Shapes of molecules
Electrons will repel each other to maximum separation and minimum repulsion. Greatest angles are between lone pairs of electrons, which cause more repulsion due to a larger electron cloud.
Linear - 2 bond pairs, 0 lone pairs, 180° (e.g.- BeCl2)
Trigonal Planar - 3 bond pairs, 0 lone pairs, 120° (e.g.- BF3)
Tetrahedral - 4 bond pairs, 0 lone pairs, 109.5° (e.g.- CH4)
Trigonal Pyramidial - 3 bond pairs, 1 lone pair, 107° (e.g.- NH3)
Bent - 2 bond pairs, 2 lone pairs, 104.5° (e.g.- H2O)
Trigonal Bipyramidial - 5 bond pairs, 0 lone pairs, 120° and 90° (e.g.- PF5)
Octahedral - 6 bond pairs, 0 lone pairs, 90° (e.g.- SF6)
Lone pairs of electrons count toward an area of electron density but as they repel to a much greater extent they reduce the bond angle by 2.5°
Allotropes are different forms of the same element in the same state. Carbon has three allotropes: diamond, graphite and fullerenes.
Diamond: Each carbon bonded to four other carbons, tetrahedral crystal lattice structure. High melting point, extremely hard, good thermal conductor, cannot conduct electricity. Insoluble.
Fullerene: Hollow balls or tubes. Each carbon bonded to three others = conductor. Nanoparticles. Soluble in organic solvents. Brightly coloured solutions. Delivering drug. Nanotechnology.
Graphite: Sheets of flat hexagons covalently bonded with three bonds each. Fourth electron delocalised. Weak bonds between layers,slide over each other, used as dry lubricants and pencils. Delocalised electrons allow electron flow. Less dense than diamond. High melting point. Insoluble in any solvent.
Electronegativity is the ability of an atom's nucleus to attract the bonding pair of electrons within a covalent bond. Covalent bonds can become polarised by differences in electronegativity, making the covalent bond polar.
A dipole is a difference in charge between two atoms caused by a shift in electron density in the bond. If the difference becomes large enough the bond is considered as ionic. The greater the difference in electronegativity, the more polar the bond.
Polar bonds can be cancelled out in a symmetrical molecule, so the whole molecule itself is not polar. Lone pairs of electrons on the central atom have an effect on the overall polarity and may cancel out the dipole created by the bonding pairs. If the polar bonds all point in similar directions, due to the shape of a molecule, the molecule will be polar.
The distance between the two nuclei is the distance where the attractive and repulsive forces balance each other. (Repulsion due to +vely charged nuclei) This distance is the bond length.
The stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length. A C=C has greater bond enthalpy and is shorter that a C-C bond, as electron density between carbon atoms is greater.
London forces cause all atoms and molecules to be attracted to each other.
1) Electrons in charge clouds are always moving very fast. At any particular moment, electrons in an atom are likely to be more to one side than the other, creating a temporary dipole.
2) This dipole can cause another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to each other.
3) Electrons are constantly moving, the dipoles are being created and destroyed all the time.
- Larger molecules have larger electron clouds, meaning stronger London forces.
- Molecules with greater surface areas also have stronger London forcesas they have a bigger exposed electron cloud.
- HYDROCARBONS: longer hydrocarbon chains have more molecular surface area, more electrons interact so stronger Londo forces. Branched-chain alkanes have smaller molecular surface areas and they cannot pack as closely together, so London forces are reduced. Therefore long-chain alkanes and straight-chain alkanes have higher boiling points.
Intermolecular Forces 2
Polar molecules have permanent dipole-dipole interactions. The charges on polar molecules cause weak electrostatic forces of attraction between molecules.
These permanent dipole-dipole interactions are stronger than instantaneous dipole-induced dipole interactions.
Electrostatically charged rods placed near a jet of a polar liquid, such as water, will attract the liquid molecules towards the rod.
Hydrogen bonding is the strongest intermolecular forces. Hydrogen bond only takes place when hydrogen is bonded to fluorine, nitrogen or oxygen. This is because these elements are very electronegative, so they draw the bonding electrons away from the hydrogen atom. The bond is so polarised and hydrogen has such a high charge density that hydrogen atoms form weak bonds with lone pairs of electrons on the F, N, O atoms of other molecules.
Molecules which have hydrogen bonding are usually organic, containing -OH or -NH groups. Water and ammonia both have hydrogen bonding.
Hydrogen bonds increase boiling and melting points. Ice is less dense than liquid water.
For one substance to dissolve into another, bonds in the substance have to break, bonds in the solvent have to break and new bonds have to form between the substance and solvent.
Ionic substances dissolve in polar substances such as water. Ions are pulled away from the ionic lattice by water molecules that they are attracted to. This is hydration. Some ionic substances do not dissolve as the electrostatic force of attraction between ions is too strong.
Alcohols dissolve in water due to polar -OH bond. Hydrogen bonds form between them. The carbon chain is not attracted to water so the more carbon atoms there are, the less soluble the alcohol will be.
Halogenoalkanes contain polar bonds but dipoles are not strong enough to form hydrogen bonds with water. The hydrogen bonds between water molecules is stronger than the bonds that would be formed with halogenoalkanes, so halogenoalkanes do not dissolve.
Non-polar substances dissolve best in non-polar solvents. Substances such as ethene have London forces between their molecules and form similar London forces with non-polar solvents such as hexane.
Redox is the transfer of electrons. Oxidation is the gain of oxygen and reduction is the loss of oxygen.
The oxidation state is a number assigned to an element in a compount according to some rules.
- The oxidation state of any elemnt e.g. Fe, H2, O2, Pa, S8 is zero (0).
- The oxidation state of oxygen in its compound is -2, except for peroxides like H2O2, in which the oxidation state for O is -1.
- The oxidation state for hydrogen is +1 in its compound, except for metal hydrides, such as NaH, LiH etc., in which the oxidation state for H is -1.
- The oxidation states for other element are then assigned to make the algebraic sum of the oxidation states equal to the net charge on the molecule or ion.
- The following elements usually have the same oxidation states in their compounds:
- +1 for alkali metals - Li, Na, K, Rb, Cs;
- +2 for alkaline earth metals - Be, Mg, Ca, Sr, Ba;
- -1 for halogen except when they form compounds with oxygen or one another.
Disproportionation is when the same species is both oxidised and reduced in the reaction.
Group 2 - Alkaline Metals
Group 2 + Oxygen –> Metal Oxide (solid)
Group 2 + Water –> Hydroxide + Hydrogen gas
Group 2 + Chlorine –> Metal Chloride (solid)
Oxide + Water –> Hydroxide
Oxide + Acid –> Chloride/Nitrate + H2O
Group 2 Hydroxides will do the same with 2H2O as a product
Sulphates: solubility decreases down the group. Test for sulphates: add barium chloride and a white precipitate will form. Compounds containing doubly charged negative ions tend to decrease in solubility down the group.
Hydroxides: solubility increases down the group. More strongly alkaline solutions formed as you go down the group due to increase in solubility. Singly charged negative ions tend to increase in solubility down the group.
Group 1 and 2 Compounds
Thermal decomposition is the break down of a substance into smaller molecules when heated.
Thermal stability increases down a group: Carbonate and nitrate ions are large and can be made unstable by the presence of a positively charged ion. The cation polarises the ion, distorting it. The greater the distortion, the less stable the anion. Large cations cause less distortion than small cations as they have a smaller charge density.
Group 1 compounds are more thermally stable than Group 2 compounds: Greater charge causes more distortion so anion is less stable. Group 2 cations have higher charge densities.
Carbonates: Group 1 – stable, do not decompose, except lithium carbonate
Group 2 carbonates + heat –> Group 2 oxide + CO2
Nitrates: Group 1 decompose to nitrite (e.g. KNO2) and oxygen, except lithium nitrate.
Group 2 decompose to oxygen, nitrogen dioxide (toxic, brown gas) and oxide.
Energy absorbed from the flame causes electrons to move to a higher energy level. This is unstable so electons move back to originel energy level, whilst emitting energy. The difference in energy between the higher and lower levels determines the wavelength of the light released.
Flame Tests: Mix a small amount of the compound you're testing with a few drops of hydrochloric acid. Heat a piece of nichrome or platinum wire in a hot Bunsen flame to clean it. Dip the wire into the compound mixture and hold over flame to observe colour.
Lithium/ Rubidium - Red
Sodium - Orange/Yellow
Potassium - Lilac
Strontium - Scarlet
Caesium - Blue
Calcium - Brick Red
Barium - Green
Group 7 - Halogens
Fluorine Pale Yellow, Gas
Chlorine Green, Gas, Colourless in water, Colourless in hexane, Low boiling and melting points
Bromine Orange-Brown, Liquid, Orange-Brown in water, Orange-Red in hexane
Iodine Grey, Solid, Brown in water, Pink/Violet in hexane, high boiling and melting points
Halogens act as oxidising agents. Larger atoms are less reactive due to larger radius and more shielding. They also become less oxidising and have disproportionation reactions with alkalis:
Cold: 2NaOH + Cl2 –> NaCl + NaClO + H2O (salt, chlorate I and water is formed)
Hot: 6NaOH + Cl2 –> 5NaCl + NaClO3 + H2O (salt, chlorate III and water is formed)
Halogens also oxidose metals. Fluorine and chlorine react with hot iron to form iron (III) halides as they are strong oxidising agents. Bromine is weaker and will form a mixture of iron (II) and iron (III) halides whilst iodine only forms iron (II) iodide.
All halogens, except iodine, will oxidise iron (II) to iron (III) ions in solution. The solution changes from green to orange.
Reactions of Halides
Reducing power of halides increases down the group as ionic radius and shielding increases. They give coloured precipitates with dilute nitric acid and silver nitrate solution: White chloride precipitate, cream bromide precipitate and yellow iodide precipitate.
Sulpuric Acid: Fluorine/Chlorine - Not strong enough reducing agent to reduce the sulphuric acid, so HX (X being the halide) is formed KCl + H2SO4 –>KHSO4 +HCL
Bromine - Bromine is strong enough to reduce sulfuric acid, it happens in a redox reaction. The products formed are bromine, sulphur dioxide and water 2HBr + H2SO4 -> Br2 +SO2 + 2H20
Iodine - Undergoes 2 redox reactions, first one produces Iodine, sulphur dioxide and water. Iodine reacts with sulphur dioxide to produce H2S, Iodine and water 6HI + SO2 -> H2S + 3I2 + 2H20
- Chloride Precipitate - Dissolves in dilute ammonia solution
- Bromide Precipitate - Dissolves in concentrated ammonia solution
- Iodide Precipitate - Insoluble in concentrated ammonia solution
Hydrogen halide gases are colourless and can form strong acids by dissolving in water. They react with ammonia gas to give white fumes, giving ammonium chloride.
Methyl Orange turns tellow to red when adding strong acid to alkali and end point is orange.
Phenolphthalein turns red to colourless when adding weak acid to alkali. End point is pale pink.
- You record readings to the nearest 0.05cm^3. This is to the limit of precision of the burette. The uncertainty of a reading from a burette is the maximum error of + or - 0.05cm^3.
- Percentage error can be reduced by measuring and using larger volumes.
- To improve the reliability, the experiment should be repeated until you have 2 concordant titres, within 0.20.
- Most systematic errors are from poor technique - always read the burette at eye-level. It is usually caused by equipment and is repeated throughout the experiment.
- Random errors vary with every experiment. This can be due to the subjectivity of end-point.
- Repeating the experiment improves reliability, but accuracy won't be improved. It eliminates random errors but not systematic errors.
- Using contaminated apparatus could make your results inaccurate, so burette needs to be cleaned with the substance that will go in the burette.
Reactions occur when particles of reactants collide with a minimum activation kinetic energy:
At higher concentrations and higher pressures, particles are in closer proximity to each other encouraging more frequent collisions.
At higher temperatures, a much higher proportion of colliding particles have sufficient activation energy to react and more particles are able to overcome the activation enthalpy barrier.
With smaller particles of reactant there is a larger surface area on which the reactions can take place, so the greater the chance of successful collisions
Heterogenous catalysts provide a surface where reacting particles may break and make bonds.
A homogenous catalyst is in the same state as reactants. It speeds up reactions by forming one or more intermediate compounds. The activation energy needed to form the intermediates is lower than making products directly from reactants. Energy profile will have two humps.
Chemical equilibria can only be established in closed systems. The reactions have 3 features:
- Reactions do not go to completion - there are always products and reactants.
- Reactions are reversible
- Dynamic - when a reaction appears to have finished, the chemicals continue to react, but the rate of the forward reaction = rate of the reverse reaction. This is called dynamic equilibrium.
Le Chatalier's Principle states that: when a change is imposed on a chemical equilibrium, the reaction responds in such a way to oppose the change.
As a result, the position of equilibrium changes. The conditions that change are temperature, pressure and concentration. The equlibrium will move to help counteract the effect.
Catalysts have no effect on the position of equilibrium. They cannot increase yield, but equlibrium is reached faster.
Electrophillic Addition with Halogens
- forms dihalogenoalkane, heterolytic fission, orange to colourless
- double bond repels electrons on Br2 polarising it - heterolytic fission occurs
- a carbocation is formed - an organic ion containing a positve carbon ion
- Br- moves to the positive carbon and bonds to form 1,2-dibromoethane
- C=C + Br-Br --> CH2BrCH2+ --> CH2BrCH2Br
Electrophillic Addition with Hydrogen Halides
- HBr - H turns delta positive Br delta negative, H joins one carbon, Br attracted to positive carbon = CH3CH2Br
- orange bromine water decolourises in the precence of C=C
- Br + H2O --> HOBr + HBr
- C=C + HOBr --> CH2BrCH2OH
Reduction with Hydrogen / Hydrogenation
- gain of hydrogen / loss of oxygen, forms alkane
- nickel catalyst at 150 degrees
- C=C + H2 --> CH3CH3
- long chain molecules with high Mr made from joining monomers
- monomer = unsaturated alkane with C=C, volatile liquids / gases
- polymers = saturated, solids due to increased van der waals
Addition of Acidified Manganate (VII)
- cold acidified manganate oxidises C=C to form a diol, purple to colourless
- ethene --> ethane 1,2 - diol
- C=C + [O] + H2O --> CH2OHCH2OH
Hydration of ethene
H2C=CH2 + H20 ---> CH3CH2OH Reagents = H3PO4 and 300C 60atm
Problems - uses non-renewable reactants
- uses large amounts of energy
Fermentaion of Sugars or Carbohydrates
C6H12O6 -----> 2CH3CH2OH + 2CO2 Reagents = Yeast and 37C
Problems - Produces only 14% alcohol
- Needs to be carried out in the absence of air as is anerobic
Halogenoalkanes from Tertiary Alcohol
Tertiary alcohols are more reactive, so to make a chloroalkane, shake the alcohol with hydrochloric acid in a separating funnel. Keep releasing the pressure as the product is volatile.
Allow mixture to settle into layers. Run off aqueous layer, leaving the impure hydrogenoalkane.
Neutralise the excess acid with sodium hydrogencarbonate solution, until no more gas is produced. Run the lower layer off.
Add distilled water, shake, and run off lower layer to remove inorganic impurities.
Remove any remaining water with anhydrous sodium sulphate.
Remove organic impurities by distilling the mixture. Collect the fraction that boils between 49 and 53 C which is chloroalkane.
Halogenoalkanes From Alcohols Using Hydrogen Halid
Primary and secondary alcohols react too slowly by just shaking.
Bromoalkanes: Rather than using aqueous hydrobromic acid, you treat alcohol with a mixture of sodium or potassium bromide and concentrated sulphuric acid, producing hydrogen bromide which reacts with the alcohol. The mixture is warmed to distil off the bromoalkane.
CH3CH2OH + HBr --> CH3CH2Br + H2O
Iodoalkanes: The alcohol is reacted with a mixture of sodium or potassium iodide and concentrated phosphuric (V) acid, H3PO4, and the iodoalkane is distilled off. The mixture produces hydrogen iodide which reacts with the alcohol.
CH3CH2OH + HI --> CH3CH2I + H2O
Phosphoric (V) acid is used instead of concentrated sulphuric acid because sulphurc acid oxidises iodide ions to iodine, producing hardly any hydrogen iodide.
Making Halogenoalkanes from Alcohols using Phospho
React an alcohol with liquid phosphorus (III) halide, PCl3: 3R-OH + PCl3 --> 3R-Cl + H3PO3
The Phosphorus first reacts with the bromine or iodine to give the phosphorus (III) halide:
2P + 3Br2 --> 2PBr3 ; 2P + 3I2 --> 2PI3
These then react with the alcohol to give the corresponding halogenoalkane which be distilled off
3CH3-CH2-OH + PBr3 --> 3CH3-CH2-Br + H3PO3
3CH3-CH2-OH + PI3 --> 3CH3-CH2-I + H3PO3
Chloroalkanes can be made using Phosphorus (V) Chloride. ROH + PCl5 -> RCl + HCl + POCl3 Test for hydroxyl group: add phosphurus (V) chloride to the unknown liquid. If -OH is present, steamy fumes of HCl gas can be seen, which dissolve in water to form chloride ions.
Alcohol and Sodium
Sodium metal reacts gently with ethanol, breaking the -OH bonds to produce ionic sodium ethoxide and hydrogen.
A small chunk of sodium can be placed in an alcohol.
The result is bubbles, as hydrogen gas is produced.
The sodium will slowly disappear too.
The result is the formation of an alkoxide and hydrogen gas.
Reactions are slower with longer chains because there are more van de waal's forces in a longer chain.
The reaction is also slower than it is with water.
The equation for a reaction is: 2CH3CH2CH3 + 2Na => 2CH3CH2CH2O-Na+ + H2
- Combustion: longer chains give more sooty flame
- CH3CH2CH2OH (l) + 4O2 (g) --> 3CO2 (g) + 4H2O
- gives sodium alkoxide, effervesence, white film forms, redox reaction, room temp
- 2CH3CH2CH2OH + 2Na --> 2CH3CH2CH2O-Na+ + H2
- Phosphorus (V) Chloride
- gives chloroalkane, white flame (HCl) given off, more dense with ammonia fumes
- HCl + NH3 --> NH4Cl - used to test for OH group
- nucleophillic subsitution - OH for Cl
- CH3CH2CH2OH + PCl5 --> CH3CH2CH2Cl + PCl3 + HCl
- gives chloroalkane, white flame (HCl) given off, more dense with ammonia fumes
- Bromide Ions
- KBr and 50% H2SO4, relfux reaction - the acid oxidises, OH replaced with Br
- CH3CH2CH2OH + KBr --> CH3CH2CH2Br + KOH
Oxidation of Alcohols
Aldehydes and ketones are carbonyl compounds, with C=O functional group, CnH2nO. Aldehydes have a hydrogen and one alkyl group. Ketones have two alkyl groups attached to the carbonyl carbon atom.
Primary alcohols can be both partially and fully oxidised. Partial oxidation takes place when the products are immediately distilled off so that complete oxidation cannot take place. The product is an aldehyde. Full oxidation takes place when the reactants are refluxed under heating in a vertical reflux condenser. Then the products are distilled off. The product is a carboxylic acid.
Refluxing a secondary alcohol with acidified dichromate (VI) will produce a ketone.
Fehling's solution and Benedict's solution are both deep Cu2+ complexes, wgich reduce to brick-red Cu2O when warmed with an aldehyde, but stay blue with a ketone.
Tollen's reagent is reduced to silver when warmed with an aldehyde, but not with a ketone. The silver will coat the inside of the apparatus to form a silver mirror.
Oxidation of Alcohols 2
- Aldehyde (from primary alcohol)
- add K2Cr2O7 in conc H2SO4, distillation reaction - H2 lost double bond formed
- orange --> green (Cr ions 6+ to 3+) Fehlings blue to brick red
- CH3CH2CH2OH + [O] --> CH3CH2CHO + H2O
- Carboxylic Acid (from primary alcohol)
- excess oxidising agent so further oxidation, heat under reflux
- Fehlings stays blue, test for by using Na2CO3 - bubbles of CO2 form
- CH3CH2CH2OH + 2[O] --> CH3CH2COOH + H2O
- Ketone (from secondary alcohol)
- orange to green, heat under reflux
- Fehlings stays blue
- CH3CH(OH)CH3 + [O] --> CH3COCH3 + H2O
- Phosphorus and Iodine
- gives iodoalkane, reflux reaction, substitution
- CH3CH2CH2OH + PI --> CH3CH2CH2I
Reactions of Halogenoalkanes
Potassium hydroxide: Halogenoalkanes can be reacted with KOH under different conditions in order to make different products.
- In aqueous conditions, a halogenoalkane can be refluxed under heating with KOH to make an alcohol. This is a nucleophilic substitution reaction. Heterolytic fission takes place.
Hydrolysis occurs through a nucleophilic substitution reaction.
- OH- ion has a lone pair of electrons which is attracted to the electron deficient carbon atom this is also known as nucleophilic attack.
- Donating the lone pair leads to a new covalent bond between the ‘O’ molecule and the carbon atom as a result this causes a break in the carbon – halogen bond by heterolytic fission forming a halide ion.
- Water molecule can also act as a weak nucleophile.
Reactions of Halogenoalkanes 2
In ethanoic (alcoholic) conditions, a halogenoalkane can be refluxed under heating with KOH to make an alkene. This is an elimination reaction.
The mixture of warm alkali dissolved in alcohol must be heated under reflux or the volatile substances will be lost. CH3-CH2-Br + KOH --> CH2=CH2 + H2O + KBr
Conditions: Concentrated solution of ammonia at room temperature, or heated in a sealed tube.
Mechanism: Nucleophilic Susbtitution.
R-Hal + NH3 -> R-NH2 + (H+) + (Hal-)
Nucleophilic Substitution Mechanisms
- SN2 - substitution nucleophillic 2nd order (2 species in slow step)
- CH3Br + OH- --> CH3OH + Br-
- nucleophile attacks from behind due to lack of space
- transition state - Br bond weakened
- SN1 - substitution nucleophillic 1st order (1 specie in slow step)
- C(CH3)3Br + OH- --> C(CH3)3OH + Br-
- large CH3 causes steric hindrance - OH cant fit
- CH3 stabalise C+ charge by positive induction effect, electrons pushed toward C so C-Br weak enough to drop
Halogenalkanes are generally unreactive, have low boiling points and are non-flammable, so they are used as fire-retardants and refrigerants.
Chlorofluorocarbons (CFCs) were found to deplete the ozone layer in the atmosphere, allowing harmful UV radiation to reach the earth surface, increasing the risk of skin cancer and eyecateracts.
Ozone is formed in the stratosphere by free radical reactions.
-Ordinary stable oxygen O2 (dioxygen) is split (dissociates) into two by high energy ultraviolet electromagnetic radiation from the Sun forming free radicals. A 'free' oxygen atom combines with an oxygen molecule (dioxygen) to form ozone (trioxygen).
O2 + uv ==> O. + O.
O. + O2 ==> O3
The ozone is a highly reactive and unstable molecule and decomposes into dioxygen when hit by other uv light photons. O3 + uv ==> O2 + O.
The chemically stable CFCs diffuse up into the stratosphere and decompose when hit by ultraviolet light (uv), producing free radicals, including free chlorine atoms, which themselves are highly reactive free radicals.
CCl2CF2 ==> CClF2. + Cl.
Chlorine free radicals readily react with ozone and change it back to much more stable ordinary oxygen. O3 + Cl. ==> O2 + ClO. and then: ClO + O ==> Cl + O2 - Cl free radical is regenerated and acts as a catalyst.
The two reactions above involving chlorine atoms are known as a catalytic cycle because the chlorine free radicals from CFC's act as a catalyst in the destruction of ozone.
NO. free radicals also destroy ozone. They are formed from nitrogen oxides, which are produced by car and aircraft engines and thunderstorms.
The greenhouse effect keeps us alive. The sun emits electromagnetic radiation. Most UV and infrared is absorbed by atmospheric gases, some of which is reflected back into space from clouds. Radiation is absorbed by the Earth, causing it to heat up. The earth then radiates energy back towards space as infrared radiation. Greenhouse gases absorb some of this radiation and re-emit it in all directions.
Main greenhouse gases are water vapour, carbon dioxide and methane. The molecules absorb IR radiation to make the bods in the molecule vibrate more. This extra vibrational energy is passed to other molecules by collisions, giving other molecules kinetic energy, raising the overall temperature.
Contribution of any particular gas to the greenhouse effect depends on how much radiation one molecule of the gas absorbs and concentration of that gas in the atmosphere in ppm.
Industrialisation, increase in human pouplation, burning of fossil fuels, deforestation are all anthropogenic changes, i.e. caused by human activities.
There is a definite correlation between CO2 and temp. rise but some argue that this does not mean rise in CO2 levels is the cause.
Types of Reactions
Addition: joining two or more molecules together to form a larger molecule.
Polymerisation: joining together the same monomer to form a polymer.
Elimination: small group of atoms break away from a larger molecule.
Substitution: one species is replaced by another.
Hydrolysis: splitting a molecule into two new molecules by adding H+ and OH- derived from water.
Oxidation: any reaction in which an atom loses electrons.
Reduction: any reaction in which an atom gains electrons.
Redox: any reaction where electrons are transferred between two species.
Free Radical Substitution: formation of free radicals, which attack and substitute.
Nucleophiles: electron pair donors. They are electron rich and are attracted to electron deficient species such as positive ions, polar bonds and carbocations (e.g. polar carbon-halogen bond).
Electrophiles: electron pair acceptors. THey are attracted to electron rich species such as negative ions, polar bonds and C=C bonds.
Free radicals: species with unpaired electrons. Cl free radicals are produced when UV light splits a Cl2 molecule. They are very reactive and react with any species.
Heterolytic Fission: bonds break unevenly. Both electrons from the shared electron pair move to one atom. This forms two different species: a positively charged cation/electrophile and a negatively charged anion/nucleophile. Electron transfer is shown by a double-headed arrow.
Homolytic Fission: bond breaks evenly. One electron moves to each atom. forming two electrically uncharged free radicals. Both atoms have an unpaired electron. Free radicals are very reactive. Electron transfer is shown by a single-headed arrow.
Mass spectrometry can be used to:
- identify unknown compounds,
- determine the abundance of each isotope in an element,
- gain further information about the structure and chemical properties of molecules.
Mass spectrometry also has a wide range of industrial and medical applications today, for example:
- monitoring the breath of patients during surgery whilst under anaesthetic,
- analysing molecules in space,
- detecting traces of toxic chemicals in contaminated marine life,
- detecting banned substances such as steroids in athletics.
A mass spectrometer determines the mass of a molecule or isotope by measuring the mass-to-charge ratio of ions. Although mass spectrometers differ considerably in their operation, the same basic principles occur inside the instrument.
VIADD: Vapourisation, Ionisation, Acceleration, Deflection, Detection
Mass Spectrometry 2
The molecular ion, M+, is the positive ion formed in mass spectrometry when a molecule loses an electron. It produces the peak with the highest m/z value in the mass spectrum
When a molecule is ionised by being impacted on by an electron, excess energy from the ionisation process can be transferred to the molecular ion, making it vibrate. This causes bonds to weaken and the molecular ion can split into pieces by fragmentation. Fragmentation results in apositive fragment ion and a neutral species.
Fragment ions are often broken up further into smaller fragments. The molecular ion and fragment ions are detected in the mass spectrometer.
Mass spectrometry can be used to determine the structure of an unknown compound and give its precise identity.
Although the molecular ion peak of two isomers will have the same m/z value, the fragmentation patterns will be different.
Each organic compound produces a unique mass spectrum, which can be used as a fingerprint for identification.
Some bonds in some molecules absorb IR radiation. This will make the bond vibrate.
The amount of energy needed to make them vibrate is dependant on the length of the bond, the strength of the bond and the mass of the atoms involved.
- For stronger bonds, higher frequency IR radiation will be required for the molecule to vibrate.
- If there are 2 atoms in the molecule, then only a stretch can occur.
- If there are more than this there can be a bend or a stretch. A shift of the partial charges is required for the bending/stretching to occur.
- The stretch can be assymetric or symmetric.
- Longer peaks are caused by strong absorption of IR.
- Smaller peaks are caused by weaker absorption of IR.
- Wider peaks are caused by absorption of IR over a wide range of frequencies.
- The unit of measurement for absorbance iswavenumber, a receprical of wavelength.
- Only molecules that change polarity as they vibrate (due to dipole movement in polar bonds) absorb IR.
- Only molecules that can absorb IR are greenhouse gases.