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 spectatorsAg+(aq) + Br-(aq) goes to AgBr(s)

o Always include state symbolso 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 Insoluble saltsAll sodium, potassium and ammonium salts.

All carbonates are insoluble except for sodium, potassium and ammonium carbonates.

Nitrates are always solubleMost chlorides, bromides and iodides.

Silver and lead chlorides, bromides and iodides.

Most sulphates. Calcium, strontium and barium sulphate.

Charges on ions

Positive ions (cations) Negative ions (anions)Hydrogen H+ Chloride Cl-Sodium Na+ Bromide Br-Silver Ag+ Fluoride F-Potassium K+ Iodide I-Lithium Li+ Hydrogencarbonate HCO3-Ammonium NH4+ Hydroxide OH-Barium Ba2+ Nitrate NO3-Calcium Ca2+ Oxide O2-Copper(II) Cu2+ Sulphide S2-Zinc Zn2+ Sulphate SO42-Magnesium Mg2+ Carbonate CO32-Lead Pb2+

Iron(II) Fe2+

Iron(III) Fe3+

Aluminium Al3+

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) + e-o 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) + e-Second IE: X+(g) goes to X2+(g) + e-Third IE: X2+(g) goes to X3+(g) + e-Fourth IE: X3+(g) goes to X4+(g) + e-Redox

OILRIGOxidation is losing (electrons) reduction is gaining (electrons).

Oxidation is an increase in oxidation statesReduction 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 righto 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

D isproportionation

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 2X-o 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 sub-shells, s, p, d and f.

S sub-shells have 1 orbital and can hold 2 electronsP sub-shells have 3 orbitals and can hold 6 electronsD 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 Cost effective for small quantities o Slow reactions are catered foro A range of products can be made in the same vesselo 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 continuouso Contamination is more likely as the same vessel is

used for different reactions.o Exothermic processes are difficult to control

Continuous

Advantages

o Suited to high tonnageo Greater throughput and no need to shut downo More controlled as very fine adjustment is possibleo Contamination is less likely as is only used for one

reactiono Consistent quality is ensuredo Minimal workforce and labour neededo More easily automated

Disadvantages

o Much larger capital cost (start up costs) than batcho Not cost effective if run below capacityo 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 cello No need to remove toxic mercury from the products.o Less environmental pollution.

Reaction at the positive anode: 2Cl- goes to Cl2 +2e-Reaction at negative cathode: 2H2O +2e- goes to 2OH- +H2Overall reaction: 2Cl- +2H2O goes to Cl2 +2OH- +H2

C V

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 OH-o Minimise Cl- ions diffusing into the negative electrodeo Minimise the OH- ions from diffusing into the positive

electrodeo 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

1dm3 = 1000cm3 so to get from dm3 from cm3, divide cm3 by 1000.

Steps for titrations calcs:1. write out balanced equation

N

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 + Hal-o 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+ + Hal-

o This reaction can also be done in reverse to produce halogenoalkanes.

ConditionsRHal + OH- Heat under reflux with

NaOH(aq). Often ethanol is added as a solvent

RHal + H2O Heat under reflux, sometimes called hydrolysis.

RHal + NH3 The halogenoalkane is heated with concentrated ammonia in a sealed tube.

ROH + Hal- 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 HClo 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 Formula StructureHydroxide ion OH-

Cyanide ion CN-

Ethanoate ion CH3COO-

Ethoxide ion C2H5O-

Water molecule H2O

Ammonia molecule

NH3