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GCSE-KS4-IGCSE Science-CHEMISTRY revision notes on

 Reversible Reactions, Chemical Equilibrium, Ammonia and nitric acid, their manufacture and uses   

This page covers: 1.Reversible Reactions * 2.Reversible reactions and Equilibrium * 3.The Haber Synthesis of ammonia * 4.The Uses of ammonia-nitric acid-fertilisers * 5.Fertilisers-environmental problems * 6.The nitrogen cycle * Advanced Level Notes on Equilibrium

  Multiple choice Quiz:Foundation or Higher * EMAIL query?comment

• A reversible reaction is a chemical change in which the products can be converted back to the original reactants under suitable conditions.
• In other words, you can change the position of the chemical
• A reversible reaction is shown by the sign ,
  • a half-arrow to the right (direction of forward reaction),
  • and a half-arrow to the left (direction of backward reaction).
• Most reactions are not reversible (irreversible) and have the usual complete arrow only pointing to the right.

Five examples 1a to 1e of reversible reactions are described below:

1a The thermal decomposition of ammonium chloride

• On heating strongly above 340^oC, the white solid ammonium chloride, thermally decomposes into a mixture of two colourless gases ammonia and hydrogen chloride.
• On cooling the reaction is reversed and solid ammonium chloride reforms.
  • This is an example of sublimation but here it involves both physical and chemical changes.
  • When a substance sublimes it changes directly from a solid into a gas without melting and on cooling reforms the solid without condensing to form a liquid.

  • Ammonium chloride + heat ammonia + hydrogen chloride

  • NH₄Cl_(s) NH_3(g) + HCl_(g)

  • so the thermal decomposition of ammonium chloride is the forward reaction, and the formation of ammonium chloride is the backward reaction.

• Note:

  • Reversing the reaction conditions reverses the direction of chemical change, typical of a reversible reaction.
  • Thermal decomposition means using 'heat' to 'break down' a molecule into smaller ones.
  • The decomposition is endothermic (heat absorbed or heat taken in) and the formation of ammonium chloride is exothermic (heat released or heat given out).
  • This means if the direction of chemical change is reversed, the energy change is also reversed.
  • Ammonium fluoride (>?^oC), ammonium bromide (>450^oC) and ammonium iodide (>550^oC), with a similar formula, all sublime in a similar physical-chemical way when heated, so the equations will be similar i.e. just swap F, Br or I for the Cl.
    • Similarly, ammonium sulphate also sublimes when heated above 235^oC and thermally decomposes into ammonia gas and sulphuric acid vapour.
      • (NH₄)₂SO_4(s) NH_3(g) + H₂SO_4(g)
    • Ammonium nitrate does not undergo a reversible sublimation reaction, it melts and then decomposes into nitrogen(I) oxide gas (dinitrogen oxide) and water vapour.
    • NH₄NO_3(s) N₂O_(g) + 2H₂O_(g)
    • This is very different reaction, in fact its an irreversible redox reaction. The nitrate ion, NO₃^-, or any nitric acid formed, HNO₃, act as an oxidising agent and oxidise the ammonium ion. If the products are cooled, ammonium nitrate is NOT reformed.
  • For more on sublimation, see the States of Matter webpage.

1b The thermal decomposition of hydrated copper(II) sulphate

• On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed.
• When the white solid is cooled and water added, blue hydrated copper(II) sulphate is reformed.

  • blue hydrated copper(II) sulphate + heat white anhydrous copper(II) sulphate + water

  • CuSO₄.5H₂O_(s) CuSO_4(s) + 5H₂O_(g)

  • The dehydration decomposition to give the white solid is the forward reaction and the 're-hydration' to reform the blue crystals is the backward reaction.

• Note:
  • The 5H₂O in the formula of hydrated copper(II) sulphate is called the water of crystallisation and forms part of the crystal structure when copper(II) sulphate solution is evaporated and crystals form.
  • This crystal structure is broken down on heating and the water is given off.
  • The thermal decomposition is endothermic as heat is absorbed to drive off the water.
  • The reverse reaction is exothermic i.e. on adding water to white anhydrous copper(II) sulphate the mixture heats up as the blue crystals reform.
  • The reverse reaction is used as a simple chemical test for water i.e. white anhydrous copper(II) sulphate turns blue.

1c The reaction of bismuth chloride with water

• When bismuth chloride (BiCl₃) is added to water, it dissolves and then reacts with water to form a white precipitate of bismuth oxychloride (BiOCl) and a colourless solution of hydrochloric acid (HCl).

  • forward reaction

  • bismuth chloride + water ==> bismuth oxychloride + hydrochloric acid

  • BiCl_3(aq) + H₂O_(l) ==> BiOCl_(s) + 2HCl_(aq)

• If hydrochloric acid is added to the mixture, the bismuth oxychloride dissolves to reform the bismuth chloride solution.

  • backward reaction

  • bismuth oxychloride + hydrochloric acid ==> bismuth chloride + water

  • BiOCl_(s) + 2HCl_(aq) ==> BiCl_3(aq) + H₂O_(l)

• Therefore the reaction is reversible and what is formed depends on the relative amounts of hydrochloric acid and water, and so the reaction equation should be written as:

  • BiCl_3(aq) + H₂O_(l) BiOCl_(s) + 2HCl_(aq)

1d The formation and decomposition of ammonia

• The synthesis of ammonia from nitrogen and hydrogen is a reversible reaction and by changing the temperature and pressure of the reacting gases you can make the reaction go one way more than another,

• e.g. high pressure makes more ammonia (forward reaction) and higher temperature causes ammonia to decompose into hydrogen and nitrogen (backward reaction).

  • nitrogen + hydrogen ammonia

  • N_2(g) + 3H_2(g) 2NH_3(g)

1e The formation and hydrolysis of an ester

• Ethyl ethanoate, an ester,  is formed by the reaction of ethanoic acid with ethanol e.g.
• ethanoic acid + ethanol ethyl ethanoate + water
• + + H₂O
• Its an equilibrium, and starting with the pure acid plus pure alcohol you get about ²/₃rds conversion* to the ester, and the reaction is catalysed by a few drops of concentrated sulphuric acid.
  • *This means a yield of about 67% in the 'atom economy'.
• If the ester is warmed with water or any dilute acid (faster), it changes back into the original acid and alcohol. This reverse reaction is called hydrolysis (backward reaction) ...
  • ethyl ethanoate + water ==> ethanoic acid + ethanol
  • whereas esterification (forward reaction) is
  • ethanoic acid + ethanol ==> ethyl ethanoate + water
• Esters are usually sweet smelling and widely used as fragrances (e.g. perfumes) and food flavourings.

• When a reversible reaction occurs in a closed system an equilibrium is formed, in which the original reactants and products formed coexist.
• In an equilibrium there is a state of balance between the concentrations of the reactants and products and once a state of chemical equilibrium is reached there is no further change in concentrations BUT the reactions don't stop!
• At equilibrium the rate at which the reactants change into products is exactly equal to the rate at which the products change back to the original reactants.
• However the final relative equilibrium amounts of the reactants and products depend on the reaction conditions e.g. the temperature and pressure.

For industrial processes, it is important to maximise the concentration of the desired products and minimise the 'leftover' reactants. A set of rules can be used to predict the best reaction conditions to give the highest possible yield of product.

Rule 1a: If the forward reaction forming the product is endothermic, raising the temperature favours its formation increasing the yield of product (lowering the temperature decreases the yield).

Rule 1b: If the forward reaction forming the product is exothermic, decreasing the temperature favours its formation (increasing temperature decreases the yield).

Rule 2a: Increasing the pressure favours the side of the equilibrium with the least number of gaseous molecules as shown by the balanced symbol equation.

Rule 2b: Decreasing the pressure favours the side of the equilibrium with the most number of gaseous molecules as shown by the balanced symbol equation.

Rules 1 above, and rule 3, below, apply to any reaction, BUT rule 2 above, ONLY applies to a reaction with one or gaseous reactants or products.

Rule 3a: If the concentration of a reactant (on the left) is increased, then some of it must change to the products (on the right) to maintain a balanced equilibrium position.

Rule 3b: If the concentration of a reactant (on the left) is decreased, then some of  the products (on the right) must change back to reactants to maintain a balanced equilibrium position.

• e.g. nitrogen + hydrogen ammonia
  • or N_2(g) + 3H_2(g) 2NH_3(g)
  • If the nitrogen or hydrogen concentration was increased, some of this extra gas would change to ammonia.
  • If the nitrogen or hydrogen concentration was decreased, some of ammonia would change back to nitrogen and hydrogen.
    • At AS-A2 advanced level things can get more complicated e.g. can you figure out why in terms of concentration to maintain the equilibrium balance? (and if a gcse student, don't worry if you can't) ...
    • So in terms of enforced change ==> system response:
      • Increasing nitrogen ==> decreases hydrogen and increases ammonia.
      • Increasing hydrogen ==> decreases nitrogen and increases ammonia.
      • Increasing ammonia ==> increases nitrogen and hydrogen.
      • Decreasing ammonia ==> decreases nitrogen and hydrogen.
      • Decreasing nitrogen ==> increases hydrogen and decreases ammonia.
      • Decreasing hydrogen ==> increases nitrogen and decreases ammonia.

Rule 4: A catalyst does NOT affect the position of an equilibrium, you just get there faster! A catalyst usually speeds up both the forward and reverse reaction but there is no way it can influence the final 'balanced' concentrations. However, the importance of a catalyst lies with economics e.g. (i) bringing about reactions with high activation energies at lower temperatures and so saving energy or (ii) saving time is saving money.

Advanced Level Notes on Chemical Equilibrium

Applying the rules 1 to 3

The forward reaction is endothermic, 178kJ of heat energy is absorbed (taken in) for every mole of calcium oxide formed.

One mole of gas is formed in the process, so there is a net increase in the moles of gas in lime formation, since there are no gaseous reactants.

From rule 1: increasing the temperature will increase the yield of calcium oxide or lime, CaO which is endothermically formed.

From rule 2: decreasing the pressure will favour the formation of more gas molecules if possible, so more carbon dioxide formed, and hence more lime.

Lime is made commercially by heating limestone to a high temperature (e.g. 1000^oC) in a limekiln that is well ventilated (this reduces the carbon dioxide pressure and so reduces the un-desired backward reaction).

The forward reaction is exothermic, 184kJ of heat energy is given out in forming hydrogen bromide according to the above equation (184/2 = 92kJ per mole of HCl formed).

There is no net change in the moles of gas (2 moles reactants 2 moles of product)

From rule 1: decreasing the temperature favours the exothermic formation of hydrogen chloride, so the equilibrium moves proportionately to the right-hand side (more HCl, less H₂ or Cl₂).

From rule 2: since there is no net change in the number of moles of gas on reaction, pressure has no effect on the yield of hydrogen chloride and the proportions of HCl, H₂ or Cl₂ stay the same.

• Ammonia gas is synthesised in the chemical industry by reacting nitrogen gas with hydrogen gas in what is known as the Haber-Bosch Process, named after two highly inventive and subsequently famous chemists.
• The nitrogen is obtained from liquified air (80% N₂). Air is cooled and compressed under high pressure to form liquid air (liquefaction). The liquid air is fractionally distilled at low temperature to separate oxygen (used in welding, hospitals etc.), nitrogen (for making ammonia), Noble Gases e.g. argon for light bulbs, helium for balloons).
• The hydrogen is made by reacting methane (natural gas) and water or from cracking hydrocarbons (both reactions are done at high temperature with a catalyst).
  • CH₄ + H₂O ==> 3H₂ + CO
  • e.g. C₈H₁₈ ==> C₈H₁₆ + H₂
• The synthesis equation for this reversible reaction is ...

  • N_2(g) + 3H_2(g) 2NH_3(g)

• .. which means an equilibrium will form, so there is no chance of 100% yield even if you use, as you actually do, the theoretical reactant ratio of nitrogen : hydrogen of 1 : 3 !
• In forming ammonia 92kJ of heat energy is given out (i.e. exothermic, 46kJ of heat released per mole of ammonia formed).
• Four moles of 'reactant' gas form two moles of 'product' gas, so there is net decrease in gas molecules on forming ammonia.
• So applying the equilibrium rules from section 2 above, the formation of ammonia is favoured by  ...
  • (a) Using high pressure because you are going from 4 to 2 gas molecules (the high pressure also speeds up the reaction because it effectively increases the concentration of the gas molecules), but higher pressure means more dangerous and more costly engineering.
  • (b) Carrying out the reaction at a low temperature, because it is an exothermic reaction favoured by low temperature, but this may produce too slow a rate of reaction,
  • So, the idea is to use a set of optimum conditions to get the most efficient yield of ammonia and this involves getting a low % yield (e.g. 8% conversion) but fast. Described below are the conditions to give the most economic production of ammonia.
  • these arguments make the point that the yield* of an equilibrium reaction depends on the conditions used.
    • * The word 'yield' means how much product you get compared to the theoretical maximum possible if the reaction goes 100%.
    • For more on chemical economics see Extra Industrial Chemistry page.
• In industry pressures of 200 - 300 times normal atmospheric pressure are used in line with the theory.
• Theoretically a low temperature would give a high yield of ammonia BUT ...
  • Nitrogen is very stable molecule and not very reactive i.e. chemically inert, so the rate of reaction is too slow at low temperatures.
  • To speed up the reaction an iron catalyst is used as well as a higher temperature (e.g. 400-450^oC).
  • The higher temperature is an economic compromise, i.e. it is more economic to get a low yield fast, than a high yield slowly!
  • Note: a catalyst does NOT affect the yield of a reaction, i.e. the equilibrium position BUT you do get there faster!
• At the end of the process, when the gases emerge from the iron catalyst reaction chamber, the gas mixture is cooled under high pressure, when only the ammonia liquefies and is so can be removed and stored in cylinders.
• Any unreacted nitrogen or hydrogen (NOT liquified), is recycled back through the reactor chamber, nothing is wasted! [nitrogen (-196^oC) and hydrogen (-252^oC) have much lower boiling points than ammonia (-33^oC). Boiling points increase with pressure, but these normal atmospheric pressure values offer a fair comparison].
• To sum up: A low % yield of ammonia is produced quickly at moderate temperatures and pressure, and is more economic than getting a higher % equilibrium yield of ammonia at a more costly high pressure and a slower lower temperature reaction.
• Detailed notes on "Rates of Reaction" for further reading.
• AND there are some more general notes on Chemical Economics on the Industrial Chemistry page.

4(a) Ammonia is used to manufacture nitric acid

• Ammonia is oxidised with oxygen from air using a hot platinum catalyst to form nitrogen monoxide and water.
• 4NH_3(g) + 5O_2(g) ==> 4NO_(g) + 6H₂O_(g)
• The gas is cooled and reacted with more oxygen to form nitrogen dioxide.
• 2NO_(g) + O_2(g) ==> 2NO_2(g)
• This is reacted with more oxygen and water to form nitric acid.
• 4NO_2(g)+ O_2(g) + 2H₂O_(l) ==> 4HNO_3(aq)
• Nitric acid is used to make nitro-aromatic compounds from which dyes are made.
• It is also used in the manufacture of artificial nitrogenous fertilisers (like ammonium nitrate, see below).

4(b) Ammonia is used to manufacture 'artificial' nitrogenous fertilisers

• Ammonia is a pungent smelling alkaline gas that is very soluble in water.

• The gas or solution turns litmus or universal indicator blue because it is a soluble weak base or weak alkali and is neutralised by acids to form salts. (more on theory on the Acids, Bases Salts page or on the Extra Aqueous Chemistry page).

• Ammonium salts are used as 'artificial' or 'synthesised' fertilisers i.e. nitrogenous fertilisers 'man-made' in a chemical works, and used as an alternative to natural manure or compost etc.

• The fertiliser salts are made by neutralising ammonia solution with the appropriate acid (more method details on Acids, Bases and Salts, but the equations are given below).

• The resulting solution is heated, evaporating the water to crystallise the salt e.g.

ammonia + sulphuric acid ==> ammonium sulphate

2NH_3(aq) + H₂SO_4(aq) ==> (NH₄)₂SO_4(aq)

AND

ammonia + nitric acid ==> ammonium nitrate

NH_3(aq) + HNO_3(aq) ==> NH₄NO_3(aq)

• These equations are sometimes written in terms of the fictitious 'ammonium hydroxide' (shown below). The  above equations are however, more correct! Quite simply, we are dealing with an aqueous solution of ammonia NH_3(aq), but NH₄OH is used in some textbooks! Only about 2% of the dissolved ammonia forms ammonium and hydroxide ions (more on this on Extra Aqueous Chemistry).

  • ammonium hydroxide + sulphuric acid ==> ammonium sulphate + water

    • 2NH₄OH_(aq) + H₂SO_4(aq) ==> (NH₄)₂SO_4(aq) + 2H₂O_(l)

  • ammonium hydroxide + nitric acid ==> ammonium nitrate + water

    • NH₄OH_(aq) + HNO_3(aq) ==> NH₄NO_3(aq) + H₂O_(l)

• The salt Ammonium chloride is used in zinc-carbon dry cell batteries. The slightly acid paste, made from the salt, slowly reacts with the zinc to provide the electrical energy from the chemical reaction.

  • ammonia + hydrochloric acid ==> ammonium chloride

  • NH_3(aq) + HCl_(aq) ==> NH₄Cl_(aq) 

    • or NH₄OH_(aq) + HCl_(aq) ==> NH₄Cl_(aq) + H₂O_(l)

• If ammonium salts are mixed with sodium hydroxide solution, free ammonia is formed (detected by smell and damp red litmus turning blue).

  • e.g. ammonium chloride + sodium hydroxide ==> sodium chloride + water + ammonia

  • NH₄Cl + NaOH ==> NaCl + H₂O + NH₃

• Ammonium sulphate or nitrate salts are widely used as 'artificial or synthetic fertilisers (preparation reactions above). There are several advantages to using artificial fertilisers in the absence of sufficient manure-silage etc. e.g. relatively cheap mass production, easily used to make poor soils fertile or quickly enrich multi-cropped fields.

• Artificial fertilisers are important to agriculture and used on fields to increase crop yields but they should be applied in a balanced manner (see 'problems' in section 5.).

  • Fertilisers usually contain compounds of three essential elements for healthy and productive plant growth to increase crop yield. They replace nutrient minerals used by a previous crop or enriches poor soil and more nitrogen gets converted into plant protein. 

    • Nitrogen (N) e.g. from ammonium or nitrate salts like ammonium sulphate, ammonium sulphate or ammonium phosphate or urea (e.g. look for the N in the formula of ammonium salts)

    • Phosphorus (P) e.g. from potassium phosphate or ammonium phosphate

    • Potassium (K) e.g. from potassium phosphate, potassium sulphate.

    • The fertiliser is marked with an 'NPK' value, i.e. the nitrogen : phosphorus : potassium ratio

  • Fertilisers must be soluble in water to be taken in by plant roots.

• Overuse of ammonia fertilisers on fields can cause major environmental problems as well as being uneconomic.
• Ammonium salts are water soluble and get washed into the groundwater, rivers and streams by rain contaminating them with ammonium ions and nitrate ions.
• This contamination causes several problems.
• Excess fertilisers in streams and rivers cause eutrophication.
  • Overuse of fertilisers results in appreciable amounts of them dissolving in rain water.
  • This increases levels of nitrate or phosphate in rivers and lakes.
  • This causes 'algal bloom' i.e. too much rapid growth of water plants on the surface where the sunlight is the strongest.
  • This prevents light from reaching plants lower in the water.
  • These lower plants decay and the active aerobic bacteria use up any dissolved oxygen.
  • This means any microorganisms or higher life forms relying on oxygen cannot respire.
  • All the eco-cycles are affected and fish and other respiring aquatic animals die.
  • The river or stream becomes 'dead' below the surface as all the food webs are disrupted.
• Nitrates are potentially carcinogenic (cancer or tumour forming).
  • The presence in drinking water is a health hazard.
  • Rivers and lakes can be used as initial sources for domestic water supply.
  • You cannot easily remove the nitrate from the water, it costs too much!
  • So levels of nitrate are carefully monitored in our water supply.
• More on water pollution on the Extra Aqueous Chemistry page acid rain on Oil Products page.

6. The Nitrogen Cycle for the gaseous element N_2(g)

• Nitrogen is an extremely important element for all plant or animal life! It is found in important molecules such as amino acids, which are combined to form proteins. Protein is used everywhere in living organisms from muscle structure in animals to enzymes in plants/animals.
• Nitrogen from the atmosphere:
  • Nitrifying bacteria, e.g. in the root nodules of certain plants like peas/beans (the legumes), can directly convert atmospheric nitrogen into nitrogen compounds in plants e.g. nitrogen => ammonia => nitrates which plants can absorb.
    • However, most plants can't do this conversion from nitrogen => ammonia, though they can all absorb nitrates, so the 'conversion' or 'fixing' ability might be introduced into other plant species by genetic engineering.
  • The nitrogen from air is converted into ammonia in the chemical industry, and from this artificial fertilisers are manufactured to add to nutrient deficient soils. However, some of the fertiliser is washed out of the soil and can cause pollution (see above).
  • The energy of lightning causes nitrogen and oxygen to combine and form nitrogen oxides which dissolve in rain that falls on the soil adding to its nitrogen content.
    1. N_2(g) + O_2(g) ==> 2NO_(g), then 
    2. then 2NO_(g) + O_2(g) ==> 2NO_2(g) 
    3. NO_2(g) + water ==> nitrates_(aq) in rain/soil
    4. Incidentally, reactions 1. and 2. can also happen in a car engine, and NO₂ is acidic and adds to the polluting acidity of rain as well as providing nutrients for plants!
• Nitrogen recycling apart from the atmosphere:
  • Nitrogen compounds, e.g. protein formed in plants or animals, are consumed by animals higher up the food chain and then bacterial and fungal decomposers break down animal waste and dead plants/animals to release nitrogen nutrient compounds into the soil (e.g. in manure/compost) which can then be re-taken up by plants. 
• Nitrogen returned to the atmosphere:
  • However, denitrifying bacteria will break down proteins completely and release nitrogen gas into the atmosphere.

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