Chemical change, or reaction, is said to occur when new substances are formed.
As each substance is unique in terms of its arrangement and types of particles, chemical reactions must involve some rearrangement of the fundamental particles that make up a substance.
This chemical change may be observed by a change in colour or appearance of the material, the evolution of a gas or some other factor that provides evidence that something is happening to the very nature of the substance in question.
Equations representing chemical reactions are called chemical equations.
A chemical reaction is described according to the substance or substances that are present at the beginning - the reactants, and the substances that are present having been formed in the process at the end of the reaction - the products.
Traditionally, the transition in time between the reactants and the products is shown using an arrow. The arrow is NOT an equals sign, but rather an indication that the process takes place over a period of time.
Word equations show the names of the reacting species on the left hand side, and the products of the reaction on the right hand side, both sides linked by an arrow to indicate that change has occurred.
The reactants and the products are described by their names. These equations give no indication as to the relative masses or numbers of moles reacting and produced and, consequently, they are of limited, or no use in calculations.
Sodium hydroxide + nitric acid sodium nitrate + water
sulfuric acid + sodium chloride sodium hydrogen sulfate + hydrogen chloride
Calcium hydroxide + ammonium chloride ammonia + calcium chloride + water
Although not specifically mentioned in the Syllabus, it is extremely useful to be familiar with many of the common types of reactions.
Reaction types - Inorganic
Chemical reactions are defined according to the nature of the reactants, products, or the process occuring.
The basic reactions of inorganic chemistry can be summarised as:
This is not intended to be an exhaustive list there are many other, less common, reaction types such as for example, disproportionation.
Two simple substances, usually elements, combine to form a more complex compound. For example, the reaction between iron and sulfur.
Example: Synthesis of iron(II) sulfide
iron + sulfur iron(II) sulfide
The term may also be applied to more complex reactions which are designed to manufacture a specific compound, such as a pharmaceutical. In organic chemistry, for example, we may talk about aspirin synthesis:
Example: Synthesis of aspirin
ethanoic anhydride + 2-hydroxy benzene carboxylic acid aspirin + ethanoic acid
Synthesis is a rather general term that can be applied to almost any chemical change and as such, not very meaningful.
Example: Synthesis of ammonia
nitrogen + hydrogen ammonia
Decomposition is when a compound is broken down, usually by heat, to produce simpler substances. An example is the thermal decomposition of calcium carbonate, producing calcium oxide and carbon dioxide.
Example: Thermal decomposition of calcium carbonate
calcium carbonate calcium oxide + carbon dioxide
Definition of thermal stability
Stability to heat is a relative issue. Enough heat will break even the strongest chemical bonds. Definitions of thermal stability usually depend on the method of heating being used. In general, a chemistry laboratory's heat supply is provided by a Bunsen burner. If a compound is not decomposed by the hottest bunsen flame then we can say that it is thermally stable.
An example of the weakness of this definition is the case of iron(II) sulfate (green vitriol). This is generally considered to be thermally stable. However, this substance was known in the Middle Ages and used to make sulfuric acid (oil of vitriol) by heating strongly until decomposition. This produced sulfur trioxide (sulfur(VI) oxide) fumes. These fumes were lead into cold water giving the acid.
FeSO4 FeO + SO3
SO3 + H2O H2SO4
The following table shows the accepted thermal stabilities of some common inorganic compounds.
Thermal stability of common inorganic compounds
|Type of anion||Type of metal ion||Effect of heat|
|Oxide (O2-)||K,Na,Ca,Mg,Al,Zn,Fe,Pb,Cu||Stable (no decomposition)|
|Ag2O||decomposes to metal|
|Ca,Mg,Al,Zn,Fe,Pb,Cu||decompose metal oxide + water|
|Ca,Mg,Al,Zn,Fe,Pb,Cu||decompose metal oxide + CO2|
|sulfate (SO42-)||K,Na,Ca,Mg,Al,Zn,Fe,Pb,Cu||All stable|
|Chloride (Cl-)||K,Na,Ca,Mg,Al,Zn,Fe,Pb,Cu||All stable|
|Nitrate (NO3-)||K,Na||decompose nitrite + oxygen|
|Ca,Mg,Al,Zn,Fe,Pb,Cu||decompose oxide + nitrogen dioxide + oxygen|
Hydrated compounds (compounds that contain water of crystallisation) also decompose when heated, losing some, or all, of their associated water molecules.
Example: Thermal decomposition of hydrated magnesium sulfate
magnesium sulfate heptahydrate Anhydrous magnesium sulfate + water
Some ionic compounds may undergo reaction between their water of crystallisation and the ions present, complicating this process. This is the case for magnesium chloride crystals, which decompose on heating to give a basic chloride salt containing chloride ions and hydroxide ions. This is not on the Syllabus.
Neutralisation is possibly the most common type of reaction encountered in schools. It is easy to carry out, there are many examples from everyday sources and it is easy to demonstrate using pH or litmus indicator papers. Acids occur in many foods such as fruits (lemons, oranges, apples, etc) vegetables (rhubarb), vinegar, etc.. Stomach acid is a reasonably strong solution of hydrochloric acid.
Neutralisation is the reaction of an acid and a base producing a salt and water. For example, the neutralisation of sulfuric acid by sodium hydroxide.
Example: Neutralisation of sulfuric acid by sodium hydroxide
sulfuric acid + sodium hydroxide sodium sulfate + water
Neutralisation can be extended to any compound that react with acids neutralising them. Such substances are called bases.
- Metal hydroxides
- Metal oxides
- Metal carbonates
- Metal sulfites
- Ammonia solution (ammonium hydroxide)
Metal hydroxides react with acids producing the corresponding salt and water.
Example: Neutralisation of hydrochloric acid by sodium hydroxide
hydrochloric acid + sodium hydroxide sodium chloride + water
Metal oxides also react with acids producing the corresponding salt and water.
Example: Neutralisation of hydrochloric acid by calcium oxide
hydrochloric acid + calcium oxide calcium chloride + water
Metal carbonates react with acids producing the corresponding salt, water and carbon dioxide.
Example: Neutralisation of hydrochloric acid by sodium carbonate
hydrochloric acid + sodium carbonate sodium chloride + water + carbon dioxide
Metal sulfites (sulfate(IV) ) react with acids producing the corresponding salt, water and sulfur dioxide (sulfur(IV) oxide).
Example: Neutralisation of hydrochloric acid by sodium sulfite (sulfate IV)
hydrochloric acid + sodium sulfite sodium chloride + water + sulfur dioxide
Nothing to do with the weather, but it does have a link in the idea of matter 'falling'.
Rain and snow falls from the sky because the atmosphere cannot support the mass of particles accumulating together. Substances 'fall out' of solution because the water cannot dissolve them.
Precipitation (called double decomposition in old text books) is a reaction where two ionic solutions are mixed together, bringing ions that form an insoluble substance in contact with one another. Once the insoluble ionic compound is formed it falls out of the solution as a precipitate.
positive ion (aq) + negative ion (aq) insoluble salt precipitate (s)
Much use is made of this type of reaction in 'wet' analysis of inorganic ions, such as the use of silver nitrate solution to test for the presence of chloride ions. A white precipitate of silver chloride appears from the solution of the suspected chloride on addition of silver nitrate solution.
Example: Precipitation of silver chloride
silver nitrate(aq) + sodium chloride(aq) silver chloride(s) + sodium nitrate(aq)
Likewise, the presence of sulfate ions can be detected by the addition of barium ions in the form of the chloride or nitrate. On this occasion, a white precipitate betrays the presence of sulfate ions.
Example: Precipitation of barium sulfate
iron(II) sulfate(aq) + barium chloride(aq) barium sulfate(s) + iron(II) chloride(aq)
Note: Other ions may interfere with the above tests and their presence must be previously discarded by prior tests. For example, carbonate ions also gives a white precipitate with barium ions. This may be prevented by the addition of a small quantity of hydrochloric acid, in conjuction with the barium chloride, to eliminate any carbonate ions present, by neutralisation.
How do I know which ionic compounds are insoluble?
The short answer is - you don't! The solubility of ionic compounds can't be predicted, you have to learn which ones are, and aren't, soluble. The good news is that the majority of ionic compounds are fairly soluble, so it makes sense to learn the insoluble ones.
| In reality, there is no such thing as soluble or insoluble,
rather that everything has different degrees of solubility. Even so-called
insoluble substances dissolve to a certain degree.
Barium sulfate is generally said to be insoluble, but some does dissolve in water at room temperature, albeit only a tiny amount. The solubility is given by a concept known as the solubility product, ksp.
Similarly a soluble substance may dissolve a large mass per litre or not much at all. Solubility curves show the change in solubility with temperature and these are different for each different substance.
The following table gives a brief outline of the insoluble common ionic compounds.
|Type of anion||Cation||Solubility|
|Oxide (O2-)||K+,Na+,Ca2+,Mg2+||React with water making metal hydroxides|
|Hydroxide (OH-)||K+,Na+||Very soluble|
|Ca2+,Mg2+||Partially or slightly soluble|
|Chloride (Cl-)||K+,Na+,Ca2+,Mg2+,Al3+,Zn2+,Fe2+,Cu2+||All soluble|
|Nitrate (NO3-)||K+,Na+,Ca2+,Mg2+,Al3+,Zn2+,Fe2+,Pb2+,Cu2+,Ag+||All nitrates are soluble|
Facts to remember
- All sodium, potassium and ammonium (NH4+) compounds are soluble.
- Silver and lead compounds are nearly always insoluble, the exceptions being nitrates and ethanoates because...
- All nitrates and ethanoates (CH3COO-) are soluble.
Lead(II) chloride, bromide and iodide are soluble in HOT water, but not in cold water.
Lead(II) iodide is bright yellow and is sometimes used to produce a sparkley gold crystal effect as the lead iodide deposits in the form of golden crystals from a hot solution as it cools down.
Displacement reactions occur in several areas of chemistry.
- Displacement of a weak (or volatile) acid from its salt by a stronger acid.
- Displacement of a weak (or volatile) base from its salt by a stronger base.
- Displacement of a metal ion from its compound by a more reactive metal (this is also a redox reaction)
- Displacement of hydrogen gas from water or acids by reactive metals (this is also redox)
Example: Displacement of nitric acid from potassium nitrate
potassium nitrate + sulfuric acid potassium hydrogen sulfate + nitric acid
Example: Displacement of ammonia from ammonium chloride
ammonium chloride + calcium hydroxide calcium chloride + ammonia + water
Example: Displacement of copper from copper(II) sulfate solution
copper(II) sulfate + zinc zinc sulfate + copper
Example: Displacement of hydrogen gas from water by calcium metal
calcium + water calcium hydroxide + hydrogen
Redox is short for Reduction and Oxidation.
These are reactions that involve a transfer of electrons from one species (the reducing agent) to another species (the oxidising agent). The terms used are wonderfully confusing as the reducing agent gets oxidised and the oxidising agent gets reduced during the course of the reaction. Extraction of metals falls into this category, as the metal ions in the metal ore need electrons to become metal atoms. For example, the following reaction occurs in the extraction of iron from its ore, haematite.
Example: The extraction of iron (blast furnace)
In this equation the iron 3+ ions get reduced by gaining three electrons from the carbon to become iron atoms.
For a full treatment of reduction and oxidation see 'Colourful Solutions 8 - Reduction and Oxidation'.
Reaction types - Organic
In organic chemistry there are also many different types of reaction. How many of the following reaction types do you know? Hover your mouse over each type for a definition.
And more specific terms for some of the general terms above as applied to other systems such as:
Specific addition reactions:
Specific elimination reactions
Specific condensation type reactions
These organic reaction terms go beyond the scope of this volume and are dealt with in Colourful Solutions 9 - 'Organic Chemistry'