Writing Equations (Cambridge O Level Chemistry)

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
• An example of a word equation for neutralisation is: 

sodium hydroxide + hydrochloric acid  →  sodium chloride + water 

• The reactants are sodium hydroxide and hydrochloric acid
• The products are sodium chloride and water 

 Names of compounds

• For compounds consisting of 2 atoms:
  • If one is a metal and the other a non-metal, then the name of the metal atom comes first and the ending of the second atom is replaced by adding -ide
    • E.g. NaCl which contains sodium and chlorine thus becomes sodium chloride 
  • If both atoms are non-metals and one of those is hydrogen, then hydrogen comes first
    • E.g. Hydrogen and chlorine combined is called hydrogen chloride
• For other combinations of non-metals as a general rule, the element that has a lower group number comes first in the name
  • E.g. 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
    • E.g. KOH is potassium hydroxide, CaCO₃ is calcium carbonate

Writing and balancing chemical equations

• Chemical equations 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 non-metals must be written as diatomic molecules (i.e. molecules that contain two atoms): 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. 
  • Examples of chemical equations:
    • Acid-base neutralisation reaction:
      NaOH (aq) + HCl (aq)  ⟶ NaCl (aq) + H₂O (l) 
    • Redox reaction:
      2Fe₂O₃ (s) + 3C (s) ⟶ 4Fe (s) + 3CO₂ (g)
    • In each equation there are equal numbers of each atom on either side of the reaction arrow so the equations are balanced

• The best approach is to practice lot of examples of balancing equations
• By trial and error change the coefficients (multipliers) in front of the formulae, one by one checking the result on the other side
• Balance elements that appear on their own, last in the process

State symbols

• State symbols are written after each formula in chemical equations to show which physical state each substance is in
• Brackets are used and they are not usually subscripted although you may come across them written in this way
• Aqueous should remind you of the word 'aqua' and means the substance is dissolved in water
  • In other words it is a solution

The four state symbols show the physical state of substances at normal conditions

• Symbol equations should be included when writing chemical equations.
• An example of a reaction with state symbols is the reaction of copper carbonate with hydrochloric acid:

CuCO₃ (s) + 2HCl (aq) ⟶ CuCl₂ (aq) + CO₂ (g) + H₂O (l)

• Sometimes it can be hard to know what the correct state symbol is and we have to look for clues in the identity of substances in a reaction
• Generally, unless they are in a solution:
  • Metal compounds will always be solid, although there are a few exceptions
  • Ionic compounds will usually be solids
• Non-metal compounds could be solids, liquids or gases, so it depends on chemical structure
• Precipitates formed in solution count as solids
• In the worked examples above the final equations with the state symbols would be
  • 2Al (s) + 3CuO (s) ⟶ Al₂O₃ (s) + 3Cu (s)
  • MgO (s)  + 2HNO₃ (aq)  ⟶ Mg(NO₃)₂ (aq)  + H₂O (l)

Balancing Ionic Equations

• In aqueous solutions ionic compounds dissociate into their ions, meaning they separate into the component ions that formed them
• E.g. hydrochloric acid and potassium hydroxide dissociate as follows:

HCl (aq) →  H⁺ (aq) + Cl^-(aq) 

KOH (aq)  → K⁺ (aq)  + OH^- (aq) 

• 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

• For some reactions, you will not be given the unbalanced equation but you will be expected to use your knowledge learnt throughout the course to know or deduce the formula of compounds and then balance the equations