There are also many simple indicators which are substances which have more than one colour form, depending on the pH. These will tell you if the solution is alkaline or acidic, but not the strength of the solution.
An acid can be defined as a proton donor, as acids produce H+ ions. The H+ ion is a proton. In water this proton is hydrated and is represented as H+ (aq).
A base can be defined as proton acceptor. Alkalis, in aqueous solution, produce OH- (aq) ions. So the process of neutralisation is the reaction between H+ (aq) and the OH- (aq) ions to give water.
H+ (aq) + OH- (aq) → H2O (l)
A strong acid or alkali is one that is fully ionised in aqueous solution. A weak acid or alkali is one that is only partially ionised in aqueous solution and is mostly in the form of molecules rather than ions.
Strong acids have lower pH than weaker acids of the same concentration. Weaker acids have slower reactions than stronger acids of the same concentration.
Concentrated and dilute tell you about the amount of acid which has gone into solution. Strong and weak tell you about the proportion of the acid which has reacted with the water to form ions.
Summary of neutralisation reactions
-
Metal + acid → salt + hydrogen
-
Metal oxide + acid → salt + water
-
Metal hydroxide + acid → salt + water
-
Metal carbonate + acid → salt + carbon dioxide + water
Salts
All ammonium, potassium and sodium salts are soluble.
Most carbonates and hydroxides are insoluble except, of course, ammonium, potassium and sodium carbonates and hydroxides.
All nitrates and ethanoates are soluble.
Most chlorides are soluble (except lead and silver chlorides).
Most sulphates are soluble (except lead, barium and calcium).
Salts have different methods of production depending on whether the reagents needed are soluble or not.
In the first method a metal or an insoluble base or insoluble carbonate with an appropriate dilute acid. This method cannot be used with unreactive metals such as copper or with very reactive metals such as sodium.
Excess insoluble solid is added to the acid, this ensures that all the acid reacts and none remains to contaminate the salt. The excess solid reactant is then removed by filtration. The solution of the required salt is then evaporated to the point of crystallisation.
In the second method a soluble base or carbonate is reacted with the appropriate acid.
Since one cannot remove any excess by filtration, one needs to use a burette and an indicator. The alkali or carbonate (usually in aqueous solution) is placed in a wide necked flask with the added indicator. Acid is added from the burette until the end point is reached, identified by a sudden change in colour.
The indicator is removed by boiling the solution with charcoal. The suspension is then filtered to remove the charcoal and the salt is evaporated from the solution.
Alternatively, the titration is carried out initially above using indicator, but not charcoal is added, as the whole procedure us repeated again, using the exact quantities as before, but with no indicator present. The final solution can be heated to obtain the salt as before.
yes
Heat the metal in dry chlorine
Do you want to make an anhydrous chloride?
no
no
Use a precipitation method. Mix two solutions, one containing the correct positive ion and the other containing the correct negative ion.
Is the salt soluble?
yes
Is the salt soluble?
no
React an acid with an excess of a solid metal (if suitably reactive), metal oxide, hydroxide or carbonate.
Is it a sodium, potassium or ammonium salt?
yes
Use a titration method. React with an acid with a solution of sodium, ammonium hydroxide or carbonate.
The Periodic Table
The Periodic table is an arrangement of all the hundred-plus elements in order of increasing atomic number, with elements with similar properties in the same vertical column.
The vertical columns in the periodic table are called groups. The elements in a group have similar properties.
The elements between the two main blocks are the transition metals. The bold stepped line on the table divides metals on the left hand side from non-metals on the right.
For any element in the main block of the periodic table, it is easy to work out the electron arrangement in the atoms.
- The number of energy levels or shells is the same as the period in which the element is found.
- The number of electrons in the outer energy level is the same as the group number (except for elements in group 0 which have 8 electrons, apart from helium which has 2 electrons).
Characteristics of metals
Almost all metals have certain properties in common:
- They conduct heat well (e.g. calcium)
- They conduct heat well (e.g. magnesium, copper)
- They are shiny (clean surface); this is known as metallic lustre (e.g. polished calcium or aluminium)
- They are malleable, i.e. they can be rolled into thin sheets (e.g. magnesium ribbon, copper foil)
- They are ductile, i.e. they can be pulled into thin wires (e.g. copper, silver, gold)
- They are sonorous, i.e. they make a sounds when hit (e.g. copper, iron)
Metals react with oxygen to form metal oxides.
Metals react with water and steam to give metal oxides or metal hydroxides.
Metals react with acids to form a salt and hydrogen gas. Hydrochloric acid gives chloride salts and sulphuric acid gives sulphate salts.
Many metal compounds are changed by heat. We call it thermal decomposition.
Metal carbonates high in the Reactivity Series do not change, their carbonates are not decomposed. Metals lower in the Series produce a metal oxide and carbon dioxide gas.
Example equation:
Calcium carbonate → calcium oxide + carbon dioxide
CaCO3 (s) → CaO (s) + CO2 (g)
Most carbonates of metals lower in the Reactivity Series than calcium behave the same way.
Metal hydroxides at the top of the Reactivity Series have hydroxides that are unchanged by heat, such as potassium and sodium. Metal hydroxides lower in the Series are decomposed to an Oxide and water.
Example equation:
Copper hydroxide → copper oxide + water
Cu (OH)2 (s) → CuO (s) + H2O (l)
The Noble Gases
The noble gases are all colourless gases. There are all very unreactive and monatomic- their molecules consist of single atoms. Their densities and boiling points increase as you go down the Group, although they are very low in the first place. The boiling points also increase as you go down the Group. This is because the attractions between one molecule and its neighbours get stronger as the atoms get bigger. More energy is needed to break the stronger attractions. The noble gases don’t form stable ions, as so don’t produce ionic compounds.
Group 1- the alkali metals
This entire group are soft metals, with very low melting points and densities. They react rapidly with air to form coating of the metal oxide. They react with water to produce the metal hydroxide and hydrogen. They increase in reactivity as you go down the group. They form compounds in which the metal has a 1+ ion. They have oxides that react with water to produce soluble alkaline hydroxides. They have white compounds that dissolve to produce colourless solutions.
To stop them reacting with oxygen or water vapour in the air, lithium, sodium and potassium are stored under oil. Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility of oxygen getting at them.
The difference in the reactions depends in part on how easily the outer electron of the metal is lost. This depends on how strongly it is attracted to the nucleus in the original atom. In all of these reactions, the metal atoms are losing electrons and forming metal ions in solution. The nucleus of an atom is positive because it contains protons; therefore it attracts the negative electrons. The outer electron is shielded from the full attraction of the nucleus by all the inner electrons. In every single atom in the element of Group 1, the outer electron will feel an overall attractive force 1+ from the nucleus, but the effect of the force falls very quickly as the distance increases.
All Group 1 metal ions are colourless. This means that their compounds will be colourless or white unless they are combined with a coloured negative ion. Potassium dichromate (VI) is orange, because the dichromate (VI) ion is orange.
The hydroxides all form strongly alkaline solutions. This is the origin of the name ‘alkali metals. They are strong bases and in solution, their pH values are very high.
The chlorides and sulphates are colourless salts, which dissolve easily in water giving neutral solutions. Sodium chloride is the most common salt in seawater.
All the nitrates decompose on heating to give the nitrate salt and oxygen gas, except lithium.
Lithium nitrate and lithium carbonate both behave in the same way as group 2 nitrates and carbonates.
The carbonates of group 1 metals are soluble and not decomposed by heating in a Bunsen flame. Sodium carbonate is sold as washing soda and is effective as a degreasing agent.
Group 2
Group 2 are all metals with low melting points and densities. They react with air, water or steam to form metal oxides or hydroxides. They show the same reactivity trend as Group 1, increasingly reactive as you go down the Group. They form compounds in which the metal has a 2+ ion and their compounds are white or colourless which give colourless solutions.
Magnesium and calcium oxides are strong bases and can be prepared by burning the elements in air or heating the carbonates.
Both magnesium and calcium hydroxides from alkaline solutions in water and are both strong bases. They can neutralise acids to give salts.
Magnesium and calcium chlorides are colourless salts, which crystallise from solution with water of crystallisation. Anhydrous calcium chloride is a useful drying agent for gases, except for ammonia with which it reacts. Calcium oxide is known as quicklime, and can be converted into calcium hydroxide by adding water; this reaction is very exothermic. Calcium hydroxide is known as slaked lime and is used in agriculture to neutralise acidic soils and so improve crop yields.
Group 7 – the Halogens
The Halogens are a family of salt- producing non-metals. They are diatomic. As the molecules get larger towards the bottom of the group, the melting and boiling points increase. They are oxidising elements that lose their oxidising capability as you go down- become less reactive as you go down. They will displace elements lower down the group from their salts. They form ionic salts with metals and molecular compounds with non-metals.
Fluorine is dangerously reactive that so we would never come across it in a school lab. All the elements have poisonous fumes and have to be handled in a fume cupboard. Liquid bromine is very corrosive.
The halogens react with hydrogen to form hydrogen halides. The hydrogen halides are all steam, acidic, poisonous gases. They are also covalently bonded, very soluble in water, reacting with it to produce solutions of acids. For example, hydrochloric acid is a solution of hydrogen chloride in water.
Hydrochloric acid is essentially hydrogen chloride dissolved in water.
Preparation of chlorine from hydrogen chloride
-
Add a few grams of solid, block powder potassium permanganate together with about 10 cm3 of water.
- Slowly add enough concentrated hydrochloric acid to start the reaction, look for the green colour of the gas.
- Test the gas using damp blue litmus paper, held near the top of the jar. The chlorine gas turns it red and then bleaches it.
Explaining the trend of reactivity
When a halogen oxidises something, it does so by removing electrons from it. (Oxidation is the loss of electrons). Each halogen has the ability to oxidise the ions of those underneath it in the group, but not those above it. Chlorine can remove electrons from bromide or iodide ions, and bromine can remove electrons from iodide ions.
Chlorine is a strong oxidising agent because its atoms readily attract an extra electron to make chloride ions. Bromine is less successful at attracting electrons, and iodine less successful still. In chlorine, there are 17 positively charged proton offset by 10 negatively charged electrons in the inner shells. That means that the new electron fells an overall pull of 7+, but in the case of bromine, this pull is further away. The incoming electron is further and further away from the nucleus as you go down the Group, and so it is les strongly attracted. This means an ion is less easily formed. It also means that it is easier to remove the extra electron from bromide ion than chloride ions. Iodide ions hold the ions weakly enough that they are easily oxidised.
Uses of the halogens
Fluorides such as sodium fluoride are added to toothpaste or dinking water to help prevent tooth decay.
Chlorine is added to the water supply in very small quantities to kill bacteria. It is also powerful bleach. Potassium iodide solution is used as an antiseptic.
Silver halides
The silver halides are important in photography.
The test for chlorine gas is damp blue litmus paper does red and then is bleached.
The Transition metals
They are all typical metallic elements. They are good conductors of heat and electricity, workable, strong and mostly high densities. Apart from mercury, they all have reasonably high to very high melting points. They are much less reactive than Groups 1 and 2, and don’t react as rapidly with air or water. They usually have a shiny appearance. They often form more than one positive ion. For example, iron forms iron (II) ions, Fe2+, and iron (III) ions, Fe3+. This property is called variable valency. Their compounds are often coloured. Transition meals and their compounds are often good catalysts. For example, iron in the manufacture of ammonia, the Haber process, and nickel in the manufacture of margarine.
Transition metals have a wide range of uses, either as pure metals or in alloys (mixtures of metals). Some have strong magnetic properties.
Steel is an alloy of iron with a small percentage of carbon. It is used from making car bodies, ships and bridges. Steel rods are also used to reinforce concrete. Stainless steel contains other transition metals such as nickel and chromium. It is more resistant to corrosion than ordinary steel. Brass is an alloy of copper and tin. Gold and silver are used in jeweller, but they are sometimes hardened by alloying with other metals.
Copper does not react with dilute hydrochloric or sulphuric acids but it does react with the oxidising acid, nitric acid. When reacting with concentrated nitric acid, the toxic brown fumes of nitrogen dioxide are released. This reaction is a redox reaction.
When copper is heated in a flame, it turns black with a coating of black copper (II) oxide. This oxide can also be produced by heating the green copper (II) carbonate. This is an example of thermal decomposition.
Copper (II) oxide is a base; it reacts with acids to form salts.
Blue copper (II) sulphate solution reacts with alkalis to form copper (II) hydroxide. When it reacts with ammonia solution the deep blue complex ion, [Cu (H20) 2 (NH3)4] 2+
Reactivity series
Elements can be arranged in order of their reactivity. Elements with atoms that have a full outer shell of electrons are very unreactive as these arrangements are stable. Elements with atoms contain one or two; or six or seven electrons in their out energy level are very reactive, as they want to lose or gain these extra electrons to finish with a stable electronic arrangement. In the middle, with elements contain three, four or five electrons in their outer shell, are less reactive, usually.
The order of reactivity amongst the most common elements is as follows:
Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, (Carbon), Zinc, Iron, Tin, Lead, (Hydrogen), Copper, Silver, Gold and Platinum. This order is with decreasing reactivity.
Any metal higher in the series will displace one lower down from its compound. For example, if you heated magnesium and copper (II) oxide, you would get magnesium oxide and copper as the magnesium displaced the copper out of the oxide compound.
This is an example of a displacement reaction. However if you heated copper with magnesium oxide, nothing would happen as copper is less reactive than magnesium.
Metals above hydrogen in the Reactivity Series react with water (or steam) to produce hydrogen. If the metal reacts with cold water, the metal hydroxide is made and hydrogen formed. If the metal reacts with steam, the metal oxide is made and hydrogen is formed.
AS you go down the Series, the reactions become less rigorous. Metals below hydrogen do not react with water or steam. This is why copper is used in rum stills and for bother hot and cold water pipes.
Metal + cold water → metal hydroxide + hydrogen
Metal + steam → metal oxide + hydrogen
Magnesium doesn’t react with cold water because it gets coated with insoluble magnesium hydroxide which prevents water from coming into contact with the magnesium. However, with steam, in a test tube, enough heat spreads back along the test tube to vaporise the water. The magnesium burns with a bright white flame, producing hydrogen.
Analysis
Gases
Gases can be collected in downwards into a test tube or gas jar, upwards into a test tube or gars jar or over water. Gases can also be collected in a gas syringe if you want to measure the volume of gas concerned.
But hydrogen forms explosives mixtures with and air, and chlorine, sulphur dioxide, hydrogen chloride and ammonia are all poisonous. Sulphur dioxide can trigger asthma attacks.
Hydrogen
Since hydrogen is less dense than air, and almost insoluble in water, you can collect it over water or upwards into a test tube. Hydrogen pops when a lighted splint is exposed to the gas. It also combines explosively with oxygen in the air to make water.
2H2 (g) + O2 (g) → 2H2O (l)
Oxygen
Oxygen has almost the same density as air and is almost insoluble in water, so it is normally collected over water. Oxygen relights a glowing splint.
Carbon dioxide
Carbon dioxide is denser than air so must be collected downwards, and as it is only slightly soluble in water; it can be collected over water as well. Carbon dioxide forms a white precipitate in lime water.
Ca(OH)2 (g) + CO2 (g) → CaCO3 (s) + H2O (l)
Chlorine
Chlorine is denser than air and so is usually collected downwards. As it is a green gas, it is easy to see when the collecting vessel is full. It cannot be collect over water but it more insoluble over a concentrated salt solution. Chlorine bleaches damp blue litmus.
Sulphur dioxide
Sulphur dioxide is denser than air and can be collected downwards. As gases become less soluble as the temperature of the liquid creases, you can collect sulphur dioxide over hot water (although not cold water). Sulphur dioxide gas is acidic and it turns blue litmus red and orange potassium dichromate (VI) paper green.
Hydrogen chloride
Hydrogen chloride must be collected downwards into a test tube. It cannot be collected over water. Hydrogen chloride turns damp blue litmus paper red.
Ammonia
Ammonia is less dense than air and can be collected upwards. It is also very soluble. It is an alkaline gas and turns damp red litmus paper blue.
Nitrogen dioxide
Nitrogen dioxide is a poisonous brown gas, and it is also soluble in water. It turns damp blue litmus paper red.
Water
Water turns white anhydrous copper (II) sulphate powder blue and blue cobalt chloride paper pink.
Flame tests
A platinum (or nichrome) wire is dipped into concentrated hydrochloric acid and when held to a Bunsen flame. This step is repeated until the wire does not impart any colour to the flame. It is then dipped into the acid once more and then a tiny sample of the solid is put into the flame.
If the flame is lilac, the metal is Potassium
If the flame is blue-green, the metal is Copper
If the flame is red, the metal is Lithium
If the flame is yellow-orange, the metal is Sodium
If the flame is orange-red, the metal is Calcium
If the flame is pale green, the metal is Barium
Hydroxides
Because most hydroxides are insoluble (except sodium, potassium and ammonium hydroxides), when you add sodium hydroxide solution to a solution containing metal ions, you will get a precipitate of the metal hydroxide. The colour of these precipitates will help you to identify the metal ion.
A blue precipitate shows the presence of copper (II) ions. The precipitate is copper (II) hydroxide.
Cu2+ (aq) + 2OH- (aq) → Cu(OH)2 (s)
Any copper (II) salt in solution will react with sodium hydroxide solution in this way.
CuSO4 (aq) + 2NaOH (aq) → Cu(OH)2 (s) + Na2SO4 (aq)
An orange brown precipitate shows the presence of iron (III) ions
Fe3+ (aq) + 3OH- (aq) → Fe(OH)3 (s)
A green precipitate shows the presence of iron (II) ions.
Fe2+ (aq) + 2OH- (aq) → Fe(OH)2 (s)
However, iron (II) hydroxide darkens on standing and turns orange around the tope of the tube. This is because the iron (II) hydroxide is being oxidised back to iron (II) hydroxide by the oxygen in the air.
A white precipitate which doesn’t dissolve when you add sodium hydroxide is either magnesium or calcium hydroxide.
Mg2+ (aq) + 2OH- (aq) → Mg(OH)2 (s)
Ca2+ (aq) + 2OH- (aq) → Ca(OH)2 (s)
You can differentiate by doing a flame test from the original compound. Calcium will give an orange-red flame, whiles magnesium has no flame colour.
A white precipitate that does dissolve when you add excess sodium hydroxide is aluminium hydroxide (or possible zinc hydroxide).
Al3+ (aq) + 3OH- (aq) → Al(OH)3 (s)
If no precipitate is formed, but there is a smell of ammonia, this shows the presence of ammonium ions.
NH4+ (s or aq) + OH- (aq) → NH3 (g) + H2O (l)
Carbonates
Most carbonates split up to give a metal oxide and carbon dioxide when you heat them. You can test the carbon dioxide given off using limewater and there may be some helpful colour changes as well.
If you add a dilute acid to a solid carbonate, carbon dioxide is produced in the cold. It is best to use nitric acid, as all nitrates are soluble.
CO32- (s) + 2H+ (aq) → CO2 (g) + H2O (l)
Sulphites
Sulphites contain the ion SO32- and so have similar formulae to carbonates. Sulphites react with dilute acids to give off SO2.
SO32- (s) + 2H+ (aq) → SO2 (g) + H2O (l)
Sulphites usually need warming with the acid before you produce enough sulphur dioxide gas to test adequately. Sulphur dioxide turn orange potassium dichromate (VI) paper green.
Sulphates
Make a solution of your suspected sulphate, and then add enough dilute hydrochloric acid to make it acidic (to destroy other compounds which might produce white precipitates) and then add some barium chloride solution.
As barium sulphate is insoluble, of there are sulphate ions are present, a white precipitate will form.
Ba2+ (aq) + SO42- (aq) → BaSO4 (s)
Chlorides, bromides and iodides
Make a solution of your suspected chloride, bromide or iodide and then add enough nitric acid to make it acidic. Then ad some silver nitrate solution
A white precipitate (of silver chloride) shows the presence of chloride ions.
Ag+ (aq) + Cl- (aq) → AgCl (s)
A pale cream precipitate (of silver bromide) shows the presence of bromide ions.
Ag+ (aq) + Br- (aq) → AgBr (s)
A yellow precipitate (of silver iodide) shows the presence of iodide ions.
Ag+ (aq) + I- (aq) → AgI (s)
Non-metals: Nitrogen, Oxygen, Sulphur and Hydrogen
The composition of air is as follows:
To show that the air contains about 1/5 oxygen:
Set up you apparatus- two gas syringes, one with 100 cm3 of air, and the other empty. Then connect them with a silica tube packed with copper.
Have a Bunsen burner ready to heat the copper.
Then push the air backwards and forwards over the heated copper, which will turn black as copper (II) oxide is being formed.
The volume of air decreases as oxygen is used up.
Make sure as the copper reacts, the Bunsen burner is moved along the tube, so that it is always heating fresh copper.
Eventually all the oxygen in the air will be used up. The volume stops contracting and the copper stops turning black. Once the apparatus has cooled, a final reading can be taken, and it should be around 79cm3 of gas left in the syringe.
When a compressed gas escapes, its temperature falls; the expanding gas thus feels colder than when it was compressed. This is the effect used to cool air until it liquefies at around -200°C. By recycling the cooled air several times, liquid air is produced as the temperature continues to fall. Since both carbon dioxide and water vapour turn to solids at low temperatures, they must first be removed; otherwise the ice would block the pipes, thus stopping the process. The gases that remain are nitrogen, oxygen and traces of the noble gases (main argon).
Liquid air can then be separated into separate gases through fractional distillation. The temperature begins to rise (from around -200°C); nitrogen boils first at -196°C, followed by oxygen at -186°C.
Laboratory preparation of oxygen
- Assemble the apparatus: A conical flask with a dropping funnel above it and a delivery tube leading to a trough with a beehive shelf and an upturned gas jar.
- Add about 1g of the catalyst, manganese (IV) oxide to the flask.
-
Add 25cm3 from the dropping funnel into the flask.
- Collect several jars of oxygen gas over the trough.
The gas should:
- Relight a glowing splint
- When magnesium ribbon is placed in it, produce magnesium oxide
Reactions with oxygen
Magnesium catches fire in oxygen easily. It burns with a blinding white flame to form magnesium oxide (MgO) a white powder.
Iron does not burn in oxygen, but the hot metal glows brightly in oxygen and gives off yellow sparks. It produces iron oxide, Fe3O4, a black powder.
Copper does not burn, but the hot metal becomes coated with a black substance, copper oxide (CuO).
Sulphur burns in oxygen to give the acidic gas, sulphur dioxide.
Iron (and steel) react, in the atmosphere, to produce reddish-brown rust. Iron rusts in the presence of oxygen and water. The formula of rust is Fe2O3.xH2O (x is a variable number). The rusting process is an oxidation reaction can be represented by the equation:
Fe → Fe3+ + 3e-
The iron loses electrons to form iron (II) ions, Fe2+, which are then oxidised by the air to iron (III) ions, Fe3+. Reactions involving water form the actual rust.
Rusting is speeded up when acid (carbon dioxide or sulphur dioxide) or salt are present.
Rusting can be prevented if either oxygen or water can be kept away from the iron. There are many ways of achieving this. They include:
-
Painting. A coat of paint prevents oxygen and water coming into contact with the iron. But this only works while the coating of paint is unbroken. This type of rust prevention is used to prevent car bodies or iron railings rusting.
-
Oil or grease. These also prevent oxygen or water coming into contact with the iron. This method is particularly useful when parts are moving.
-
Galvanising (coating with zinc). This is often used with fences, as it cannot be used for food cans, as zinc compounds are poisonous. For food cans, tin plate is used instead. As long as the zinc layer is unscratched, the zinc serves as a barrier to air and water. Even if it is scratched, the iron will not rust, as zinc is more reactive than iron and so corrodes instead of iron. As it loses its electrons, any iron atom that loses any electrons can immediately regain them.
-
Sacrificial protection. Magnesium or zinc blocks can be strapped to the hulls of ships and will corrode in preference to iron. As long as they remain, no rusting will take place. These blocks can easily be replaced.
-
Electroplating. This can also be called alloying. When iron is alloyed with chromium and nickel to produce stainless steel, this prevents the iron from rust. However this is expensive.
Oxidation and reduction
Oxidation and reduction reactions occur together. If one substance is oxidised, another substance is reduced. A reaction where oxidation and reduction take place is called a redox reaction.
If a substance has been oxidised, it has gained oxygen. Oxidation is gain of oxygen.
If a substance has been reduced, it has lost oxygen. Reduction is loss of oxygen.
Magnesium + copper (II) oxide → magnesium oxide + copper
The magnesium has been oxidised. It is therefore the reducing agent (something that reduces something else). The copper has been reduced. It is therefore an oxidising agent (something that oxidises something else).
No reaction would have taken place if the copper (II) oxide was heated alone as the magnesium was necessary as it removes the oxygen. Likewise, heating magnesium alone would not have worked, as the copper (II) oxide was needed, to supply the oxygen.
Common reducing agents include hydrogen, carbon and carbon monoxide. Common oxidising agents include oxygen, chlorine, concentrated sulphuric acid and concentrated nitric acid.
Hydrogen peroxide is unusual as it can act as both an oxidising and reducing agent. It can oxidise lead sulphide to lead sulphate.
PbS + 4H2O → PbSO4 + 4H2O
But if it reacts with an even stronger oxidising agent, such as the purple manganate ion, then it acts as a reducing agent.
Electron transfer
If we look careful at the reaction taking place,
Magnesium + copper (II) oxide → magnesium oxide + copper
Mg(s) + CuO(s) → MgO(s) + Cu(s)
The magnesium and the copper are both metals and are made up of metal atoms but the copper (II) oxide and the magnesium oxide are both ionic compounds. The copper (II) oxide contains Cu2+ and O2- ions whilst the magnesium oxide contains Mg2+ and O2- ions. Writing this in an equation gives:
Mg(s) + Cu2+(s) + O2-(s) → Mg2+(s) + O2-(s) +Cu(s)
The oxide ion is completely unaffected by the reaction. It is a spectator ion. The equation showing things being changed is this:
Mg(s) + Cu2+(s) → Mg2+(s) +Cu(s)
This is known as an ionic equation and shows what is actually happening is that magnesium atoms are becoming magnesium ions. The magnesium atoms lose electrons to form magnesium ions.
Mg (s) → Mg2+ (s) + 2e-
Those electrons have been gained by the copper (II) ions to make the atoms present in the metallic copper.
Cu2+(s)+ 2e- → Cu(s)
Oxidation is thus defined as the loss of electrons and reduction as the gain of electrons.
Laboratory preparation of carbon dioxide
- Prepare the apparatus in the same was as oxygen
- Place about 10g of marble chips in the flask
-
Add about 25cm3 dilute hydrochloric acid from the dropping tube to the flask
- Collect several jars of gas
When you bubble carbon dioxide through limewater, a white precipitate forms.
Carbon dioxide as a gas is colourless, denser than air and fairly soluble in water. Carbon dioxide is a solid below -78°C.
Reactions of carbon dioxide
Carbon dioxide reacts with water to give a solution of a weak acid, carbonic acid. The salts of carbonic acid are carbonates. Carbonic acid decomposes if you warm the solution, and carbon dioxide is released. So the reaction is reversible.
If carbon dioxide is bubbled into an alkali, such as sodium hydroxide, it forms the salt sodium carbonate.
With the alkali solution of the alkali, calcium hydroxide (limewater), a solid carbonate is precipitated. This is why it goes cloudy in the presence of carbon dioxide. If more carbon dioxide is bubbled through the limewater, the precipitate re-dissolves to form soluble calcium hydrogen carbonate.
Uses of carbon dioxide
Carbon dioxide should be used in place of water to extinguish fires, as it is a very dense gas. It is also used to make fizzy (carbonated) drinks. The gas is dissolved under pressure and is released when the bottle or can is opened. The sharp taste of fizzy drinks is caused by the carbonic acid, but as the carbon dioxide bubbles out, the drink tends to taste sweeter.
Nitrogen
Nitrogen is a very unreactive gas, usually under normal conditions; it does not react with the oxygen in the air. But during thunderstorms, the energy released by lightening is enough to make the two gases react. The product of the reaction is a mixture of nitrogen oxides.
High temperatures and pressures in car engines can also make nitrogen and oxygen react.
N2 + O2 → 2NO
Followed by…
2NO + O2 → 2NO2
Nitrogen dioxide is a brown acidic gas that dissolves in water to give an acidic solution; this is one of the causes of acid rain pollution.
Liquid nitrogen is a very cold substance! So cold, it can rapidly freeze fresh food to preserve it. This food can be stored for extended periods of time without it going bad. As nitrogen is gas is so unreactive, it is often used to fill food containers instead of air. The absence of oxygen stops the food from oxidising and reduces the chances of bacteria growth and thus spoilt food.
Sulphur
Sulphur is a solid, yellow element. It is found in many places of the world and is often released by volcanoes. Many elements form compounds with sulphur. These are called sulphides. Many of the most important metal ores are sulphides, such as galena and lead sulphide. Sulphur exists in 3 main forms (2 crystalline forms and one plastic form).
Rhombic sulphur is the form of sulphur that is stable at room temperature. They occur in rhombus-shaped crystals. The individual molecules are in rings of 8 atoms. Rhombic crystals can be obtained by dissolving powdered sulphur in the toxic solvent carbon disulphide, followed by evaporation.
Monoclinic sulphur is obtained when rhombic sulphur is heated up and then cooled. It is in the form of long, needle-shaped crystals. The sulphur is still in S8 but arranged in a different crystal structure to rhombic sulphur. At room temperature, the crystals will gradually change back to rhombic form.
Plastic sulphur is formed when hot molten sulphur is poured suddenly into cold water. It makes a flexible brown filament (like chewing gum) as the sulphur has cooled too fast for crystals to form. But if the plastic sulphur is left for a few days it will turn back into rhombic sulphur.
Sulphur dioxide reacts with water to form an acidic solution containing sulphurous acid (H2SO3). This is also called sulphuric (VI) acid. If it is bubbled through an alkali such as sodium hydroxide, it forms a hydrogen sulphite salt. But if still more alkali is added, the sulphite salt is produced.
Hydrogen
Hydrochloric acid reacts vigorously with magnesium, although a strong hydrochloric acid is needed to react with aluminium and the reaction between zinc and iron and hydrochloric acid is quite slow. Hydrogen gas is produced in all of the above reactions.
Laboratory preparation of hydrogen gas
- Set up the apparatus: a conical flask (containing granulated zinc) fitted with a delivery tube and funnel as well as a burette (containing dilute hydrochloric acid). The delivery tube should go to a trough filled with water with a beehive shelf as support for an overturned gas jar.
- Add the acid slowly until a steady stream of bubbles can be seen. Do not collect the gas immediately since it will be mixed with air from the flask and hydrogen-air mixtures are explosive.
- Collect several small gas jars of hydrogen.
The gas collected can be tested:
- With indicator. Hydrogen is neutral.
- With a burning splint, a small pop sounds should be heard.
- Igniting a small sample of hydrogen, the gas explodes to form water.
If a jet of burning hydrogen comes into contact with a cold surface, a colourless liquid condenses. This liquid is pure water. Two volumes of hydrogen combine with one volume of water, explaining the formula H2O.
Water turns white anhydrous copper (II) sulphate blue. Anhydrous copper (II) sulphate lacks water of crystallisation and by adding water; you add the water of crystallisation.
In the absence of water, cobalt chloride paper is blue. Hydrous cobalt chloride paper is pink.
Pure water boils at 100°C under normal conditions and has a density of 1 g/cm3.
Rates of reaction
Chemical reactions occur at different rates. Some are very fast like that between hydrogen and air (where a squeaky pop is heard), while others take hundreds of years before an effect can be seen, for example, the reaction between a limestone building and acid gases in the air.
In a laboratory, for the experiment to be useful, the experiment must not be either too slow or too fast, as it is necessary to find a change that can be observed during a reasonable amount of time.
On a graph, when the line becomes horizontal, the reaction that has been taking place has finished. When the graph is at its steepest, usually at the start of the reaction, the rate of reaction is at its fastest (greatest).
There are many changes in a reaction that be measured; they include colour changes, formation of precipitate, pH change, temperature change and volume change/ loss of mass.
Everything in the world is made of particles, which can move. These particles must collide before a reaction can take place. But not every collision leads to a reaction, as the particles must have a sufficient amount of energy, (activation energy) for a reaction to take place. Sometimes the particles just bounce off each other, for a fast reaction, there needs to be lots of collisions, enough energy for success and correct orientation.
In a concentrated solution there will be lots of particles colliding at the beginning of a reaction, but as the reaction proceeds, the rate of reaction will decrease as particles will have reacted and been used up, leading to fewer collisions.
When you compare the size of solid, always the smaller sized particles will react faster as they have an increased surface area, and the reaction must take place on the surface of the solid, so there is a greater area for the collisions to take place. But the mass of gas produced at the end will eventually be the same as the same quantities of everything (i.e. solid and acid) and conditions (temperature, concentration etc.) would be used to ensure fairness. High surface areas are frequently used to speed up reactions outside the lab.
Increasing the temperature means that particles are moving faster and so will collide more often, this will speed up the reaction.
By increasing the pressure on gases, you are forcing the particles closer together, so they are also more likely to collide. This is the same as increasing the concentration of the gas.
Another way of increasing the rate of reaction is to use a catalyst. Catalysts provide an alternative route with lower activation energy. This means that many more collisions are likely to be successful. However the catalysts themselves remain chemically unchanged (unless they are enzymes, which can be denatured).
Energetics
Chemical reactions are accompanied by an energy change, which, in solution, may be detected as a temperature change.
Exothermic reactions are reactions where heat energy is released and given out to the surroundings. Endothermic reactions are reactions where energy is taken in fro the surroundings. An example of an exothermic reaction is the combustion of methane with oxygen. An endothermic reaction is thermal decomposition. Neutralisation reactions are also exothermic.
Electrolysis
Electrolysis is a chemical change caused by passing electricity through a compound which is either molten or in solution. An electrolyte is a substance that undergoes electrolysis. All electrolytes contain ions. Electricity passes in and out of the electrolytes via electrodes.
In a metal or graphite, electricity is simply a flow of electrons. The movement of the electrons doesn’t produce any chemical change in the metal or graphite. Graphite and metal contain free-moving (delocalised) electrons. In an electrical circuit, the battery or power pack is a type of ‘electron pump’, pushing the electrons through the metal or graphite. Even liquid metal (mercury) conduct electricity; no decomposition takes place.
Hardly any solid compounds conduct electricity, yet, lots of compounds will conduct electricity if they are molten or when they are dissolved in water, this is because their structure is now broken up and there are ions free to move.
Carbon is frequently used for electrodes as it conducts electricity (when solid) and it is fairly chemically inert. Platinum is also used, as it is also fairly inert. The positive electrode is called the anode. The negative electrode is called the cathode. Positive Anode; Negative Cathode. Ions move (or migrate) towards the electrode of opposite charge. When the ions reach the electrode they may be discharged. This involves either a transfer of charge to or from the electron.
There are some rules for predicting products of electrolysis for aqueous solutions. At the cathode, if the metal in the salt is low in the activity series (e.g. silver, copper or lead), the metal is deposited. If the metal in the salt is high in the activity series (e.g. potassium or sodium), hydrogen is evolved from the water.
At the anode, if the non-metal ion in the salt is a halide ion (e.g. chloride, bromide or iodide), then the halogen is released. If the non-metal ion in the salt is not a halogen, oxygen from the water is given off.
4OH-(aq) – 4e- → O2 + 2H2O
Electrolysis of molten zinc chloride
A bulb won’t light up while zinc chloride is solid. But as soon as the zinc chloride melts, the bulb lights. Zinc starts to form around the cathode (the negative electrode). Chlorine gas is produced at the anode (the positive electrode). The bulb goes out when the heat is removed and the metal solidifies. The overall reaction is ZnCl2 → Zn + Cl2.
Zinc chloride is an ionic compound. The solid consists of a giant structure of zinc and chloride ions packed regularly in a crystal lattice. It doesn’t have any free electrons, and the ions are locked tightly in the lattice and aren’t free to move. The solid zinc chloride doesn’t conduct electricity. As soon as the solids melts, the ions become free to move and this movement, which enables the electrons to the electrons to flow in the external circuit.
As soon as you connect the power source, it pumps any mobile electrons away from the anode and towards the cathode.
The excess of electrons at the cathode makes it negatively charged, and the anode becomes positively charged, as has a shortage of electrons. There is a limit to how many extra electrons the power source can squeeze into the cathode because of the repulsion by the electrons already there.
When the zinc chloride melts, the ions are free to move. The positive zinc ions are attracted to the cathode, as the electrons are transferred from the cathode to the ion and the ion is changed into a zinc atom. At the anode, electrons are transferred from the chlorine ions to the anode and the ions are changed into chlorine atoms. Two chlorine atoms combine to form a chlorine molecule. The new electrons on the anode are pumped away by the power source to help fill the spaces being created at the cathode. Because the electrons are flowing in the external circuit, the bulb lights up.
Ions are discharged at electrodes. Discharging an ion means that it loses its charge. Electrons can flow in the external circuit because of the chemical changes to the ions arriving at the electrodes.
The zinc ions gain electrons at the cathode:
Pb2+ (l) + 2e- → Pb (l)
Gain of electrons is reduction. The zinc ions are reduced to zinc atoms.
The chlorine ions lose electrons at the anode:
2Br- (l) → Br2+ (g) + 2e-
Loss of electrons is oxidation. Bromide ions are oxidised to bromine molecules.
Reduction happens at the cathode and oxidation happens at the anode.
Positive ions are cations because they are attracted to the cathode. The negative ions move to the anode, where they give electrons to the electrodes. Negative ions are known as anions.
Aluminium extraction
Aluminium metal is always covered in a thin layer of aluminium oxide. This oxide layer can be made thicker to protect the metal from corrosion. The aluminium object is made the anode of an electrolytic cell. Oxygen is produced at the anode and this combines with the metal to give a thicker oxide layer. This is called anodising.
Aluminium is the most common metal in the Earth’s crust; its main ore is bauxite (impure aluminium oxide). The bauxite is first treated to give pure aluminium oxide. Because aluminium is a fairly reactive metal it has to be extracted using electrolysis, Aluminium oxide has a very high melting point and it isn’t practical to electrolyse molten aluminium oxide.
Instead aluminium oxide is dissolved in molten cryolite. This is another aluminium compound that melts at a lower temperature. The electrolyte is a solution of aluminium oxide in molten cryolite at a temperature of around 1000°C.
The electrolysis cell has a carbon lining of a steel tank (with heat resistant bricks) as the cathode and carbon anodes.
The molten aluminium is siphoned off from time to time, and a fresh aluminium oxide is added to the cell. The heat operates at about 5 volts but with currents of up to about 100,000 amps. The heat generated by the huge current keeps the electrolyte molten.
The two major costs of extracting aluminium are the costs of electricity and also the replacing of the anodes as they burn away. At the working temperature of the cell, the oxygen reacts with the carbon of the anode to produce carbon dioxide. Smelters are situated where there is good access to low cost electricity.
Industrial electrolysis
The copper extracted by heating the ore is not pure enough to as an electrical conductor. This impure copper is purified using electrolysis.
The purification of copper uses the electrolysis of copper (II) sulphate solution. The anode is impure copper and the cathode is pure copper. The copper from the anode goes into solution as Cu2+ ions.
Cu(s) → Cu2+ (a) + 2e-
At the cathode, Cu2+ ions are deposited as copper.
Cu2+ (aq) + 2e- → Cu(s)
The impure copper at the anode gradually disappears, and pure copper is plated on the cathode.
The impurities collect in the anode mud. As copper ores have become rare and expensive to use, much of the new copper needed is obtained by recycling old copper wires and pipes.
Uses of copper
- Electrical wiring as it is a good conductor of electricity and it is easily drawn into wires.
- Domestic plumbing as it doesn’t react with water and is easily bent into shape.
- Making brass. Brass is an alloy of copper and zinc in various proportions.
- Making coins. ‘Silver’ coins are made from an alloy of copper and nickel (‘cupronickel’).
Electrolysis of brine
The electrolysis of brine (sodium chloride solution) produces sodium hydroxide, hydrogen and chlorine.
The salt solution is purified to remove ions other than sodium and chloride ions and is then electrolysed to produce three useful chemicals. The electrolysis can be carried out in a membrane cell. The cell is designed to keep the products apart. If chlorine comes into contact with sodium hydroxide, it reacts to make bleach (a mixture of sodium chloride and sodium chlorate (I) solution. If chlorine comes into contact hydrogen, it produces a mixture that would explode violently on exposure to sunlight or heat.
The anode is made of titanium and the cathode of steel. The diaphragm in the middle is made of asbestos; its function is to let sodium ions through but to keep the gases apart. Chloride ions cannot go through, as they are larger than sodium ions.
At the anode, chloride ions are discharge to produce chlorine gas.
2Cl-(aq) → Cl2 (g) + 2e-
The sodium ions are attracted into the right hand compartment by the negatively charged electrode. It is too difficult to discharge sodium ions, so hydrogen ions from the water are discharged instead to produce hydrogen gas.
2H+ (aq) = 2e- → H2 (g)
More and more water keeps splitting up to replace the hydrogen ions as soon as they are discharged. Each time a water molecule splits up it produces a hydroxide ion as well. This means that there is a build up of sodium ions and hydroxide ions in the right hand compartment – sodium hydroxide solution is produced.
Uses of products
- The plastic PVC
- Bleaches
- Paints and dyestuffs
- Medical drugs
- Sterilising agent
- Sodium hydroxide solution
- Soaps
- Detergents
- Dyes
- Ceramics
- As fuel in hydrogen fuel cells
- To harden vegetable oils to make margarine.
- To make hydrogen peroxide
Hydrocarbons
Hydrocarbons are compounds of carbon and hydrogen only. There are two common families of hydrocarbons, the alkanes and the alkenes. Members of family have similar chemical properties and physical properties gradually from one member to another. Hydrocarbons are organic molecules, they can exist in chains, branched chains or in rings of carbon atoms with hydrogens attached.
Crude oil is a mixture of hydrocarbons. Crude oil is finite and non-renewable.
As molecules get bigger (due to the increased number of carbon atoms), many of properties change in a regular pattern. As the molecules get bigger, the attractions increase between neighbouring molecules. When these intermolecular attractions increase it becomes harder for one molecule to pull away from its neighbours.
As the molecule size increases:
-
Boiling point increases, larger molecules are attracted to each other more strongly than smaller ones, so more heat is needed to break those attractions.
-
Liquids become less volatile. The bigger the hydrocarbon, the more slowly it evaporates at room temperature. This is also because the bigger molecules are more strongly attracted and so do not release each other and turn into a gas easily.
-
Liquids become more viscous (flow less easily). Liquids become stickier as the forces of attraction are larger.
-
Bigger molecules become less flammable. This is because the ratio of carbon to hydrogen increases and more oxygen is needed for the molecule to burn.
3-dimensional formula
H
C
H H H
Hydrocarbons are often drawn using 3-D formula. The wedge shows the bond pointing towards you. The dotted line shows the bond pointing away from you. The normal lines are bonds in the plane of the paper.
The bond shown is methane; methane has a bond angle of around 109°. This is the angle in tetrahedral molecules.
Names for organic compounds can look complex but each part of the name tells you some specific about the molecule.
A family of hydrocarbons is called a homologous series. Each member of a homologous series contains the same functional group, differing form the one before by one carbon atom.
Alkanes
Alkanes are saturated hydrocarbons; this means that they contain only single carbon-carbon bonds. The first members of the family are used extensively as fuels, but other than that have few other reactions. All alkanes fit the formula CnH2n+2. Many alkanes are obtained from crude oil by fractional distillation.
Alkanes burn in oxygen to form CO2 and H2O. They give off CO if there is only a low supply of oxygen.
They are quite unreactive and there is no chemical test for an alkane. Their main usage is fuels.
Structural isomerism
H
I
H—C—H
H H
I I
H—C—C—C—H
I I I
H H H
The carbon atoms in a hydrocarbon molecule can be arranged in different ways. For example, in butane, C4H10, the carbon atoms can be positioned in two ways. Structural isomerism is the existence of two or more different structures with the same molecular formula.
H H H H
I I I I
H—C—C—C—C –H
I I I I
H H H H
The second isomer shown is 2-methylpropane, because the longest straight chain is propane, the branch is methane, and the branch is on the second carbon from the end.
The various isomers will have slightly different physical properties because of their different intermolecular forces. Branched chains have weaker intermolecular attractions that straight ones. Intermolecular forces are only effective over very short distances. The more branching there is in a chain, the more difficult it is for the molecules to get close to each other.
Large unbranched alkanes have higher boiling points than branched isomers because there are larger attractive intermolecular forces between the unbranched isomer molecules. The higher boiling points are a result of the large amount energy needed to separate one molecule from another.
As the branching increases, the boiling points fall.
The chlorination of methane
Halogenation is the replacement of one or more hydrogens in an organic compound by halogen atoms. This is called a substitution reaction. The reactions use light energy to start.
When methane is reacted with chlorine the products of the react depend on whether there is an excess of methane or an excess of chlorine.
methane + chlorine → methyl chloride + hydrogen chloride
CH4 + Cl2 → CH3Cl + HCl (all gases)
In sunlight, chlorine reacts so fast with alkanes that it can explode.
Alkenes
Alkenes are another homologous series, so there have similar chemical and physical properties that gradually change from one member to the next.
Alkenes are unsaturated hydrocarbons, which mean that they contain at least one carbon-carbon double bond. They burn well and are reactive in other ways too. Their reactivity is fur to the carbon-carbon double bond.
The position of the double bond can vary and this is reflected in the naming of the alkene. For example, but-1-ene has the double chain starting on the 1st carbon, but but-2-ene has the double bond starting on the 2nd carbon.
The general formula for alkenes is CnH2n. There cannot be an alkene with only one carbon atom as alkenes all have a carbon-carbon double bond.
In alkenes the bonds are arranged symmetrically around the carbon. The bond angel is 120° and the bonds point to the corners of an equilateral triangle.
Alkenes burn in carbon dioxide to give carbon dioxide and water. But more importantly, they undergo addition reactions. Part of the double bond breaks and the electrons are used to join other atoms are the carbon atoms. After undergoing an addition reaction, the unsaturated alkene becomes a saturated alkane.
The test to distinguish an alkene from an alkane is an addition reaction. If an alkene, such as ethane, is bubbled through bromine water, the solution changes from red-brown to colourless.
Ethane + bromine → 1, 2-dibromoethane
Separating crude oil
Crude oil was formed millions of years ago from the remains of animals that were pressed together under layers of rock. It is usually found underground and is a finite fuel. But crude oil is an extremely valuable resource which must be used efficiently. As it contains many useful chemicals and we need to separate these chemicals so they are not wasted.
The chemicals in crude oil are separated unto useful fractions by a process known as fractional distillation. The crude oil is separated into separate fractions in an oil refinery. Each fraction contains hydrocarbons which boil within a temperature range, and has a different use.
The fractionating column, in which the crude oil is heated, is cooler at the top and hotter at the bottom. The crude oil is heated in a furnace and the vapour mixture given off rises up the column and the different fractions condense out at different parts of the column, depending on their boiling points. The lower the boiling point, the higher the vapour condenses in the column. For example, if a hydrocarbon boils at 120 °C, the bottom of the column will be much hotter than 120 °C, so the hydrocarbon will remain as a gas. But as it travels up the column, the temperature lowers, until it falls to around 120 °C. Then the hydrocarbon will turn into a liquid; it condenses and can be tapped off.
Those hydrocarbons removed near the bottom of the column are dark and viscous, whereas the hydrocarbons in petroleum gases have boiling points so low that the temperature in the fractionating column never cools enough for them to condense to liquids.
Use of the fractions
All hydrocarbons burn in air to form carbon dioxide and water. They also release a lot of heat in the process; therefore they can be used as fuels.
If there isn’t enough air or oxygen, you get an incomplete combustion, this leads to the formation of carbon monoxide instead of carbon dioxide. Carbon monoxide is colourless, odourless and very poisonous as it combines with haemoglobin, preventing it from carry the oxygen. You may become seriously ill and die, if you cannot get enough oxygen to your cells.
A flame from an incomplete combustion will be blue, but where the combustion is incomplete, the flame is likely to be yellow.
Cracking
Although the fractions from crude oil distillation are useful, the amount of each fraction you get depends on the proportions of the various hydrocarbons available in the crude oil, not the amount in which they are need. More petrol is needed than is found in crude oil. The hydrocarbons in crude oil are also fairly unreactive; to make organic chemicals they must be converted into something more reactive.
Larger molecules can be broken down into smaller ones by cracking. Cracking involves passing the vapour of the high boiling point fraction over a catalyst at a high pressure with a high temperature. It is a very useful process; the bog hydrocarbon molecules in gas oil can be broken down into the smaller ones needed for petrol.
The majority of hydrocarbons found in crude oil have single carbon-carbon bonds, but during the process, new molecules are formed which have double carbon-carbon bonds between their atoms. These molecules are much more reactive and can be used to make lots more things.
The gas oil or naphtha fractions are heated to give a gas and then passed over a catalyst of mixed silicon dioxide and aluminium oxide at about 500°C. Cracking is just an example of thermal decomposition which produces a mixture of alkanes and alkenes.
Problems from transporting crude oil
It is still cheap to transport crude oil by sea and the numbers of trade ships continues to rise but occasionally, ships have accidents and shed their cargo. When a tanker carrying crude oil has a spill, it causes enormous environmental damage, especially if it is near land and washes ashore and affects the birds and other sea life.
After an oil spill, a thin layer of oil covers a large area of sea. If allowed to spread, it will prevent the evaporation of water, affecting the water cycle.
Burning hydrocarbons produces carbon dioxide, water and nitrogen oxides. These contribute to global warming and acid rain. In major cities, the main pollutant is unburned fuel from motor vehicles, this leads to smog and a low air quality which is hazardous to people living there.
Global warming is a result of the carbon dioxide produced from burning fuels. Acid rain is produced by sulphur dioxide in fuels, or by nitrogen oxides produced when they burn in air. The acid rain damages trees and makes lakes acidic, killing the fish. Buildings also corrode. The holes in the ozone layer are also an effect of the hydrocarbons reacting with the ozone.
Ethanol
Ethanol is a member of the alcohol family. Alcohols contain an –OH group covalently bonded onto a carbon chain. Their general formula is CnH2n+1OH.
The boiling points of alcohols increase as they get bigger.
Uses of alcohol
Alcohol is found in alcoholic drinks. Alcohol depressed the activity of some of the higher parts of the brain. This releases inhibitions. Alcohol is poisonous in large quantities. Even small quantities can reduce concentration.
It is also used as a solvent in perfumes as the –OH group allows it to dissolve in water and it also dissolves in organic compounds.
In an increasing amount of countries, alcohol is being used as a fuel, as it only releases CO2 and H2O and not the pollutant gases of petrol. It is also renewable because it comes from plants.
Production of ethanol
Ethanol can be produced in two ways, either by fermentation or using ethene and steam.
When ethanol is made by fermentation, yeast is added to a sugar solution and left in the warm in anaerobic conditions. Ethanol can be prepared by the fermentation of sugar solutions using enzymes in yeast.
The chemical reaction for fermentation is:
Glucose (sugar) → ethanol + carbon dioxide
C6H12O6 (s) → 2C2H5OH (aq) + 2CO2 (g)
Yeast is killed by more than 15% of alcohol in the mixture so it is impossible to make pure alcohol by fermentation.
At the end of the fermentation process, which takes time as it is an enzymic reaction and a batch process, a more concentrated solution of ethanol is produced by fractional distillation. This takes advantage of the different in boiling points in ethanol and water. Water boils at 100 °C whereas ethanol boils at 78 °C.
The source of the sugar determines the type of alcoholic drink produced, for example, grapes produce wine and barley for beer.
Ethanol can also be made from alkenes produced by refine crude oil. The reaction is:
Phosphoric acid
ethene + steam → ethanol
C2H4 (g) + H2O (g) → C2H5OH (g)
It is made by reacting ethene with water, a process known as hydration. There are quite extreme conditions in terms of energy (300 °C) and 60-70 atmospheres of pressure, so the cost is high.
Only a small amount of ethene reacts, the unreacted ethene is recycled through the process.
The method used to manufacture ethanol depends upon the materials available. Countries, usually developed countries) with easy access to crude oil, often produce ethanol from ethene; whereas counties which do not have crude oil but have sugar, produce ethanol from sugar cane, preferring fermentation.
There are disadvantages and advantages for each method. The advantages for fermentation are that it uses renewable resources, like sugar cane, and the process also utilises otherwise waste materials. The disadvantages are that a large volume of starting materials is needed to produce a small amount of ethanol. Large reaction vessels are also needed. Not only is the fermentation slow, fractional distillation is expensive. But the advantages of the ethene method are that it produces pure ethanol and does not need large reaction vessels. It is also a continuous process. However, ethene is a non-renewable resource and energy is needed to produce steam. A high percentage of ethene remains unreacted and must be recycled.
Reaction with sodium
Alcohols react gently with sodium to produce hydrogen. The product is sodium ethoxide and this reaction is often used to treat small sodium spills or dispose of small amount of sodium ethoxide.
2Na (s) + 2C2H5OH (l) → 2C2H5ONa (s) + H2 (g)
Oxidation of ethanol
Ethanol can be oxidised to make ethanoic acid (a carboxylic acid) using an oxidising agent, e.g. potassium dichromate (VI) solution in sulphuric acid. The acidic solution of potassium dichromate is orange in colour, but we know that a reaction has occurs where its colour has changed to green,
Ethanoic acid is the acid in vinegar. If wine or beer is not properly sealed, the ethanol is oxidised over the oxygen in the air and the wine or beer becomes acidic, so spoiling their taste.
Dehydration of ethanol
Ethanol can be dehydrated to give ethene by heating it with an excess of concentrate sulphuric acid at about 170°C. Alternatively, you can soak mineral wool in ethanol and heat it over a catalyst of aluminium oxide and collect the ethene over cold water.
Ethanol → ethene + water
Ethanol and carboxylic acids
Ethanol reacts with carboxylic acids (with a concentrated sulphuric acid) from esters. Carboxylic acids have the general formulae of RCOOH, for example, CH3COOH is ethanoic acid. The beginning of the name, ‘R’ or ‘eth’, etc, includes all the carbon atoms in the –COOH group.
Alcohol + acid → ester + water
Esters often have distinct pleasant smells. They are often used as flavourings for perfumes; small esters can be used as solvents.
The name of the ester is the alcohol first, and then the acid used –‘oate’.
Extraction and uses of metals
The method used to extract the metal from the ore depends on the position of the metal in the Reactivity Series. If the metal is high in the series, its ores are stable and the metal can be obtained only through electrolysis. These metals include potassium, sodium, calcium, magnesium and aluminium.
Metals in the middle of the Reactivity Series do not form very stable ores and they can be extracted by reduction reactions, often with carbon. Examples of metal extracted by reduction are zinc, iron and lead.
Metals low in the Reactivity Series, if at all present in ores can be extracted by heating because their ores are unstable.
There are also a few metals such as gold that are found un-combined in the Earth.
Iron
Iron is an example of a metal extracted by reduction. The reducing agent is carbon monoxide. This removes oxygen from the iron oxide to leave iron alone. The extraction is carried out in blast furnace.
The furnace is loaded with iron ore, coke and limestone and is heated by blowing hot air into the base. Hot waste gases at the top of the furnace are also piped away and used to heat the air blast at the bottom. Inside the furnace lots of reactions are taking place.
Coke is impure carbon and it burns in the hot air blast to form carbon dioxide. This is a strongly exothermic reaction and so raises the temperature to around 1500°C.
C (s) + O2 (g) →CO2 (g)
At the high temperatures in the furnace, the carbon dioxide is reduced by more carbon to give carbon monoxide.
CO2 (g) + C (s) → 2CO (g)
The carbon monoxide then reduces the iron ore to iron. The iron ore is usually haematite.
F2O3 (s) + 3CO (g) → 2Fe (l) + 3CO2 (g)
The iron melts and flows to the bottom of the furnace where it can be tapped off. In hotter parts of the furnace some of the iron oxide can be reduced by the carbon itself.
F2O3 (s) + 3C (s) → 2Fe (l) + CO2 (g)
The limestone is added to the furnace to remove impurities in the ore which would otherwise close the furnace with solid material. Because eth furnace is so hot, the limestone decomposes into calcium oxide and carbon dioxide. Since this is an endothermic reaction it is important to avoid using too much limestone as this would cool the furnace.
CaCO3 (s) → CaO (s) + CO2 (g)
The basic calcium oxide can then react with the acidic silicon dioxide and form slag which will float on the surface of the molten iron. This can be tapped off separately and is used for road making.
CaO (s) + SiO2 (s) → CaSiO3 (l)
Chromium
The major ore of chromium is chromite, FeCr2O4. Some chromium is extracted directly from this ore to produce an alloy of iron and chromium and is used to prepare stainless steel. But if pure chromium is needed, the chromite must first be converted to chromium oxide, Cr2O3. The chromium is extracted by heating the oxide with aluminium as aluminium is higher in the Reactivity Series than chromium.
Chromium oxide + aluminium → chromium + aluminium oxide
This is a spectacular reaction that gives out lots of heat and is called the thermite reaction.
Zinc
Most zinc is produced by electrolysis, using the electrolyte of zinc sulphate, prepared form the oxide. The anode is made out of lead and the cathode is of aluminium. The zinc builds up a layer around the cathode.
Zn2+ (aq) + 2e- → Zn (s)
However, zinc can also be extracted using thermal extraction when zinc oxide is heated with coke in a furnace.
Zinc oxide + carbon → zinc + carbon monoxide
Zinc oxide + carbon monoxide → zinc + carbon dioxide
Zinc boils at about 900°C, a lower temperature than inside the furnace. So the zinc distils and can be condensed by cooling.
Synthetic polymers
A polymer is formed by joining up many smaller sub-units called monomers. These smaller molecules are identical.
Under the right conditions, molecules contain carbon-carbon double bonds can join together to produce very long chains as part of the double bond is broken and the electrons in it are used to join to neighbouring molecules. When monomers add together like this the material produced is called a addition polymer. Poly(ethene) is made this way from lots of ethene molecules. It is made using heat, high pressures and an initiator which is used to get the process started.
Uses
Low density poly(ethene) is used as a thin film to make polythene bags. It is very flexible and not very strong.
High density poly(ethene) is used where greater strength and rigidity is needed for example with plastic milk bottles.
Poly(propene) is used to make ropes and crates.
Poly(chloroethene) is quite strong and rigid and can be used to make things such as drain pipes.
Condensation polymerisation
Polymers can also be made by joined two different monomers so that they react together. When they react, they expel a small molecule which is usually water, hence the name.
The manufacture of chemicals
The Haber process
The Haber process is the process for the manufacture of ammonia. Ammonia is a compound of nitrogen and hydrogen. The nitrogen needed is taken from the air using fractional distillation and the hydrogen is obtained from cracking and natural gas.
The equation for the reaction is:
N2 (g) + 3H2 ↔ 2NH3 (g)
Since the reaction is reversible, the products can decompose, reforming the reactants. But by choosing the best conditions, the highest yield of ammonia can be produced.
These conditions are one part nitrogen to 3 parts hydrogen, a high pressure and low temperature. But using a low temperature reduces the rate of reaction. This can be speeded up by using a catalyst.
The temperature is 450°C and the pressure used is 200 atmospheres. An iron catalyst is used. 450°C is a compromise temperature- producing reasonable yield reasonably quickly.
Each time the gases pass through the reaction vessel, only about 1% of the gases combine to make ammonia. The reaction mixture is cooled and the ammonia condenses as liquid. The unreacted nitrogen and ammonia are recycled.
The ammonia is used to make fertilisers and nitric acid.
Nitric acid
Ammonia is mixed with air and passed over a red-hot platinum-rhodium catalyst. The ammonia is oxidised by oxygen in the air to form nitrogen monoxide.
4NH3 (g) + 5O2 (g) → 4NO (g) + 6H2O (g)
Once cooled, the nitrogen monoxide reacts with more oxygen to form nitrogen dioxide.
2NO (g) + O2 (g) → 2NO2 (g)
This is finally absorbed in water in the presence of still more oxygen to give nitric acid.
2H2O (l) + 4NO2 (g) + O2 → 4HNO3 (aq)
The Contact process
Sulphuric acid made in industry in a three-stage process.
The first stage is burning sulphur in air to make sulphur dioxide.
S + O2 → SO2
The second step is to convert the sulphur dioxide with oxygen to form sulphur trioxide. An excess of oxygen so that the maximum amount of sulphur dioxide is converted, to avoid waste and pollution. The conditions used are 450°C. 1-2 atmospheres is the pressure used as well as V2O5 catalyst.
2SO2 (g) + O2 (g) ↔ 2SO3 (g)
In principle the third stage could be reacting sulphur trioxide with water to make sulphuric acid, but this produces an uncontrollable fog of concentrated sulphuric acid. Instead the sulphur trioxide is absorbed is in concentrated sulphuric acid to give oleum.
H2SO4 (l) + SO3 (g) → H2S2O7 (l)
This is then converted to concentrated sulphuric acid by the careful addition of water.
H2S2O7 (l) + H2O (l) → 2H2SO4 (l)