, slightly soluble in water and organic solvents
to give a pale green solution. Forms a white precipitate of silver
chloride with silver ions.
Bromine: dark red liquid, forms a brown gas on warming. Soluble in
water and organic solvents to give red-brown and red solutions
respectively. Forms a cream precipitate of silver bromide with silver
ions.
Iodine: Shiny black solid, sublimes on warming to give a purple vapour,
barely soluble in water to give brown solution, soluble in organic
solvents to give violet solution. Forms a yellow precipitate of silver
iodide with silver ions.
o The intramolecular bonds in the halogens are covalent and the
intermolecular bonds are ID-ID bonds.
o As the size of the molecule increases, so does the strength of the
ID-ID bonds, as there are more electrons.
o This explains why the halogens change from gas to liquid to solid
down the group.
o Reactivity decreases down the group because the halogens get
bigger down the group sot the outer electrons are less attracted
to the nucleus.
o Halogens only displace (oxidise) lower halides.
o The general reaction of a halogen with silver ions is:
Ag+(aq) +X-(aq) goes to AgX(s)
o In redox, halogens act as oxidising agents (decrease oxidation
states – remove electrons from other elements)
X2 + 2e- goes to 2Xo Chlorine is used in water treatment, to make bleach, HCl and
plastics. It is highly toxic so is transported road or rail tanker as a
liquid.
o Bromine is used to make flame retardants, agricultural fumigants
and in photography. It’s transported in lead-lined steel tanks in
metal frames. Transport routes avoid residential areas and it is
often transported at night.
SPDF
Elements are split into number shells 1-4, these are split up into subshells, s, p, d and f.
S sub-shells have 1 orbital and can hold 2 electrons
P sub-shells have 3 orbitals and can hold 6 electrons
D sub-shells have 5 orbitals and can hold 10 electrons.
o Orbitals fill with one electron in each before going back and
putting 2 electrons in each orbital.
o Orbitals with 1 electron in have parallel spins.
o Orbitals with 2 electrons in have opposite spins.
1
1s
2
2s
3
4
4s is always filled before 3d because it has a lower energy and the sub
shells are filled in order of increasing energy.
Also, the d shells prefer to be exactly half full or completely full eg. 3d5
or 3d10 rather than 3d7 because this is lower energy, this therefore
sometimes leaves the 4s shell half full.
If the last electron in the element entered the s sub shell, it is said to be
in the ‘s-block’ of the periodic table. This is also true of the other 3
blocks, p, d and f.
Chemical Manufacturing
Extracting Bromine from sea water
o Partially evaporated, acidified sea water is warmed and chlorine is
added. It needs to be acidified as sea water is naturally slightly
alkaline and both chlorine and bromine dissolve in alkali.
o Steam is blown through the product. Bromine is given off as it is
volatile.
o The vapours are condensed and as bromine is not very soluble in
water, two layers are formed.
o The bromine is distilled and dried.
Batch Vs Continuous Process?
Batch. The starting materials are put into a vessel and allowed to react
together. The reaction is monitored and when complete, is terminated.
The product is then separated from the mixture. The process is repeated
batch by batch.
Continuous. The starting materials are fed in at one end and withdraw
at the other in a continuous flow.
Batch
Advantages
o
o
o
o
Cost effective for small quantities
Slow reactions are catered for
A range of products can be made in the same vessel
A greater % conversion for the same amount in the same time
than with continuous
Disadvantages
o Charging and emptying the vessel is time consuming as is hut
down time.
o Larger workforce needed than continuous
o Contamination is more likely as the same vessel is used for
different reactions.
o Exothermic processes are difficult to control
Continuous
Advantages
o
o
o
o
o
o
o
Suited to high tonnage
Greater throughput and no need to shut down
More controlled as very fine adjustment is possible
Contamination is less likely as is only used for one reaction
Consistent quality is ensured
Minimal workforce and labour needed
More easily automated
Disadvantages
o Much larger capital cost (start up costs) than batch
o Not cost effective if run below capacity
o Contamination risk is very high if used for two or more products.
Green Chemistry aims to reduce the use of feedstocks to a minimum eg.
Recycle unused reactants and solvents and increase atom economy.
Reduce energy consumption eg. Use enzymes to reduce temperature
needed or keep no. of steps to a minimum.
Minimise waste – recycle or find ways to make use of waste products.
Co-products are produced at the same time as the desired product via
the same reaction. As the amount of desired product increases, so does
the amount of co-products. Co-products can be sold for further profit.
By-products are the result of unwanted side reactions. The conditions of
any chemical process are designed to increase the amount of desired
product and decrease the amount of by-products.
The Chloralkali Process
The mercury cell
This is being phased out by 2020 because it very expensive to run and
produces toxic mercury emissions.
The membrane cell
Advantages:
o lower running costs than mercury as less energy used per tonne of
chlorine produced.
o Larger capacity for chlorine in the same space than mercury cell
o No need to remove toxic mercury from the products.
o Less environmental pollution.
Reaction at the positive anode: 2Cl- goes to Cl2 +2eReaction at negative cathode: 2H2O +2e- goes to 2OH- +H2
Overall reaction: 2Cl- +2H2O goes to Cl2 +2OH- +H2
The membrane is made from the PTFE polymer and has negatively
charged side chains which the Na+ ions are attracted to. These side
chains repel Cl- and OH- ions. Some water molecules are attached to
the sodium ions and pass through but the membrane is impermeable to
free water molecules.
The cell is designed to:
o Prevent chlorine from reacting with OHo Minimise Cl- ions diffusing into the negative electrode
o Minimise the OH- ions from diffusing into the positive electrode
o Prevent hydrogen and chlorine from mixing because the mixture
is explosive.
% yield and atom economy
% yield =
actual mass of product
Theoretical maximum mass of product
x100
This measures the efficiency of the reaction in terms of waste and how
much of the theoretical product is actually produced.
% Atom economy = Mr of useful product
Mr of all reactants
x100
This measures the efficiency of the reactants in the process and how
much ends up in the products.
Titration Calculations
N
1dm3 = 1000cm3 so to get from dm3 from cm3, divide
cm3 by 1000.
C
V
Steps for titrations calcs:
1. write out balanced equation
2. work out no. of moles (n=c x v) of one you know c and v for.
Remember to divide v by 1000 to change to dm3.
3. write down mole ratio, if 1:1 skip to next step. If the ‘n’ you just
worked out in step 1 is the smaller ratio then times ‘n’ by the
bigger mole ratio for the next purposes. If ‘n’ is the larger of
the 2 then divide by the larger mole ratio.
4. carry out c=n/v using your previously worked out ‘n’ value and
the ‘v’ given in the question (remember to change it into dm3
so 25 cm3 is 0.025 dm3)
Intermolecular Bonds
o bond polarity depends on the electronegativity difference
between the atoms
o molecular polarity depends on the electronegativity difference
and the shape of the molecule. If molecule contain a NOF atom
and are not symmetrical, they are likely to contain a PD-PD bond.
ID- ID bonds
These occur in all molecules and are the weakest type of intermolecular
bonds.
Bond enthalpy
A Dipole occurs when a molecule has a positive end and a negative end
When a molecule has a dipole we say it is polarised. If a molecule has a
permanent dipole it is called a polar molecule.
The constant movement of electrons in a molecule means that at any
one time the electron density may be unevenly distributed, this creates
an instantaneous dipole. The polarity of the molecule may change
because the electron density is constantly moving. If other molecules
are close to the molecule with a dipole then electrons get attracted to
the positive end of the original molecule, this induces a dipole in the
new molecule. The instantaneous dipole and the induced dipole attract
each other.
ID-ID bonds are continuously forming and breaking because of the
constant movement of electrons within the molecules.
The bigger the molecule, the more electrons it has so the greater the
attraction and the strength of the ID-ID bond so the higher the boiling
point.
o A bigger molecule also has a larger surface area so has more
points of contact with neighbouring molecules.
o A long chain molecule has stronger ID-IDs because it has more
points of contact.
o A straight chain molecule has stronger ID-IDs than a branched
chain because it can line up more closely so has more points of
contact.
Poly(ethene) contains only ID-ID bonds but is solid at room
temperature. This is because the chains are long and can pack close
together meaning that there are many ID-ID bonds which compensates
for their weakness.
PD-PD bonds
o Molecules with PD-PD bonds contain atoms with different
electronegativity values.
o
The partially positive end of one molecule attracts the
partially negative end of the other.
o PD-PD bonds are stronger than ID-ID bonds but weaker than
hydrogen bonds.
o PD-PD binds hold polyester molecules together.
o Occurs if the molecule is not symmetrical and the dipole
moments don’t cancel each other out.
o PD-PD bonds can also induce dipoles in neighbouring molecules.
Electronegativity is the ability of an atom to attract the bonding
electron in a covalent bond
The bigger the difference in electronegativity a molecule has, the more
ionic it is.
Halogenoalkanes
o The homologous series of a halogenoalkane is R-X or R-Hal where
X (Hal) is a halogen – Cl, Br or I.
o All halogenoalkanes are immiscible with water.
o The larger the halogen the higher the boiling point because there
will be stronger ID0ID bonds between the molecules.
Homolytic Fission
o This forms radicals of both the alkane and the halogen.
o The condition is radiation of the right frequency (visible or UV)
which is absorbed by the halogenoalkanes.
o This kind of reaction occurs when halogenoalkanes reach the
stratosphere which is what forms chlorine radicals.
o CH3-Cl + hv goes to CH3. +Cl.
o One electron goes to each atom
Heterolytic Fission
o In Heterolytic fission, the halogenoalkanes tend to react in a polar
solvent such as ethanol, or ethanol and water. This is the
condition.
o No radiation is needed in this process and no radicals are formed.
o Instead a negative halide ion and a positive carbocation are
formed by the complete breakage of the C-Hal bond.
o Here, both electrons go to the halogen to form a halide ion.
o The reaction is as follows:
Reactivity trends
o Strength of C-Hal bond decreases down the group because the
size of the halogen atom increases.
o Therefore reactivity increases.
o Bond strength rather than polarity has the greatest effect on the
reactivity of the halogenoalkanes.
o This is why Br bonds break down in the troposphere whereas Cl
bonds only break down in the stratosphere, wreaking havoc on
the ozone layer.
Nucleophilic substitution reactions
o The general equation where X- is a nucleophile is as follows:
R-Hal + X- goes to R-X + Halo The C-Hal bond breaks and the halogen atom is replaced by the
nucleophile.
o The mechanism for the Nucleophilic substitution of a
halogenoalkanes is as follows:
o The C-I bond is polar so the carbon atom has a partially + charge
and the iodine atom has a partially – charge.
o The OH- ion is negatively charged so the oxygen is attracted to
the partially positive C atom.
o A lone pair of electrons move from the oxygen atom to the
carbon atom as the C-I bond breaks.
o Iodine ions are formed, however a free carbocation is not formed
because the OH- attacks at the same time that the C-Br bond
breaks.
A nucleophile is a molecule or negatively charged ion with a lone pair
of electrons that it can donate to a positively charged atom to form a
covalent bond.
All halogens have 3 lone pairs
The negatively charged atom in the nucleophile always donates the lone
pair.
Water as a Nucleophile
o The reaction with water as nucleophile is slower because water
does not have a full negative charge.
o The reaction happens in two steps; first the water attacks the
halogenoalkane in much the same way as a normal nucleophile.
o The second step involves the resulting positively charged ion from
the first step, loses H+ to form an alcohol.
o The overall equation for this reaction is:
R-Hal + H2O goes to ROH + H+ + Halo This reaction can also be done in reverse to produce
halogenoalkanes.
RHal + OHRHal + H2O
RHal + NH3
ROH + Hal-
Conditions
Heat under reflux with NaOH(aq).
Often ethanol is added as a solvent
Heat under reflux, sometimes
called hydrolysis.
The halogenoalkane is heated with
concentrated ammonia in a sealed
tube.
In presence of a strong acid.
Preparation of halogenoalkanes
Carrying out the reaction
o Weigh 5g of 2-methylpropan-2-ol and pour into a separating
funnel.
o Gradually add 20cm3 of concentrated HCl
o Shake the mixture in the separating funnel for 20 minutes
releasing the pressure from time to time.
Separating the required product from the reaction mixture
o Allow the mixture to stand until the 2 layers have separated. The
alcohol (CH3)3COH will be on the bottom and (CH3)3CCl will be
on the top layer.
o To remove the excess acid, add sodium hydrogen carbonate
solution but be careful to release the pressure as carbon dioxide is
formed. Run off the lower aqueous layer. Repeat until no more
gas is given off.
o Add 10cm3 distilled water and run off aqueous layer again.
o Run organic (bottom) layer (CH3)3CCl into a clean conical flask.
o Add anhydrous sodium sulphate (a drying agent) to remove any
traces of water.
Purifying the product
o The mixture is transferred to a flask and then distilled.
Name
Hydroxide ion
Formula
OH-
Cyanide ion
CN-
Ethanoate ion
CH3COO-
Ethoxide ion
C2H5O-
Water molecule
H2O
Ammonia molecule
NH3
Structure