Writing Equations

Writing Word Equations & Balanced Equations

Word equations

• These show the reactants and products of a chemical reaction using their full chemical names
• The arrow (which is spoken as “goes to” or “produces”) implies the conversion of reactants into products
• Reaction conditions or the name of a catalyst can be written above the arrow

Names of compounds

For compounds consisting of 2 atoms:

• If one is a metal and the other a nonmetal, then the name of the metal atom comes first and the ending of the second atom is replaced by adding –ide
  • Eg. NaCl which contains sodium and chlorine thus becomes sodium chloride
• If both atoms are nonmetals and one of those is hydrogen, then hydrogen comes first
  • Eg. hydrogen and chlorine combined is called hydrogen chloride
• For other combinations of nonmetals as a general rule, the element that has a lower Group number comes first in the name
  • Eg. carbon and oxygen combine to form CO₂ which is carbon dioxide since carbon is in Group 4 and oxygen in Group 6

For compounds that contain certain groups of atoms:

• There are common groups of atoms which occur regularly in chemistry
• Examples include the carbonate ion(CO₃^2-), sulfate ion (SO₄^2-), hydroxide ion (OH^–) and the nitrate ion (NO₃^–)
• When these ions form a compound with a metal atom, the name of the metal comes first
• Eg. KOH is potassium hydroxide, CaCO₃ is calcium carbonate

Writing and balancing chemical equations

• These use the chemical symbols of each reactant and product
• When balancing equations, there needs to be the same number of atoms of each element on either side of the equation
• The following nonmetals must be written as molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂
• Work across the equation from left to right, checking one element after another
• If there is a group of atoms, for example a nitrate group (NO₃^–) that has not changed from one side to the other, then count the whole group as one entity rather than counting the individual atoms. For example:
  • NaOH + HCl → NaCl + H₂O
• There are equal numbers of each atom on either side of the reaction arrow so the equation is balanced

Extended Only

Equations with State Symbols & Deducing Balanced Equations

Using state symbols:

State symbols are written after formulae in chemical equations to show which physical state each substance is in:

Example 1

Aluminium (s)  +   Copper (II) Oxide (s)  →

Aluminium Oxide (s)  +   Copper (s)

Unbalanced symbol equation:    

Al     +     CuO     →     Al2O3     +     Cu

ALUMINIUM: There is 1 aluminium atom on the left and 2 on the right so if you end up with 2, you must start with 2. To achieve this, it must be 2Al

2Al     +     CuO     →     Al2O3     +     Cu

OXYGEN: There is 1 oxygen atom on the left and 3 on the right so if you end up with 3, you must start with 3. To achieve this, it must be 3CuO

2Al     +     3CuO     →     Al2O3     +     Cu

COPPER: There is 3 copper atoms on the left and 1 on the right. The only way of achieving 3 on the right is to have 3Cu

2Al     +     3CuO     →     Al2O3     +     3Cu

Example 2

Magnesium Oxide (s)  +  Nitric Acid (aq)  →

Magnesium Nitrate (aq)  +   Water (l)

Unbalanced symbol equation:  

MgO   +   HNO3   →   Mg(NO3)2   +   H2O

MAGNESIUM: There is 1 magnesium atom on the left and 1 on the right so there are equal numbers of magnesium atoms on both sides so these are kept the same

MgO   +   HNO3   →   Mg(NO3)2   +   H2O

OXYGEN: There is 1 oxygen atom on the left and 1 on the right so there is an equal number of oxygen atoms on both sides. It is therefore kept the same (remember that you are counting the nitrate group as a separate group, so do not count the oxygen atoms in this group)

MgO   +   HNO3   →   Mg(NO3)2   +   H2O

HYDROGEN: There is 1 hydrogen atom on the left and 2 on the right. Therefore you must change HNO₃ to 2HNO₃

MgO   +   2HNO3   →   Mg(NO3)2   +   H2O

Balancing ionic equations

• In aqueous solutions ionic compounds dissociate into their ions, meaning they separate into the component atoms or ions that formed them
• g. hydrochloric acid and potassium hydroxide dissociate as follows:
  • HCl →H⁺ + Cl^–
  • KOH → K⁺ + OH^–
• It is important that you can recognise common ionic compounds and their constituent ions
• These include:
  • Acids such as HCl and H₂SO₄
  • Group I and Group II hydroxides e.g. sodium hydroxide
  • Soluble salts e.g. potassium sulfate, sodium chloride
• Follow the example below to write ionic equations

Example

Write the ionic equation for the reaction of aqueous chlorine and aqueous potassium iodide.

Step 1:   Write out the full balanced the equation:     

2KI(aq)     +     Cl2(aq)     →

2KCl(aq)    +     I2(aq)

Step 2:   Identify the ionic substances and write down the ions separately:

2K+ (aq) + 2I-(aq)     +     Cl2(aq)     →

2K+(aq) + 2Cl-(aq)    +     I2(aq)

Step 3:   Rewrite the equation eliminating the ions which appear on both sides of the equation (spectator ions ) which in this case are the K⁺ ions:

2I-(aq)     +     Cl2(aq)     →

2Cl-(aq)    +     I2(aq)

Exam Tip

When balancing equations you cannot change any of the formulae, only the amount of each atom or molecule. This is done by changing the numbers that go in front of each chemical species.

You need to be able to identify the products which are not ions in ionic equations. These are usually molecules such as water or bromine but they may also be precipitated solids.