Formulae Revision Notes

Each chemical element is represented by a unique symbol, usually derived from its English or Latin name. The symbol consists of one or two letters, with the first letter always capitalised and the second letter, if present, in lowercase. For example, hydrogen is represented as H, oxygen as O, sodium as Na (from Latin 'Natrium'), and aluminium as Al.

These symbols are universal and used internationally to identify elements in chemical formulas and equations. Memorising the symbols of elements listed in the syllabus is essential for recognising substances in chemistry.

• Symbols are one or two letters, e.g., H, O, Na, Al.
• First letter capitalised, second letter lowercase if present.
• Symbols are derived from English or Latin names.
• Used to represent elements in formulas and equations.

Compounds consist of two or more different elements chemically combined in fixed proportions. Their chemical formulas show which elements are present and the number of atoms of each element in one molecule or formula unit of the compound.

Formulas are written by placing the symbols of the constituent elements side by side. Subscripts (small numbers written to the lower right of a symbol) indicate the number of atoms of that element in the compound. If there is only one atom of an element, no subscript is written.

For example, water is H₂O, indicating two hydrogen atoms and one oxygen atom per molecule. Carbon dioxide is CO₂, with one carbon atom and two oxygen atoms per molecule. Sodium chloride is NaCl, showing one sodium ion and one chloride ion per formula unit.

• Compounds combine element symbols to show composition.
• Subscripts indicate the number of atoms of each element.
• No subscript means one atom of that element.
• Examples: H₂O, CO₂, NaCl.

It is important to memorise the chemical formulas of common elements and compounds specified in the syllabus. Below are examples with their correct formulas:

Elements: Hydrogen (H₂), Oxygen (O₂), Nitrogen (N₂), Chlorine (Cl₂), Sulfur (S₈), Iron (Fe), Copper (Cu), Aluminium (Al).

Compounds: Water (H₂O), Carbon dioxide (CO₂), Ammonia (NH₃), Methane (CH₄), Sodium chloride (NaCl), Calcium carbonate (CaCO₃), Sulfuric acid (H₂SO₄), Ethanol (C₂H₅OH).

• Elements may exist as molecules (e.g., O₂, N₂) or single atoms (e.g., Fe).
• Compounds have fixed formulas showing element ratios.
• Memorise formulas for syllabus-listed substances.
• Examples include H₂O, CO₂, NH₃, NaCl, CaCO₃, H₂SO₄.

When writing chemical formulas, place the symbols of elements in the order given by convention or syllabus lists. Use subscripts to indicate the number of atoms. For polyatomic ions or groups, brackets may be used if more than one group is present, e.g., calcium carbonate is CaCO₃ (one calcium ion, one carbonate ion).

Interpreting a formula involves identifying each element and the number of atoms present. For example, in H₂SO₄, there are 2 hydrogen atoms, 1 sulfur atom, and 4 oxygen atoms.

Understanding formulas allows prediction of the composition and properties of substances, and is essential for writing balanced chemical equations.

• Write element symbols with correct capitalization.
• Use subscripts to show atom numbers.
• Use brackets for multiple polyatomic groups if needed.
• Interpret formulas by counting atoms of each element.
• Example: H₂SO₄ has 2 H, 1 S, 4 O atoms.

Water is a compound made of hydrogen and oxygen. Each water molecule contains two hydrogen atoms and one oxygen atom.

Step 1: Identify the symbols for hydrogen (H) and oxygen (O).

Step 2: Determine the number of atoms of each element in one molecule: 2 hydrogen atoms, 1 oxygen atom. This is known from experimental data and molecular structure studies, but for this objective, the formula is given or memorised from the syllabus list.

• Step 1: Hydrogen symbol is H; oxygen symbol is O.
• Step 2: Two hydrogen atoms and one oxygen atom per molecule.
• Step 3: Write formula as H₂O, with subscript 2 for hydrogen.
• This formula shows the composition of water molecules.

A molecular formula shows the exact number of atoms of each element present in one molecule of a compound. It provides detailed information about the composition of the molecule, unlike empirical formulas which only give the simplest ratio of atoms.

For example, the molecular formula of water is H₂O, indicating that each molecule contains two hydrogen atoms and one oxygen atom.

• Specifies types of atoms in a molecule.
• Indicates the exact number of each atom in one molecule.
• Represents the actual composition of the molecule.

Each element in a molecular formula is represented by its chemical symbol, followed by a subscript number indicating how many atoms of that element are present in one molecule. If no subscript is shown, it means there is only one atom of that element.

For example, in glucose, C₆H₁₂O₆, there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms in each molecule.

• Subscripts show the number of atoms of each element.
• No subscript means one atom of that element.
• Example: C₆H₁₂O₆ means 6 C, 12 H, 6 O atoms per molecule.

The molecular formula gives the actual number of atoms in a molecule, while the empirical formula gives the simplest whole-number ratio of atoms.

For example, hydrogen peroxide has the molecular formula H₂O₂ but the empirical formula HO, showing the simplest ratio of hydrogen to oxygen atoms is 1:1.

Understanding this difference is important for correctly identifying compounds and their properties.

• Molecular formula = actual number of atoms.
• Empirical formula = simplest ratio of atoms.
• Example: H₂O₂ (molecular) vs HO (empirical).

Consider the molecular formula ethanol, C₂H₆O. To find the number of atoms:

Carbon (C): 2 atoms per molecule, from the subscript 2.

Hydrogen (H): 6 atoms per molecule, from the subscript 6, not just 1 or 6 molecules of hydrogen atoms separately but 6 atoms bonded within the molecule itself. Oxygen (O): 1 atom per molecule, as no subscript means one atom.

• C₂H₆O means 2 carbon, 6 hydrogen, 1 oxygen atoms per molecule.
• Subscripts directly indicate atom counts.
• No subscript means one atom of that element.

Molecular formulas are used in balanced chemical equations to show the reactants and products involved in a reaction.

For example, the combustion of methane is represented as:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) This shows one molecule of methane (CH₄) reacts with two molecules of oxygen (O₂) to produce one molecule of carbon dioxide (CO₂) and two molecules of water (H₂O).

• Use molecular formulas to write balanced equations.
• Subscripts indicate atom numbers in molecules.
• Coefficients indicate molecule numbers in reactions.

A chemical formula represents the types and numbers of atoms in a compound. It shows which elements are present and in what proportions. For example, water has the formula H₂O, indicating two hydrogen atoms and one oxygen atom per molecule.

Chemical formulas are written using element symbols from the periodic table, with numbers (subscripts) to show how many atoms of each element are present. If only one atom of an element is present, no subscript is written.

Determining the formula of a compound involves analysing the relative quantities of atoms and expressing these as the simplest whole number ratio.

• Chemical formula = element symbols + subscripts showing atom counts.
• Subscripts indicate the number of atoms of each element.
• No subscript means one atom of that element.

Models or diagrams of compounds show atoms as spheres or symbols, often in different colours to represent different elements. For example, a model of water might show two small white spheres (hydrogen) attached to one larger red sphere (oxygen).

To determine the formula, first identify the different types of atoms by their colour or label. Common colour conventions include: hydrogen (white), oxygen (red), carbon (black), nitrogen (blue), chlorine (green).

Count the number of atoms of each type shown in the model or diagram carefully.

• Use colour or labels to identify atom types.
• Count each type of atom accurately.
• Ignore bonds for counting atoms; focus on atom numbers.

Once the numbers of each type of atom are counted, express these numbers as a ratio. For example, if a model shows 4 carbon atoms and 10 hydrogen atoms, the ratio is 4:10.

Simplify the ratio to the smallest whole numbers by dividing both numbers by their greatest common divisor. For 4:10, divide both by 2 to get 2:5.

This simplest ratio corresponds to the empirical formula of the compound, which is the formula with the smallest whole number ratio of atoms.

• Write atom counts as a ratio.
• Divide both numbers by the greatest common divisor.
• Use the simplest whole number ratio for the formula.

Write the chemical formula by placing the element symbols side by side, usually with the more metallic element first or according to standard conventions.

Add subscripts to each symbol to show the number of atoms from the simplest ratio. For example, C₂H₅ means 2 carbon atoms and 5 hydrogen atoms.

If the ratio is 1 for an element, do not write the subscript '1'. For example, H₂O means 2 hydrogen atoms and 1 oxygen atom.

• Write element symbols in correct order.
• Add subscripts from simplest ratio.
• Omit subscript '1' for single atoms.

A model shows 3 atoms of nitrogen (blue spheres) and 9 atoms of hydrogen (white spheres).

Step 1: Identify atoms — nitrogen (N) and hydrogen (H).

Step 2: Count atoms — N = 3, H = 9. Write ratio 3:9. Step 3: Simplify ratio by dividing both by 3: 3 ÷ 3 = 1, 9 ÷ 3 = 3, so ratio is 1:3. Step 4: Write formula using symbols and subscripts: NH₃. This is the formula for ammonia, showing one nitrogen atom bonded to three hydrogen atoms.

• Identify atom types and count atoms.
• Write ratio of atoms.
• Simplify ratio to smallest whole numbers.
• Write formula with element symbols and subscripts.

A chemical reaction involves the transformation of reactants into products. To represent this change, chemists use chemical equations. These equations provide a concise way to describe what substances react and what substances are formed.

There are two main types of chemical equations: word equations and symbol equations. Word equations use the names of substances, while symbol equations use chemical formulas and symbols to represent the substances involved.

Including physical states in equations is important because it gives additional information about the form of each substance during the reaction. The common physical state symbols are: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (a substance dissolved in water).

• Chemical reactions convert reactants into products.
• Word equations use names; symbol equations use formulas.
• Physical states provide context about the form of substances.

Word equations list the names of reactants on the left side and products on the right side, separated by an arrow (→) indicating the direction of the reaction.

For example, when hydrogen gas reacts with oxygen gas to form water, the word equation is:

Hydrogen + Oxygen → Water. This shows the substances involved but does not indicate quantities or physical states explicitly.

• Use correct chemical names for reactants and products.
• Separate reactants and products with an arrow (→).
• Word equations do not show quantities or states.

Symbol equations use chemical formulas to represent substances and must be balanced to obey the law of conservation of mass, meaning the number of atoms of each element is the same on both sides.

Physical state symbols are added in parentheses immediately after each formula to indicate the state of the substance.

For example, the reaction of hydrogen gas with oxygen gas to produce liquid water is written as: 2H₂(g) + O₂(g) → 2H₂O(l). This equation is balanced and shows that hydrogen and oxygen are gases, while water is a liquid.

• Use correct chemical formulas for all substances.
• Balance the equation so atoms are conserved.
• Add physical state symbols: (s), (l), (g), (aq).
• Example: 2H₂(g) + O₂(g) → 2H₂O(l)

Consider the reaction of magnesium metal with oxygen gas to form magnesium oxide. The word equation is:

Magnesium + Oxygen → Magnesium oxide. To write the symbol equation, first write the formulas: Mg + O₂ → MgO.

Next, balance the equation by ensuring the same number of atoms of each element on both sides. Oxygen appears as O₂ on the left and as O in MgO on the right, so balance oxygen by placing a coefficient 2 before MgO: Mg + O₂ → 2MgO. Now balance magnesium by placing 2 before Mg: 2Mg + O₂ → 2MgO. This is the balanced symbol equation with physical states: 2Mg(s) + O₂(g) → 2MgO(s).

• Write unbalanced symbol equation: Mg + O₂ → MgO.
• Balance oxygen atoms by adjusting coefficients: Mg + O₂ → 2MgO.
• Balance magnesium atoms: 2Mg + O₂ → 2MgO.
• Add physical states: (s) for solids, (g) for gases.

(s) Solid: A substance with a fixed shape and volume, e.g., magnesium metal, sodium chloride crystals.

(l) Liquid: A substance with fixed volume but no fixed shape, e.g., water at room temperature.

(g) Gas: A substance with neither fixed shape nor volume, e.g., oxygen gas, hydrogen gas, carbon dioxide gas at room temperature and pressure (RTP). Gas state is often indicated when a substance is a gas under RTP conditions or when it is evolved in a reaction (e.g., bubbles). (aq) Aqueous: A substance dissolved in water, forming a solution, e.g., sodium chloride dissolved in water is NaCl(aq). This is important to distinguish from pure solids or liquids.

• (s) = solid, fixed shape and volume.
• (l) = liquid, fixed volume, no fixed shape.
• (g) = gas, no fixed shape or volume.
• (aq) = aqueous solution, dissolved in water.

The empirical formula of a compound is the simplest whole-number ratio of the different atoms or ions present in the compound. It shows the relative number of each type of atom but does not indicate the actual number of atoms in a molecule or the structure of the compound.

For example, the empirical formula of hydrogen peroxide (H₂O₂) is HO, indicating a 1:1 ratio of hydrogen to oxygen atoms, even though each molecule contains two hydrogen and two oxygen atoms.

• Represents simplest ratio of atoms or ions.
• Does not show actual number of atoms in a molecule.
• Does not provide structural information.

The molecular formula shows the actual number of atoms of each element in a molecule of the compound, whereas the empirical formula shows only the simplest ratio.

For example, glucose has the molecular formula C₆H₁₂O₆ but the empirical formula CH₂O, reflecting the 1:2:1 ratio of carbon, hydrogen, and oxygen atoms.

Knowing the empirical formula is a step towards determining the molecular formula when the relative molecular mass is known.

• Molecular formula shows actual atom counts.
• Empirical formula shows simplest ratio.
• Empirical formula is foundational for molecular formula determination.

Empirical formulas can be calculated from experimental data such as mass percentages or masses of elements in a compound.

The general steps are: convert masses to moles by dividing by atomic masses, find the simplest mole ratio by dividing all mole values by the smallest mole number, and express the ratio in whole numbers.

If ratios are not whole numbers, multiply all by the smallest integer to obtain whole numbers.

• Convert element masses to moles: moles = mass / atomic mass.
• Divide all mole values by smallest mole value.
• Multiply to get whole-number ratios if necessary.

A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass. Calculate its empirical formula.

Step 1: Assume 100 g sample, so masses are 40.0 g C, 6.7 g H, 53.3 g O.

Step 2: Calculate moles: C: 40.0 g ÷ 12.0 g/mol = 3.33 mol; H: 6.7 g ÷ 1.0 g/mol = 6.7 mol; O: 53.3 g ÷ 16.0 g/mol = 3.33 mol; Step 3: Divide by smallest mole number (3.33): C: 3.33/3.33 = 1; H: 6.7/3.33 = 2.01 ≈ 2; O: 3.33/3.33 = 1; Step 4: Empirical formula is CH₂O.

• Use atomic masses: C=12.0, H=1.0, O=16.0 g/mol.
• Divide moles by smallest value to find ratio.
• Round to nearest whole number for formula.

Empirical formulas provide essential information about the elemental composition of compounds, which is critical in chemical analysis, synthesis, and identification.

They are particularly useful when the molecular formula is unknown or when compounds do not exist as discrete molecules, such as ionic compounds.

Empirical formulas are used in stoichiometric calculations and to understand the proportions of elements in reactions.

• Indicate elemental composition ratios.
• Useful for ionic and molecular compounds.
• Foundation for stoichiometry and chemical analysis.

Ionic compounds consist of positive ions (cations) and negative ions (anions) held together by strong electrostatic forces of attraction called ionic bonds.

Each ion has a specific charge, for example, sodium ion (Na⁺) has a +1 charge, and chloride ion (Cl⁻) has a -1 charge.

The overall compound must be electrically neutral, meaning the total positive charge must exactly balance the total negative charge. This principle is fundamental when determining the formula of an ionic compound.

• Ionic bond = electrostatic attraction between oppositely charged ions.
• Compound charge neutrality: total positive charge + total negative charge = 0.

The charge on an ion depends on how many electrons are lost or gained compared to the neutral atom.

Group 1 metals form +1 ions, Group 2 metals form +2 ions, and Group 3 metals form +3 ions.

Non-metals typically form negative ions with charges equal to 8 minus their group number (e.g., Group 7 elements form -1 ions).

• Na (Group 1) → Na⁺
• Mg (Group 2) → Mg²⁺
• Al (Group 3) → Al³⁺
• Cl (Group 7) → Cl⁻
• O (Group 6) → O²⁻

To write the formula of an ionic compound, balance the total positive and negative charges so the compound is neutral.

This often involves finding the lowest whole-number ratio of ions that balances the charges.

For example, aluminium ions (Al³⁺) and oxide ions (O²⁻) combine so that total positive charge equals total negative charge: 2 Al³⁺ ions give +6 charge, 3 O²⁻ ions give -6 charge, so formula is Al₂O₃.

• Write the ions with their charges.
• Balance total positive and negative charges.
• Use the smallest whole-number ratio of ions.
• Write the formula with cation first, then anion.

Magnesium forms Mg²⁺ ions and chlorine forms Cl⁻ ions.

To balance charges, two chloride ions are needed to balance one magnesium ion: total positive charge = +2, total negative charge = 2 × (-1) = -2.

Therefore, the formula is MgCl₂. This means one magnesium ion combines with two chloride ions to form a neutral compound.

• Mg²⁺ and Cl⁻ ions.
• Charge balance: +2 and 2 × (-1) = -2.
• Formula: MgCl₂.

Models or diagrams may show the relative numbers of ions in an ionic compound.

Count the number of each type of ion and write the formula using the simplest whole-number ratio.

If charges are given, use them to confirm the ratio balances the overall charge to zero.

• Count ions shown in the model.
• Write formula using simplest ratio.
• Confirm charge neutrality.

A chemical symbol equation represents a chemical reaction using chemical symbols and formulae for reactants and products. It shows how substances change during the reaction.

State symbols indicate the physical state of each substance: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (dissolved in water).

Including state symbols is important because the physical state affects reaction conditions and product properties. For example, sodium chloride dissolved in water is written as NaCl(aq), while solid sodium chloride is NaCl(s).

• (s) = solid, e.g. NaCl(s)
• (l) = liquid, e.g. H2O(l)
• (g) = gas, e.g. O2(g)
• (aq) = aqueous solution, e.g. Na+(aq), Cl−(aq)

To write a balanced symbol equation: first identify the correct chemical formulae of all reactants and products, then include their state symbols, and finally balance the number of atoms of each element on both sides.

Balancing ensures the law of conservation of mass is obeyed — atoms are neither created nor destroyed.

Example: The reaction of hydrogen gas with oxygen gas to form liquid water is written as: 2H2(g) + O2(g) → 2H2O(l)

• Identify correct formulae for reactants and products.
• Add state symbols based on physical states.
• Balance atoms by adjusting coefficients (numbers before formulae).

Ionic equations show the ions involved in reactions occurring in aqueous solutions. Ionic compounds dissolve in water to give their constituent ions, which are shown with charges and state symbol (aq).

In ionic equations, strong electrolytes are written as separate ions, while solids, liquids, and gases remain as formula units.

Example: When aqueous sodium chloride reacts with aqueous silver nitrate, a precipitate of silver chloride forms: NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq) can be written as the ionic equation: Na+(aq) + Cl−(aq) + Ag+(aq) + NO3−(aq) → AgCl(s) + Na+(aq) + NO3−(aq)

• Write soluble ionic compounds as separate ions with charges.
• Include state symbols, especially (aq) for ions in solution.
• Keep solids, liquids, and gases as formula units.
• Balance charges and atoms on both sides.

Zinc metal reacts with hydrochloric acid to produce zinc chloride and hydrogen gas.

Step 1: Write formulae and states: Zinc is solid (Zn(s)), hydrochloric acid is aqueous (HCl(aq)), zinc chloride is aqueous (ZnCl2(aq)), hydrogen is gas (H2(g)).

Step 2: Write unbalanced equation: Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g)

• Balance chlorine atoms: ZnCl2 has 2 Cl atoms, so need 2 HCl molecules.
• Balance hydrogen atoms: 2 HCl gives 2 H atoms, which form H2 gas.
• Balanced equation: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

Always include state symbols in symbol equations when required, as marks are often awarded for this.

Use correct chemical formulae, including charges for ions in ionic equations.

Balance equations fully, ensuring atom and charge balance in ionic equations, and atom balance in molecular equations.

• State symbols: (s), (l), (g), (aq)
• Correct formulae and charges for ions
• Balanced atoms and charges
• Show ions explicitly in ionic equations

A symbol equation represents a chemical reaction using chemical formulas for the reactants and products. It shows the substances involved and their physical states, providing a concise and precise description of the reaction.

Each chemical species is represented by its formula, such as H2 for hydrogen gas or NaCl for sodium chloride. The equation uses an arrow (→) to separate reactants on the left from products on the right.

State symbols indicate the physical state of each substance: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (dissolved in water). Including these symbols helps clarify reaction conditions and product forms.

• Symbol equation format: reactants → products
• Use chemical formulas, not names
• Include state symbols: (s), (l), (g), (aq)

To write a symbol equation, first identify the correct chemical formulas of all reactants and products. This may be given directly or inferred from a word equation or experimental description.

For example, if the reaction is between hydrochloric acid and sodium hydroxide, the formulas are HCl (aq) and NaOH (aq). The products are sodium chloride (NaCl) and water (H2O).

Ensure the formulas reflect the correct chemical species and their ionic or molecular nature.

• Use correct element symbols and subscripts
• Recognise common compounds and ions
• Distinguish between molecular and ionic formulas

Chemical equations must obey the law of conservation of mass: the number of atoms of each element must be the same on both sides.

Balance equations by adjusting coefficients (whole numbers placed before formulas) to equalise atom counts.

Never change the subscripts in formulas to balance equations, as this changes the substances involved.

• Count atoms of each element on both sides
• Adjust coefficients to balance atoms
• Do not alter chemical formulas

State symbols provide information about the physical form of each reactant and product at the reaction conditions.

Common state symbols are: (s) solid, (l) liquid, (g) gas, and (aq) aqueous solution.

For example, in the reaction of zinc metal with hydrochloric acid, zinc is solid (Zn(s)) and hydrochloric acid is aqueous (HCl(aq)). The products include hydrogen gas (H2(g)) and aqueous zinc chloride (ZnCl2(aq)).

• Identify physical states from context or experimental data
• Use correct symbols after each formula
• State symbols are enclosed in parentheses

Word equation: zinc + hydrochloric acid → zinc chloride + hydrogen

Step 1: Write formulas with states: zinc is solid (Zn(s)), hydrochloric acid is aqueous (HCl(aq)), zinc chloride is aqueous (ZnCl2(aq)), hydrogen is gas (H2(g)).

Step 2: Write unbalanced symbol equation: Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g)

• Step 3: Count atoms: Left: Zn=1, H=1, Cl=1; Right: Zn=1, Cl=2, H=2
• Step 4: Balance chlorine and hydrogen by placing coefficient 2 before HCl: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
• Step 5: Check atom balance: Zn=1, H=2, Cl=2 on both sides
• Final balanced equation with states: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)