Chemistry 2 AQA

• Thermochemistry- the study of heat changes during chemical reactions
• Energy must be put in to break bonds and energy is given out when bonds are formed.
• Exothermic- energy given out (e.g. neutralising an acid with an alkali)
• Endothermic- energy taken in (e.g. the break down of limestone to lime and carbon dioxide)
• A reaction that is exothermic in one direction will be endothermic in the other.

• When we measure a heat change at constant pressure it is called an enthalpy change (ΔH)
• Standard conditions-100kPa at 298k
• A reaction is not over until the products have cooled back to 298k
• In an exothermic reaction the products end up with less energy that the starting materials because they have lost heat energy when they heated up their surroundings. This means that ΔH is negative.
• In an endothermic reaction the products end up with more energy that the starting materials, so ΔH is positive.

• Standard molar enthalpy of formation- the enthalpy change when one mole of compound is formed from its constituent elements under standard conditions, all reactants and products in their standard states.
• Standard molar enthalpy of combustion- the enthalpy change when one mole of compound is completely burned in oxygen under standard conditions, all reactants and products in their standard states.
• To measure enthalpy change we need to know- the mass of the substance being heated/cooled, the temperature change, the SHC of the substance.
• enthalpy change= mass of substance x SHC x temperature change(q=mc x ΔT)
• acid + alkali →salt + water
• A metal that is more reactive than another with displace the less reactive one from the compound.

• Hess's Law- the enthalpy change for a chemical reaction is the same, whatever route is taken from reactants to products.
• We can work out the energy change on a thermochemical cycle or on an enthalpy diagram.
• The enthalpies of all elements in their standard states (at 298k and 100kPa) are taken as zero.

• Bond Dissociation Enthalpy- the enthalpy change required to break a covalent bond with all species in gaseous state
• Mean Bond Enthalpy- the average value of the bond dissociation enthalpy for a given type of bond taken from a range of different compounds

• For a reaction to take place between two molecules they must collide hard enough with enough energy to break bonds.
• Most collisions between molecules or other particles do not lead to a reaction. they either don't have enough energy or are in the wrong orientation.
• Factors that effect the rate of reaction:
• Increasing Temperature: Increases speed of molecules and therefore increases their energy and number of collisions.
• Increasing concentration of a volume:more particles in a given volume therefore more collisions likely, therefore faster reaction rate. However, as reactants used up conc falls and reaction goes on. (increasing pressure has same effect.)
• Increasing surface area of solid reactant: more particles able to collide at one time and so more site for reaction.
• Using a catalyst: changes rate of reaction without being used up itself.
• Activation energy: minimum energy needed to start a reaction
• The top of an enthalpy diagram is called the transition state or activated complex.

• x axis- fraction of particles with energy E. y axis- energy E
• The MBD shows that:
• -No particles have zero energy
• -Most particles have intermediate energies- around the peak of the curve
• -Very few have very high energies (RHS of curve). There is no upper limit.
• -The average energy is not the same as the most probable energy.
• At a higher temperature the curve is lower and moves to the right. The number is particles with very high energy increases. The area beneath the curve is the same is it represents the total number of particles.
• The graph shows that at higher temperatures more of the molecules will have an energy greater than the activation energy so more collisions will result in a reaction.

• Catalysts work because they provide a different pathway for the reaction, one with a lower Ea. Catalysts do not affect the enthalpy change of the reactions, nor do they affect the position of equilibrium (=m)
• Heterogeneous catalysts- where the catalyst is in a different phase to the reactants- usually a solid catalyst and liquid or gaseous reactants.
• Homogeneous catalysts- where catalyst and reactants are in the same phase.
• Catalytic converters are honeycombs, made of ceramic material coated with platinum and rhodium metals (the catalysts). Honeycomb shape provides enormous surface area.
• carbon monoxide + nitrogen oxides→nitrogen + carbon dioxide
• hydrocarbons + nitrogen oxides→nitrogen + carbon dioxide + water
• On the catalyst- gases form weak bonds with metal atoms of catalyst (adsorption). This holds gases in correct position to react. Gases react on the surface.
• Products break away from metal atoms (desorption). This frees up room on catalyst for more gases to react.

• Zeolites- minerals that have open pore structure that ions/molecules can fit into. It confines molecules in small spaces so that changes occur in their structure and reactivity. 150 have been synthesized and 48 naturally occurring zeolites are known.
• Hardening fats- Unsaturated fats are 'hardened' when hydrogen is added across some of the double bonds. This is done by bubbling hydrogen into the liquid fat which has a nickel catalyst mixed with it. Nickel is filtered off after the reaction.
• CFC's were previously widely used until they realised that they escaped high into the atmosphere where they stay as they are so nonreactive (due to strength of C-H bond). They eventually decompose to produce Cl atoms which act as catalysts to the destruction of the ozone layer. Ozone is important is it prevents too much UV reaching the Earth's surface.
• O3(g) + O(g) chlorine atom catalyst→2O2(g)
• At the moment HCFC's and HFC's are temporarily being used as a substitute.

• Equilibrium mixture- the mixture of reactants and products formed when a reversible reaction is allowed to proceed in a closed container until no further change occurs. The forward and backward reactions are still proceeding but at the same rate.
• Dynamic Equilibrium- a situation in which the composition of a reaction mixture does not changes because both forward and backward reactions are proceeding at the same rate.
• The conditions for equilibrium (=m):
• =m can only be reached in a closed system.
• =m can be approached from either direction and the final =m position will be the same.
• =m is a dynamic process. It is reached with the rates of the two opposing processes, which are going on at the same time, are the same.
• We know that =m has been reached when the macroscopic properties (density, conc., colour and pressure) of the system do not change with time.

Note- an =m mixture can have any proportions of reactants and products. It does not necessarily have to be half and half.

• If the proportion of products in the =m mixture is increased then we say =m has moved to the right (or forwards).
• If the proportion of reactants in the =m mixture is increased then we say =m has moved to the left (or backwards).
• Le Châtelier's Principle- If a system at =m is disturbed, the =m moves in the direction that tends to reduce the disturbance
• Changing conc.- increasing the conc. of one of the reactants will cause =m to shift in the directing that will reduce the conc. of this reactant. The was to do this would be to reactant the reactant with another reactant thus producing more product
• Changing pressure-this only affects reactions with gases. It is equivalent to increasing the conc. in a solution.
• Changing temperature- LCP tells us that if we increase the temperature, =m move in the direction to cool the system down so it will move in the direction to absorb heat (endothermic). So in an exothermic reaction it will move to the left and the =m mixture will contain a greater proportion of reactants. If we cooled the mixture =m would move to the right and we would have a greater proportion of products.
• Catalysts- Catalysts have no effect on the position of =m so they do not alter the composition of the =m mixture.

• The Haber Process- makes ammonia:
• The raw materials for the process are air, which provides nitrogen, water and methane. These provide hydrogen:
• CH4(g) + H20(g)→CO(g) + 3H2(g)
• N and H are fed into a converter in a 3:1 ratio and passed over an iron catalyst (20000kPa and 670k are compromise conditions). Nitrogen and hydrogen are continuously passed over the catalyst until =m is reached and a 15% conversion to ammonia is reached. The NH3 is cooled so that it becomes liquid and is piped off.
• 80% of NH3 is used to make fertilisers (ammonium sulfate, ammonium nitrate and urea). It is also used in the manufacture of nylon, explosives and dyes.
• Ethanol (C2H5OH): made by fermentation from sugars with a yeast catalyst but the main source for it is from crude oil by fractional distillation and then cracking.
• Ethanol is made by the hydration of ethene, the reaction is reversible, it is speeded up by a catalyst is phosphoric acid absorbed on silica.
• H2C=CH2(g) + H20(g) ↔ CH3CH2OH(g)
• All reactants and products are gaseous and compromise conditions are used of about 570k and 6500kPa giving a conversion to ethanol of about 5% but unreacted ethene is recycled over the catalyst again.

• Methanol (CH3OH): used as chemical feedstock and is used in the manufacture of plastics (perspex and terylene) and methanal (formaldehyde).
• CO(g) + 3H2(g) ↔ CH3OH(g)
• The reaction uses a copper catalyst and uses compromise conditions of 500k and 10000kPa and produces around 5-10% yield.

• OIL RIG- oxidation is loss, reduction is gain
• reducing agents give away electrons- they are electron donors
• oxidising agents accept electrons
• Half equations are used to show the electron transfer in a redox reaction
• Spectator ions take no part in a reaction and so are not put in a half equation

• In an ionic compound the oxidation state simply tells us how many electrons is has lost or gained compared to the element in its uncombined state.
• In a molecule the oxidation state tells us about the distribution of electrons between elements of different electronegativities.
• Every element in it uncombined state have an oxidation state of zero.
• A positive number shows that the element has lost electrons and has therefore been oxidised.
• A negative number shows that the element has gained electrons and has therefore been reduced.
• H= +1, Group 1= +1, Group 2= +2, Al= +3, O= -2, F= -1 and Cl= -1
• The sum of all oxidation states in a compound is 0.
• The sum of oxidation states of a complex ion equals the charge on the ion.
• In a compound the most electronegative element always has a negative oxidation state.

• Elements exist as diatomic molecules
• Fluorine= pale yellow gas, Chlorine= greenish gas, Bromine= red-brown liquid, Iodine= black solid so get darker and denser as we go down the group.
• All have characteristic 'swimming bath' smell.
• Fluorine has a number of untypical properties that stem from the unexpectedly weak F-F bond. The small size of the fluorine atom leads to repulsion between the non-bonding electrons because they are so close together.
• Size of atoms gets bigger as we go down the group as there are added electron shells.
• Electronegativity decreases down the group despite increasing nuclear charge because of shielding by inner shells.
• Melting and boiling points increase down the group because the larger atoms have more electrons and this makes the van der Waals forces between the molecules stronger. The lower the boiling point the more volatile the liquid.

• Halogens usually react by gaining electrons to become negative ions (-1 charge). The oxidising ability of the halogens increases as we go up the group.
• Halogens will react with metal halides in solution in such a way that the halide in the compound will be displaces by a more reactive halogen (displacement reaction).
• The oxidation of a halide by a halogen is the basis of a method for extracting bromine from sea water which contains small amounts of bromide ions which can be oxidised by chlorine to produce bromine.
• Cl2(aq) + 2Brˉ(aq) → Br2(aq) + 2Clˉ(aq)

• Halide ions can act as reducing agents, they lose elections to become halogen molecules.
• The larger the ion the more easily it loses an electron therefore their reducing power increases down the group.
• Reactions with conc. sulfuric acid:
• Sodium Chloride:Observation- steamy fumes of hydrogen chloride and solid sodium hydrogensulfate. NaCl(s) + H2SO4(l) → NaHSO4 + HCl(g)This is not a redox reaction as no oxidation state has changed. The chloride ion is too weak a reducing agent to reduce the sulfur to sulfuring acid. It is an acid-base reaction.
• Fluoride ions are even weaker reducing agents that Chloride so the reaction would be similar.
• Sodium Bromide: Observation- steamy fumes of hydrogen bromide and brown fumes of bromine (colourless sulfur dioxide also formed). 1) similar acid base reaction: NaBr(s) + H2SO4(l) → NaHSO4(s) + HBr(g) 2) bromide ions are strong enough reducing agents to reduce sulfuric acid to sulfur dioxide 2H+ + 2Br- + H2SO4(l) → SO2(g) + 2H2O(l) + Br2(l)This is an exothermic redox reaction and some of the bromine vaporises.

• Sodium Iodide: Observations- steamy fumes of hydrogen iodide and black solid iodine plus the bad egg smell of hydrogen sulfide gas. Yellow solid sulfur may also be seen (colourless sulfur dioxide is also formed). 1) Acid-Base reaction: NaI(s) + H2SO4(l) → NaHSO4(s) + HI(g) 2) Iodide ions are even better reducing agents that bromide so they reduce the sulfur in sulfuric acid so that sulfur dioxide, sulfur and hydrogen sulfide gas are produced 8H+ +8I- + H2SO4(l) → H2S(g) + 4H2O(l) + 4I2(s)
• All metal halides react with silver ions in aqueous solution. The reaction can be used as a test for halide ions because of the colour of the precipitate that is formed and its solubility in ammonia.
• Silver fluoride: no precipitate, Silver chloride: white precipitate which dissolves in dilute ammonia, Silver bromide: cream precipitate which dissolves in conc. ammonia, Silver iodide: pale yellow precipitate which is insoluble in conc. ammonia

• Chlorine is poisonous and soluble in water.
• Cl2(g) + H2O(l) ↔ HClO(aq) + HCl(aq) In this reaction one of the chlorine atoms oxidation number increases from 0 to +1 and the other decreases from 0 to -1, this is called disproportionation. This reaction takes place when chlorine is used to purify water.
• Chlorine is rapidly lost from pool water in sunlight so it needs topping up.
• An alternative is to add solid sodium/calcium chlorate which dissolves in water to form chloric acid. NaClO(s) + H2O ↔ Na+(aq) + OH- + HClO(aq) In alkaline solution =m moves to the left and HClO is removed as ClO- ions therefore pools need to be kept slightly acidic.
• Chlorine reacts with cold NaOH to form NaClO in a disproportionation reaction. Cl2(g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H2O(l) Other halogens react similarly.

• As we go down the group the electrons in the 'sea' of delocalized electrons are further from the positive nuclei and so the strength of the metallic bonding decreases we go down the group. This is why the melting points decrease slightly down the group.
• Metals get more reactive as we go down the group.
• Reaction with water (with M as any group two metal): M(s) + 2H2O(l) → M(OH)2(aq) + H2(g)
• Magnesium hydroxide is used in indigestion remedies to neutralise excess stomach acid. Calcium hydroxide is used to treat acidic soil.
• Magnesium reacts very slowly with cold water but rapidly in steam to form an alkaline oxide and hydrogen. Mg(s) + H2O(g) → MgO(s) + H2(g)
• Ca reacts the same way but more vigorously, and Sr and Ba even more vigorously.
• All Group 2 metal hydroxides become more soluble down the group. The hydroxides are all white solids.
• Barium hydroxide dissolves to produce a strongly alkaline solution. Ba(OH)2(s) + aq → Ba2+(aq) + 2OH-(aq)

• Sulfates become less soluble down the group.
• Therefore Barium Sulfate is insoluble and so can be taken as a 'barium meal' to line the gut for x-rays.
• It is also used for a test for sulfate ions. Solution is acidified with nitric or hydrochloric acid and then barium chloride solution is added to the solution under test. If a sulfate is present then a white precipitate of barium sulfate is formed. Ba2+(aq) + SO4 2-(aq) → BaSO4(s)
• The addition of acid removes carbonate ions as carbon dioxide since barium carbonate is also a white solid and therefore indistinguishable against barium sulfate.

• Metals are usually found in ores combined with oxygen or sulfur. The oxides and sulfides of the metals have positive oxidation states ( the metals have a state of 0). So to separate the metal from its ore it has to be reduced.
• Before reduction sulfide ores are converted to oxides by heating them in air (roasting). e.g. Zns(s) + 1.5O2(g) → ZnO(s) + SO2(g) The by-product of sulfur is a problem as it contributes to acid rain if it escapes and it converted to sulfuric acid. But it can be collected and converted to H2SO4 in controlled conditions and sold. SO2(g) + H2O(g) + 0.5O2(g) → H2SO4(l)
• Possible reducing agents: Coke; cheaply obtained by burning coal in the absence on air, but some metals need a high temperature for the reaction and if it is a reactive metal then they will react with carbon to form carbides, hydrogen; made from methane and water, used to extract tungsten, electrolysis; for metals high up in the reactivity series (Na or Al), or a more reactive metal.

• Iron is found in the ores magnetite (Fe3O4) and haematite (Fe2O3) with the major impurity silica (SiO2). When extracted coke is used in blast furnaces up to 70m high and is run continuously. The hopper is charged with iron, coke and limestone. Coke burns in the base in a blast of hot air (generated by exothermic process so temperature is around 2000k where melting point of iron is 1808k). C(s) + O2 → CO2(g) then CO2(g) + C(s) → 2CO(g) Carbon monoxide is the reducing agent and so reacts with iron (Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)) to produce molten iron.

• Aluminium is extracted from purified bauxite (largely aluminium oxide). Oxide is dissolved in molten cryolite to form a solution which melts around 1240k compared to Al2O3 at 2345k. The solution is electrolysed in rows of cells using currents up to 300000A. Aluminium is produced at the negative electrode (the steel casing of the shell) and oxygen at the positive electrode (made of carbon). The overall process is the decomposition of aluminium oxide. Al203 → 2Al + 1.5O2. The process runs continuously and the solution is kept molten by the temperature of the reaction. Al is formed as a liquid and is siphoned off while the oxygen burns the carbon to form CO2 and therefore the carbon electrodes have to be regularly replaced.
• Titanium is strong, low density, resistant to corrosion and its ores are abundant. It can't be reduced with carbon because it makes the metal brittle. The ore (rutile) is converted titanium (IV) chloride by reacting it with coke and chlorine. TiO2(S) + 2C(s) + 2Cl2(g) → TiCl2(l) + 2CO(g) The liquid titatnium is purified by distillation. TiCl2 is reduced with molten Na under an inert argon atmosphere at 1300k. TiCl4(l) + 4Na(l) → Ti(l) + 4NaCl(l). It is a batch process.
• Tungsten is rare and is useful in light bulb filaments as it has high melting points. It can't be reduced by carbon because a carbide is formed so it is extracted by reduction with hydrogen. WO3 + 3H2 → W + 3H2O

• Metals are easy to recycle because they can be melted and reformed:
• -This reduces the scrap iron that is put in refill sites
• -Scrap iron has already been extracted from its ore and it easy to extract from other materials as its magnetic
• -Extracting iron from its ore contribute to the greenhouse effect but energy to melt the metal produced CO2.
• -Recycled aluminium uses only about 5% of the amount of energy to extract it.
• -Recycling avoids production of CO2. However, this gives rise to financial and energy costs involved in sorting and transporting recycled material.

• Halogens are more electronegative than carbon and so the bond is polar, Cδ+ - Xδ-. As we go down the group the bonds become less polar.
• The C-X bonds are not polar enough to make the haloalkanes soluble in water. The main force of attractions are dipole-dipole (DD) and van der Waals (VDW). Haloalkanes can mix with hydrocarbons so they can remove oily stains.
• Boiling point increases with chain length and as we go down the halogen group. Both effects are due to increased VDW forces. Increased branching decreases the boiling point. Haloalkanes have higher boiling points that alkanes of similar chain lengths because a) they have higher RAM and b) they are more polar.
• When haloalkanes react the C-X usually breaks. How readily they react depends on the bond enthalpy and polarity.
• Bond polarity: Since the carbon is partially positive it can be attacked by nucleophiles (electron pair donor). The polarity of the bond determines how reactive it'll be. C-F is most reactive and C-I is the least reactive.
• Bond enthalpy: Bonds get weaker down the group. The weakest bond is the most reactive. So C-F is the strongest (and the least reactive) and C-I is the weakest (and the most reactive).

• A nucleophile is negatively charged and is a reagent that attacks and forms bonds with positive carbon atoms. It has a lone pair which it can use to form a covalent bond. Common nucleophiles are hydroxide ions, ammonia and cyanide ions.
• When reacting with aqueous sodium/potassium hydroxide the haloalkane can't mix with water so ethanol is used as a solvent (hydrolysis reaction).
• The product of substitution with a cyanide ion is called a nitrile

• In elimination the OH- can be used as a base, removing the H+ ion from the haloalkane.
• Process of elimination: the OH- ion uses its lone pair to form a bond with one of these hydrogen atoms on the carbon next to the C-X bond. The electron pair from the C-H bond now becomes part of the C-C double bond. The halogen takes the pair fo electrons from the C-X bond and leaves as an ion.
• Hydroxide ions are room temperature, dissolved in water, favour substitution.
• Hydroxide ions at high temperature, dissolved in ethanol, favour elimination.
• Primary haloalkanes tend to react by substitution and tertiary ones by elimination. Secondary do both.

• Haloalkanes form when you put a mixture of an alkane and a halogen under UV light.
• Free radical substitution: a chain reaction which takes place in initiation, propagation and termination.
• Initiation: Chlorine molecules absorb the energy of one quantum of UV light and the energy of the UV is greater than the bond so it breaks to form free radicals which are highly reactive.
• Propagation: 1) The Chlorine free radical takes a H atom from methane to form HCL (stable) which leaves a methyl free radical. 2) The methyl free radical is also reactive and reacts with a chlorine molecule. This produces a chlorine free radical which is ready to react with more methane. This continues thousands of times over until all the chlorine and methane is used up.
• Termination: All the free radicals are removed by all the free radicals reacting with each other no matter whether they are identical or not.

• They have the general formula CnH2n.
• Alkenes can form position isomers and geometrical isomers.
• Position isomers: Isomers where the double bond is in a different position.
• Geometrical isomers: Isomers have the same structural formula but the bonds are arranged differently in space.
• The physical properties of the alkenes are very similar to the properties of alkanes i.e, VDW are the only intermolecular forces, melting and boiling points increase with the number of carbon atoms present, are insoluble in water.
• Alkenes are more reactive that alkanes because of the electron rich area of the double bond which can be easily attacked by positively charged reagents.

• Electrophillic addition: The electrophile is attracted to the double bond. Electrophiles are positively charged and accept a pair of electrons from the double bond.A carbocation is formed (a positive ion). A negatively charged ion forms a bond with the carbocation.
• This addition reaction is used to test for C-C double bonds. When a few drops of bromine solution are added to an alkene, the solution is decolorized because the products are colourless.

• When alkenes, such as ethene, polymerise to form poly(ethene) despite the 'ene' ending, it is actually an alkane and unreactive. The reaction is addition polymerisation.
• Polymers are made from joining up monomers.
• Uses of poly(ethene): washing up bowls, plastic bags
• Uses of poly(propene): rope
• Poly(propene) is a thermoplastic polymer and so will soften when heated so it can be melted and reused but can only be done a limited number of times because when it is heated some of the chains break and become shorter thus degrading the plastics properties.

• The OH group in alcohols means that hydrogen bonding occurs between the molecules and so alcohols have a higher melting and boiling point than alkanes. Hydrogen bonding also makes the shorter chain alcohols soluble in water because hydrogen bond can form between the alcohol and the water.

• From crude oil: when the fractions are cracked ethanol is made from ethane (and a little from the naphtha fraction). Ethene is then hydrated.
• By fermentation: Carbs from plants are broken down into sugars and then converted by yeast into ethanol. Carbs come from sugar cane and sugar beet. The process is anaerobic respiration: C6H12O6(aq) → 2C2H5OH(aq) + 2CO2(g) The rate is affected by temperature so a compromise temperature is used of 35˚. Air is kept out to prevent oxidation of ethanol to ethanoic acid. Once the fermenting solution contain 15% ethanol the enzymes can't function and the reaction stops.
• Ethanol is used as a motor fuel mixed with petrol. It is made by renewable crops and so it a bio fuel because when the ethanol is burned the same amount of CO2 goes back into the atmosphere as was locked in the plant so it is carbon neutral.

• Alcohols burn completely to CO2 and H20. Combustion is usually complete oxidation. Primary alcohols are oxidised to aldehydes, RCHO, and can then be further oxidised to carboxylic acids, RCOOH. Secondary alcohols are oxidised to ketones, R2CO. Tertiary alcohols are not easily oxidised because oxidation would need a C-C bond to break rather than a C-H bond.
• A solution of potassium dichromate, acidified with dilute sulfuric acid is used to oxidise alcohols to aldehydes and ketones. In the reaction the orange dichromate (VI) ions are reduced to green chromium (III) ions.
• To oxidise ethanol to ethanal: use dilute acid and less potassium dichromate (VI) than is needed. Heat the mixture gently. CH3CH2OH(l)+[O]→CH3CHO(g)+H2O(l)
• To oxidise ethanol to ethanoic acid: use conc. sulfuric acid and mroe than enough potassium dichromate (VI) and reflux the mixture in the apparatus. CH3CH2OH(l) + 2[O] → CH3COOH(g) + H20(l)
• Secondary alcohols are oxidised to ketones by acidified dichromate.
• Aldehydes and ketones both have a C=O group called a carbonyl group. In aldehydes it is at the end of the hydrocarbon chain and in ketones it is in the body of the chain. Aldehydes have the suffix -al and ketones have the suffix -one.

• Tests to distinguish between aldehydes and ketones:
• The Tollens Test: Tollens reagent is a oxidising agent made from silver nitrate in aqueous ammonia. It oxidises aldehydes but has no effect on ketones. It contains colourless silver (I) complex ions which are reduced to metallic silver.
• The Fehling's/Benedict's test: The are both reagents which contain blue copper (II) complex ions which will oxidise aldehydes but not ketones. During the reaction both reagents gradually change from blue to a brick red precipitate of copper (I) oxide. Cu2+ + e- → Cu+
• When alcohols are eliminated a water molecule always leaves as the leaving group. It is made from the OH group and the H is from the hydrogen next to the OH group.
• Dehydration: alcohols can be dehydrated with excess hot conc. H2SO4 or by passing their vapour over heated aluminium oxide. An alkene is formed. Phosphoric (V) acid is an alternative dehydrating agent.

• When ethanol is ionised it forms the ion C2H5Oh+ which is the molecular ion. These ions will break up as they are ionised so we have other ions of smaller molecular mass (fragmentation). The main peak furthest to the right of the mass spectrum is the molecular ion.
• High resolution mass spectrometry allows us to work out the molecular formulae of a parent ion. e.g.

A parent ion of mass 200 (3s.f) could have the molecular formulae C10H1604, C11H4O4, C11H20O3

Adding up their accurate atomic masses gives

C10H16O6= 200.1049 C11H4O4= 200.0110 C11H20O3= 200.1413

These can easily be distinguished by high resolution mass spectrometry

• A pair of atoms joined by a bond are always vibrating. Stronger bonds vibrate faster and heavier atoms make the bond vibrate slowly. Every bond has its own unique frequency which can absorb infra-red radiation of the same frequency. Therefore the radiation that emerges from the sample will be missing the frequencies that correspond to the bonds in the sample. An instrument then plots a graph of the intensity of the radiation emerging from the sample (transmittance) against the frequency of radiation. The frequency is expressed as a wave number.
• On the spectrum, usually below 1500cm-1, there is an area with many peaks caused by complex vibrations of the whole molecule. It can be used to identify the chemical and is called the fingerprint region.
• Infra-red spectra can also be used to show up the presence of impurities as you can see it by changes in graphs.