Chemistry Unit 2

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  • Created by: Leah Gray
  • Created on: 28-05-13 09:33

Shapes of Molecules: THE PROCESS

Molecular shape depends on the electron pairs around the central atom.

Electron pairs repel one another. The shape of the molecule is where the bonds and lone pairs are arranged in a way that minimum repusion and maximum distance between the bonds and lone pairs.

The size of repulsion leading to the size of angle is like this:

Lone -pair/Lone-pair angles > Lone-pair/Bonded-pair angles > Bonded-pair/Bonded-pair angles

2 electron pairs on a central atom: Linear: 180 degree bond angle

4 electron pairs on a central atom: Tetrahedral: 109.5 degree bond angle

3 electron pairs and one lone pair on a central atom: Trigonal pyramidal107  degree bond angle

5 electron pairs on a central atom: Trigonal bypyramidal: 90 degree and 120 degree bond angle

6 electron pairs on a central atom: Octahedral: 90 degree bond angle

 

 

 

 

 

 

 

 

 

 

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Allotropes of Carbon

Allotropes are different forms of the same element in the same state (for example: solid). Carbon forms 3 allotropes: diamond, graphite and fullerenes.

Diamond: It sublimes when its temperature is over 3800K. It is extremely hard. Vibrations can travel easily through the stiff lattice, so it is a good thermal conductor. It cannot conduct electrisity as its outer electrons are held in localised bonds. It won't dissolve in any solvant.

Graphite: Has a different macromolecular structure to diamond. Weak bonds between layers mean it can be easily broken. The delocalised electrons are free to move along sheets, so an electric current can flow. The layers are quite far apart relative to the covelant bonds, so graphite is less dense than diamond. In a hexagonal shape, due to stong covelant bonds has high melting point (it sublimes over 3900K). Like diamond, graphite is insoluble.

Fullerenes:Molecules of carbon that form hollow balls or tubes. Each carbon atom forms three covelant bonds, leaving free electrons that can conduct electrisity. Fullerenes are nanoparticles. Nanoparticles are generally up to 100 nanometres across. Many fullerenes are soluble in organic solvants and form brightly coloured solutions. Due to their hollow structure they can be used to 'cage' other molecules.

 

 

 

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Electronegativity

There is a gradual transition from ionic to covelant bonding. Very few atoms have close purely ionic bonding.

Diatomic gases have the closest bonding topurely covelant. Most compounds have some ionic and some covelant properties (e.g. hydrogen chloride has covelant bonding but dissolves in water to form hydrochloric acid, which is a ionic solution).

The ability to attact the bonding electrons in a covelant bond is called electronegativity.

Covelant bonds may be polarised by differences in electronegativity. In a polar bond the difference in electronegativity causes a dipole to form. A dipole is a difference in charge caused by a shift in electron density in a bond.

The length of bonds is also important. The shorter the bond lenght the greater the bond enthalpy. Covelant bonds aren't just the attraction between the nuclei and the shared pair of electrons, the positive nuclei also repel one another. The stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length. If there is more attraction the nuclei will pull closer.

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Oxidisation Rules

1. All atoms are treated as ions for this, even if they are covelantly bonded.

2. Uncombined elements have an oxidisation number of 0.

3. Diatomic atoms have an oxidisation number of 0.

4.Oxidisation number of a simple monatomic ion is the same as it's charge would be if it were an ion.

5.In compounds the oxidisation number is just the ion charge.

6. The sum of the oxidisation numbers for a neuteral compound is 0.

7. Combined oxygen is nearly always 2-, except for in peroxides.

8. Combined hyrogen is nearly always 1+, except in metal hydrides where it is -1 (and the diatomic version of hydrogen where the oxidisation number is 0)

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Group 2 (Part 1)

Ionisation energy decreases down the group.

Group 2 elements react with water to make hydroxides.

Burn in oxygen to form solid white metal oxides. They burn with their charectistic flame colours.

The metals react with chlorine to form white solid chlorides.

The oxides and hydorxides are bases (Magnesium oxide being an exception. It only reacts slowly and is insoluble.) Therefore they form alkaline solutions and neutralise acids.

Solubility of compounds depends on the anion. Hydroxides get more soluble down group 2 and sulfates get less soluble down group 2.

Thermal stability increases down a group . The sulfate and nitrate ions are large and so can be made unstable by the presence of postively charged ion, as the cation polarises the anion, distorting it. The greater the distortion the less stable the anion. Large cations cause less distortion to the anion than small cations. Therefore, further down the group the less distorted the anion, the more thermally stable the compound is.

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Group 2 (Part 2)

Group 2 compounds are less thermally stable than group 1 compunds. The greater the carge on the cation the grater the polarising power of the cation and therefore, group 1 element polarise the anion less and so the anion is more stable.

Lithium: Red

 Sodium: Orange/Yellow

 Potassium: Lilac

Rubidium: Red

Calcium: Brick-Red,

Barium: Green.

Flame colours: The energy absorbed from the flame causes electrons to move in to higher energy levels. The colours are seen as the electrons fall back down into lower energy levels, releasing the energy in the form of light. The difference in energy between the higher and lower levels determins the wavelength of light released which in turn determines the colour of the light.

 

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The Halogens (Part 1)

Chlorine: virtually colourless in water, virtually colourless in hexane.

Bromine: yellow/orange in water, orange/red in hexane.

Iodine: brown in water, pink/violet in hexane.

Halogens get less reactive down the group. Halogens react by gainig electrons, so they are oxidising agents (as they reduce they oxidise other substances). As the atoms become larger, the outer electrons get further from the neucleus. There is also more shielding. This means the larger the halogen atom, the harder it is for the halogen atom to attract the electron it needs to form an ion.

Melting and boiling points increases down the group.

Halogens undergo disproportionation with alkalis.

The oxidising power of halogens decreases down the group.

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The Halogens (Part 2)

Halogens oxidise metals. For example, fluorine and chlorine react with iron to form iron (iii) halides. Iron is taken to it's higest oxidisation state because these halogens are very  strong oxidising agents. Because bromine is a weaker oxidising agent you get a mixture of iron (ii) and iron (iii) bromide when bromine reacts with iron.

Halogens also oxidise non-metals. For example: chlorine reacts with sulfur to form sulfur (i) chloride.

And all halogens can oxidise some ions (except for iodine as it is only a weak oxidising agent.) Halogens will oxidise iron (ii) to iron (iii) in solution. The solution will turn from green to orange.

The reducing power of halides increase down the group.

Hydrogen halides are acidic, colourless gases. They are very soluble and dissolve to make strong acids.

Halogens relative oxidising strengths can be seen in their displacement reactions with hydrogen halides. For example if you mix bromine water with potassium iodide, the bromine displaces the iodide ions (as it oxidises them) making KBr and iodine.

Chlorine water is colourless, bromine water is orange and iodine solution is brown.

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The Halogens (Part 3)

Halides give coloured precipitates when mixed with silver nitrate.

Floride: no precipitate, Chloride: white precipitate, Bromide: cream precipitate, Iodide: yellow precipitates.

Silver halides react with sunlight. Silver halides decompose when light shines on them.

Iron (iii) iodide is unstable because  iodide ions are strong enough to reduce iron (iii) to iron (ii). However, iron (iii) chloride is stable becaise chloride ions are not strong enough to reduce iron (ii).

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Iodine-Sodium Thiosulfate Titrations

Iodine-sodium titrations are a way of finding the concerntration of an oxidising agent.

Stage 1: Use a sample of oxidising agent to oxidise as much iodine as possible. Measure out a certain volume of potassium iodate (V). Add this to an excess of acidic potassium iodide solution. The iodate (V) ions in the potassium iodate (V) solution oxidises some of the iodide ions to iodine.

Stage 2: Find out how many moles of iodine have been produced. Titrate the resulting mixture from the last step with (a known concertrated) sodium thiosufate. Use starch solution as indicator. When the blue colour dissapears all the iodine has gone.

Stage 3: Calculate the concertration of the oxidising agent. Use the ratio to work out how much potassium iodide the is.

Poblems with Titration:

Contaminated apparatus, reading the burrett reading from the bottom of the miniscus, repeat the titration to eliminate random errors, wash flask between titrations.

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Making Haloalkanes from Alchols

An alchol can be primary (when the carbon that the -OH group is bonded to is only boned to one carbon), secondary (when the carbon that the -OH group is bonded to is bonded to 2 other carbons) or tertairy (when the carbon that the -OH group is attatched to is bonded to 3 other carbons).

-OH can be swapped for a halogen to make a halogenoalkane. Tertairy alchols are more reactive than primary and secondary alchols. Primary and secondary alchols react too slowly to be made this way but can be made using phosphorus (iii) halides. The general formula is: 3ROH + PX3 -> 3RX + H3PO3. PCl3 is straightforward but PBr3 and PI3 are not straight forward and are usually made in situ by refluxing the alchol with 'red phosphorus' and either bromine or iodine.

Chloroalkanes can also be made using phosphorus (V) chloride. The general equation is: ROH + PCl5 -> 3RX + H3PO3.

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Reactions of Alcohols

Alcohols react with sodium to produce alkoxides. The sodium metal reacts gently with ethanol, breaking the O-H bond to produce ionic sodium ethoxide and hydogen. The longer the hydrocarbon chain of the alcohol, the less reactive it is with sodium.

The hydroxyl group can form hydrogen bonds. Hydrogen bonds are the strongest type of intermolecular bond so it gives alcohols high boiling points compared to non-polar compounds of a similar size. Alchols also are miscible with water as hydrogen bonds can be formed between -OH and water. If it is a short alcohol, hydrogen bonding lets it mix freely with water. But larger alchols, most of the molecule is non-polar and so there is less attraction to the water per molecule. Therefore as alcohols increase in size they become less miscible in water.

The test for the hydroxyl group is to add phosporus (V) chloride to the unknown liquid. If -OH is present, you will get steamy fumes of HCl gas, which dissolves in water to form chlroide ions. You can then test for chloride ions using silver nitrate. The steamy fumes of blue litmus paper can also turn moist blue litmus paper red.

Alcohols burn to produce carbon dioxide and water. Ethanol is easy to alight and burns with a pale blue colour. The C-C and C-H bonds are broken and completely oxidised to make carbon dioxide and water.

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Oxidisation of Alcohols

Primary alcohols are oxidised to aldehydes and then to carboxylic acids, by using: potassium dichromate and sulfuric acid

Secondary alcohols are oxidised to ketones only.

Tertiary alcohols will not oxidise.

Aldehydes and Ketones both have C=O functional groups. Aldehydes have a hydrogen attatched to the carbonyl carbon atom. Whereas ketones have two alkyl groups attactched to the carbonyl carbon atom.

To control how far an alchol will oxidise can be controlled by the reaction conditions. Potassium docromate is used. To produce an aldehyde it must be removed from the oxidising solution quickly. To do this it must be oxidised under distillation. However when making a carboxylic acid must be formed under reflux.

Secondary alcohols will oxidise to ketones. To produce a ketone, a secondary alcohol must be refluxed with acidified potassium dichromate.

To distinguish between an aldehyde and a ketone, benedicts solution can be used: turns brick-red when warmed with a aldehyde but stays blue when heated with a ketone. Tollens reagent is reduced to silver when it is heated with warmed with an aldehyde but doesn't when warmed with a ketone.

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Halogenoalkanes (Part 1)

Halogenoalkanes can be primary, secondary or tertiary.

When you mix a halogenoalkane with water it will react to form an alcohol.

When you mix a halogenoalkane with silver nitrate they form a silver halide precipitate.

Teriary haolgenoalkanes are the most reactive and primary halogenoalkanes are the least reactive.

Chloroethene can be use to make PVC and tetrafluoroethene can be used to make the non-stick coating on pans.

CHlorofluorocarbons are very unreactive and non-toxic. They have been used a solvants, refridgerants or as fire retardants.

Hydrochlorofluorocarbons can be used instead of carbonfluorocarbons as they are better for the enviroment. They are less stable and so decompose in the lower atmosphere.

Halogenoalkanes can react by nucleophilic substitution. Halogens are more polar that carbons and so a carbon halogen bond is polar. The delta postive carbon needs some more carbons and so is suseptable to a neuclephilic attack. OH- and ammonia are examples of neucleophiles. The halogen is subsituted for the neucleophile.

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Halogenalkanes (Part 2)

Halogenoalkanes react with aqueous alkalis to form alcohols, for example aqueous KOH. It is a hydrolysis reaction. To react bromoethane with water to form ethanol you have to use warm aqueous sodium under reflux.

Water is a weak neucleophile, but it is a neucleophile and will eventually react to form an alcohol. When water reacts with a halogenoalkane there is an itimediate where water is bonded to carbon but also has 2 bonds to hydrogen.

Haloalkanes react with ammonia to form amines. You must warm the haloalkane with excess ethanolic ammonia under reflux conditions. This is also a neucleophilic substitution reaction. The ammonia acts as a neucleophile as it has a pair of unbonded electrons.

If you react a haloalkane with a warm alkali dissolved in alcohol, you get an alkene. This must be heated under reflux. The OH- acts as a base and takes a proton, H+, from the carbon. This makes water. The carbon now has a spare electron, so it forms a pi bond with the other carbon. To form the pi bond the carbon releases the Br as Br-.

 

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Chlorine and Ozone

The ozone layer is in the layer of the atmosphere called the stratosphere. The ozone layer removes UV raditation, too much of which can cause skin cancer as it damages DNA in cells.

O2 + O. -> O3 is a reversible reaction. Therefore ozone is always being replaced.

Chlorine free radicals are made from CFC's which are broken down by UV rays. These Chlorine radicals are catalystes.

The reaction is:

Cl. + O3 -> O2 + ClO.

ClO. + O3 -> 2O2 + Cl.

Therefore the overall reaction is 2O3 -> 3O2

The montreal protocol of 1989 was an international treaty to ban CFC's.

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Nitrogen Oxides and Ozone

Nitric oxide, NO., free radicals destroy ozone as well as chlroine radicals.

NO. free radicals come from nitrogen oxides, which are made by aircraft engines and thunderstorms.

NO. free radicals act in the same way as Cl. radicals do. They are produced at the end of the reaction and therefore NO free radicals are catalystes.

NO. + O3 -> O2 + NO2.

No2. + O3 -> 2O2 + NO.

A free radical is......

Free radicals are particles with an unpaired electron, written as Cl., NO. or CH3. - you make free radicals when bonds split homolytically (or evenly).

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Infrared Spectra

Some molecules absorb energy from infrared radiation. This extra energy makes covelant bonds vibrate. Only molecules of different elements can absord infrared radiation because the polarity of bonds changes as the vibrate. Gases that absord infrared radiation are called greenhouse gases becauses they stop the infrared radiation from escaping from the atmosphere.

Gas molecules have certain fixed energy levels. These are called quantised levels. This means a bond's energy can only jump form one enrgy level to the other. It is not a contiuum but steps. This means that only frequencies which are related to a certain bond can be absorbed.

Different molecules absord different levels of infrared radiation.

The infrared radiation is absorbed by the covelant bond in the molecule, making them vibrate more. Bonds between different atoms absorb different frequencies of infrared radiation. As well bonds in different positions in molecule absorb different frequencies. For example the O-H group in alcohols and the O-H groups in carboxylic acids absorb different frequencies.

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Reaction Rates

Particles in liquids and gases are always colliding, however for a reaction to take place the collision must be productive. Meaning the atoms must have sufficent energy to react and must be the correct atoms to make the atoms take place. The minimun amount of enery needed to begin a reaction is called the activation energy and is the minimum amount of energy needed to break bonds and so begin the reaction.

Increasing temperature makes the reaction take place faster as the atoms will have more kinetic enery and so will move fast, therefore the frequency of collisions will increase and so the chance of a productive collision is increased. Also more atoms will have the activation energy to begin the reaction.

Increasing the concertration in a solution or the pressure of a gas means the atoms are on average closer together and so they will collide more often. The more collisions means the more chance that a reaction will take place.

Increasing the surface area speed up the reaction. If a reactant is in a big lump most of these particles will not collide with other reactants. However if you crush the reactants the reactants more of the particles come into contact with the other reactants. The smaller the pieces, the larger the surface area, the speedier the reaction.

Catalyste provide another route for the reaction to take place which requires a lower activation energy. With a lower activation energy, more particles will have enough energy to react.

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Solubility (Polar and Non-Polar)

For one substance to dissolve another, all these things have to happen: bonds in the substance must break, bonds in the solvant must break and new bonds must form between substance and solvant.

Ionic substances dissolve in polar solvants such as water. The ions are attracted to oppisitely charged ends of water moelcules. The ions are pulled away from the ionic lattice by the water molecules, which surround the ions. However some ionic substances do not diisolve because their ionic bonding is too strong, like aluminium oxide.

Alcohols dissolve in polar substances such as water although they are covelantly bonded. The polar -OH group is attracted to the polar water molecules, forming hydrogen bonds. The carbon chain isn't attracted to the water, so the longer the carbon chain, the less soluble alcohol will be.

Haloalkanes contain polar bonds but the dipoles are not strong enough to form hydrogen bonds with water. Because the bonds between water molecules are stronger than the bonds which would form if the haloalkane did dissolve.

Non-polar substances disssolve best in non-polar solvants because non-polar sunstances only have Van der Waals forces holding the molecules together. Molecules of polar solvants such as water are attracted to oneanother more strongly than they are to non-polar substances and so  non polar substances do not dissolve in polar solvants.

Usually LIKE dissolves LIKE - substances with similar bonds dissolve best.

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SN1 Mechanism

Here is the SN1 mechanism for tertiary haloalkanes:

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SN2 Mechanism

SN2 mechanism for primary haloalkanes:

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Water Neucleophile Mechanism

The water molecule is a weak neucleophile, but even water will substitute for a halogen. Here is the mechanism:

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Amine Mechanism

Warm haolgenoalkanes react with ammonia to form amines, here is the mechanism:

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Chlorinisation of Methane

Here is the free radical reaction of chlorine and methane:

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