• The molecular formula is the formula that shows the number and type of each atom in a molecule

  • E.g. the molecular formula of ethanoic acid is C₂H₄O₂

  
  

• The empirical formula is the simplest whole number ratio of the atoms of each element present in one molecule or formula unit of the compound

  • E.g. the empirical formula of ethanoic acid is CH₂O

  
  

• Organic molecules often have different empirical and molecular formulae

• The formula of an ionic compound is always an empirical formula

• It is calculated from knowledge of the ratio of masses of each element in the compound

Example: A compound that contains 10 g of hydrogen and 80 g of oxygen has an empirical formula of H₂O. This can be shown by the following calculations:

Amount of hydrogen atoms = mass in grams ÷ A_r of hydrogen = (10 ÷ 1) = 10 moles
Amount of oxygen atoms = mass in grams ÷ A_r of oxygen = (80 ÷ 16) = 5 moles

• Divide the relative formula mass of the molecular formula by the relative formula mass of the empirical formula

• Multiply the number of each element present in the empirical formula by this number to find the molecular formula

Nothing created - nothing destroyed

• New substances are made during chemical reactions

  • However, the same atoms are always present before and after reaction

  • They have just joined up in different ways

  • Atoms cannot be created or destroyed, so if they exist in the reactants then they absolutely must be in the products!

• Because of this the total mass of reactants is always equal to the total mass of products

• This idea is known as the Law of Conservation of Mass

Conservation of Mass

• The Law of Conservation of Mass enables us to balance chemical equations, since no atoms can be lost or created

• You should be able to:

  • Write word equations for reactions outlined in these notes

  • Write formulae and balanced chemical equations for the reactions in these notes

• These show the reactants and products of a chemical reaction using their full chemical names

• The reactants are those substances on the left-hand side of the arrow and can be thought of as the chemical ingredients of the reaction

• They react with each other and form new substances

• The products are the new substances which are on the right-hand side of the arrow

• 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 (a substance added to make a reaction go faster) can be written above the arrow

• An example is the reaction of sodium hydroxide (a base) and hydrochloric acid produces sodium chloride (common table salt) and water:

sodium hydroxide + hydrochloric acid ⟶ sodium chloride + water

• Chemical equations use the chemical symbols of each reactant and product

• When balancing equations, there has to be the same number of atoms of each element on either side of the equation in accordance with the Law of Conservation of Mass

• A symbol equation uses the formulae of the reactants and products to show what happens in a chemical reaction

• A symbol equation must be balanced to give the correct ratio of reactants and products:   

S + O₂ → SO₂

• This equation shows that one atom of sulfur (S) reacts with one  molecule of oxygen (O₂) to make one molecule of sulfur dioxide (SO₂)

• The following non-metals must be written as molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂

• To balance an equation you 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 (aq) + 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

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