Formulae - CIE iGCSE Chemistry Notes - New Syllabus
Formulae Notes for iGCSE
Core Syllabus
- State the formulae of the elements and compounds named in the subject content
- Define the molecular formula of a compound as the number and type of different atoms in one molecule
- Deduce the formula of a simple compound from the relative numbers of atoms present in a model or a diagrammatic representation
- Construct word equations and symbol equations to show how reactants form products, including state symbols
Supplement Syllabus
- Define the empirical formula of a compound as the simplest whole number ratio of the different atoms or ions in a compound
- Deduce the formula of an ionic compound from the relative numbers of the ions present in a model or a diagrammatic representation or from the charges on the ions
- Construct symbol equations with state symbols, including ionic equations
- Deduce the symbol equation with state symbols for a chemical reaction, given relevant information
Formulae of elements and compounds
Formulae of elements and compounds
Chemical Formula
A chemical formula shows the types and numbers of atoms in a substance. It uses element symbols and numbers to represent how many atoms of each element are present.
1. Formulae of Elements:
Some elements exist as single atoms (monatomic), while others exist as molecules (diatomic or larger).
- Monatomic elements: These exist as single atoms in their natural state.
Examples: He, Ne, Ar (all noble gases), metals like Na, Fe, Al - Diatomic elements: These exist as molecules made of two atoms of the same element.
Examples: \( \text{H}_2, \text{O}_2, \text{N}_2, \text{Cl}_2, \text{F}_2, \text{Br}_2, \text{I}_2 \) - Other molecular elements: Some elements form larger molecules.
Examples: Sulfur is \( \text{S}_8 \), Phosphorus is \( \text{P}_4 \)
List of Common Elemental Formulae:
- Hydrogen – \( \text{H}_2 \)
- Oxygen – \( \text{O}_2 \)
- Nitrogen – \( \text{N}_2 \)
- Chlorine – \( \text{Cl}_2 \)
- Bromine – \( \text{Br}_2 \)
- Iodine – \( \text{I}_2 \)
- Fluorine – \( \text{F}_2 \)
- Sulfur – \( \text{S}_8 \)
- Phosphorus – \( \text{P}_4 \)
- Metals – Just the symbol: e.g. Sodium = Na, Iron = Fe, Zinc = Zn
2. Formulae of Common Compounds:
The formula of a compound depends on the elements involved and their valencies (combining powers).
- Water: \( \text{H}_2\text{O} \)
- Carbon dioxide: \( \text{CO}_2 \)
- Ammonia: \( \text{NH}_3 \)
- Methane: \( \text{CH}_4 \)
- Sodium chloride: \( \text{NaCl} \)
- Hydrochloric acid: \( \text{HCl} \)
- Sulfuric acid: \( \text{H}_2\text{SO}_4 \)
- Nitric acid: \( \text{HNO}_3 \)
- Sodium hydroxide: \( \text{NaOH} \)
- Calcium carbonate: \( \text{CaCO}_3 \)
- Copper(II) sulfate: \( \text{CuSO}_4 \)
- Magnesium oxide: \( \text{MgO} \)
- Zinc chloride: \( \text{ZnCl}_2 \)
Important Tips:
- The number of atoms is shown by a subscript (e.g. the “2” in \( \text{H}_2 \)).
- If there is only one atom, no number is shown (e.g. NaCl, not Na1Cl1).
- Use parentheses when multiple polyatomic ions are needed (e.g. \( \text{Ca(OH)}_2 \)).
Example
Write the correct formula for sulfur dioxide and magnesium chloride.
▶️Answer/Explanation
Sulfur dioxide: One sulfur atom + two oxygen atoms = \( \text{SO}_2 \)
Magnesium chloride: Magnesium ion is \( \text{Mg}^{2+} \), chloride is \( \text{Cl}^- \). To balance the charges, we need 2 chloride ions:
Formula = \( \text{MgCl}_2 \)
Molecular formula of a compound as the number and type of different atoms in one molecule
The molecular formula of a compound shows the actual number and types of atoms present in a single molecule of that compound. It tells you:
- Which elements are present
- How many atoms of each element are in one molecule
Key Characteristics:
- It applies only to covalent compounds (which form molecules)
- It is written using element symbols and numerical subscripts
- The order of elements usually follows a standard (e.g. carbon before hydrogen in organic compounds)
Examples of Molecular Formulae:
- Water (H2O): 2 hydrogen atoms and 1 oxygen atom per molecule
- Carbon dioxide (CO2): 1 carbon atom and 2 oxygen atoms
- Ammonia (NH3): 1 nitrogen atom and 3 hydrogen atoms
- Methane (CH4): 1 carbon atom and 4 hydrogen atoms
- Ethene (C2H4): 2 carbon atoms and 4 hydrogen atoms
- Glucose (C6H12O6): 6 carbon, 12 hydrogen, and 6 oxygen atoms
Comparison with Empirical Formula:
The molecular formula is different from the empirical formula, which shows the simplest whole-number ratio of atoms in a compound.
Example:
For glucose: – Molecular formula = \( \text{C}_6\text{H}_{12}\text{O}_6 \) – Empirical formula = \( \text{CH}_2\text{O} \) (simplest ratio 1:2:1)
Example
State the molecular formula of a compound that contains 3 carbon atoms and 8 hydrogen atoms per molecule.
▶️Answer/Explanation
There are 3 carbon atoms and 8 hydrogen atoms.
So, the molecular formula is: \( \text{C}_3\text{H}_8 \)
Deducing a Formula from a Model or Diagram
Deducing a Formula from a Model or Diagram
When given a diagram or model of a simple compound, you can determine its formula by counting how many atoms of each type are shown. Each atom is represented either as a colored sphere (in models) or as an atomic symbol with bonding in diagrams.
Steps:
- Identify the elements in the compound (e.g. carbon, hydrogen, oxygen).
- Count the number of atoms of each element in the molecule.
- Write the chemical formula by writing each element symbol with its corresponding number as a subscript.
Examples:
- Model shows: 1 carbon atom, 4 hydrogen atoms → \( \text{CH}_4 \)
- Model shows: 2 hydrogen atoms, 1 oxygen atom → \( \text{H}_2\text{O} \)
- Model shows: 1 nitrogen atom, 3 hydrogen atoms → \( \text{NH}_3 \)
- Model shows: 1 carbon atom, 2 oxygen atoms → \( \text{CO}_2 \)
Tips:
- Don’t write the number 1 as a subscript – it’s assumed (e.g. write CO, not C1O1).
- Learn common color codes in ball-and-stick models:
- Hydrogen = white
- Oxygen = red
- Carbon = black or grey
- Nitrogen = blue
- Chlorine = green
Example
A ball-and-stick model shows 2 carbon atoms, 6 hydrogen atoms, and 1 oxygen atom.
▶️Answer/Explanation
Count the atoms:
Carbon = 2
Hydrogen = 6
Oxygen = 1
So the formula is: \( \text{C}_2\text{H}_6\text{O} \)
Constructing word equations and symbol equations
Constructing word equations and symbol equations
Chemical Equation
A chemical equation represents a chemical reaction. It shows the reactants (starting substances) and the products (new substances formed).
There are two main types:
1. Word Equations
These describe the reaction using full names of chemicals.
Example:
Magnesium + hydrochloric acid → magnesium chloride + hydrogen
2. Symbol Equations
A symbol equation shows a chemical reaction using chemical formulae instead of words.
Example:
\( \text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2 \)
Balancing Symbol Equations:
The number of atoms of each element must be the same on both sides (conservation of mass).
- Only change the number in front of a formula (called a coefficient)
- Never change subscripts (numbers inside a formula)
Deducing a Symbol Equation for a chemical reaction
Step 1: Identify the Reactants and Products
Read the description carefully and pick out all substances involved. Look for keywords like:
- “reacts with” → helps you find the reactants
- “forms” or “produces” → shows the products
Step 2: Write the Chemical Formulae
Use your knowledge of element symbols, compound rules (ionic or covalent), and polyatomic ions to write each formula correctly.
Step 3: Add State Symbols
State symbols show the physical state of each substance:
- \( (s) \) for solids (e.g. metals, precipitates)
- \( (l) \) for pure liquids (e.g. water, bromine)
- \( (g) \) for gases (e.g. O₂, H₂, CO₂)
- \( (aq) \) for solutions (aqueous) (substances dissolved in water)
Example with state symbols:
\( \text{NaOH (aq)} + \text{HCl (aq)} \rightarrow \text{NaCl (aq)} + \text{H}_2\text{O (l)} \)
Step 4: Balance the Equation
Make sure the number of atoms of each element is the same on both sides. Do not change formulas – only add coefficients in front.
Tips:
- Watch out for acids, gases, and precipitates – they help you decide correct state symbols.
- Check solubility rules if unsure whether something is a solid or aqueous.
Example
“Calcium carbonate reacts with hydrochloric acid to form calcium chloride, water, and carbon dioxide.” Give symbol equation.
▶️Answer/Explanation
Step 1 – Identify reactants and products:
Reactants: calcium carbonate (CaCO₃), hydrochloric acid (HCl)
Products: calcium chloride (CaCl₂), water (H₂O), carbon dioxide (CO₂)
Step 2 – Write the unbalanced equation:
\( \text{CaCO}_3 + \text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2 \)
Step 3 – Balance the equation and add state symbols:
\( \text{CaCO}_3 (s) + 2\text{HCl} (aq) \rightarrow \text{CaCl}_2 (aq) + \text{H}_2\text{O} (l) + \text{CO}_2 (g) \)
Explanation:
Calcium carbonate reacts with 2 moles of hydrochloric acid to produce an aqueous salt (calcium chloride), water, and carbon dioxide gas. This is an example of a neutralisation and decomposition reaction.
Example
“Iron(III) oxide reacts with carbon monoxide to form iron and carbon dioxide.” Give symbol equation.
▶️Answer/Explanation
Reactants: Iron(III) oxide = Fe₂O₃, Carbon monoxide = CO
Products: Iron = Fe, Carbon dioxide = CO₂
Balanced symbol equation:
\( \text{Fe}_2\text{O}_3 (s) + 3\text{CO} (g) \rightarrow 2\text{Fe} (s) + 3\text{CO}_2 (g) \)
Empirical formula of a compound
Empirical formula of a compound
The empirical formula is the simplest whole-number ratio of atoms (or ions) of each element in a compound.
It does not show the actual number of atoms in a molecule – it shows only the ratio.
Empirical vs Molecular Formula:
- Molecular formula – actual number of atoms in a molecule
e.g. \( \text{C}_6\text{H}_{12}\text{O}_6 \) for glucose - Empirical formula – simplest ratio
e.g. glucose becomes \( \text{CH}_2\text{O} \)
Why use empirical formula?
- Used to represent ionic compounds (which don’t form molecules)
- Used when the molecular structure is unknown or too large
How to find the empirical formula from molecular formula:
- Find the greatest common factor for all subscripts
- Divide all subscripts by that number
Examples:
- \( \text{C}_2\text{H}_6 \) → divide by 2 → \( \text{CH}_3 \)
- \( \text{N}_2\text{O}_4 \) → divide by 2 → \( \text{NO}_2 \)
- \( \text{P}_4\text{O}_{10} \) → divide by 2 → \( \text{P}_2\text{O}_5 \)
Note: If the molecular formula is already in simplest form, the empirical formula is the same.
Example
What is the empirical formula of ethene, \( \text{C}_2\text{H}_4 \)?
▶️Answer/Explanation
Divide both subscripts by 2:
\( \text{C}_2\text{H}_4 \rightarrow \text{CH}_2 \)
Empirical formula = \( \text{CH}_2 \)
Ionic Compound and Equation
Deducing the Formula of an Ionic Compound
Ionic compounds are made of positive ions (cations) and negative ions (anions). Their formulas are written so that the total positive charge equals the total negative charge – the compound must be electrically neutral.
Step-by-step method using charges:
- Write the symbol and charge of the cation and anion.
- Balance the total charges by choosing the correct number of each ion.
- Write the formula without showing charges. Use subscripts to show how many ions are present.
Common ion charges to remember:
- Group 1 metals: +1 (e.g. Na⁺, K⁺)
- Group 2 metals: +2 (e.g. Mg²⁺, Ca²⁺)
- Group 17 non-metals: -1 (e.g. Cl⁻, Br⁻)
- Group 16 non-metals: -2 (e.g. O²⁻, S²⁻)
- Transition metals often form +2 or +3 ions (e.g. Fe²⁺, Fe³⁺)
- Some common polyatomic ions:
- Ammonium = NH₄⁺
- Sulfate = SO₄²⁻
- Nitrate = NO₃⁻
- Hydroxide = OH⁻
- Carbonate = CO₃²⁻
Example using charges:
- Sodium ion = Na⁺, Chloride ion = Cl⁻ → Charges balance 1:1 → Formula: \( \text{NaCl} \)
- Magnesium ion = Mg²⁺, Chloride ion = Cl⁻ → Need 2 Cl⁻ to balance 1 Mg²⁺ → Formula: \( \text{MgCl}_2 \)
- Aluminum ion = Al³⁺, Oxide ion = O²⁻ → LCM of 3 and 2 is 6 → Need 2 Al³⁺ and 3 O²⁻ → Formula: \( \text{Al}_2\text{O}_3 \)
If using a diagram or model:
Count how many positive and negative ions are shown. Use those numbers to write the formula.
Example
A diagram shows 1 calcium ion (Ca²⁺) and 2 nitrate ions (NO₃⁻). Deduce the formula.
▶️Answer/Explanation
Ca²⁺ = +2
Each NO₃⁻ = -1, and there are two → -2
Charges balance: +2 and -2
Formula = \( \text{Ca(NO}_3)_2 \)
Ionic Equation
An ionic equation only shows the ions that actually change in the reaction. Spectator ions (ions that remain unchanged) are left out.
Steps to write an ionic equation:
- Write a full balanced symbol equation with state symbols.
- Split all aqueous ionic compounds into ions.
- Cancel ions that appear unchanged on both sides (spectator ions).
→ 
Tips:
- Only aqueous ionic compounds can be split into ions.
- Solids, liquids, and gases are written as whole formulas in ionic equations.
Example
Write the ionic equation for the reaction between sodium hydroxide and hydrochloric acid.
▶️Answer/Explanation
Step 1 – Full equation:
\( \text{NaOH (aq)} + \text{HCl (aq)} \rightarrow \text{NaCl (aq)} + \text{H}_2\text{O (l)} \)
Step 2 – Write aqueous ionic compounds as ions:
\( \text{Na}^+ (aq) + \text{OH}^- (aq) + \text{H}^+ (aq) + \text{Cl}^- (aq) \rightarrow \text{Na}^+ (aq) + \text{Cl}^- (aq) + \text{H}_2\text{O (l)} \)
Step 3 – Cancel spectator ions (Na⁺ and Cl⁻):
\( \text{OH}^- (aq) + \text{H}^+ (aq) \rightarrow \text{H}_2\text{O (l)} \)
Final Ionic Equation:
\( \text{OH}^- (aq) + \text{H}^+ (aq) \rightarrow \text{H}_2\text{O (l)} \)
Example
Write the ionic equation for the reaction between barium nitrate and sodium sulfate, both aqueous.
▶️Answer/Explanation
Full equation:
\( \text{Ba(NO}_3)_2 (aq) + \text{Na}_2\text{SO}_4 (aq) \rightarrow \text{BaSO}_4 (s) + 2\text{NaNO}_3 (aq) \)
Ions involved:
\( \text{Ba}^{2+} (aq) + 2\text{NO}_3^- (aq) + 2\text{Na}^+ (aq) + \text{SO}_4^{2-} (aq) \rightarrow \text{BaSO}_4 (s) + 2\text{Na}^+ (aq) + 2\text{NO}_3^- (aq) \)
Cancel spectator ions (Na⁺, NO₃⁻):
Ionic equation:
\( \text{Ba}^{2+} (aq) + \text{SO}_4^{2-} (aq) \rightarrow \text{BaSO}_4 (s) \)