C2 - Elements, compounds and mixtures

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Dmitri Mendeleev made the first periodic table

  • In 1869, Dmitri Mendeleev took the 50 or so elements known at the time and arranged them into his table of elements - with various gaps
  • Mendeleev put the elements in order of atomic mass. To keep elements with similar properties in the same vertical groups, he had to swap one or two elements round and leave a few gaps. He was prepared to leave some very big gaps in the first two rows before the transition metals come in on the third row. 
  • The gaps were the really clever bit because they predicted the properties of so far undiscovered elements. When they were found and they fitted the pattern, it helped confirm Mendeleev's ideas. For example, Mendeleev made really good predictions about the chemical and physical properties of an element he called ekasilicon, which we know today as germanium
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How the periodic table looks today

  • Once protons and electrons were discovered, the atomic number of each element could be found, based on the number of protons in its nucleus. The modern periodic table shows the elements in order of ascending atomic number - and they fit the same patterns that Mendeleev worked out. 
  • The periodic table is laid out so elements with similar chemical properties form columns - these are called groups. (Elements with similar chemical properties react in similar ways.)
  • The group to which the element belongs corresponds to the number of electrons it has in its outer shell. E.g. Group 1 elements have 1 outer shell electron, Group 7 elements have 7, etc. Group 0 elements are the exception - they have full outer shells of 8 electrons (or 2 in the case of helium). 
  • The rows are called periods. Each new period represents another full shell of electrons.
  • The period to which the element belongs corresponds to the number of shells of electrons it has. 
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Electron shell rules

  • Electrons occupy shells (sometimes called energy shells).
  • The lowest energy levels are always filled first

Only a certain number of electrons are allowed in each shell:

  • 1st shell - 2 electrons 
  • 2nd shell - 8 electrons 
  • 3rd shell - 8 electrons
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Working out electronic structures

Electronic structure for Nitrogen:

  • The periodic table tells you that nitrogen has seven protons, so it must have seven electrons. 
  • The first shell can only take 2 electrons and the 2nd shell can take a maximum of 8 electrons. 
  • So the electronic structure for nitrogen must be 2.5.

You can also work out the electronic structure of an element by its period and group.

  • The number of shells which contain electrons is the same as the period of the element. 
  • The group number tells you how many electrons occupy the outer shell of the element. 
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Simple ions form when atoms lose/gain electrons

  • Ions are charged particles - they can be single atoms (e.g. CL-) or groups of atoms (e.g. N03-)
  • When atoms lose or gain electrons to form ions, all they're trying to do is get a full outer shell (also called a "stable electronic structure"). Atoms like full outer shells. 
  • When metals form ions, they lose electrons to form positive ions
  • When non-metals form ions, they gain electrons to form negative ions
  • The number of electrons lost or gained is the same as the charge on the ion. E.g. if 2 electrons are lost the charge is 2+. If 3 electrons are gained the charge is 3-. 
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Groups 1&2 and 6&7 are most likely to form ions

  • The elements that most readily form ions are those in Groups 1, 2, 6 and 7
  • Group 1 and 2 elements are metals. They lose electrons to form positive ions (cations). 
  • Group 6 and 7 elements are non-metals. They gain electrons to form negative ions (anions). 
  • Elements in the same group all have the same number of outer electrons. So they have to lose or gain the same number to get a full outer shell. And this means that they form ions with the same charges
  • Group 1 elements form 1+ ions
  • Group 2 elemens form 2+ ons
  • Group 6 elements form 2- ions.
  • Group 7 elements form 1- ions
  • As you go down each group you add electron shells, so the outer electrons get further from the nucleus. 
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Groups 1&2 and 6&7 are most likely to form ions

  • For Groups 1 and 2, this means that it gets easier to remove the outer electrons to form ions - so the elements get more reactive as you go down the groups. 
  • But for Groups 6 and 7, it means that it gets harder for the nucleus to attract extra electrons to form ions - so the elements get less reactive as you go down the groups. 
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Ions with opposite charges form ionic bonds

  • When a metal and a non-metal react together, the metal can lose electrons to form a positively charged ion and the non-metal can gain electrons to form a negatively charged ion. These oppositely charged ions are then strongly attracted to one another by electrostatic forces and form an ionic bond

To find the formula of an ionic compound, you balance the positive and negative charges. For example:

  • Sodium chloride - Na+ + Cl- →  NaCl - The sodium ion has a 1+ charge and the chloride ions has a 1- charge, so they balance
  • Magnesium chloride - Mg2+ + 2Cl-  MgCL2 - The magnesium ion has a 2+ charge and the chloride ion has a 1- charge, so you need two Cl- ions to balance the Mg2+ ion. 
  • Potassium oxide - 2K+ + O2- → K20 - The potassium ion has a 1+ charge and the oxygen ion has a 2- charge, so you need two K+ ions to balance the O2- ion. 
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Show ionic bonding using dot and cross diagrams

  • Dot and cross diagrams show the arrangement of electrons in an atom or ion. They can also show what happens to the electrons when atoms react with each other. Each electron is represented by a dot or a cross diagram. 
  • Dot and cross diagrams are really useful for showing how ionic compounds are formed, but they don't show the structure of the compound. 

Image result for bbc bitesize dot and cross diagrams

  • Soidum chloride (NaCl) - The sodium atom gives up its outer electron, becoming an Na+ ion. The chlorine atom picks up the electron, becoming a Cl- (chloride) ion.
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Ionic compounds have a regular lattice structure

  • Ionic compounds always have giant ionic lattice structures. The ions form a closely packed regular lattice. There are very strong electrostatic forces of attraction between oppositely charged ions, in all directions
  • A single crystal of sodium chloride (salt) is one giant ionic lattice. The Na+ and Cl- ions are held together in a regular lattice. This model shows the scale of the ions, but it only lets you see the outer layer of the compound
  • The ball and stick method shows how the ions are arranged, but the scale is wrong - in reality, there aren't gaps between the ions, and the ions are different sizes. 
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Lattice structure and ball and stick model

Image result for sodium chloride ball and stick 

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Ionic compounds all have similar properties

  • Ionic compounds have high melting and boiling points due to the strong attraction between the ions. It takes a large amount of energy to overcome this attraction. 
  • Solid ionic compounds don't conduct electricity because the ions are fixed in place can't move. But when an ionic compound melts. the ions are free to move and will carry an electric current
  • Many also dissolve easily in water. The ions seperate and are all free to move in the solution, so they'll carry an electric current
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Covalent bonds - sharing electrons

  • When non-metal atoms combine together they form covalent bonds by sharing pairs of electrons. 
  • This way both atoms feel that they have a full outer shell
  • Each covalent bond provides one extra shared electron for each atom. 
  • Covalent bonds are strong because there's a strong electrostatic attraction between the positive nuclei of the atoms and the negative electrons in each shared pair.  
  • Usually, each atom involved makes enough covalent bonds to fill up its outer shell. 
  • You can use dot and cross diagrams to show covalent bonds. 
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Examples of covalent bonds

  • Hydrogen Gas - Hydrogen atoms have just one electron. They need one more to complete the first shell, so they form a single covalent bond to achieve this. 

Image result for hydrogen gas covalent bond

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Simple molecules - low melting/boiling points

  • Substances formed with covalent bonds usually have simple molecular structures, like carbon dioxide and water. 
  • The atoms within the molecules are held together by very strong covalent bonds
  • The forces of attraction between these molecules are very weak. It's these feeble intermolecular forces that you have to overcome to melt or boil a simple covalent compound. 
  • So the melting and boiling points are very low, because the molecules are easily parted from each other. 
  • Most simple molecular substances are gases or liquids at room temperature. 
  • Simple molecular substances don't conduct electricity, because they don't have free electrons or ions. 
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Ball and stick model

Ball and stick models show how the atoms in covalent molecules are connected. You can make them with plastic molecular model kits, or as computer models. 

  • They're great for helping to visualise the structure of molecules, as they show you the shape of the molecule in 3D
  • They're more realistic than 2D drawings, but they're still a bit misleading. They make it look like there are massive gaps between the atoms - in reality this is where the electron clouds interact. 
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Giant covalent structures contain covalent bonds

  • Giant covalent structures are similar to giant ionic lattices except that there are no charged ions.
  • The atoms are bonded to each other by strong covalent bonds. 
  • This means that they have very high melting and boiling points. 
  • They don't conduct electricity - not even when molten (except for graphite, graphene and fullerenes).
  • The examples of giant covalent structures you need to know about are made from carbon atoms
  • Carbon can form loads of different types of molecules, because carbon atoms can form up to four covalent bonds, and bond easily to other carbon atoms to make chains and rings
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Diamond

  • Pure diamonds are lustrous (sparkly) and colourless. Ideal for jewellery. 
  • Each carbon atom forms four covalent bonds in a very rigid giant covalent structure, which makes diamond really hard. This makes diamond ideal as cutting tools. 
  • All those strong covalent bonds take a lot of energy to break and give diamond a very high melting point, which is another reason diamond is a good cutting tool
  • It doesn't conduct electricity because it has no free electrons or ions. 
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Graphite and graphene

  • Graphite is black and opaque, but still kind of shiny
  • Each carbon atom only forms three covalent bonds, creating sheets of carbon atoms which are free to slide over each other.
  • The layers are held together weakly so they are slippery and can be rubbed off onto paper to leave a black mark - that's how a pencil works. This also makes graphite ideal as a lubricating material
  • Graphite's got a high melting point - the covalent bond needs loads of energy to break.
  • Since only three out of each carbon's four outer electrons are used in bonds, there are lots of delocalised electrons that can move. This means graphite conducts electricity
  • A single sheet of graphite is called graphene. Graphene's covalent bonds make it strong and sheet of graphene is so thin that its transparent and light. Its delocalised electrons are completely free to move about, which makes it even better at conducting electricity than graphite. 
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Fullerenes are large carbon molecules

  • Fullerenes are another form of carbon. They aren't giant covalent structures, they're large molecules shaped like hollow balls or tubes. Different fullerenes contain different numbers of carbon atoms. 
  • The carbon atoms in fullerenes are arranged in rings, similar to those in graphite. And like graphite, they have delocalised electrons so they can conduct electricty
  • Their melting and boiling points aren't anything like as high as those of diamond and graphite, but they're pretty high for molecular substances because they're big molecules (and bigger molecules have more intermolecular forces). 
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Plastics are long-chain molecules called polymers

  • Polymers are formed when lots of small molecules called monomers joined together. This reaction is called polymerisation - and it usually needs high pressure and a catalyst. 
  • Plastics are polymers. They're usually carbon based and their monomers are often alkenes (a type of hydrocarbon containing a carbon-carbon double bond). 
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Forces between molecules determine the properties

  • Strong covalent bonds hold the atoms together in polymer chains. But it's the forces between the different chains that determine the properties of the plastic.

Weak forces:

  • If the plastic is made up of chains that are only held together by weak intermolecular forces, then the chains will be free to slide over each other. This means that the plastic can be stretched easily, and will have a low melting point. 

Strong forces:

  • Some plastics have stronger bonds between the polymer chains - these might be covalent bonds (sometimes called cross-links). These plastics have higher melting points, are rigid and can't be stretched, as the cross links hold the chains firmly together. 
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Properties of metals depend on structure & bonding

  • All the different types of materials in this topic have their own special properties. What they've all got in common is the fact that their properties are down to the structures and bonding in the material. 
  • The individual atoms in the material don't have these properties themselves - it's the type and strength of the bonds in a material that determines its properties.
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Examples of structure and bonding

For example, chlorine is pretty good at forming both ionic and covalent bonds.

  • It's found in many common ionic compounds like sodium chloride. The ionic bonds in sodium chloride are really strong, because there's strong electrostatic attraction between the ions, which acts in all directions within the lattice structure. The strength of these bonds is what gives sodium chloride its high melting and boiling point
  • Chlorine also forms simple molecular substances such as chloromethane. Although the covalent bonds which hold together the atoms in each molecule of chloromethane are very strong, the intermolecular forces which attract the molecules to each other are weak and easily overly overcome. So chloromethane has a low melthing and boiling point
  • Some polymers, such as polyvinyl chloride (PVC), also contain chloride. PVC is strong and rigid because the intermolecular forces between the polymer chains in PVC are relatively strong
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Metals have a crystal structure

  • All metals have the same basic properties, due to the special type of bonding that exists in metals.
  • In metals, the outer electrons of each atom can move freely. The atoms become positive ions in a 'sea' of delocalised electrons
  • Metallic bonding is the electrostatic attraction between these ions and electrons. The ions are surrounded by the electrons, so the attraction acts in all directions
  • This bonding is what gives rise to many of the properties or metals. 
  • Metals are generally on the left-hand side of the periodic table. This explains the bonding in solid metals - elements on the left of the table normall gets a full outer shell by losing electrons, so metal atoms find it easy to become positive ions. The electrons they give up form the electron 'sea'. 
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Metals have high melting/boiling points & density

  • Metals are very hard, dense and lustrous (i.e. shiny).
  • There's a strong attraction between the delocalised electrons and the closely packed positive ions - causing very strong metallic bonding
  • Metals generally have high melting and boiling points because of these strong metallic bonds. You need to use a lot of energy to break them apart. 
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Metals are strong, but bendy and malleable

  • Metals have a high tensile strength - they're strong and hard to break
  • But they can also be hammered into different shapes (they're malleable). 
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Metals are good conductors of heat and electricty

  • This is entirely due to the sea of delocalised electrons which move freely through the metal, carrying the electrical current.
  • They can also carry heat energy through the metal. 
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Metals react with oxygen to form metal oxides

  • Most metals react with oxygen to form metal oxides. Most metal oxides are solid at room temperature and form basic solutions when you dissolve them in water. 
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Metals can be mixed with elements to make alloys

  • Pure metals often aren't quite right for certain jobs. You can change their properties by mixing them with other elements (either metals or non-metals) to make alloys
  • Alloys have different properties from the main metal (or metals) they contain. For example, they may be stronger, more malleable or more corrosion resistant

Non-metal elements generally have different properties from metals. Non-metals usually have low melting and boiling points. When solid, they tend to be weak and brittle. They have lower densities than metals and don't conduct electrictity. (But there are exceptions, e.g. carbon breaks some of these rules).

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Structure & bonding affect melting/boiling points

  • The type of bonding in a substance affects its melting point and boiling point. It's all to do with how much energy you need to put in to get the atoms, ions or molecules apart.
  • The stronger the bonds are that keep the particles together in a solid or liquid, the more heat energy you need to put in to overcome those bonds and seperate the particles.
  • Simple covalent substances have strong bonds within each molecule, but only weak intermolecular forces between the molecules. It doesn't take much energy to overcome these forces, so simple covalent substances melt and boil at fairly low temperatures.
  • Most metals have high melting and boiling points because the metal ions are very strongly attracted to the delocalised electrons 'sea'. 
  • The positive and negative ions in ionic lattices are strongly attracted to each other. This strong electrostatic attraction means ionic substances have high melting and boiling points.  
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Structure & bonding affect melting/boiling points

  • Giant covalent lattices are held together by strong covalent bonds. These bonds take a lot of energy to break, so giant covalent substances have very high melting and boiling points. (Some giant covalent substances sublime instead - that means they go straight from solid to gas.
  • Covalent bonds - low melting/boiling points
  • Metallic - high melting/boiling points
  • Ionic - high melting/boiling points 
  • Giant covalent - high melting/boiling points 
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Pure substances contain only one thing

  • A substance is pure if it's completely made up of a single element or compound
  • If you've got more than one compound present, or different elements that aren't all part of a single compound, then you've got a mixture, not a pure substance.
  • So, for example, fresh air might be thought of as a nice and 'pure', but it's chemically impure, because it's a mixture of nitrogen, oxygen, argon, carbon dioxide, water vapour and various other gases. 
  • Lots of mixtures are really useful - alloys are a great example. But sometimes chemists need to obtain a pure sample of a substance.
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Test for purity using melting/boiling points

  • Every pure substance has a specific melting point and boiling point. For example, pure ice melts at 0°C, and pure water boils at 100°C.
  • So you can test the purity of a sample of a substance by comparing the actual melting or boiling point of the sample to the expected value
  • If a substance is impure, the melting point will be too low. So if some ice melts at -2°C, it's probably got an impurity in it (e.g. salt)
  • The boiling point of an impure substance will be too high. For example, seawater contains salt (and other impurities). It's boiling point tends to be around 100.6°C.

You can also sometimes tell if a sample of a solid or liquid is a mixture by heating it up.

  • In a mixture, the different components will melt or boil at different temperatures, so part of the mixture will melt or boil first, while the rest will stay in its original state. Mixtures will often melt over a range of temperatures.
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Simple distillation is used to seperate solutions

Simple distillation is used for seperating out a liquid from a solution. Here's how to use simple distillation to get pure water from seawater:

  • Pour your sample of seawater into the distillation flask
  • Set up the apparatus as shown in the diagram. Connect the bottom end of the condenser to a cold tap using rubber tubing. Run cold water through the condenser to keep it cool. 
  • Gradually heat the distillation flask. The part of the solution that has the lowest boiling point will evaporate (e.g. water)
  • The water vapour passes into the condenser where it cools and condenses (turns back into a liquid). It then flows into the beaker where it is collected
  • Eventually you'll end up with just the salt left in the flask. 

The problem with simple distillation is that you can only use it to seperate things with very different boiling points. 

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Fractional distillation - mixture of liquids

If you've got a mixture of liquids you can seperate it using fractional distillation. A lab demonstartion that can be used to model fractional distillation of crude oil at a refinery:

  • Put your mixture in a flask. Attach a fractioning column and condenser above the flask as shown. 
  • Gradually heat the flask. The different liquids will all have different boiling points - so they will evaporate at different temperatures
  • The liquid with the lowest boiling point evaporates first. When the temperature on the thermometer matches the boiling point of this liquid, it will reach the top of the column.
  • Liquids with higher boiling points might also start to evaporate. But the column is cooler towards the top, so they will only get part of the way up before condensing and running back down towards the flask. 
  • When the first liquid has been collected, raise the temperature until the next one reaches the top. 
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Filtration -separate insoluble solid from a liquid

  • If the product of a reaction is an insoluble solid, you can use filtration to seperate it out from the liquid reaction mixture
  • It can be used in purification as well. For example, solid impurities can be seperated out from a reaction mixture using filtration.
  • All you do is put the filter paper into a funnel and pour the mixture into it. The liquid part of the mixture runs through the paper, leving behind a solid residue
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Crystallisation - a soluble solid from a solution

Here's how you crystallise a product:

  • Pour the solution into an evaporating dish and gently heat the solution. Some of the solvent (which will usually be water) will evaporate and the solution will get more concentrated
  • Once some of the solvent has evaporated, or when you see crystals start to form (the point of crystallisation), remove the dish from the heat and leave the solution to cool.
  • The salt should start to form crystals as it becomes insoluble in the cold, highly concentrated solution.
  • Filter the crystals out of the solution, and leave them in a warm place to dry. You could also use a drying oven or a disiccator (a desiccator contains chemicals that remove water from the surroundings).
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Choose the right purification method

  • You might have to pick one of the techniques (i.e. Distillation, fractional distillation, crystallisation and filtration) to seperate a mixture. The best technique to use will depend on the properties of the substances in the mixture.
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Chromatography uses two phases

Chromatography is a method used to seperate and identify the substances in a mixture. There are lots of different types of chromatography -but they all have two phases.

  • A mobile phase - where the molecules can move. This is always a liquid or a gas.
  • A stationary phase - where the molecules can't move. This can be a solid or a really thick liquid.
  • The components in the mixture seperate out as the mobile phase moves over the stationary phase - they all end up in different places in the stationary phase
  • This happens because each of the chemicals in a mixture will spend different amounts of time dissolved in the mobile phase and stuck to the stationary phase.
  • How fast a chemical moves through the stationary phase depends on how it distributes itself between the two phases.
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In TLC the mobile phase is a solvent

  • In thin-layer chromatography (TLC), the stationary phase is a thin layer of a solid (e.g. alicia gel or aluminium oxide powder) on a glass or plastic plate. The mobile phase is a solvent (e.g.ethanol). Here's the method for setting it up :
  • Draw a line near the bottom of the plate. (Use a pencil to do this - pencil marks are insoluble and won't react with the solvent.) Put a spot of the mixture to be seperated on the line.
  • Put some of the solvent into a beaker. Dip the bottom of the plate (not the spot) into the solvent.
  • Put a watch glass over the beaker to stop any solvent from evaporating away.
  • The solvent will start to move up the plate. When the chemicals in the mixture dissolve in the solvent, they will move up the plate too.
  • You will see the different chemicals in the sample seperate out, forming spots at different places on the place.
  • Remove the plate from the beaker before the solvent reaches the top. Mark the distance the solvent has moved (the solvent front) in pencil.
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In TLC the mobile phase is a solvent

The amount of time the molecules spend in each phase depends on two things:

  • How soluble they are in the solvent.
  • How attracted they are to the stationary phase

Molecules with a higher solubility in the solvent (and which are less attracted to the stationary phase) will spend more time in the mobile phase than the stationary phase - so they'll be carried further up the plate.

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Paper chromatography is similar to TLC

  • Paper chromatography is very similar to TLC, but the stationary phase is a sheet of chromatography paper (often filter paper).
  • The mobile phase is a solvent such as ethanol.
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You can calculate the Rf value for each chemical

  • The result of chromatography analysis is called a chromatogram.
  • Sometimes, the spots on a chromatogram might be colourless. If they are, you'll need to use a locating agent to show where they are (e.g. you might have to spray the chromatogram with a special reagent).
  • You need know how to work out the Rf values for spots (solutes) on a chromatogram.
  • An Rf values is the ratio between the distance travelled by the dissolved substance (the solute) and the distance travelled by the solvent.
  • Rf = distance travelled by solute ÷ distance travelled by solvent
  • To find the distance travelled by the solute, measure from the baseline to the centre of the spot.
  • Chromatography is often carried out to see if a certain substance is present in a mixture. You run a pure sample of the substance alongside the unknown mixture. If the Rf values match, the substances  may be the same.
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You can calculate the Rf value for each chemical

  • Chemists sometimes run samples of pure substances called standard reference materials (SRMs) next to a mixture to check the identities of its components. SRMs have controlled concentration and purities.
  • You can also use chromatography to do a purity test. A pure substance won't be seperated by chromatography - it'll always move as one blob, while a mixture can produce multiple blobs.
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Gas chromatography is a bit more high-tec

Gas chromatography (GC) is used to analyse unknown substances too. If they're not already gases, then they have to be vaporised. The mobile phase is an unreactive gas, such as nitrogen and the stationary phase is a viscous (thick) liquid, such as oil.

The process is quite different from paper chromatography and TLC:

  • The unknown mixture is injected into a long tube coated on the inside with the stationary phase.
  • The mixture moves along the tube with the mobile phase until it comes out the other end. As in the other chromatography methods, the substances are distributed between the phases (so each substance spends different amounts of time dissolved in the mobile phase and stuck to the stationary phase).
  • The time it takes a chemical to travel through the tube is called the retention time.
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Gas chromatography is a bit more high-tec

  • The retention time is different for each chemical - that's what's used to identify it.

The chromatogram from GC is a graph. Each peak on the graph represents a different chemical.

  • The distance along the x axis is the retention time - which can be looked up to find out what the chemical is.
  • The relative areas under the peaks show you the relative amounts of each chemical in the sample.
  • There's one peak for each chemical, which means a sample of a pure substance will produce a single peak.
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Relative atomic mass, A

  • In the periodic table, the elements all have two numbers next to them. The bigger one is the relative atomic mass (A,) of the element.
  • The relative atomic mass of an element is the average mass of one atom of the element, comapred to 1/12 of the mass of one atom of carbon-12.
  • If an element only has one isotope, its A, will be the same as its mass number.
  • If an element has more than one isotope, its A, will be the average of the mass numbers of all the different isotopes, taking into account how much there is of each one,
  • Example: Chlorine has two stable isotopes, chlorine-35 and chlorine-37. There's quite a lot of chlorine-35 around and not so much chlorine-37 - so chlorine's, A, works out as 35.5.
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Relative formula mass, Mr,

  • The relative formula mass of a compound is all the relative atomic mass in its formula added together.

Example:

  • MgCl2 = 24.3 + (2 x 35.5) = 95.3 = Relative formula mass
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Molecular and displayed formula - No. of atoms

  • You can work out how many atoms of each type there are in a substance when you're given its formula.
  • CH4 - This is called a molecular. It shows the number and type of atoms in a molecule.
  • O=O - This is called a displayed formula. It shows the atoms and the covalent bonds in a molecule as a picture. 
  • If you have the displayed formula of a molecule, you can use it to write the molecular formula (and vice versa). 
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Empirical formula is the simplst ratio of atoms

  • An empirical formula of a compound tells you the smallest whole number ratio of atoms in the compound. 

For example:

  • C2H= CH
  • Empirical formula = CH3
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