Compounds, formulae and equations — OCR A-Level Chemistry
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Compounds, formulae and equations explained
Compounds, formulae and equations covers 5 specification points, including Formulae and equations and (a) the writing of formulae of ionic compounds….
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Formulae and equations
This section heading introduces the working tools of quantitative chemistry: writing formulae and balancing equations. Read it as a study guide rather than a single testable fact. You should learn to build the formula of an ionic compound from the charges on its ions, so that the total positive charge equals the total negative charge; for example, calcium ions (2⁺) and chloride ions (1⁻) combine as CaCl₂. You should also learn to construct and balance symbol equations, adding coefficients so that every element and every charge is conserved, and to write state symbols where they are needed. Treat the heading as a map: it tells you that formulae and equations are linked skills, and that accuracy in one supports accuracy in the other.
(a) the writing of formulae of ionic compounds from ionic charges, including:
To write an ionic formula, start from the charge on each ion and combine them so the total positive charge equals the total negative charge. A group 1 ion such as Na⁺ pairs one-to-one with a 1⁻ ion, giving NaCl. A 2⁺ ion such as Mg²⁺ needs two 1⁻ ions, giving MgCl₂, while a 2⁺ ion with a 2⁻ ion such as O²⁻ pairs one-to-one, giving MgO. When both charges are equal in magnitude, the ions combine in a 1:1 ratio. Use the lowest whole-number ratio, and place polyatomic ions such as sulfate (SO₄²⁻) or nitrate (NO₃⁻) in brackets when more than one is needed, for example Al₂(SO₄)₃. Check that the charges cancel before you finalise the formula.
(i) prediction of ionic charge from the position of an element in the periodic table
Ions form when atoms gain or lose electrons to reach a stable noble-gas configuration. For main-group elements the charge of the common ion follows directly from the group number. Group 1 metals lose one electron and form 1⁺ ions, such as Na⁺. Group 2 metals lose two electrons and form 2⁺ ions, such as Mg²⁺. Group 3 metals such as Al lose three electrons and form Al³⁺. Non-metals gain electrons: Group 5 elements form 3⁻ ions such as N³⁻, Group 6 elements form 2⁻ ions such as O²⁻, and Group 7 elements form 1⁻ ions such as Cl⁻. Group 4 elements are not usually described this way. To predict a charge, locate the element's group, decide whether it is metal or non-metal, and count electrons gained or lost to reach the nearest noble-gas shell. Transition metals are variable and are not predicted this way.
(ii) recall of the names and formulae for the following ions: NO₃⁻, CO₃²⁻, SO₄²⁻, OH⁻, NH₄⁺, Zn²⁺ and Ag⁺
You must know both the name and the formula of seven common ions. The nitrate ion is NO₃⁻, the carbonate ion is CO₃²⁻, the sulfate ion is SO₄²⁻ and the hydroxide ion is OH⁻; all four are negative. The ammonium ion is NH₄⁺, a positive ion made of nitrogen and hydrogen. Zinc forms the Zn²⁺ ion and silver forms the Ag⁺ ion. Learn each formula with its charge attached, because the charge controls how ions combine in a compound. For example, calcium ions (2⁺) balance carbonate ions (2⁻) one-to-one to give CaCO₃, while sodium ions (1⁺) need two to balance one carbonate, giving Na₂CO₃. Practise writing each formula from its name and each name from its formula until recall is automatic.
(b) construction of balanced chemical equations (including ionic equations), including state symbols, for reactions studied and for unfamiliar reactions given appropriate information.
This criterion asks you to build equations, not merely to read them. For a familiar reaction you write the correct formulae for every reactant and product, then adjust coefficients so that each element and each charge balances. For an unfamiliar reaction you use the information given, such as the names or formulae of the species involved, to deduce the products and then balance. Ionic equations show only the species that change: you split soluble ionic compounds into their ions, keep insoluble solids, gases and water together, then cancel spectator ions. State symbols (s), (l), (g) and (aq) must match the physical state under the stated conditions. Check both atom balance and, in ionic equations, charge balance.
Your focus
- Identify the sub-skills covered by the formulae and equations section and plan study around them.
- Construct the formula of an ionic compound so that total positive and negative charges balance.
- Balance a symbol equation by adjusting coefficients while keeping formulae unchanged.
Show all 15 objectives
- Deduce the formula of an ionic compound from the charges on its ions.
- Apply the lowest whole-number ratio when combining cations and anions.
- Use brackets correctly when a polyatomic ion appears more than once in a formula.
- Predict the charge on the common ion of a main-group element from its group number
- Distinguish positive ions formed by metals from negative ions formed by non-metals
- Write ionic charges using correct notation with the number before the sign
- Recall the name and formula of each of the seven specified ions
- State the charge on each specified ion correctly
- Use these ions to construct the formula of a compound from its constituent ions
- Construct a balanced full equation for a familiar reaction, including state symbols
- Derive an ionic equation by splitting soluble ionic compounds and cancelling spectator ions
- Construct a balanced equation for an unfamiliar reaction from the information supplied
Compounds, formulae and equations flashcards
Card 1 of 11
Group 1 metals
Compounds, formulae and equations exam tips
Marking Points
- Balances the total positive charge against the total negative charge so the compound is electrically neutral.
- Uses the lowest whole-number ratio of ions when combining a cation with an anion.
- Places a polyatomic ion in brackets when more than one unit is required in the formula.
- Applies the correct charge to each ion, including group numbers and common polyatomic ions, before combining them.
- Identifies the group number of the element from its position in the periodic table
- States that metals lose electrons to form positive ions and non-metals gain electrons to form negative ions
- Applies the pattern that Group 1 gives 1⁺, Group 2 gives 2⁺ and Group 3 gives 3⁺ ions
- Applies the pattern that Group 5 gives 3⁻, Group 6 gives 2⁻ and Group 7 gives 1⁻ ions
- Writes the charge with the number before the sign, for example Mg²⁺ and O²⁻
- Recognises that transition metals do not follow a single predictable charge
- Recalls NO₃⁻ as the nitrate ion
- Recalls CO₃²⁻ as the carbonate ion
- Recalls SO₄²⁻ as the sulfate ion
- Recalls OH⁻ as the hydroxide ion
- Recalls NH₄⁺ as the ammonium ion
- Recalls Zn²⁺ as the zinc ion and Ag⁺ as the silver ion
- Matches each formula to its correct name and charge without mixing them up
- Writes correct formulae for all reactants and products before adjusting any coefficients
- Balances each element in turn and confirms the final atom count on both sides
- Splits soluble ionic compounds into their constituent ions when constructing an ionic equation
- Cancels spectator ions so that only the species that change appear in the ionic equation
- Adds state symbols that match the physical state of each species under the stated conditions
- Checks that total charge balances on both sides of an ionic equation
Examiner Tips
- 💡Use the heading to plan revision: list the sub-skills (ionic formulae from charges, balancing equations, state symbols) and check each one separately.
- 💡When you meet a worked example, cover the answer and rebuild it yourself, then compare your method step by step rather than only the final result.
- 💡Keep a running check that every equation you write conserves both atoms and charge, because examiners reward correct, balanced equations.
- 💡Write the charge above each ion first, then find the ratio that makes the charges cancel before writing the final formula.
- 💡Simplify subscripts to the lowest whole-number ratio, since a formula such as Mg₂O₂ should be reduced to MgO.
- 💡Check polyatomic ions carefully: count how many are needed and add brackets when the number is greater than one.
- 💡Read the question to see whether it asks for the group, the ion charge or the ion symbol, and answer exactly that.
- 💡Check the sign and the size of the charge against the group number before selecting an option.
- 💡Eliminate options that give transition metals a single fixed charge, since these are usually wrong.
- 💡Learn the seven ions as name–formula–charge triplets so recall works in either direction.
- 💡Check the charge on each ion before combining ions in a formula, since the ratio depends on it.
- 💡Use the number of oxygen atoms and the overall charge to distinguish similar-looking ions.
- 💡Write the full balanced equation first, then derive the ionic equation by splitting and cancelling
- 💡Underline each element in turn and count atoms on both sides before moving on
- 💡For unfamiliar reactions, list the species named in the question and predict products from the information given
- 💡Recheck state symbols against any conditions stated in the question stem
Common Mistakes
- Reads the heading as a single fact to memorise; instead, treat it as a guide to two linked skills, formula writing and equation balancing, and practise both.
- Assumes a formula can be written without checking charge balance; instead, confirm that total positive charge equals total negative charge before finalising any ionic formula.
- Balances an equation by changing subscript numbers inside formulae; instead, adjust only the coefficients placed in front of each species so the formulae stay correct.
- Writes a formula without checking charge balance; instead, confirm that total positive charge equals total negative charge.
- Uses a ratio that is not the lowest whole-number ratio; instead, simplify the subscripts to their smallest whole numbers.
- Omits brackets around a repeated polyatomic ion; instead, enclose the polyatomic ion in brackets when more than one is needed.
- Writes the charge with the sign before the number; place the number first, as in Ca²⁺
- Assumes every metal forms a 1⁺ ion; match the charge to the group, so calcium in Group 2 forms Ca²⁺
- Predicts a fixed charge for a transition metal such as iron; treat transition-metal charges as variable
- Gives a Group 7 element a positive charge; halogens gain one electron to form 1⁻ ions such as Br⁻
- Writes the carbonate ion without its 2⁻ charge; include the charge, as in CO₃²⁻
- Gives the ammonium ion a negative charge; ammonium is positive, NH₄⁺
- Confuses the sulfate and sulfite formulae; the required ion here is sulfate, SO₄²⁻
- Writes the zinc ion as Zn⁺; zinc forms Zn²⁺
- Changes a subscript inside a formula to balance atoms; adjust only the coefficients in front of formulae
- Leaves spectator ions in an ionic equation; cancel ions that appear unchanged on both sides
- Omits or misassigns state symbols; match each symbol to the actual physical state described
- Balances atoms but ignores charge in an ionic equation; confirm the charge total on each side