Formulae and Equations Revision Notes for OCR A-Level Chemistry A

Formulae and equations

Writing chemical formulae and constructing balanced chemical equations are essential skills in A-Level chemistry. This topic builds on your understanding of ionic bonding to show you how to write correct formulae for ionic compounds and how to represent chemical reactions using balanced equations.

Ionic charges and the periodic table

Understanding simple ions

Atoms can lose or gain electrons to achieve a stable electron configuration similar to the nearest noble gas (helium through to radon). This process creates charged particles called ions:

Metal atoms (found on the left side of the periodic table) lose electrons to form cations (positively charged ions)
Non-metal atoms (found on the right side of the periodic table) gain electrons to form anions (negatively charged ions)

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The key difference: metals form positive ions (cations) by losing electrons, while non-metals form negative ions (anions) by gaining electrons. This fundamental pattern underlies all ionic compound formation.

Predicting ionic charges

For many elements, you can work out the likely charge on the ion by looking at the element's position in the periodic table:

The diagram shows a clear pattern:

Group 1 metals form 1+ ions (e.g., $\text{Li}^{+}$ , $\text{Na}^{+}$ , $\text{K}^{+}$ )
Group 2 metals form 2+ ions (e.g., $\text{Mg}^{2+}$ , $\text{Ca}^{2+}$ , $\text{Sr}^{2+}$ )
Aluminium (Group 3) forms 3+ ions ( $\text{Al}^{3+}$ )
Group 5 non-metals form 3- ions (e.g., $\text{N}^{3-}$ , $\text{P}^{3-}$ )
Group 6 non-metals form 2- ions (e.g., $\text{O}^{2-}$ , $\text{S}^{2-}$ )
Group 7 non-metals (halogens) form 1- ions (e.g., $\text{F}^{-}$ , $\text{Cl}^{-}$ , $\text{Br}^{-}$ , $\text{I}^{-}$ )

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Exceptions to memorise: Zinc ( $\text{Zn}^{2+}$ ) and silver ( $\text{Ag}^{+}$ ) are important exceptions that don't follow the general transition metal pattern. You need to memorise these specific charges as they appear frequently in exam questions.

Transition metals and variable charges

Many transition metals can form ions with different charges. The ionic charge is indicated using Roman numerals in the name of the ion:

Copper forms two ions: copper(I) ( $\text{Cu}^{+}$ ) and copper(II) ( $\text{Cu}^{2+}$ )
Iron forms two ions: iron(II) ( $\text{Fe}^{2+}$ ) and iron(III) ( $\text{Fe}^{3+}$ )

The Roman numeral tells you the charge on the metal ion.

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When you see a compound name like "iron(III) oxide", the Roman numeral (III) immediately tells you that iron has a 3+ charge in this compound. This systematic naming convention helps you work out the correct formula.

Binary compounds

A binary compound contains exactly two different elements.

To name a binary ionic compound:

Write the name of the metal element first
Change the ending of the non-metal element's name to -ide
For ionic compounds, the metal ion always comes first in the formula

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Example: Sodium Oxide Formation

Sodium and oxygen combine to form sodium oxide ( $\text{Na}_{2}\text{O}$ )

Sodium = metal (comes first)
Oxygen → oxide (change ending to -ide)
Formula shows 2 sodium ions for every 1 oxide ion

Polyatomic ions

What are polyatomic ions?

Sometimes ions contain atoms of more than one element bonded together. These are called polyatomic ions. Unlike simple ions which consist of single atoms, polyatomic ions are groups of atoms that carry an overall charge.

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Memorisation Required: You need to memorise the names and formulae of all common polyatomic ions shown in the table. There are no shortcuts or patterns to help you work these out, so regular practice is essential. These ions appear frequently in exams and are fundamental to writing correct formulae.

Common polyatomic ions include:

1+ ions: Ammonium ( $\text{NH}_{4}^{+}$ )
1- ions: Hydroxide ( $\text{OH}^{-}$ ), nitrate ( $\text{NO}_{3}^{-}$ ), nitrite ( $\text{NO}_{2}^{-}$ ), hydrogencarbonate ( $\text{HCO}_{3}^{-}$ ), manganate(VII) or permanganate ( $\text{MnO}_{4}^{-}$ )
2- ions: Carbonate ( $\text{CO}_{3}^{2-}$ ), sulfate ( $\text{SO}_{4}^{2-}$ ), sulfite ( $\text{SO}_{3}^{2-}$ ), dichromate(VI) ( $\text{Cr}_{2}\text{O}_{7}^{2-}$ )
3- ions: Phosphate ( $\text{PO}_{4}^{3-}$ )

Writing formulae from ions

The charge balancing rule

An ionic compound contains both a cation and an anion. You can work out the correct formula by balancing the charges on each ion.

For a correct formula:

The overall charge must be zero
Sum of positive charges = sum of negative charges

This means you need to find the ratio of ions that makes the charges cancel out.

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Think of charge balancing like a see-saw that must be perfectly balanced. The positive charges on one side must exactly equal the negative charges on the other side. This is the fundamental rule for all ionic compounds.

Worked examples

Let's look at how to balance charges systematically:

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Worked Example 1: Zinc Chloride

Ions present: $\text{Zn}^{2+}$ and $\text{Cl}^{-}$
One zinc ion has a 2+ charge
Each chloride ion has a 1- charge
To balance: you need two chloride ions for every one zinc ion
Charge check: $(1 \times 2+) = (2 \times 1-)$ ✓
Formula: $\text{ZnCl}_{2}$

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Worked Example 2: Aluminium Sulfate

Ions present: $\text{Al}^{3+}$ and $\text{SO}_{4}^{2-}$
Each aluminium ion has a 3+ charge
Each sulfate ion has a 2- charge
To balance: you need 2 aluminium ions (total 6+) and 3 sulfate ions (total 6-)
Charge check: $(2 \times 3+) = (3 \times 2-)$ ✓
Formula: $\text{Al}_{2}(\text{SO}_{4})_{3}$

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Tip: The charges must balance through multiplication. Think of it like finding the lowest common multiple of the two charge numbers. For aluminium (3+) and sulfate (2-), the LCM of 3 and 2 is 6, so you need charges totaling 6+ and 6-.

Writing the formula correctly

When writing the final formula, follow these rules:

The number of each ion is shown as a subscript after the ion symbol or formula
The ionic charges are not written in the final completed formula
Brackets must be used if there is more than one polyatomic ion

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Example: Aluminium Sulfate Notation

Aluminium sulfate contains 2 $\text{Al}^{3+}$ ions and 3 $\text{SO}_{4}^{2-}$ ions, so the formula is $\text{Al}_{2}(\text{SO}_{4})_{3}$ .

Note the brackets around $\text{SO}_{4}$ because we need three sulfate ions, and the subscript 3 goes outside the bracket.

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Common Mistake: Writing magnesium hydroxide as $\text{MgOH}_{2}$ is incorrect because it means one oxygen and two hydrogens. The correct formula is $\text{Mg}(\text{OH})_{2}$ , which means two complete hydroxide ions. Always use brackets for multiple polyatomic ions!

Writing and balancing chemical equations

Representing elements in equations

In chemical equations, elements are shown using their chemical symbols, but there's an important exception:

Diatomic molecules: Some elements exist naturally as molecules containing two atoms bonded together. These are:

$\text{H}_{2}$ (hydrogen)
$\text{N}_{2}$ (nitrogen)
$\text{O}_{2}$ (oxygen)
$\text{F}_{2}$ (fluorine)
$\text{Cl}_{2}$ (chlorine)
$\text{Br}_{2}$ (bromine)
$\text{I}_{2}$ (iodine)

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Memory aid: Remember "BrINClHOF" or "Have No Fear Of Ice Cold Beer"

These mnemonics help you remember all seven diatomic elements that must always be written with a subscript 2 in equations.

The only other elements that exist as small molecules are phosphorus ( $\text{P}_{4}$ ) and sulfur ( $\text{S}_{8}$ ). However, sulfur is usually written simply as S in equations (otherwise you'd need to multiply everything by 8).

Representing compounds in equations

Covalent compounds exist as molecules with a fixed number of atoms bonded together. In equations, you use the molecular formula (e.g., $\text{CO}_{2}$ , $\text{H}_{2}\text{O}$ ).

Ionic compounds don't exist as molecules - they form giant lattice structures. In equations, you use the formula unit worked out from the ionic charges (e.g., $\text{NaCl}$ , $\text{Al}_{2}\text{O}_{3}$ ).

State symbols in equations

State symbols are written in brackets after each formula to indicate the physical state:

(g) = gas
(l) = liquid
(s) = solid
(aq) = aqueous (dissolved in water)

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State symbols are an essential part of a complete chemical equation. Many exam questions will specifically ask you to include state symbols, and marks can be lost if they are omitted or incorrect.

Balancing chemical equations

A balanced equation shows the same number of atoms of each element on both sides, representing the conservation of mass during a chemical reaction.

Key rules for balancing:

Never change a chemical formula - the formulae are fixed by the rules above
Balancing numbers go in front of formulae (on the line, not as subscripts)
Multiply each formula by a balancing number until the atoms of each element match on both sides
The equation is balanced when atom counts are equal on both sides

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Critical Rule: Never change the subscripts in a chemical formula when balancing. For example, if you need more oxygen atoms, write $2\text{O}_{2}$ not $\text{O}_{4}$ . Changing subscripts changes the identity of the substance completely!

Worked example: constructing a balanced equation

Reaction: Aluminium reacts with oxygen to form aluminium oxide

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Worked Example: Balancing Aluminium Combustion

Step 1: Work out the formula of the product

Aluminium oxide contains $\text{Al}^{3+}$ and $\text{O}^{2-}$ ions
Balancing charges: 2 aluminium ions (6+) and 3 oxide ions (6-)
Formula: $\text{Al}_{2}\text{O}_{3}$

Step 2: Write an equation using formulae for all reactants and products

Aluminium is a metal element, so just Al
Oxygen exists as diatomic molecules, so $\text{O}_{2}$
Unbalanced equation: $\text{Al} + \text{O}_{2} \rightarrow \text{Al}_{2}\text{O}_{3}$

Step 3: Balance the equation

Start with the formulae of compounds and count atoms
Left side: 1 Al, 2 O
Right side: 2 Al, 3 O
The key is to get oxygen atoms equal on both sides
Try: $4\text{Al} + 3\text{O}_{2} \rightarrow 2\text{Al}_{2}\text{O}_{3}$

Check:

Left side: 4 Al and $3 \times 2 = 6$ O
Right side: $2 \times 2 = 4$ Al and $2 \times 3 = 6$ O ✓

Step 4: Add state symbols

$4\text{Al(s)} + 3\text{O}_{2}\text{(g)} \rightarrow 2\text{Al}_{2}\text{O}_{3}\text{(s)}$

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Alternative approach: Some students find fractions easier initially. For the worked example, you could write:

$2\text{Al} + 1\frac{1}{2}\text{O}_{2} \rightarrow \text{Al}_{2}\text{O}_{3}$

Then multiply through by 2 to remove the fraction, giving the same final answer. This method works well when dealing with diatomic elements.

Balancing formulae with brackets

Take extra care when balancing formulae containing brackets - the balancing number multiplies the entire formula.

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Example: Counting Atoms in Zinc Nitrate

$3\text{Zn}(\text{NO}_{3})_{2}$ means $3 \times \text{Zn}(\text{NO}_{3})_{2}$

This gives:

3 Zn atoms
$3 \times 2 = 6$ N atoms
$3 \times 3 \times 2 = 18$ O atoms

The 3 multiplies everything, and the subscript 2 outside the bracket multiplies everything inside the bracket (one N and three O atoms per nitrate ion).

Remember!

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Key Points to Remember:

Metal atoms lose electrons to form positive ions (cations); non-metal atoms gain electrons to form negative ions (anions)
You can predict many ionic charges from the element's group in the periodic table, but transition metals form ions with variable charges shown using Roman numerals
Polyatomic ions are groups of atoms bonded together with an overall charge - you must memorise the common ones
To write ionic formulae correctly, balance the charges so the sum of positive charges equals the sum of negative charges, giving an overall charge of zero
When balancing equations, never change the chemical formulae - only add balancing numbers in front of formulae
Seven elements exist as diatomic molecules ( $\text{H}_{2}$ , $\text{N}_{2}$ , $\text{O}_{2}$ , $\text{F}_{2}$ , $\text{Cl}_{2}$ , $\text{Br}_{2}$ , $\text{I}_{2}$ ) and must be written this way in equations

Exam focus checklist

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Self-Assessment Questions:

Can you predict ionic charges from the periodic table position?
Can you recall the formulae of all common polyatomic ions?
Can you balance ionic charges to write correct formulae?
Do you remember to use brackets when there is more than one polyatomic ion?
Can you write and balance equations for unfamiliar reactions?
Do you always include state symbols in equations?
Can you recognise and correctly represent diatomic elements?

Formulae and Equations (OCR A-Level Chemistry A): Revision Notes