ESO 2 Science 6

Chemical Changes

What to Learn

  • Physical changes vs. chemical changes.
  • Chemical equations. Reactants vs. products.
  • Conservation of mass in chemical reactions: the law of constant proportions; balancing chemical equations.
  • Energy changes during chemical reactions.
  • Simple chemical reactions: metals with non-metals, oxidation, combustion, formation of acids, formation of salts.

Key Information

Physical Changes

Physical changes are those in which the substances that make up a physical body remain the same when the body undergoes a change.

A physical change can be produced…

  • When you apply a force to an object, as when you crumble a piece of paper with your hand;
  • When you give or take energy to/from an object, as when you heat up or cool down water.

Some examples of physical changes are…

  • A change of state, such as the condensation of water;
  • A change of temperature, as the cooling down of food inside the fridge;
  • A change in the shape, as when you shatter mom's favourite china vase;
  • A change in the amount or type of energy in an object: so happens when you drop a stone to the floor.
Chemical Changes

Chemical changes (or chemical reactions) are those that transform some chemical/s into new other chemical/s.

They normally occur when different substances, prone to react between them, are put into contact. And so, when iron (Fe) and oxygen (O2) are put together, they will combine and disappear to form a new chemical: rust (Fe2O3). But the wine and the nitrogen (kept in the upper part of the bottles of wine) will never react, allowing the wine to preserve its properties over time. Not all chemicals will react when put into contact.

The initial molecules that react in a chemical reaction are called reactants, and the final resulting molecules are called products. In the previous reaction, the oxygen and the iron are the reactants, whereas the rust is the only product.

Some chemical changes imply a loss of energy, i.e., they are exothermic, because they release energy (usually heat) to the environment. Some other chemical changes are endothermic, they absorb energy from the environment, and the products will have a greater amount of energy than the reactants.

Nevertheless, most chemical reactions, either exothermic or endothermic, require an initial contribution of external energy to take place. If you want a piece of paper to combine with the atmospheric oxygen and go into combustion, you need to heat up the paper (what you usually do with a flame). But during the reaction, the paper releases a lot of energy (heat, light and even sound), so the reaction is exothermic: the paper and the oxygen had more energy than the ashes, the smoke, the CO2 and the water vapour produced.

Quite often, chemical changes can be noticed by some conspicuous events, such as…

  • A change in the temperature, as in any combustion;
  • A change in the colour, as in the rusting of iron;
  • The formation of bubbles, as when the calcite (CaCO3) reacts with hydrochloric acid (HCl);
  • A change in the volume, as in baking bread;
  • Etc.

Most chemical changes are irreversible: they can't be undone (e.g. a combustion), whereas most physical changes are reversible (e.g. the condensation of water).

Chemical Equations

Chemical equations are the way by which chemical reactions are represented. To the left you write the formulas of the reactants (either elements, molecules or crystals) and to the right, the formulas of the products. In between, you draw an arrow that represents the direction of the change: from the reactants to the products. In case of a reversible reaction, you must draw a double arrow.

One example would be the following:

  • 3 H2 + N2 → 2 NH3

The number of molecules of each type is represented with a coefficient to the left of the molecule. The number of atoms for each kind of molecule is represented by a subindex to the right of its chemical symbol (nothing when it's just one). Thus, you can see that in the equation above we have six atoms of hydrogen and two of nitrogen in each side. This is because in a chemical reaction, neither the chemical elements, nor the number of atoms of each chemical element change. What changes is the arrangement of those atoms because a chemical change actually consists of the breaking down of some chemical bonds and the formation of some new chemical bonds.

The following chemical equation does not represent any real chemical reaction, because there aren't exactly the same atoms in each side of the equation:

  • H2 + N2 → NH3

Also, as the atoms of the reactants have to be the same as the atoms of the products, the overall mass does not change during a chemical reaction.

Balancing a chemical equation means writing the right coefficients to the right of each molecular species, to ensure that the number of each type of atoms (and so, the overall mass) remains the same along the equation.

Chemical Reactions of Metals
Metals with oxygen and water

Metals react with oxygen and water depending on their reactivity. This means that some are highly reactive, some others are moderately reactive, and some others are very little reactive or even nothing at all. The following list shows the reactivity of some common metals: K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > Cu > Ag > Au > Pt.

This way, while potassium reacts violently with O2, iron forms rust (Fe2O3), silver slowly tarnishes (forming Ag2O), and gold does not react at all and preserves its properties forever.

Likewise, as combustions are oxidations boosted by heat, the behaviour of metals when they are heated varies: some burn (Mg) while others just rust (Cu).

Metals with acids

Metals react with acids forming hydrogen gas (H2) and a salt. This reaction consists of the displacement of the hydrogen of the gas by the metal, as this one is more reactive than the hydrogen.

One example is sodium reacting with hydrochloric acid to form sodium chloride (table salt) and hydrogen:

  • Na + HCl → NaCl + H2

This type of reactions can be tested by testing the formation of H2, which is a gas that is able to put out the flame of a match.

Displacement of metals by metals

As some metals are more reactive than others, the more reactive metals can displace the less reactive metals from their compounds. For instance: magnesium is more reactive than iron, and so it can displace the iron from the ferric oxide:

  • 3 Mg + Fe2O3 → 3 MgO + 2 Fe

For the same reason, platinum would never displace any other metal from its compounds, as it is the least reactive metal of all.

These type of reactions are useful to purify metals that are rarely found in pure state in nature. This is the case of iron, that usually occurs forming the minerals hematite (Fe2O3) and magnetite (Fe3O4). Aluminium can be used to displace iron from its oxides, and thus, purify it:

  • 2 Al + Fe2O3 → Al2O3 + 2 Fe
Acids, Alkalis and Salts
Acids and alkalis

Acids and alkalis are common everyday substances, such as acetic acid (in vinegar), citric acid (in lemon juice), caustic soda (lye, NaOH) or bicarbonate of soda (baking soda).

Acids and alkalis have a very different behaviour in solution: while acids release protons (H+), alkalis sequestrate the protons. The amount of free protons is measured with the pH scale: the greater the amount of free protons, the lower the pH, and the lower the amount of free protons, the higher the pH. Very acidic solutions have a pH of 1, whilst very alkaline solutions have a pH of 14. Neutral solutions and distilled water, neither acidic nor alkaline, have a pH of 7. Salts make neutral solutions; such is the case of a solution of table salt in water.

Indicators, such as litmus paper or a solution of phenolphthalein, can tell us whether a solution is acidic or alkaline: litmus paper or litmus solution will turn red in an acidic solution, and blue in an alkaline solution. Universal indicator is more precise, though: it is a mixture of dyes that can be present in paper strips or in solutions, and it yields a wide range of colours for the whole pH scale: from red (very acidic solutions) to green (neutral solutions) and to purple (very alkaline solutions).

Reactions between acids and bases

Alkalis are soluble substances that belong to a group of chemicals called bases. Bases react with acids producing a salt and water. These are called neutralisation reactions, as both the acid and the base "disappear" to form a salt, which gives a neutral solution

Some examples of everyday neutralisation reactions are the following:

  • Treating bee stings (acidic) with sodium bicarbonate (a base);
  • Treating wasp stings (alkaline) with vinegar (that contains acetic acid);
  • Upping the pH of an acidic soil with lime (CaOH, an alkali);
  • Neutralising the excess of HCl that your stomach produces during digestion with antacids such as baking soda (a base).

The reactions between acids and bases can produce a wide range of different salts. For example:

  • Salts of sulphuric acid are called sulphates:
  •     H2SO4 (sulphuric acid) + 2 NaOH (caustic soda) → Na2SO4 (sodium sulphate) + 2 H2O
  • Salts of nitric acid are called nitrates:
  •     HNO3 (nitric acid) + NaOH (caustic soda) → NaNO3 (sodium nitrate) + H2O
  • Salts of hydrochloric acid are called chlorides:
  •     2 HCl (hydrochloric acid) + CaCO3 (calcium carbonate) → CaCl2 (calcium chloride) + H2O + CO2


Chemical Reactions
Chemical party

Chemical party

Watch almost real chemical elements interacting live.




Exercises: Balance the following chemical equations
  1. C4H10 + O2 → CO2 + H2O
  2. H2O → O2 + H2
  3. H2S + O2 → SO2 + H2O
  4. HCl + Al → AlCl3 + H2
  5. NO → N2 + O2
  6. N2 + H2 → NH3
  7. Fe + O2 → Fe2O3
  8. KClO3 → KCl + O2
  9. Mg + Fe2O3 → MgO + Fe
  10. Na + Ag2O → Ag + Na2O
  11. CO + O2 → CO2
  12. C2H5OH + O2 → H2O + CO2
  13. SO + O2 → S2O3
  14. Mg + Ag2O → Ag + MgO
  15. Na + Fe2O3 → Na2O + Fe
  16. H2SO4 + Zn → ZnSO4 + H2
  17. Zn + HCl → ZnCl2 + H2
  18. CaCO3 → CaO + CO2
  19. N2O5 + H2O → HNO3
  20. HNO3 + NaOH → NaNO3 + H2O
  21. H2SO4 + KOH → K2SO4 + H2O
  22. H2CO3 + NaOH → Na2CO3 + H2O
  23. H2S + CaCO3 → CaS + H2O + CO2
  24. NaCl + AgNO3 → NaNO3 + AgCl
  25. HCl + CaCO3 → CaCl2 + H2O + CO2